Plus One Chemistry Chapter Wise Questions and Answers Chapter 3 Classification of Elements and Periodicity in Properties

Students can Download Chapter 3 Classification of Elements and Periodicity in Properties Questions and Answers, Plus One Chemistry Chapter Wise Questions and Answers helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Chemistry Chapter Wise Questions and Answers Chapter 3 Classification of Elements and Periodicity in Properties

Plus One Chemistry Classification of Elements and Periodicity in Properties One Mark Questions and Answers

Plus One Chemistry Chapter 3 Questions And Answers Question 1.
Which of the following is not a Dobereiner triad?
a) Cl, Br, I
b) Ca, Sr, Ba
c) Li, Na, K
d) Fe, Co, Ni
d) Fe, Co, Ni

Plus One Chemistry Chapter 3 Question 2.
The elements of s-block and p-block are collectively called ___________
Representative elements

Classification Of Elements And Periodicity In Properties Questions And Answers Question 3.
The cause of periodicity of properties is
b) Increasing atomic weights
c) Number of electrons in the valence shell
d) The recurrence of similar outer electronic configuration
d) The recurrence of similar outer electronic configuration

An online valence electrons calculator finds the abbreviated or condensed electron configuration of an element with these instructions.

Classification Of Elements And Periodicity In Properties Important Questions Question 4.
Halogen with highest ionization enthalpy is ___________ .
Fluorine

Classification Of Elements And Periodicity In Properties Question 5.
Which of the following represents the most electropositive element?
a) [He]2s1
b) [He]2s2
c) [Xe]6s1
d) [Xe]6s2
c) [Xe]6s1

Periodic Classification Of Elements Class 10 Extra Questions With Answers Question 6.
Second electron gain enthalpy is
always positive

Classification Of Elements And Periodicity In Properties Class 11 Pdf Question 7.
Correct order of polarising power is
a) Cs+ < K+ < Mg2+ < Al3+
b) Al3+ < Mg2 + K+ < Cs+
c) Mg2+ < Al3+ < K+ < Cs+
d) K+ < Cs+ < Mg2+ < Al3+
a) Cs+ < K+ < Mg2+ < Al3+

Chemistry Chapter 3 Class 11 Important Questions Question 8.
The IUPAC name of the element with atomic number is 109 is ___________
Une

Important Questions Of Periodic Classification Of Elements Class 11 Question 9.
The size of iso electronic species F~, Ne and Na+ is affected by
a) Nuclear charge
b) Principal quantum number.
c) Electron – electron interaction in outer orbitals.
d) None of the factors because their size is the same.
Nuclear charge as nuclear charge is high the size is small.

Chapter 3 Chemistry Class 11 Question 10.
In transition elements the differentiating electron occupies (n-1)d sublevel in preference to ______________
np level

Plus One Chemistry Classification of Elements and Periodicity in Properties Two Mark Questions and Answers

Important Questions For Class 11 Chemistry Chapter 3 Question 1.
The arguments made by two students are as given:
Student 1: ‘Hydrogen belongs to Group 1.’
Student 2: ‘Hydrogen belongs to Group 17.’
1. Who is right?
1. Nobody is right.

2. Hydrogen has a one s-electron and hence can be placed in group 1 (alkali metals). It can also • gain an electron to achieve a noble gas arrangement and hence it can behave similar to a group 17 (halogen family) element. Because it a special case, hydrogen is placed separately at the top of the Periodic Table.

Class 11 Chemistry Chapter 3 Important Questions With Answers Question 2.
Match the following:

 Sodium f-block Oxygen s-block Uranium d-block Silver p-block

Sodium – s-block
Oxygen – p-block
Uranium – f-block
Silver – d-block

Periodic Classification Of Elements Important Questions Question 3.

1. Which one has greater size, Na or K?

1. K

2. K comes below Na in the Periodic Table. The atomic size increases down the group due to the fact that the inner energy levels are filled with electrons, which serve to shield the outer electrons from the pull of the nucleus.

Questions On Periodic Classification Of Elements Class 11 Question 4.
The general characteristics of a particular block of elements is as given:
They are highly electropositive, soft metals. They are good reducing agents. They lose the outermost electron(s) readily to form 1+ ion of 2+ ion.

1. Which block has this general characteristics?
2. Write down two general characteristics of p-block.

1. s-block

2. The p-block contains metals, non-metals and metalloids. They form ionic as well as covalent compounds.

Class 11 Chemistry Chapter 3 Important Questions Question 5.
The variation in electron gain enthalpies of elements is less systematic than for ionization enthalpies. Out of oxygen and sulphur which has greater negative value for electron gain enthalpy? Justify.
Sulphur has greater negative value (-200 kJ mol1) for electron gain enthalpy compared to that of oxygen (-141 kJ mol1). This is because, due to smaller size of oxygen the added electron experiences much repulsion from the electrons present in the shell. Due to large size, the electron-electron repulsion is much less in sulphur.

Question 6.
Some elements are given. Li, Cs, Be, C, N, O F, I, Ne, Xe.

1. Arrange the above elements in the increasing order of ionization enthalpy.
2. Arrange the given elements in the decreasing order of negative electron gain enthalpy.

1. Cs < I < Li < Xe < Be < C < N < O < F < Ne
2. Cl > F > O > N > C > Be > Li > I >C s > Xe > Ne

Plus One Chemistry Classification of Elements and Periodicity in Properties Three Mark Questions and Answers

Question 1.
Analyse the given figure and answer the questions that follow.

1. What is meant by atomic radius?
3. Write down another two types of terms expressed as atomic size.

1. Atomic radius is defined as the distance from the centre of the nucleus of the atom to the outermost shell of electrons.
2. Covalent radius is defined as one half of the distance between the centre of nuclei of two similar atoms bonded by a single covalent bond.

Question 2.
Consider the statement: The element with 1s2
configuration belongs to the p-block.’

1. Identify the element.
2. Do you agree with this statement?
3. Justify.

1. Helium
2. Yes
3. Strictly, helium belongs to the s-block but its positioning in the p-block along with other group 18 elements is justified because it has a completely filled valence shell (1s²) and as a result, exhibits characteristic of other noble gases.

Question 3.
The properties of elements are a periodic function of
their atomic weights.

1. Who proposed this law?
2. Can you see anything wrong in this law? If yes, justify your answer.
3. State modem periodic law.

1. Mendeleev
2. Yes, atomic number is the more fundamental property of an element than atomic mass.
3. The physical and chemical properties of the elements are periodic functions of their atomic numbers.

Question 4.
1. Define ionisation enthalpy.
2. IE1 <IE2 <IE3
What is meant by this? Justify.
1. Ionisation enthalpy is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

2. It shows the increasing order the magnitude of successive ionization enthalpies. As the positive charge of the ion increases, it becomes more difficult to remove the valence electron due to increase in effective nuclear charge.

Question 5.
Say whether the following are true or false:

1. On moving across a period ionization enthalpy decreases.
2. Mg is biggerthan Cl.
3. Ionization enthalpy of Li is less than that of K.

1. False
2. True
3. False

Question 6.
Analyze the graph given below:

1. Identify the graph.
2. Account forthe following observations:
i) ‘Ne’ has the maximum value of ∆iH.
ii) In the graph from Be to B, ∆iH decrease.
1. Graph showing the variation of ionisation enthalpy (∆iH) with atomic number (Z) of the elements of second period.

2. i) ‘Ne’ is an inert gas and it has closed electron shell with stable octet electronic configuration. Hence it has the maximum ionisation enthalpy in the second period.
ii) Be has completely filled electronic configuration and a more stable. So ionization enthalpy is high

Question 7.
Study the graph and answer the questions that follow:

1. On moving down a group what happens to electron gain enthalpy?
2. Why chlorine shows more negative electron gain enthalpy than fluorine? ,

1. On moving down a group,electron gain enthalpy becomes less negative because the size of the atom increases and the added electron would be farther from the nucleus.

2. Due to small size of F, the added electron goes to the smaller 2p quantum level and suffers significant replusion from the other electrons present in this level. But due to big size of Cl, the added electron goes to the n = 3p quantum level and occupies a larger region of space and the electron-electron repulsion is much less.

Question 8.
Electron gain enthalpy is an important periodic property.
1. What is meant by electron gain enthalpy?
2. What are the factors affecting electron gain enthalpy?
3. How electron gain enthalpy varies on moving across a period? Justify.

1. Electron gain enthalpy(∆egH) is the enthalpy change accompanying the process of addition of an electron to a neutral gaseous atom to convert it into a negative ion.

2. Effective nuclear charge, atomic size, electronic configuration

3. Electron gain enthalpy becomes more negative with increase in the atomic number across a period. The effective nuclear charge increases from left to right across a period and consequently, it will be easier to add an electron to a smaller atom since the added electron on an average would be closer to the nucleus. Thus more energy is released.

Question 9.
1. The electron gain enthalpies of Be and Mg are positive. What is your opinion? Justify.
2. Electron gain enthalpies of nobles gases have large positive values. Why?
1. The statement is correct. Electron gain enthalpies of Be (+240 kJ/mol) and Mg (+230 kJ/mol) are positive because they have stable electronic configurations with fully filled s-orbitals (Be – 2s², Mg – 3s²). Hence, the gain of electron is highly endothermic.

2. Noble gases have stable octet electronic configuration of ns2 np6 (except He -1 s²) and they have practically no tendency to accept additional electron. They have large positive electron gain enthalpies because the electron has to enter the next higher principal quantum level leading to a very unstable electronic configuration. Hence, energy has to be supplied to add an extra electron.

Question 10.
During a group discussion, a student argued that ionization enthalpy depends only upon electronic configuration.
1. Do you agree? Comment.
2. Define shielding effect/screening effect.
3. Is there any relation between ionization enthalpy and shielding effect/screening effect? Explain.
1. No. In addition to electornic configuration, ionization enthalpy depends upon other factors like atomic size, nuclear charge and shielding effect/screening effect.

2. In multi-electron atoms, the nuclear charge experienced by the valence electron will be less than the actual charge on the nucleus because it is shielded by inner core of electrons. This is called shielding effect or screening effect.

3. Yes. Ionization enthalpy decreases with increase in shielding effect/screening effect of inner electrons. This is because as a result of shielding of the valence electron by the intervening core of electrons it experiences a net positive charge which is less than the actual charge of the nucleus. In general, shielding is effective when the orbitals in the inner shells are completely filled.

Question 11.
a) Which of the following has higher first ionization enthalpy, N or O? Justify.
b) Which one is bigger, For F ? Why?
1. N. This is because in N, three 2p-electrons reside in different atomic orbitals in accordance with Hund’s rule ($$2p_{ x }^{ 1 }2{ p }_{ y }^{ 1 }2{ p }_{ z }^{ 1 }$$) whereas in O, two the four 2p-electrons must occupy the same 2p-orbital resulting in an increased electron-electron repulsion ($$2p_{ x }^{ 2 }2{ p }_{ y }^{ 1 }2{ p }_{ z }^{ 1 }$$). Consequently, it is easier to remove the fourth 2p-electron from O than it is, to remove one of the three 2p-electrons from N.

2. F (136 pm) is bigger than F (72 pm). An anion is bigger than its parent atom. Addition of one electron in F results in increased repulsion among the electrons and a decrease in effective nuclear charge. Thus, attraction between nucleus and the electrons decreases and hence size increases.

Question 12.
Electronegativity differs from electron gain enthalpy.

1. Do you agree?
2. What do you mean by electronegativity?

1. Yes.
2. Electronegativity is defined as the tendency of an atom in a chemical compound to attract the shared pair of electrons to itself.

Question 13.
Ionization enthalpy is an important periodic property.
1. What is the unit in which it is expressed?
2. What are the factors influencing ionization enthalpy?
3. How ionisation enthalpy varies in the periodic table?
1. kJ mol-1.

2. Atomic/ionic radius, nuclear charge, shielding effect/screening effect, penetration effect and electronic configuration.

3. Ionization enthalpy generally increases with increase in atomic number across a period due to regular increase in nuclear charge and decrease in atomic size. Thus, the attractive force between the nucleus and the electron cloud increases. Consequently, the electrons are more and more tightly bound to the nucleus.

Ionisation enthalpy generally decreases from top to bottom a group: This is because the effect of increased nuclear charge is cancelled by the increase in atomic size and the shielding effect. Consequently, the nucelar hold on the valence electron decreases gradually and ionization enthalpy decreases.

Question 14.
The second period elements show anomalous ‘ behaviour.
1. Give reason.
2. What are the anomalous properties of second period elements?
1. The first element is each group belong to the second period. The difference in behaviour of these elements from the other elements of the same group can be attributed to the following factors:

• Small atomic size
• High electronegativity
• Absence of d-orbtials in the valence shell

2. The important anomalous properties of second period elements are: diagonal relationship, maximum covalence of four and ability to form pπ —pπ multiple bonds.

Question 15.
Atomic size, valency, ionization enthalpy, electron gain enthalpy and electronegativity are the important periodic properties of elements.
1. What do you mean by periodicity?
2. Periodic properties are directly related to ___________
1. The periodical repetition of elements with similar properties after certain regular intervals when the elements are arranged in the order of increasing atomic numbers is called periodicity,

2. Electronic configuration.

Question 16.
1. Which is the element among alkali metals having lowest ionization enthalpy?
2. What is meant by valence of an element? How it varies in the periodic table?
3.

Identify the elements A, B, C and D, if the graph represents halogens.
1. Fr

2. Valence of an element is the combining capacity of that element. In the case of representative elements the number of valence electrons increases from 1 to 8 on moving across a period, the valence to the element with respect to H and Cl increases from 1 to 4 and then decreases from 4 to zero. On moving down a group, the number of valence electrons remains same and, therefore, all the elements in a group exhibit same valence.

3. A = F, B = Cl, C = Br, D = I

Question 17.
Li and Mg belonging to first and second group in periodic table respectively resemble each other in many respects.
1. Name the relationship.
2. B can only form [BF4] ion while Al can form [AIF6]3-, though both B and Al belong to group 13. Justify.
1. Diagonal relationship.

2. This is because B, being a second period element has a maximum covalence of 4. It cannot expand its covalence beyond 4 due to absence of d-orbitals. But Al, being a third period element has vacant d-orbitals in its valence shell and hence can expand its covalence beyond 4.

Question 18.
During a group discussion a student argues that both oxidation state and valence are the same.
1. Do you agree?
2. Justify taking the case of [AICI(H2O)5]2+.
1. No. Valence refers to the combining capacity of an element whereas oxidation state is the charge assigned to an element in a compound based on the assumption that the shared electron in a covalent bond belongs entirly to the more electronegative element.

2. In [AICI(H2O)5]2+the valence of Al is 6 while its oxidation state is +3.

Question 19.
Among the elements of the third period, identify the element

1. With highest first ionization enthalpy.
2. That is the most reactive metal.
3. With the largest atomic radius.

1. Ar
2. Na
3. Na

Question 20.
A cation is smaller than the corresponding neutral atom while an anion is larger. Justify.
A cation is smaller than its parent atom because it has fewer number of electrons while its nuclear charge remains the same. The size of an anion will be larger than that of the parent atom because the addition of one or more electrons would result in increased . repulsion among the electrons and a decrease in effective nuclear charge.

Question 21.
1. How the metallic character varies in the periodic table?
2. Categorize the following oxides into acidic, basic, neutral and amphoteric:
Al2O3, Na2O, CO2, Cl2O7, MgO, CO, As2O3, N2O
1. The metallic character decreases from left to right across the period due to increase in ionization enthalpy along a period which makes loss of electrons difficult. From top to bottom a group metallic character increases due to decrease in ionization enthalpy. Thus, metallic character decreases diagonally from left bottom to right top of the periodic table.

2. Acidic oxides: CO2, Cl2O7
Basic oxides: Na2O, MgO
Neutral oxides: CO, N2O
Amphoteric oxides: Al2O3, As2O3

Question 22.
A group of ions are given below:
Na+, Al3+, O2-, Ca2+, Mg2+, F, N3-, Br
1. Find the pair which is not isoelectronic.
2. Arrange the above ions in the increasing order of size.
1. Ca2+ and Br
2. Al3+ < Mg2+ < Na+ < F < O2- < N3- < Ca2+ < Br

Plus One Chemistry Classification of Elements and Periodicity in Properties Four Mark Questions and Answers

Question 1.
Statement 1: ‘Atomic mass is the fundamental property of an element.’
Statement 2: ‘Atomic number is a more fundamental property of an element than its atomic mass.’
2. Name the scientist who proposed this statement? What observation led him to this conclusion?
1. Statement 2. Atomic number indicates the number of electrons present in an element. Most of the chemical properties of an element depend on its electronic configuration,

2. Henry Moseley. He observed regularities in the characteristic X-ray spectra of the elements. A plot of √υ (where u is the frequency of the X-rays emitted) against atomic number (Z) gave a straight line and not the plot of √υ vs atomic mass.

Question 2.
Atoms possessing stable configuration have less tendency to loss electrons and consequently will have high value of ionization enthalpy.
1. Justify this statement by taking the case of half-filled and completely filled electronic configurations.
2. The noble gases have highest ionization enthalpies in each respective periods. Why?
1. Atoms with half-filled and completely filled electronic configurations have extra stability due to symmetric distribution of electrons and maximum exchange energy. Hence, more energy is required for the removal of their electrons. Elements like N (1s² 2s² $$2p_{ x }^{ 1 }2{ p }_{ y }^{ 1 }2{ p }_{ z }^{ 1 }$$), P (1s² 2s² 2p6 3s² $$3p_{ x }^{ 1 }3{ p }_{ y }^{ 1 }3{ p }_{ z }^{ 1 }$$) etc. possessing half-filled shells have high ionization enthalpies.
Elements like Be ((1s² 2s²), Mg(1s² 2s² 2p6 3s²) etc. having completely filled shells show high values of ionization enthalpy.

2. Noble gases have closed electron shells and very stable octet electronic configurations (except He). Hence, maximum amount of energy is required to remove their valence electron.

Question 3.
Mendeleev arranged the elements in the order of increasing atomic weights.
a) Write down the merits of Mendeleev’s periodic table.
b) What are the demerits of Mendeleev’s periodic table?
a) Merits of Mendeleev’s periodic table:

• Study of elements – Elements are classified into groups with similar properties, thus facilitating the study of properties of elements.
• Prediction of new elements – Mendeleev left certain vacant places in his table which provided a clue for the discovery of new elements, e.g. Eka-AI (for Ga), Eka-Si (for Ge).
• Determination of correct atomic weights – With the help of this table, doubtful atomic weights of some elements were corrected.

b) Demerits of Mendeleev’s periodic table:

• Position of hydrogen was not certain.
• Anomalous pairs of elements – Certain elements of higher atomic weight preceed those with lower atomic weight, e.g. I (at.wt. 127) was placed after Te (at.wt. 128).
• Lanthanides and Actinides are not given proper places in this periodic table.
• No proper position for isotopes.

Question 4.
1. Name any three numerical scales of electronegativity.
2. How electronegativity varies in the periodic table? Justify.
1. Pauling scale, Mullimen-Jaffe scale, Allred- Rochow scale.
2. Electronegativity generally increases from left to right a period and decreases from top to bottom in a group. This is because, from left to right across a period atomic size decreases and attraction between the valence electrons and the nucleus increases. From top to bottom in a group atomic size increases and attraction between the valence electrons and the nucleus decreases.

Question 5.
1. First ionization enthalpy of Na is lower than that of Mg. But its second ionization enthalpy is higher than that of Mg. Explain.
2. Which one is smaller, Na or Na+? Give reason.
1. By the removal of one electron from Na it gets
the stable octet configuration of Ne. But when the first electron is removed from Mg it gets the unstable configuration of Na. It requires more energy due to small size and greater nuclear charge of Mg. In the case of Na the second electron is to be removed from a stable octet configuration which requires more energy than the removal of second electron from Mg.

2. Na+ (95 pm) is smaller than Na (186 pm). Acation is smaller than its parent atom. Na+ has fewer number of electrons (10 electrons) compared to Na (11 electrons). But the nuclear charge remains the same in both. Thus, effective nuclear charge per electron is greater in Na+. Thus, the attraction between nucleus and the remaining electrons increases and size decreases.

Question 6.
Removal of electron becomes easier on moving down the group.
1. Comment the above statement based on ionization enthalpy.
2. How electronic configuration influences the ionization enthalpy value?
1. On moving from top to bottom in a given group, size of the atom increases and ionisation enthalpy decreases. Hence, it becomes easier to remove the valence electron.

2. Atoms with octet configuration, half-filled and completely filled configurations have extra stability and hence have higher values of ionization enthalpy.

Question 7.
The energy released during the addition of an electron to an isolated neutral atom is called electron gain enthalpy.
1. Explain how electron gain enthalpy differ from electronegativity.
2. The second ionisation enthalpy of an element is always greater than the first ionisation enthalpy. Give reason.
1. Electron gain enthalpy(AegH) is the enthalpy change accompanying the process of addition of an electron to a neutral gaseous atom to convert it into a negative ion. It is a quantitative property of an isolated gaseous atom, which can be measured. Whereas, electronegativity is a qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself. It is not a measureable quantity.

2. This is because to remove second electron from a positively charged ion more amount of energy is required due to increase in effective nuclear charge.

Question 8.
The physical and chemical properties of elements are periodic functions of their atomic numbers.
1. The atomic number of an element ‘X’ is 19. Write the group number, period and block to which X’ belong in the periodic table.
2. Name the element with
i) highest electronegativity and
ii) highest electron gain enthalpy
1. The element is K.
19K= 1s² 2s² 2p6 3s² 3p6 4s1
Group number = 1
Period number = 4
Block = s-block

2. i) Fluorine
ii) Chlorine

Plus One Chemistry Classification of Elements and Periodicity in Properties NCERT Questions and Answers

Question 1.
What is the basic theme of organisation in the periodic table? (2)
The basic theme of organisation of elements in the periodic table is to facilitate the study of the properties of all the elements and their compounds. On the basis of similarities in chemical properties, the various elements have been divided into different groups. This has made the study of elements simple because their properties are now studied in the form of groups rather than individually.

Question 2.
What is the basic difference in approach between Mendeleev’s Periodic Law and the Modern Periodic Law? (2)
Mendeleev Periodic Law states that the properties of the elements are a periodic function of their atomic weights whereas Modern Periodic Law states that the properties of elements are a periodic function of their atomic numbers. Thus, the basic difference in approach between Mendeleev’s Periodic Law and Modern Periodic Law is the change in basis of arrangements of elements from atomic weight to atomic number.

Question 3.
Consider the following species. N3-, O2-, F, N2+, Mg2+ and Al3+ (2)
1. What is common in them?
2. Arrange them in order of increasing ionic radii.
1. Each one of these ions contains 10 electrons and
hence these are isoelectronic ions,

2. The ionic radii of isoelectronic ions decrease with the increase in the magnitude of the nuclear charge. Among the isoelectronic ions: N3-, O2-, F, Na+, Mg2+ and Al3+, nuclear charge increase in the order:
N3-< O2- < F < Na+ < Mg2+ < Al3+
Therefore, the ionic radii decrease in the order:
N3- > O2- > F > Na+ > Mg2+ > Al3+

Question 4.
Anything that influences the valence electrons will affect the chemistry of the element. Which one of the following factors does not affect the valence shell?
(a) Valence principal quantum number (n)
(b) Nuclear charge (Z)
(c) Nuclear mass
(d) Number of core electrons (1)
Nuclear mass does not affect the valence shell. Thus, option (c) is the correct answer.

Question 5.
Considering the elements F, Cl, O and N, the correct order of their chemical reactivity in terms of oxidising property is: (1)
a) F > Cl > O > N
b) F > O > Cl > N
c) Cl > F > O > N
d)0>F>N>CI
a) F > Cl > O > N
Across a period, the oxidising character increases from left to right. Therefore, among F, O and N, oxidising power decreases in the order: F > O > N. However, within a group, oxidising power decreases from top to bottom. Thus, F is a stronger oxidising agent than Cl. Thus, overall decreasing order of oxidising power is F > Cl > O > N and the choice (a) is correct.

Plus One Chemistry Chapter Wise Questions and Answers Chapter 7 Equilibrium

Students can Download Chapter 7 Equilibrium Questions and Answers, Plus One Chemistry Chapter Wise Questions and Answers helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Chemistry Chapter Wise Questions and Answers Chapter 7 Equilibrium

Plus One Chemistry Equilibrium One Mark Questions and Answers

Question 1.
Equilibrium in a system having more than one phase is called _________
heterogeneous

Question 2.
Addition of a catalyst to a chemical system at equilibrium would result in
a) Increase in the rate of forward reaction
b) Increase in the rate of reverse reaction
c) A new reaction path
d) Increase in the amount of heat evolved in the reaction
c) A new reaction path

Question 3.
With increase in temperature, equilibrium constant of a reaction
a) Always increases
b) Always decreases
c) May increase or decrease depending upon the number of moles of reactants and products
d) May increase or decrease depending upon whether reaction is exothermic or endothermic
d) May increase or decrease depending upon whether reaction is exothermic or endothermic

Question 4.
Water is a conjugate base of ____________ .
H3O+

Question 5.
Which of the following substances on dissolving in water will give a basic solution?
a) Na2CO3
b) Al2(SO4)3
c) NH4Cl
d) KNO3
a) Na2CO3

Ged exam books to download, equilibrium concentration calculator, Finding the Least Common Denominator, combining like expressions.

Question 6.
Choose the correct answer for the reaction,
N2(g) + 3H2(g) $$\rightleftharpoons$$ 2NH3(g); ∆rH = -91.8 kJ mol-1 The concentration of H2(g) at equilibrium can be increased by
1) Lowering the temperature
2) Increase the volume of the system.
3) Adding N2 at constant volume.
4) Adding H2 at constant volume.
2) and 4) are correct.

Question 7.
Conjugate base of a strong acid is a
Weak base

Question 8.
The expression forostwald dilution law is
Ka = Cα²

Question 9.
The hydroxyl ion concentration in a solution having pH = 4 will be
10-14

Question 10.
A mono protic acid in 1M solution is 0.01 % ionized the dissociation constant of this acid is
10-8

Question 11.
A certain buffer solution contains equal concentration of x and Hx. The Kafor Hx is 10-6 pH of buffer is
6

Question 12.
In the equilibrium reaction
CaCO3(s) → CaO(s) + CO2(g) the equilibrium constant is given by —–
PCO2

Question 13.
Congugate base of a strong acid is a _________ .
a) Strong base
b) Strong acid
c) Weak acid
d) Weak base
e) Salt
d) Weak base

Question 14.
The species acting both as bronsted acid and base is _________ .

b) HSO4

Question 15.
PH of .01 M KOH solution will be _________ .
12

Plus One Chemistry Equilibrium Two Mark Questions and Answers

Question 1.
“High pressure and low temperature favours the formation of ammonia in Haber’s process.” Analyse the statement and illustrate the conditions using Le-Chatliers principle?
The given statement is correct.
N2(g) + 3H2(g) $$\rightleftharpoons$$ 2NH3(g); ∆rH =-91.8 kJ mol-1 Since the number of moles decreases in the forward reaction a high pressure of ~ 200 atm is applied. Since the forward reaction is exothermic the optimum temperature of ~ 700 K is employed for maximum yield of ammonia.

Question 2.
“Chemical equilibrium is dynamic in nature”. Analyse the statement and justify your answer.
At equilibrium the reaction does not stop. Both forward and backward reactions are taking place at equal rates. Thus, at equilibrium two exactly opposite changes occur at the same rate. Hence, chemical equilibrium is dynamic in nature.

Question 3.
Pressure has no influence in the following equilibrium: N2(g) + O2(g) $$\rightleftharpoons$$ 2NO(g)

1. Do you agree with this?
2. What is the reason for this?

1. Yes.
2. Here the total number of moles of the reactants is equal to that of the products. Hence pressure is having no influence in this equilibrium.

Question 4.
During a class room discussion a student is of the view that the value of equilibrium constant can be influenced by catalyst.

1. Do you agree with the statement?
2. Justify the role of catalyst in an equilibrium reaction?

1. No.
2. Catalyst does not affect the equilibrium composition of a reaction mixture. It does not appear in the balanced chemical equation or in the equilibrium constant expression. It only helps to attain the equilibrium state in a faster rate.

Question 5.
What is the equilibrium constant (K) in the following cases?

1. Reaction is reversed.
2. Reaction is divide by 2.
3. Reaction is multiplied by 2.
4. Reaction is splitted into two.

1. 1/K
2. √K
3. K2
4. K1K2

Question 6.
1. What is homogeneous equilibrium?
2. Suggest an example for this.
1. The equilibrium in which the reactants and products are in the same phase,

2. N2(g) + 3H2(g) $$\rightleftharpoons$$ 2NH3(g)
In this equilibrium, the reactants and products are in the gaseous phase.

Question 7.
The equilibrium constants for two reactions are given. In which case the yield of product will be the maximum?
For first reaction: K1 = 3.2 × 10-6
For second reaction: K2 = 7.4 × 10-6
Higher the value of K, greater will be the yield of product. So maximum yield will be in the second case.

Question 8.
Write an expression for equilibrium constant, Kc forthe ‘ reaction, 4NH3(g) + 5O2(g) $$\rightleftharpoons$$ 4NO(g) + 6H2O(g)

Question 9.
1. State Henry’s law.
2. Suggest an example fora gas in liquid equilibrium.
1. The mass of a gas dissolved in a given mass of a solvent at any temperature is directly proportional to the pressure of the gas above the solvent.

2. Equilibrium between the CO2 molecules in the. gaseous state and the CO2 molecules dissolved in water under pressure,
CO2(g) $$\rightleftharpoons$$ CO2(in solution)

Question 10.
1. What is heterogeneous equilibrium?
2. Suggest an example forthis.
1. Equilibrium in a system having more than one phase is called heterogeneous equilibrium

2. Equilibrium between solid Ca(OH)2 and its saturated solution:
Ca(OH)2(s) + (aq) $$\rightleftharpoons$$ Ca2+(aq) + 2OH(aq)

Question 11.
For the equilibrium 2SO3(g) → 2SO2(g) + O2(g), Kc at 47 °C 3.25 × 10-9 mol per litre. What will be the value of Kp at this temperature (R = 8.314 J K-1mol-1).
R = 8.314 J K-1 mol-1 ∆n = 3 – 2 = 1
T = 47 °C = 273 + 47 = 320 K
Kc = 3.25 × 10-9
Kp = Kc (RT)∆n
= 3.25 × 10-9 (8.314 × 320)1
= 3.25 × 10-9 × 8.314 × 320 = 8.65 × 10-12

Question 12.
The pH of a sample of vinegar is 3.76. Calculate the concentration of hydrogen ions in it.
pH= -log[H+] = 3.76
log[H+] = – 3.76
[H+] = antilog of (- 3.76) = 1.738 × 10-4 mol L-1

Question 13.
The equilibrium constant can be expressed in terms of partial pressure as well as concentration.
1. Give the relation between Kp and Kc.
2. What is the relation between Kp and Kc for the reaction, N2(g) + O2(g) $$\rightleftharpoons$$ 2NO(g)?
1. Kp = Kc(RT)∆n, where ∆n = (number of moles of gaseous products) – (number of moles of gaseous reactants).

2. Here, ∆n = 2 – 2 = 0
Kp = Kc(RT)∆n , Kp = Kc (RT)°
∴ Kp = Kc

Question 14.
1. Explain Arrhenius concept of acids and bases with suitable examples.
2. How proton exists in aqueous solution? Give reason.
1. According to Arrhenius theory, acids are substances that dissociates in water to give hydrogen ions, H+(aq) and bases are substances that produce hydroxyl ions, OH(aq). Forexample, HCl is an Arrhenius acid and NaOH is an Arrhenius base.

2. In aqueous solution the proton bonds to the oxygen atom of a solvent water molecule to give trigonal pyramidal hydronium ion, H3O+(aq). This is because a bare proton, H+ is very reactive and cannot exist freely in aqueous solutions.

Question 15.
1. What is an acidic buffer?
2. Suggest an example for an acidic buffer.
1. An acidic buffer is a buffer solution having pH less than 7. It is prepared by mixing a weak acid and its salt formed with a strong base.

2. Mixture of acetic acid (CH3COOH) and sodium acetate (CH3COONa) is an example for an acidic buffer. Its pH is around 4.75.

Question 16.
1. What is a basic buffer?
2. Suggest an example for basic buffer.
1. A basic buffer is a buffer solution having pH greater than 7. It is prepared by mixing a weak base and its salt formed with a strong acid.

2. Mixture of ammonium hydroxide (NH4OH) and ammonium chloride (NH4CI) is an example for a basic buffer. Its pH is around 9.25.

Plus One Chemistry Equilibrium Three Mark Questions and Answers

Question 1.
The concentration of reactant and products for the reaction, H2(g) + l2(g) $$\rightleftharpoons$$ 2Hl(g) are recorded as follows:

 Reactant or Product Molar Concentration H2 0.080 l2 0.060 HI 0.490

a) Write down the expressions for equilibrium constant of the above reaction.
b) Calculate the equilibrium constant at the temperature 298 K if [Hl] = 0.49 M, [H2]=0.08 M and [l2]=0.06 M.

Question 2.
1. When equilibrium is reached in a chemical reaction?
2. What is the influence of molar concentration in a reaction at equilibrium?
3. Write the expression for equilibrium constant for the decomposition of NH4CI by the reaction,
NH4Cl $$\rightleftharpoons$$ NH3+HCl
1. When the rate of forward reaction is equal to rate of backward reaction, the chemical reaction is said to be in equilibrium.

2. Rate of chemical reaction is directly proportional to the product of molar concentration of the reactants.

3. $$\mathrm{K}=\frac{\left[\mathrm{NH}_{3}\right] \mathrm{HCl}}{\left[\mathrm{NH}_{4} \mathrm{Cl}\right]}$$

Question 3.
1. What is meant by Kp?
2. How Kp is related to Kc?
1. Kp is the equilibrium constant in terms of the partial pressures of the reactants and products (Pressure should be expressed in bar as standard state is 1 bar). It is used for reactions involving gases.

2. Kp = Kc (RT)∆n
where R = universal gas constant, T = absolute temperature and ∆n = number of moles of gaseous product(s) – number of moles of gasesous reactant(s).

Question 4.
a) What do you mean by equilibrium constant?
b) Write any two characteristics of equilibrium constant.
c) Write an expression for equilibrium constant of the reaction, 2SO2(g) + O2(g) $$\rightleftharpoons$$ SO3(g).
a) Equilibrium constant at a given temperature is the ratio of product of molar concentrations of the products to that of the reactants, each concentration term being raised to the respective stoichiometric coefficients in the balanced chemical equation.

b) 1. The value of equilibrium constant is independent of the initial concentrations of the reactants and products.
2. Equilibrium constant is temperature dependent having one unique value for a particular reaction represented by a balanced equation at a given temperature.

Question 5.
2NO2(g) $$\rightleftharpoons$$ N2O4(g); ∆H = -52.7 kJ mol-1
1. What change will happen if we increase the temperature?
2. What is the effect of increase in pressure in the above equilibrium?
3. What happens when N2O4 is removed from the reaction medium?
1. Since the forward reaction is exothermic, on increasing temperature the rate of backward reaction (endothermic reaction) increases.

2. Since the number of moles decreases in the forward reaction, on increasing pressure, the rate of forward reaction increases.

3. Rate of forward reaction increases.

Question 6.
Consider this reaction:
CO(g) + 2H2(g) $$\rightleftharpoons$$ CH3OH(g); ∆rH = -92 kJ mol-1
Explain the influence of the following on the basis of
Le Chatelier’s principle.
1. Decrease in pressure.
2. Increase in temperature.
3. Increase in the partial pressure of hydrogen.
1. On decreasing pressure the reaction shifts in the direction in which there is increase in the number of moles. Thus, the rate of backward reaction increases on decreasing pressure.

2. On increasing temperature, the rate of endothermic reaction increases. Here, backward reaction is endothermic. Hence, on increasing temperature the rate of backward reaction increases.

3. Hydrogen, being a reactant increase in its partial pressure increases the rate of forward reaction.

Question 7.
The equilibrium showing dissociation of phosgene gas is given below:
COCl2(g) $$\rightleftharpoons$$ CO(g) + Cl2(g)
When a mixture of these three gases at equilibrium is compressed at constant temperature, what happens to
1. The amount of CO in mixture?
2. The partial pressure of COCl2?
3. The equilibrium constant for the reaction?
1. Amount of CO decreases, because the system favours the reaction in which number of moles decreases with increase of pressure i.e., backward reaction.

2. Increases.
3. Equilibrium constant remains the same since temperature is constant.

Question 8.
The equilibrium constant of the reaction H2(g) + l2(g) $$\rightleftharpoons$$ 2Hl(g) is 57 at 700 K. Now, give the equilibrium constants for the following reactions at the same temperature:

Question 9.
1. What are buffer solutions?
2. Which of the following are buffer solutions?
NaCl + HCl
NH4Cl + NH4OH
HCOOH + HCOOK
3. What is the effect of pressure on the following equilibria?
i) Ice $$\rightleftharpoons$$ Water
ii) N2(g) + O2(g) $$\rightleftharpoons$$ 2N0(g)
1. These are solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali.

2. NH4Cl + NH4OH

3. i) When pressure is increased the melting point of ice decreases and hence the rate of forward reaction will increase.
ii) Pressure has no effect in this equilibrium because there is no change in the number of moles of the gaseous reactants and products.

Question 10.
The aqueous solution of the compounds NaCl, NH4Cl and CH3COONa show different pH.

1. Identify the acidic, basic and neutral solution among them.
2. The concentration of hydrogen ion in a soft drink is 4 × 10-4. What is its pH?

1. Acidic-aqueous solution of NH4Cl Neutral – aqueous solution of NaCl Basic – aqueous solution of CH3COONa
2. pH = – log[H+] = – log[4 × 10-4] = 3.398

Question 11.
1. What is pH? What is its significance?
2. The concentration of hydrogen ion in a sample of soft drink is 3.8 × 10-3. What is its pH?
1. pH is a logarithmic scale used to express the hydronium ion concenration in molarity more conveniently. The pH of a solution is defined as negative logarithm to the base 10 of the activity of hydrogen ion. pH = — log $${ { a }_{ { H }^{ + } } }$$ = —log[H+]

2. [H+] = 3.8 × 10-3
pH = -log[H+]
= -log [3.8 × 10-3] = -(-2.42) = 2.42

Question 12.
1. State the Le-Chatelier’s principle.
2. Apply the above principle in the following equilibrium and predict the effect of pressure.
CO(g) + 3H2(g) $$\rightleftharpoons$$ CH4(g) + H2O(g)
1. The Le Chateliers principle states that a change in any of the factors that determine the equilibrium conditions of a system will cause the system to change in such a manner so as to reduce or to counteract the effect of the change.

2. On increasing pressure the rate of forward reaction increases. This is because number of moles decreases in the forward reaction. In other words, the value of Qc decreases on increasing pressure. As Qc < Kc, the reaction proceeds in the forward direction.

Question 13.
1. Explain Lewis concept of acids and bases.
2. Why does BF3 act as a Lewis acid?
1. A Lewis acid can be defined as a species which accepts electron pair and a Lewis base is a species which donates an electron pair.

2. In BF3, the boron atom is electron deficient and it accepts a lone pair of electron. So it acts as a Lewis acid.

Question 14.
1. How the value of AG influence the direction of an equilibrium process?
2. The equilibrium constant for a reaction is 8. What will be the value of ∆G at 27 °C?
1. If ∆G is negative, then the reaction is spontaneous and proceeds in the forward direction.
If ∆G is positive, then reaction is considered non- spontaneous. instead, as reverse reaction would have a negative ∆G, the products of the forward reaction shall be converted to the reactants.
If ∆G is 0, reaction has achieved equilibrium. At this point, there is no longer any free energy to drive the reaction.
2.

Question 15.
1. What is common ion effect?
2. Suggest an example for this effect.
1. Common ion effect may be defined as the suppression of the dissociation of a weak electrolyte by the addition of some strong electrolyte containing a common ion.

2. The dissociation equilibrium of NH4OH is shifted towards left in presence of NH4Cl having the common ion, NH4+.

Question 16.
1. Predict whether an aqueous solution of (NH4)2SO4 is acidic, basic or neutral?
1. An aqueous solution of (NH4)2SC4 is acidic in nature.

2. (NH4)2SO4 is formed from weak base, NH4OH, and strong acid, H2SO4. In water, it dissociates completely
(NH4)2SO4(aq) → 2NH4+ (aq) + SO42- (aq)
NH4+ ions undergoes hydrolysis to form NH4OH and H+ ions.
NH4+(aq) + H2O(l) $$\rightleftharpoons$$ NH4OH(aq) + H+(aq)
NH4OH is a weak base and therefore remains almost unionised in solution. This results in increased H+ ion concentration in solution making the solution acidic.

Question 17.
1. What are sparingly soluble salts? Suggest an example.
2. Define solubility product constant, Ksp.
3. Obtain the relation between solubility product constant (Ksp) and solubility (S), of a solid salt of general formula Mxp+ Xyq-.
1. Sparingly soluble salts are those salts with solubility less than 0.01 M.
e.g. BaSO4

2. The solubility product of a sparingly soluble salt at a given temperature is defined as the product of the concentrations of its ions in the saturated solution, with each concentration term raised to the power equal to the number of times the ion occurs in the equation representing the dissociation of the electrolyte.

3. The equilibrium in the saturated solution of the salt can be represented as,

Question 18.
The Le Chatelier’s principle is applicable to physical and chemical equilibria.
1. What are the factors which can influence the equilibrium state of a system?
2. Explain the factors affecting the chemical equilibrium on the basis of Le Chatelier’s principle taking Haber’s process for the manufacture of ammonia as an example.
1. The following factors can influence the equilibrium state of a system:

• Change in concentration of the reactants or products.
• Change in temperature.
• Change in pressure.
• Presence of catalyst.

2. N2(g) + 3H2(g) $$\rightleftharpoons$$ 2NH3(g); ∆rH = -91.8 kJ mol-1
When concentration of N2 or H2 is increased, a good yield of NH3 can be achieved. The rate of forward reaction can also be increased by removing NH3 from the reaction mixture.

When pressure is increased, the system will try to decrease pressure and for this system will proceed in that direction where there is minimum number of moles i.e., forward reaction. Thus, a good yield of NH3 can be achieved by increasing pressure.

Since the formation of NH3 is an exothermic reaction, a good yield of NH3 can be achieved by decreasing the temperature. But if the temperature is decreased to very low value the reactant molecules do not have sufficient energy to interact. Hence, an optimum temperature of 500°C is used.

Question 19.
1. Soda water is prepared by dissolving CO2 in water under high pressure. What is the principle involved in this process?
2. At 1000 K, equilibrium constant Kc for the reaction 2SO3(g) $$\rightleftharpoons$$ 2SO2(g) + O2(g) is 0.027. What is the value of Kp at this temperature?
1. Henry’s law
2. Kp =Kc(RT)∆n
∆n = 3.2 = 1
Kp = 0.027 × (0.0831 × 1000)1 = 2.2437

Question 20.
1. For the reaction PCL $$\rightleftharpoons$$ PCl3 +Clc
i) Write the expression of Kc.
ii) What happens if pressure is increased?
2. Write the conjugate acid and base of the following species:
i) H20 ii) HCO;
3. Name the phenomenon involved in the preparation of soap by adding NaCI.
1. i) $$\kappa_{c}=\frac{\left[\mathrm{PC}_{3}\right]\left[\mathrm{Cl}_{2}\right]}{\left[\mathrm{PCl}_{5}\right]}$$
ii) If we increase the pressure the system will try to decrease the pressure. For this system will proceed in the direction where there is minimum number of moles, i.e., rate of backward reaction increases by decreasing the pressure.

2. i) Conjugate acid of H2O is H3O+
Conjugate base of H2O is OH”
ii) Conjugate acid of HCO,” is H2CO3 Conjugate base of HCO3 is CO32-

3. Common ion effect.

Question 21.
a) The pH of black coffee is 5.0. Calculate the hydrogen ion concentration.
b) The Ksp of barium sulphate is 1.5 × 10-9. Calculate the solubility of barium sulphate in pure water.

Question 22.
1. What is conjugate acid-base pair?
2. Illustrate with an example.
1. The acid-base pairthat differs only by one proton is called conjugate acid-base pair. Such acid-base pairs are formed by loss or gain of a proton.

2. Consider the ionization of hydrochloric acid in water.

HCl(aq) acts as an acid by donating a proton to H2O molecule which acts as a base because it accepts the proton. The species H2O+ is produced when water accepts a proton from HCl. Therefore, Cl is the conjugate base of HCl and HCl is the conjugate acid of Cl. Similarly, H2O is the conjugate base of H2O+ and H3O++ is the conjugate acid of the base H2O.

Question 23.
1. What are the applications of equilibrium constant?
2. What is meant by reaction quotient, Qc?
3. Predict the direction of net reaction in the following cases:
i) Qc < Kc
ii) Qc > Kc
iii) Qc = Kc
1. The applications of equilibrium constant are:
• To predict the extent of a reaction on the basis of its magnitude.
• To predict the direction of the reaction.
• To calculate equilibrium concentrations.

2. Reaction quotient, Qc at a given temperature is defined as the ratio of the product of concentrations of the reaction products to that of the reactants, each concentration term being raised to their individual stoichiometric coefficients in the balanced chemical equation, where the concentrations are not necessarily equilibrium values.

3. i) When Qc > Kc, the reaction will proceed in the
direction of reactants (reverse reaction), i.e., net reaction goes from right to left.
ii) When Qc < Kc, the reaction will proceed in the direction of products (forward reaction), i.e., net reaction goes from left to right.
iii) When Qc = Kc, the reaction mixture will be at equilibrium, i.e., no net reaction occurs.

Question 24.
Solubility product helps to predict the precipitation of salts from solution.
1. Find the relation between solubility (S) and solubility product (Ksp) of calcium fluoride and zirconium phosphate.
2. The solubility product of two sparingly soluble salts XY2 and AB are 4 × 10-15 and 1.2 × 10-16 respectively. Which salt is more soluble? Explain.
1. The equilibrium in the saturated solution of calcium fluoride can be represented as,
CaF2(s) $$\rightleftharpoons$$ Ca2+(aq) + 2F(aq)
Ksp = [Ca2+][F]2 = S.(2S)2 = 4S3
The equilibrium in the saturated solution of zirconium phosphate can be represented as,
Zr3(PO4)4(s) $$\rightleftharpoons$$ 3Zr4+(aq) + 4PO43-(aq)
Ksp = [Zr4+]3[PO43-]4 = (3S)3.(4S)4 = 6912S7

2. XY2 is more soluble than AB.

Ionic Strength Calculator … Ionic strength of a solution indicates the concentration of ionic charge in the solution.

Question 25.
a) How common ion effect can influence the solubility of ionic salts?
b) What is the application of common ion effect in gravimetric estimation?
1. In a salt solution, if we increase the concentration of any one of the ions, according to Le Chatelier’s principle, it should combine with the ion of its opposite charge and some of the salt will be precipitated till Ksp = Qsp. Similarly, if the concentration of one of the ions is decreased, more salt will dissolve to increase the concentration of both the ions till Ksp = Qsp.

2. The common ion effect is used for almost complete precipitation of a particular ion as its sparingly soluble salt, with very low value of solubility product for gravimetric estimation.

Plus One Chemistry Equilibrium Four Mark Questions and Answers

Question 1.
Match the following:

Question 2.
In Contact process, SO3 is prepared by the oxidation of SO2 as per the following reaction:
2SO2(g) + O2(g) $$\rightleftharpoons$$ 2SO3(g); ∆H = -189.4
a) What happens to the rate of forward reaction when i) temperature is increased?
ii) pressure is decreased?
iii) a catalyst V2O5 is added?
b) Calculate the pH of 0.01 M H2SO4 solution. Also, calculate the hydroxyl ion concentration in the above solution.
1. D When temperature is increased, the rate of forward reaction decreases since it is exothermic.
ii) When pressure is decreased the rate of forward reaction decreases since it is associated with decrease in number of moles.
iii) When a catalyst V2O5 is added the rate of both forward and backward reactions are increased by the same extent and equilibrium is reached earlier.
2.

The method presented in our buffer pH calculator allows you to compute the pH of both arterial and venous blood.

Question 3.
Calculate the [H+] in the following biological fluids whose pH are given in brackets.
i) Human muscle fluid (6.83)
ii) Human stomach fluid (1.22)
iii) Human blood (7.38)
iv) Human saliva (6.4)
i) pH =-log[H+] = 6.83
log[H+] = – 6.83
[H+] = antilog (- 6.83) = 1.48 × 10-7 mol L-1
ii) [H+] = antilog (- 1.22) = 6.03 × 10-2 mol L-1
iii) [H+] = antilog (- 7.38) = 4.17 × 10-8 mol L-1
iv) [H+] = antilog (- 6.4) = 3.98 × 10-7 mol L-1

Question 4.
The pH value of a solution determines whether it is acidic, basic or neutral in nature.
1. The concentration of hydrogen ion in the sample of a soft drink is 3.8 × 10-3 mol/L. Calculate its pH. Also predict whether the above solution is acidic, basic or neutral.
2. The dissociation constants of formic acid (HCOOH) and acetic acid (CH3COOH) are 1.8 × 10-4and 1.8 × 10-4 respectively. Which is relatively more acidic? Justify your answer.
1. pH = – log[H+] = – log[3.8 × 10-3] = 2.42
Since pH is less than 7, it is an acidic solution,

2. HCOOH is more acidic.
Ka value is directly proportional to the acid strength, i.e., greater the Ka value, stronger is the acid.

Question 5.
a) Write the expression for Henderson – Hasselbalch equation for i) An acidic buffer & ii) A basic buffer.
b) Calculate the pH of a solution which is 0.1 M in
CH3COOH and 0.5 M in CH3COONa. Ka for CH3COOH is 1.8 × 10-6.

Plus One Chemistry Equilibrium NCERT Questions and Answers

Question 1
A liquid is in equilibrium with its vapour in a sealed containerat a fixed temperature. The volume of the container is suddenly increased. (3)
a) What is the initial effect of the change on vapour pressure?
b) How do rates of evaporation and condensation change initially?
c) What happens when equilibrium is restored finally and what will be the final vapour pressure?
a) Vapour pressure decreases due to increase in volume.
b) Rate of evaporation remains same and rate of condensation decreases.
c) Finally the same vapour pressure is restored and the rate of evaporation becomes equal to the rate

Question 2.
The concentration of hydrogen ion in a sample of soft drink is 3.8 × 10-3 M. What is its pH? (2)
pH = -log[H+]
= -log (3.8 × 10-3) = 2.42

Question 3.
The pH of a sample of vinegar is 3.76. Calculate the concentration of hydrogen ion in it. (2)
H= -log [H+]
or log [H+] = -3.76 = -4.24
[H+] = antilog (-3.76) = 1.74 × 10-4M

Question 4.
The ionization constants of HF, HCOOH and HCN at 298 K are 6.8 × 10-4, 1.8 × 10-4 and 4.8 × 10-9 respectively. Calculate the ionization constants of the corresponding conjugate base. (3)
The relation between ionization constant of an acid and that of its conjungate base is Ka x Kb = Kw

Question 5
The ionization constant of nitrous acid is 4.5 x 10-4. Calculate the pH of 0.04 M sodium nitrite solution and also its degree of hydrolysis. (2)
Sodium nitrite is a salt of strong base and weak acid. Its degree of hydrolysis, h is given by the relation.

Question 6
A 0.02 M solution of pyridinium hydrochloride has pH = 3.44. Calculate the ionization constant of pyridine. (2)
Pyridinium hydrochloride is a salt of a weak base (pyridine) and a strong acid (HCl). The pH of an aqueous solution of this salt is given by the relation:

Plus One Chemistry Chapter Wise Questions and Answers Chapter 1 Some Basic Concepts of Chemistry

Students can Download Chapter 1 Some Basic Concepts of Chemistry Questions and Answers, Plus One Chemistry Chapter Wise Questions and Answers helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Chemistry Chapter Wise Questions and Answers Chapter 1 Some Basic Concepts of Chemistry

Plus One Chemistry Some Basic Concepts of Chemistry One Mark Questions and Answers

Plus One Chemistry Chapter Wise Questions And Answers Question 1.
Which of the following is a mixture?
a) Graphite
b) Sodium chloride
c) Distilled water
d) Steel
d) Steel

Plus One Chemistry Chapter 1 Questions And Answers Question 2.
1 µ g = __________ g
[10-3, 10-6, 10-9, 10-12] ‘
10-6

Plus One Chemistry First Chapter Questions And Answers Question 3.
The number of significant figures in 0.00503060 is __________ .
6

Plus One Chemistry Chapter Wise Questions And Answers Pdf Question 4.
The balancing of chemical equations is based on which of the following law?
a) Law of multiple proportions
b) Law of conservation of mass
c) Law of definite proportions
d) Gay-Lussac law
b) Law of conservation of mass

By using the percent/actual yield calculator, you can get the accurate amount of percent, actual, and theoretical yield produced in a chemical reaction.

Plus One Chemistry Questions And Answers Question 5.
Which among the following is the heaviest?
a) 1 mole of oxygen
b) 1 molecule of sulfur trioxide
c) 100 u of uranium
d) 44 g carbon dioxide
d) 44 g carbon dioxide

Plus One Chemistry Previous Year Question Papers And Answers Chapter Wise Question 6.
Calculate the number of atoms in 48 g of He?
Gram atomic mass of He = 4 g.
Thus, numberofatomsin4g (1 mol) He = 6.02 × 1023
So number of atoms in 48 g of He = $$\frac{48}{4}$$ × 6.02 × 1023
=12 × 6.02 × 1023
= 7.224 × 1024

Hsslive Chemistry Previous Questions And Answers Chapter Wise Plus One Question 7.
One mole of CO2 contains how many gram atoms?
3 gram atoms.

Plus One Chemistry Previous Year Questions And Answers Chapter Wise Question 8.
The ratio of gram atoms of Au and Cu in 22ct gold is __________
7 : 2

A compound contains 69.5% oxygen and 30.5% nitrogen and its molecular weight is 92. The compound will be
N2O4

Plus One Chemistry Chapter 1 Question 10.
The total number of electrons present in 1 mole of water is
6 × 1024.

Hsslive Plus One Chemistry Chapter Wise Questions And Answers Question 11.
40g NaOH is present in 100 ml of a solution. Its molarity is __________
10 M

Plus One Chemistry Some Basic Concepts of Chemistry Two Mark Questions and Answers

Plus One Chemistry Some Basic Concepts Of Chemistry Questions Question 1.
Classify the following substances into homogeneous and heterogeneous mixtures.

• Milk
• Iron
• Air
• Gasoline
• Kerosene
• Muddy water

 Homogeneous Heterogeneous Milk, Iron Gasoline Air Kerosene Muddy Water

Plus One Chemistry Previous Year Questions Chapter Wise Question 2.
Calculate the volume occupied by 4.4 g of CO2 at STP?
1 mole CO2 = 44 g
4.4 CO2 = 0.1 mole CO2
Volume occupied by 1 mol CO2 at STP = 22.4 L
∴ Volume occupied 0.1 mol CO2 at STP = 0.1 × 22.4 L
= 2.24 L

Plus One Chemistry Chapter 1 Previous Questions And Answers Question 3.
During a group discussion a student argued that “the water of sea and river should have different chemical composition”.

3. State the law.

1. I can’t join with him.
The water of sea and water of river must have the same chemical composition.
2. Law of definite proportions.
3. A given compound always contains exactly the same proportion of elements by weight.

Plus One Chemistry Question And Answer Question 4.
“When science developed some theories are also modified”.
Write the modified atomic theory.

1. Atom is no longer considered as indivisible, it has been found that atom is made up of sub atomic particles called protons, neutrons and electrons.
2. Atoms of the same element may not be similar in all respects.
3. Atoms of different elements may be similar in one or more respects.
4. The ratio in which atomic unit may be fixed and integral but may not be simple.
5. The mass of atom can be changed into energy.

Plus One Chemistry Chapter Wise Previous Questions And Answers Question 5.
Carbon combines with oxygen to form CO and CO2.

1. What is the law behind this?
2. State the law.

1. Law of multiple proportions.
2. If two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element, are in the ratio of small whole numbers.

Hss Live Plus One Chemistry Chapter Wise Questions And Answers Question 6.
Calculate the volume occupied by 6.02×1025 molecules of oxygen at STP.
Volume occupied by 1 mole of oxygen gas at STP = 22.4 l
i.e., Volume occupied by 6.02×1023 molecules of oxygen gas at STP = 22.4 l
Hence the volume occupied by 6.02 × 1025 molecules
of oxygen gas at STP = $$\frac{22.4 \times 6.02 \times 10^{25}}{6.02 \times 10^{23}}$$ = 2240 l.

Plus One Chemistry Chapter Wise Questions And Answers Hsslive Question 7.
Calculate the molality of a solution of NaOH containing 20g of NaOH in 400 g solvent.

Question 8.
Calculate the mole fraction of NaOH in a solution containing 20 g of NaOH per 360 g of water.

Question 9.
12 g of carbon reacts with 32 g of oxygen to form 44g of carbon dioxide.

1. Which law of chemical combination is applicable here?
2. State the law.

1. Law of conservation of mass.
2. Matter can neither be created nor destroyed. Or, during any physical or chemical change, the total mass of the products remains equal to the total mass of the reactants.

Question 10.
When hydrogen and oxygen combine to form water, the ratio between volume of reactants and products is 2:1:2.

1. Which law of chemical combination is applicable here?
2. State the law.

1. Gay Lussac’s law of gaseous volumes.
2. When gases combine or are produced in a chemical reaction they do so in a simple ratio by volume provided all gases are at same temperature and pressure.

Question 11.
Carbon form two oxides, the first contains 42.9% C and the second contains 27.3% carbon. Show that these are in agreement with the law of multiple proportions.
In the first compound:
C = 42.9%
O = 100-42.9 = 57.1%
So, the ratio between the masses of C and O = 42.9:57.1 = 1:1.33
In the second compound:
C = 27.3%
O= 100-27.3= 72.7%
So, the ratio between the masses of C and O = 27.3 : 72.7= 1:2.66
Hence, the ratio of masses of oxygen which combines with a fixed mass of carbon is 1.33:2.66 or 1:2, a simple whole number ratio. This illustrates the law of multiple proportions.

Question 12.
Match the following:

 A B 1 amu 1.008 x 1.66 x10–24 Mass of 1 H atom 6.02 x 1023 Molar volume of O2 at STP 11.2L Volume of 14g of N2 at STP 1.66 x 10–24 Avogadro number 22.4L

 A B 1 amu 1.66 x 10–24 Mass of 1 H atom 1.008 x 1.6 x10–24 Molar volume of O2 at STP 22.4L Volume of 14g of N2 at STP 11.2L Avogadro number 6.02 x 1023

Question 13.
Calculate the molality of a solution containing 10 g ofNaOH in 200 cm3 of solution. Density of solution is 1.4 g/mL. (Molar mass of NaOH = 40)
Mass of the solution = 200 × 1.04 = 208 g
Mass of NaOH (WB) = 10g Molar mass of NaOH (MB) = 40 g mol-1
Mass of water (WA) = (208 -10) g = 198 g = 0.198 kg

Question 14.
Calculate the mass percentage of oxygen in CaCO3.
Molecular mass of CaCO3 = 100 g mol-1
Mass of oxygen in 100 g CaCO3=3 × 16 g = 48 g
Percentage of oxygen in CaCO3 =$$\frac{48}{100}$$×100 = 48%

Question 15.
KCIO3 on heating decomposes to KCI and O2. Calculate the mass and volume of O2 produced by heating 50 g of KCIO3.
The reaction is represented as,

According to the equation, 96 g of oxygen is obtained from 245 g of KCIO3.
Hence mass of oxygen obtained from 50 g KCIO3 is $$\frac{96 \times 50}{245}=19.6 \mathrm{g}$$
According to the equation 245 g of KCIO3 gives 3 moles of O2 at STP which is 3 × 22.4 L = 67.2 L
Volume of oxygen liberated by 50g of KCIO3
= $$\frac{67.2 \times 50}{245}=13.71 \mathrm{L}$$

Question 16.
Calculate the number of molecules present in

1. 11g of CO2.
2. 56 mL Of CO2 at STP.

1. $$\frac{11}{44}$$ = 0.25 mole
1 mole of CO2 contains 6.022 × 1023 molecules.
∴ 0.25 mole CO2 contains 6.022 × 1023 × 0.25
= 1.51 × 1023 molecules.

2. 56 mL = 0.056 L
$$\frac{0.056}{22.4}$$ =0.0025 mole
= 6.022 × 1023 × 0.0025=1.5 × 1021 molecules

Question 17.
Calculate the number of moles of 02 required to produce 240 g of MgO by burning Mg metal. [Atomic mass: Mg=24, 0=16]
2 Mg + O2 → 2MgO
No. of moles of MgO = $$\frac{240}{40}$$=6
No. of moles of 02 required = 6/2 = 3

Question 18.
Arrange the following in the increasing order of their mass.
(a) 1 g of Ca
(b) 12 amu of C
(c) 6.022 × 1023 mol-ecules of CO2
(d) 11.2 L of N2 at STP
a) Mass of 1 g Ca = 1 g
b) Mass of 12 amu C = $$\frac{12}{6.022 \times 10^{23}}$$ = 2 × 10-23
c) Mass of 6.022 × 1023 molecules of CO2 = 44 g
d) Mass of 11.2 L of N2 at NTP = $$\frac{28 \times 11.2}{22.4}$$ = 14 g
(b) < (a) < (d) < (c)

Question 19.
Complete the table:

 42g N2 1.5 mole N2 33600mLN2 (STP) 16g 0: – – – – mole O2 11.2 L of O2 (STP) ….g CO2 1 mole CO2 – – – – L of C O2 (STP) 28g CO 1 mole CO – – – – mLCO (STP)

 42g N2 1.5 Mole N2 33600mLN2 (STP) 16g O2 0.5 Mole O2 11.2 L of O2 (STP) 44 g CO2 1 mole CO2 22.4 L of CO2 (STP) 28g CO 1 mole CO 22400mLCO (STP)

Question 20.
1. Irrespective of the source, pure sample of H20 always contains 88.89% by mass of oxygen and 11.11% by mass of hydrogen.
a) Which law is illustrated here?
b) State the law.
2. Complete the table by filling in the blanks:

 48 g O2 1.5 mol O2 ……mL O2 (at STP) …… g Na 2 gram atom Na 2NA Na atoms …….g CO2 2.5 mol CO2 56 L (at STP) 8.5 g NH3 ……mol NH3 11.2 L (at STP)

1. a) Law of definite proportions.
b) The same chemical compound always contains the same elements combined in the same fixed proprotion by mass.
2.

 48 g O2 1.5 mol O2 33600 mL O2 (at STP) 46g Na 2 gram atom Na 2Na Na atoms 110 g CO2 2.5 mol CO2 56 L (at STP) 8.5 g NH3 0.5 mol NH3 11.2L(atSTP)

Question 21.
Calcium carbonate reacts with aqueous HCl to give CaCl2 and CO2 according to the reaction:
CaCO3 + 2HCl → CaCl2 + CO2 + H2O
What mass of CaCO3 is required to react completely with 25 mL of 0.75 M HCl?.
Mass of HCI in 0.25 mL 0.75 M HCl $$=\frac{36.5 \times 0.75 \times 0.25}{1000}=0.6844 \mathrm{g}$$
As per reaction 100 g CaCO3 reacts with 2 × 36.5 = 73 g of HCl
∴ Mass of CaCO3 reacting with 0.6844 g HCl $$=\frac{100 \times 0.6844}{73}=0.9375 \mathrm{g}$$

Plus One Chemistry Some Basic Concepts of Chemistry Three Mark Questions and Answers

Question 1.
During a Seminar, a student remarked that “Dalton’s atomic theory has some faulty assumptions”.
a) Do you agree with him?
b) What is the present status of Dalton’s atomic theory?
c) Write any two wrong postulates of Dalton’s atomic theory.
a) I agree with him. Out of 6 Dalton’s postulates, 5 postulates are faulty and only one is correct.
b) Dalton’s atomic theory has undergone many modifications.
c)

• All substances are made up of small indivisible particles called atoms.
• Atoms of the same elements are identical in mass and other properties.

Question 2.
One gram atom of an element contains 6.023 × 1023 atoms.

1. Find the number of atoms in 8 g oxygen.
2. Which is heavier, 1 oxygen atom or 10 hydrogen atoms?
3. Define mole and Avogadro number.

1. 16 g oxygen contains 6.022 × 1023 atoms
∴ 8 g oxygen contains $$\frac{6.022 \times 10^{23}}{2}$$ = 3.011 × 1023

2.

3. One mole is the amount of a substance that contains as many particles or entities as there are atoms in exactly 12 g (or 0.012 kg) of the 12C isotope.
Avogadro number – It is the number of discrete particles present in 1 mole of any substsnce. (Avogadro number, NA = 6.022 × 1023)

Question 3.

1. Classify the following as homogeneous and heterogeneous mixtures.
Air, Smoke, Gunpowder, NaCI solution, Petrol, Bronze, Mixture of sugar and sand.
2. State and explain law of multiple proportions with example.

1. Homogeneous-Air, NaCI solution, Bronze, Gun powder, Petrol.
Heterogeneous – Mixture of sugar and sand, Smoke.

2. When two elements combine to form more than one compound the different masses of one of the elements combining with fixed mass of the other element bear a simple ratio. ,
eg. Carbon reacts with oxygen to form two compounds viz. CO and CO2. In CO mass ratio is 12:16. In CO2 mass ratio is 12:32. Then mass ratio between oxygen in the 2 compounds is 16:32 or 1:2 which is a simple whole number ratio. Hence, the law is verified.

Question 4.
1. One mole of an ideal gas occupies 22.4 L at STP
a) Calculate the mass of 11.2 L of oxygen gas at STP.
b) Calculate the number of atoms present in the above sample.
2. 21 g of nitrogen gas is mixed with 5 g of hydrogen gas to yield ammonia according to the equation.
N2 + 3H2 → 2NH3
Calculate the maximum amount of ammonia that can be formed.
1. a) Mass of 22.4 L oxygen at STP = 32 g
∴ Mass of 11.2 L oxygen at STP = 16 g
b) No. of atoms present in 16 g of O2
$$\frac{6.02 \times 10^{23}}{2}$$ ×2 = 6.02 × 1023 atoms

2. N2 + 3H2 → 2NH3
1 mole N2 + 3 mole H2 → 2moles NH3
1 mole N2 requires 3 mol H2
i.e., 28g N2 requires 6 g H2
Hence, 21 g N2 requires $$\frac{6 \times 21}{28}$$ = 4.5 g H2
21 g N2 reacts completely and 0.5g H2 remains unreacted.
Hence, N2 is the limiting reagent.
28g N2 gives 2 × 17 g NH3
∴ 21 g N2 gives $$\frac{2 \times 17 \times 21}{28}$$ = 25.5 g NH3

Question 5.
When two elements combine to form more than one compound the different masses of one of the elements combining with fixed mass of the other bear a simple ratio.
i) Name the above law.
ii) Explain the above law by taking oxides of carbon.
i) Law of multiple proportions.
ii) Carbon reacts with oxygen to form two compounds viz. CO and CO2.
In CO, mass ratio is 12:16
In CO2,mass ratio is 12:32
Ratio of the masses of oxygen combining with a fixed mass of carbon in the two compounds is 16:32 or 1:2, which is a simple whole number ratio.

Question 6.
A compound contains 80% carbon and 20% hydrogen. If the molecular mass is 30 calculate empirical formula and molecular formula.

Question 7.
A compound contains 4.07% of hydrogen, 24.27% of carbon and 71.65% of chlorine. The molar mass is 98.96. What is the empirical and molecular formula?

Question 8.
Nitrogen forms various oxides.
1. Identify the law of chemical combination illustrated here. Also state the law.
2. Determine the formula of each oxide from the given data and illustrate the law.

1. Law of multiple proportions.
When two elements combine to form more than one compound the different mass of one of the elements which combine with the fixed mass of the other element bear a simple ratio.

2.

In NO and NO2, the masses of oxygen combining with a fixed mass (14 g) of nitrogen are in the ratio, 16:32 = 1:2. Similarly, in N2O and N2O3, the masses of oxygen combining with a fixed mass (28 g) of nitrogen are in the ratio, 16:48 = 1:3. These are simple whole number ratios. Hence, the law of multiple poportions is verified.

Plus One Chemistry Some Basic Concepts of Chemistry Four Mark Questions and Answers

Question 1.
Which of the following weighs more?
a) 1 mole of glucose
b) 4 moles of oxygen
c) 6 moles of N
d) 5 moles of sodium
a) 1 mole glucose = (72 + 12 + 96) g = 180 g
b) 4 moles of oxygen = 4 × 32g = 128g
c) 6 moles of nitrogen = 6 × 14 g = 84 g
d) 5 moles of Na = 5 × 23 g = 115 g.
Thus, 1 mole glucose weighs more.

Question 2.
3 g of H2 is mixed with 29 g of O2 to yield water.
1. Which is the limiting reagent?
2. Calculate the maximum amount of water that can be formed.
3. Calculate the amount of the reactants which remains unreacted.
1.

According to the equation, 4 g H2 requires 32 g
So 3 g H2 requires $$\frac{377 \times 33}{44}$$ = 24 g O2.
Here 3 g H2 is mixed with 29 g of O2. All H2 will react. Hence H2 is the limiting reagent.

2. According to the equation, 4 g H2 gives 36 g H2O. Hence 3 g H2 will give 36 × 3/4 = 27 g H2O.

3. Amount of O2 unreacted = (29 – 24)g = 5 g

Question 3.
a) Calculate the mass of oxygen required for the complete burning of 2 g of carbon.
b) Calculate the molar mass of (i) CO2 (ii) CH4

Question 4.
One gram mole of a substance contains 6.022 x 1023 molecules.
1. 24 g of carbon is treated with 72 g of oxygen to form CO2. Identify the limiting reagent.
2. Find the number of molecules of CO2 formed in this situation.
1.

2 mol of C requires 2 mol of O2.
2 mol C completely reacts with 2 mol of O2 and 0.25 mol O2 and 0.25 mol O2 remains unreacted. Hence, C is the limiting reagent.

2. No. of moles of CO2 formed = 2
∴ No. of molecules of CO2 formed
= 2 × 6.022 × 1023 = 1.2044 × 1024

Question 5.
One gram mole of a substance contains 6.022×1023 molecules.
i) Find out the number of molecules in 2.8 g of nitrogen.
ii) Which is the heavier-one SO2 molecule or one CO2 molecule?
i) No. of molecules in1 mole of N2 = 6.022 × 1023
i.e., No. of molecules in 28 g of N2 = 6.022 × 1023
∴ No. of molecules in 2.8 g N2

Question 6.
a) How can you illustrate the law of multiple proportions by using oxides of metals containing 78.7% and 64.5% of the metal?
b) Match the following:
1/12th the mass of C12 atom – 1 mole
1 g of hydrogen atom – amu
22.4 L O2 at NTP – gram mole
180 g of glucose – gram atom
6.022 × 1023 particles – molar volume
a) In 100 g samples of the two oxides, the masses of the metal are 78.7 g and 64.5 g respectively.
First Oxide :
Mass of oxygen = 100 – 78.7 = 21.3 g
No. of parts by mass of oxygen combining with one part by mass of metal =$$\frac{78.7}{21.3}=3.7 \mathrm{g}$$

Second oxide:
Mass of oxygen = 100 – 64.5 = 35.5 g
No. of parts by mass of oxygen combining with one part by mass of metal = $$\frac{64.5}{35.5}=1.9 \mathrm{g}$$
The ratio of masses of oxygen combining with a fixed mass of metal = 3.7 : 1.9 = 2: 1, a simple whole number ratio.

b) 1/12th the mass C12 atom – amu
1 g of hydrogen atom – gram atom
22.4 L O2 at NTP – molar volume
180 g of glucose – gram mole
6.022 × 1023 particles – 1 mole

Question 7.
Calculate
1. The number of molecules present in 1 g of water.
2. The volume of 0.2 mole of sulphur dioxide at STP.
1. Number of moles in 1 g water = $$\frac{1}{8}$$
∴ No. of molecules in 1 g water
$$=\frac{1 \times 6.022 \times 10^{23}}{18}=3.35 \times 10^{22}$$

2. Volume of 0.2 mol S02 at STP = 0.2 × 22.4 litre
= 4.48 litre

Question 8.
“One mole of all substances contain the same number of specified particles.”
a) Justify the statement.
b) Howto connect mole, gram mole, and gram atom?
c) What is the relation between number of moles and volume?
d) Calculate the number of moles of a gas in 11.2 L at • STP.
a) This statement is true i.e., one mole of all sub-stances contain the same number of specified particles. According to Avogadro’s law.
1 mole of any substance contains 6.022 × 1023 specified particles.

b)

1 gram mole is the molecular mass expressed in gram. It is the mass of 1 mole molecules in gram. Thus, 1 gram mole contains 1 mole molecules.
1 gram atom is the atomic mass expressed in gram. It is the mass of 1 mole atoms in gram. Thus, 1 gram atom contains 1 mole atoms.

c) Number of moles is directly proportional to volume (according to Avogadro law).

Plus One Chemistry Some Basic Concepts of Chemistry NCERT Questions and Answers

Question 1.
Calculate the molecular mass of the following : (3)
1. H2O
2. CO2
3. CH4
1. Molecular mass of H2O = 2(1.008 u) +16.00 u
= 18.016u

2. Molecular mass of CO2 = 12.01 u + 2(16.00 u)
= 44.01 u

3. Molecular mass of CH4 = 12.01 u + 4(1.008 u)
= 16.042 u

Question 2.
Calculate the mass percent of different elements present in sodium sulphate (Na2SO4). (2)

Question 3
Calculate the amount of carbon dioxide that could be produced when (3)
1. 1 mole of carbon is burnt in air.
2. 1 mole of carbon is burnt in 16 g of dioxygen.
3. 2 moles of carbon are burnt in 16 g of dioxygen.
1. The balanced equation for the combustion of carbon dioxide in dioxygen in air is

In air combustion is complete.
Hence, Amount of CO2 produced when 1 mole of carbon is burnt in air = 44 g

2. As only 16 g dioxygen is available it is the limiting reagent.
Hence, amount of CO2 produced = $$\frac{44}{32}$$×16 = 22 g

3. Here again, dioxygen is the limiting reactant. Therefore, amount of CO2 produced from 16 g dioxygen = $$\frac{44}{32}$$×16 = 22g

Question 4.
Chlorine is prepared in the laboratory by treating manganase dioxide (MnO2) with aqueous hydrochloric acid according to the reaction,
4HCl(aq) + MnO2(s) → 2H2O(l) + MnCl2(aq) + Cl2(g)
How many grams of HCl react with 5.0 g of manganese dioxide?
(Atomic mass of Mn = 54.94 u) (2)
1 mol of MnO2, i.e., 54.94 + 32 = 86.94 g MnO2 react with 4 moles of HCl, i.e., 4 × 36.5 g = 146 g of HCl.
∴ Mass of HCl reacting with 5.0 g of MnO2
$$=\frac{146}{86.94} \times 5 \mathrm{g}=8.4 \mathrm{g}$$

Question 5.
Calculate the molarity of a solution of ethanol in water in which the mole fraction of ethanol is 0.040. (2)

Plus One Chemistry Notes Chapter 6 Thermodynamics

Students can Download Chapter 6 Thermodynamics Notes, Plus One Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Chemistry Notes Chapter 6 Thermodynamics

Introduction
The study of energy transformations forms the subject matter or thermodynamics.

Thermodynamic Terms
The system and the surroundings
A system in thermodynamics refers to that part of universe in which observations are made and remain-ing universe constitutes the surroundings. The surroundings include everything other than the system. System and the surroundings together constitute the universe.

Types Of The System
1. Open System:
In an open system, there is exchange of energy and matter between system and surroundings

2. Closed System:
In a closed system, there is no exchange of matter, but exchange of energy is possible between system and the surroundings

3. Isolated System:
In an isolated system, there is no exchange of energy or matter between the system and the surroundings. The presence of reactants in a thermos flask or any other closed insulated vessel is an example of an isolated system.

The State Of The System
The state of a system means the condition of the system when is macroscopic properties have definite values. If any of the macroscopic properties of the system changes, the state of the system will change. A process is said to occur when the state of the system changes.
The measurable properties required to describe the state of a system are called state variables or state functions. Temperature, pressure, volume, composition etc. are state variables.

The Internal Energy as a State Function
1. Work:
The system which can’t exchange heat between the system and surroundings through its boundary is called adiabatic system. The manner in which the state of such a system may be changed will be called adiabatic process. Adiabatic process is a process in which there is no transfer of heat between the system and surroundings.
Internal energy, U, of the system is a state function. The positive sign expresses that wad is positive when work is done on the system. Similarly, if the work is done by the system, wad will be negative.

2. Heat:
A system change its internal energy by ex-change of heat. The q is positive, when heat is transferred from the surroundings to the system and q is negative when heat is transferred from system to the surroundings. first law of thermodynamics, which states that The energy of an isolated system is constant.
i. e., ∆U = q + w
It is commonly stated as the law of conservation of energy i.e., energy can neither be created nor be destroyed.

Applications
Work
The work done due to expansion or compression of a gas against an opposing external pressure is called the pressure – volume type of work. It is a kind of mechanical work.

If Y is the initial volume and Vf is the final volume of a certain amount of gas and Pex is the external pressure, then the work involved in the process is given by
w = – Pex (Vf – Vi) or w = -Pex ∆V

The negative sign of this expression is required to obtain conventional sign for w.

It must be noted that the above expression gives the work done by the gas in irreversible expansion or compression

Work done in isothermal reversible expansion (or compression) of a gas is given by the relation
wrev =-2.303 nRT log $$\frac{V_{f}}{V_{i}}$$
Where n = the number of moles of the gas

Free expansion
Expansion of a gas in vaccum is called free expansion. Since P = 0 in vacuum, work done in free expansion = 0
Isothermal and free expansion of an ideal gas.
1. For isothermal irreversible change q= -w = pex (Vf – Vi)
2. For isothermal reversible change
q = -w = nRT In $$\frac{V_{f}}{V_{i}}$$
= 2.303 nRT log $$\frac{V_{f}}{V_{i}}$$

Enthalpy, H
1. A useful new state function:
We know that the heat absorbed at constant volume is equal to change in the internal dnergy i.e., ∆U= qp

We may write equation as ∆U=qp – p∆V at constant pressure, where qp is heat absorbed by the system and -pAV represent expansion work done by the system.

We can rewrite the above equation as U2 -U1 = qp – p(V2 – V1)

On rearranging, we get qp = (U2 +pV2) = (U1 +pV1) Now we can define another thermodynamic function, the enthalpy H [Greek word enthalpien, to warm or heat content] as :
H=U+PV

so, equation becomes qp = H2 – H1 = ∆H

Although q is a path dependent function, H is a state function because it depends on U, p and V, all of which are state functions.

Therefore, ∆H is independent of path. Hence,qp is also independent of path.

For finite changes at constant pressure, we can write ∆H = ∆U + ∆pV

It is important to note that when heat is absorbed by the system at constant pressure, we are actually measuring changes in the enthalpy. Remember ∆H = qp heat absorbed by the system at constant pressure.

∆H is negative for exothermic reactions which evolve heat during the reaction and ∆H is positive for endothermic reactions which absorb heat from the surroundings.

2. Extensive and Intensive Properties:
An extensive property is a property whose value depends on the quantity or size of matter present, in the system. For example, mass, volume, internal energy, enthalpy, heat capacity, etc. Those properties which do not depend on the quantity or size of matter present are known as intensive properties. For example temperature, density, pressure, etc.

3. Heat Capacity:
The heat required to rise the temperature of the system in case of heat absorbed by the system.

The increase of temperature is proportional to the heat transferred q = coeff × ∆T

The magnitude of the coefficient depends on the size, composition, and nature of the system. We can also write it as q = C ∆T

The coefficient, C is called the heat capacity. Water has a large heat capacity i.e., a lot of energy is needed to raise its temperature. C is directly proportional to amount of substance. The molar heat capacity of a substance, Cm = $$\frac{C}{n}$$, is the heat capacity for one mole of the substance and is the quantity of heat needed to raise the temperature of one mole by one degree Celsius (or one kelvin).

Specific heat, also called specific heat capacity is the quantity of heat required to raise the temperature of one unit mass of a substance by one degree Celsius (or one kelvin). q = c × m × ∆T

4. The relationship between Cp and Cv for an ideal gas:
At constant volume, the heat capacity, C is denoted by Cv and at constant pressure, this is denoted by Cp. Let us find the relationship between the two. We can write equation for heat, q at constant volume as qv=Cv ∆T = ∆U at constant pressure as qp = Cp∆T = ∆H

The difference between Cp and Cv can be derived for an ideal gas as:
For a mole of an ideal gas, ∆H = ∆U + ∆(pV)
= ∆U + ∆(RT)
= ∆U + R∆T
On putting the values ∆H of ∆H and we have
Cp∆T = Cv∆T
Cp = Cv +R
Cp – Cv = R

Measurement Of ∆U And ∆H: Calorimetry
We can measure energy changes associated with chemical or physical processes by an experimental technique called calorimetry.

1. ∆U measurements:
Here, a steel vessel (the bomb) is immersed in a water bath. The whole device is called calorimeter. The steel vessel is immersed in water bath to ensure that no heat is lost to the surroundings. A combustible substance is burnt in pure dioxygen supplied in the steel bomb. Heat evolved during the reaction is transferred to the water around the bomb and its temperature is monitored. Since the bomb calorimeter is sealed, its volume does not change i.e., the energy changes associated with reactions are measured at constant volume. Under these conditions, no work is done as the reaction is carried out at constant volume in the bomb calorimeter. Even for reactions involving gases, there is no work done as ∆V = 0.

2. ∆H measurements:
Measurement of heat change at constant pressure (generally under atmospheric pressure) can be done in a calorimeter at constant p

In an exothermic reaction, heat is evolved, and system loses heat to the surroundings.

Therefore, qp will be negative and ∆rH will also be negative. Similarly in an endothermic reaction, heat is absorbed, qp is positive and ∆rH will be positive.

Enthalpy Change, ∆rH Of A Reaction – Reaction Enthalpy
The enthalpy change accompanying a reaction is called the reaction enthalpy.
The reaction enthalpy change is denoted by ∆rH
rH = (sum of enthalpies of products) – (sum of enthalpies of reactants)

1. Standard enthalpy of reactions:
The standard enthalpy of reaction is the enthalpy change for a reaction when all the participating substances are in their standard states.
The standard state of a substance at a specified temperature is its pure form at 1 bar.

2. Enthalpy changes during phase transformations:
The enthalpy change that accompanies melting of one mole of a solid substance in standard state is called standard enthalpy of fusion or molar enthalpy of fusion ∆fusH°.

Amount of heat required to vaporize one mole of a liquid at constant temperature and under standard pressure (1bar) is called its standard enthalpy of vaporization or molar enthalpy of vaporization ∆vapH°. And that of sublimation is called Standard enthalpy of sublimation, ∆subH°.

3. Standard enthalpy of formation
The standard enthalpy change for the formation of one mole of a compound from its elements in their most stable states of aggregation (also known as reference states) is called Standard Molar Enthalpy of Formation( ∆fH°).

Hess’s Law of Constant Heat Summation Hess’s Law:
If a reaction takes place in several steps then its standard reaction enthalpy is the sum of the standard enthalpies of the intermediate reactions into which the overall reaction may be divided at the same temperature.

Enthalpy Calculator is a free online tool that displays the Enthalpy for the given equation.

Enthalpies For Different Types Of Reactions
1. Standard enthalpy of combustion
Enthalpy of combustion of a substance is defined as the enthalpy change accompanying the complete combustion of one mole of the substance in excess of air or oxygen.

Standard enthalpy of combustion is defined as the enthalpy change accompanying the complete combustion of one mole of the substance in excess of air or oxygen when all the reactants and products are tin their standard states at the specified temperature. It is denoted as ∆cH°.
For example, the complete combustion of one mole of methane evolves 890.3 kJ of heat. Thus, the enthalpy of combustion of methane is- 890.3 kJ mol-1.

Combustion reactions are always accompanied by the evolution of heat. Hence enthalpy of combustion is always negative.

2. Enthalpy of atomization (symbol: ∆aH°)
It is the enthalpy change on breaking one mole of bonds completely to obtain atoms in the gas phase.

3. Bond Enthalpy (symbol: ∆bondH°)
The bond dissociation enthalpy is the change in enthalpy when one mole of covalent bonds of a gaseous covalent compound is broken to form products in the gas phase.

Lattice Enthalpy
The lattice enthalpy of an ionic compound is the enthalpy change which occurs when one mole of an ionic compound dissociates into its ions in gaseous state.

Spontaneity
A process which has an urge or a natural tendency to occur under a given set of conditions is known as a spontaneous process.
Some of the spontaneous process need no initiation, i.e., they take place by themselves. Dissolution of common salt in water, evaporation of water in an open vessel, combination of NO and oxygen to form NO2, neutralisation reaction between NaOH and HCl, etc. are examples of such processes. But some other spontaneous processes need initiation. For example, hydrogen reacts with oxygen to form water only when initiated by passing an electric spark. Once initiated, it occurs by itself.

1. Is decrease in enthalpy a criterion for spontaneity?
By analogy, we may be tempted to state that a chemical reaction is spontaneous in a given direction, because decrease in energy has taken place, as in the case of exothermic reactions.lt becomes obvious that while decrease in enthalpy may be a contributory factor for spontaneity, but it is not true for all cases.

2. Entropy and spontaneity
Entropy(S) is the measure of the degree of randomness or disorder in the system. The greater the disorder in an isolated system, the higher is the entropy. The crystalline solid state is the state of lowest entropy (most ordered), The gaseous state is state of highest entropy. ∆S is independent of path.

Dilution factor formula. After the first tube, each tube is the dilution of the previous dilution tube.

∆S is related with q and T for a reversible reaction as: ∆S = $$\frac{q_{\text {rev }}}{T}$$

When a system is in equilibrium, the entropy is maximum, and the change in entropy, ∆S = 0.

3. Gibbs energy and spontaneity
we define a new thermodynamic function the Gibbs energy or Gibbs function, G, as G = H – TS

Gibbs energy change is a better parameter to determine the spontaneity or feasibility of a process. It can be summarised as follows.
i) If ∆G is negative (i.e., <0) the precess will be spontaneous. ii) If ∆G is zero, the precess is in equilibrium state. iii) If ∆G is positive (i.e., >0), the process is non- spontaneous in the forward direction. The reverse process may be spontaneous.

Ncert Supplementary Syllabus

Enthalpy of Dilution
It is known that enthalpy of solution is the enthalpy change associated with the addition of a specified amount of solute to the specified amount of solvent at a constant temperature and pressure. This argument can be applied to any solvent with slight modification. Enthalpy change for dissolving one mole of gaseous hydrogen chloride in 10 mol of water can be represented by the following equation.
HCl(g) + 10 aq. → HCl. 10 aq. ∆H = -69.01 kJ/mol

Let us consider the following set of enthalpy changes:
(S- 5) HCl(g) + 25 aq. → HCl.25 aq. ∆H = -72.03 kJ/mol
(S-2) HCtlgi + 40 aq. → HCl.40 aq. ∆H = -72.79 kJ/mol
(S-3) HCl(g) + ∞ aq. → HCl. ∞aq. ∆H = -74.85 kJ/mol

The values of ∆H show general dependence of the enthalpy of solution on amount of solvent. As more and more solvent is used, the enthalpy of solution approaches a limiting value, i.e, the value in infinitely dilute solution. For hydrochloric acid this value of AH is given above in equation (S-3).
If we subtract the first equation (equation S-1) from the second equation (equation S-2) in the above set of equations, we obtain

This value (-0.76kJ/mol) of ∆H is enthalpy of dilution. It is the heat withdrawn from the
HCl.25 aq. + 15 aq. → HCl.40 aq.
∆H = [ -72.79 – (-72.03)] kJ/mol
= -0.76 kJ/mol

This value (-0.76kJ/mol) of ∆H is enthalpy of dilution. It is the heat withdrawn from the surroundings when additional solvent is added to the solution. The enthalpy of dilution of a solution is dependent on the original concentration of the solution and the amount of solvent added.

Theoretical Yield Calculator is a free online tool that displays the amount of product predicted with the complete utilisation of the limiting reactant.

Entropy and Second Law of Thermodynamics
We know that for an isolated system the change in energy remains constant. Therefore, increase in entropy in such systems is the natural direction of a spontaneous change. This, in fact, is the second law of thermodynamics. Like first law of thermodynamics, second law can also be stated in several ways. The second law of thermodynamics explains why spontaneous exothermic reactions are so common. In exothermic reactions, heat released by the reaction increases the disorder of the surroundings and overall entropy change is positive which makes the reaction spontaneous.

Absolute Entropy and Third Law of Thermodynamics
Molecules of a substance may move in a straight line in any direction, they may spin like a top and the bonds in the molecules may stretch and compress. These motions of the molecule are called translational, rotational and vibrational motion respectively. When temperature of the system rises, these motions become more vigorous and entropy increases. On the other hand, when temperature is lowered, the entropy decreases. The entropy of any pure crystalline substance approaches zero as the temperature approaches absolute zero. This is called third law of thermodynamics. This is so because there is perfect order in a crystal at absolute zero.

The statement is confined to pure crystalline solids because theoretical arguments and practical evidences have shown that entropy of solutions and super cooled liquids is not zero at 0 K. The importance of the third law lies in the fact that it permits the calculation of absolute values of entropy of pure substance from thermal data alone. For a pure substance, this can be done by summing $$\frac{q_{\text {rev }}}{T}$$ increments from 0 K to 298 K. Standard entropies can be used to calculate standard entropy changes by a Hess’s law type of calculation.

Plus One Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure

Students can Download Chapter 4 Chemical Bonding and Molecular Structure Notes, Plus One Chemistry Notes helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Chemistry Notes Chapter 4 Chemical Bonding and Molecular Structure

Introduction
Matter is made up of different type of elements. The attractive force which holds the constituents together is called a chemical bond.

Kossel-Lewis Approach To Chemical Bonding
The bond between constituents are formed by the sharing of a pair of electrons or their transfer. G.N. Lewis introduced simple notations to represent these outer shell electrons in an atom. These notations are called Lewis symbols.

Significance of Lewis Symbols :
The number of dots around the symbol represents the number of valence electrons. This number of valence electrons helps to calculate the common or group valence of the element. The group valence of the elements is generally either equal to the number of dots in Lewis symbols or8 minus the number of dots or valence electrons.

How to Calculate Bond Order? Introduction to Bond Order.

Kossel, in relation to chemical bonding, drew attention to the following facts:
The bond formed, as a result of the electrostatic attraction between the positiveand negative ions was termed as the electrovalent bond. The electrovalence is thus equal to the number of unit charge (s) on the ion.
In terms of Lewis structures

1. According to electronic theory of chemical bond¬ing, atoms can combine either by transfer of valence electrons from one atom to another (gaining or losing) or by sharing of valence electrons in order to have an octet in their valence shells. This is known as octet rule.

2. Covalent Bond, Langmuir in 1919 refined the Lewis postulations by abandoning the idea of the stationary cubical arrangement of the octet, and by introducing the term covalent bond.

By Lewis – Langmuir theory the formation of chlorine molecule is as follows :

In water molecule covalent bond is as follows:

when two atoms share one electron pair they are said to be joined by a single covalent bond. If two atoms share two pairs of electrons, the covalent bond between them is called a double bond. And when combining atoms share three electron pairs as in the case of N2 molecule a triple bond a triple bond is formed.

• The total number of electrons required for writing the structures are obtained by adding the valence electrons of the combining atoms.
• For anions, each negative charge would mean addition of one electron. For cations, each positive charge would result in subtraction of one electron from the total number of valence electrons.
• The least electronegative atom occupies the central position in the molecule/ion.

Formal charge
Formal charge (F.C.) on an atom in a Lewis structure = total number of valence electrons in the free atom— total number of non bonding (lone pairjelectrons—(1/2) total number of bonding(shared)electrons.

Let us consider the ozone molecule (O3).
The Lewis structure of O3 may be drawn as:

The atoms have been marked as 1,2 and 3. The formal charge on:

• The central O atom marked 1 = 6 – 2- $$\frac{1}{2}$$(6) = +1
• The end O atom marked 2 = 6 – 4 – $$\frac{1}{2}$$(4) = o
• The end O atom marked 3 = 6 – 6 – $$\frac{1}{2}$$(2) = -1

Hence, we represent O3 along with the formal changes as follows:

Limitations of the Octet Rule
There are three types of exceptions to the octet rule. The incomplete octet of the central atom In some compounds, the number of electrons surrounding the central atom is less than eight. This is especially the case with elements having less than four valence electrons.
Some compounds are BCl3, AlCl3 and BF3.

Odd-electron molecules
In molecules with an odd number of electrons like nitric oxide, NO and nitrogen dioxide, NO2, the octet rule is not satisfied for all the atoms.

The expanded octet
In a number of compounds of these elements there are more than eight valence electrons around the central atom.Some examples are given below.

Ionic Or Electrovalent Bond
The formation of a positive ion involves ionization, i.e., removal of electrons from the neutral atom and that of the negative ion involves the addition of electron(s) to the neutral atom.

A qualitative measure of the stability of an ionic compound is provided by its enthalpy of lattice formation and not simply by achieving octet of electrons around the ionic species in gaseous state.

Lattice Enthalpy
The Lattice Enthalpy of an ionic solid is defined as the energy required to completely separate one mole of a solid ionic compound into gaseous constituentions.

Bond Parameters

Bond Length
It may be defined as the equilibrium distance between the centres of the nuclei of the two bonded i atoms in a molecule. Bond length are measured by spectroscopic, X-ray diffraction and electron diffraction techniques. It is usually expressed in Angstrom units (A°) or picometres (pm)
1 A° = 10-10m and 1 pm = 10-12m

Bond Angle
It is defined as the angle between the orbitals containing bonding electron pairs around the central atom in a molecule.

Bond Enthalpy
It is defined as the amount of energy required to break one mole of bonds of a particular type between two atoms in a gaseous state.

Bond Order:
In the Lewis description of covalent bond, the bond order is given by the number of bonds between the two atoms in a molecule. For example, the bond order in H2 is one, in O2 is two and in N2 is three. Isoelectronic molecules and ions have identical bond orders. For example N2, CO and NO+ have bond order 3. It is found that as the bond order increases, bond enthalpy increases and bond length decreases.

Resonance Structures
It is often observed that a single Lewis structure is inadequate for the representation of a molecule in conformity with its experimentally determined parameters. As in the case of O3.

O3 is represented by the above 3 structures. These are called canonical structures.

Experimentally determined oxygen-oxygen bond lengths in the O3 molecule are same (128 pm). Thus the oxygen-oxygen bonds in the O3 molecule are intermediate between a double and a single bond. According to the concept of resonance, the canonical structures of the hybrid describes the molecule accurately.

Some of the other examples of resonance structures are provided by the carbonate ion and the carbon dioxide molecule.

Polarity of Bonds
In reality no bond or a compound is either completely covalent or ionic. Even in case of covalent bond between two hydrogen atoms, there is some ionic character. As a result of polarisation, the molecule possesses the dipole moment. Which can be defined as the product of the magnitude of the charge and the distance between the centres of positive and negative charge. It is usually designated by a Greek letter Mathematically, it is expressed as follows: Dipole moment (µ) = change (Q) X distance of separation (r)

In case of polyatomic molecules, the dipole moment not only depend upon the individual dipole moments of bonds known as bond dipoles but also on the spatial arrangement of various bonds in the molecule.lt is due to the shifting of electrons to the side of more eletro negative element. For example,

The shifting of electrons is represented by an arrow. In case of H2O the resultant dipole moment is given by the following figure:

Fajans Rules:
Just as all the covalent bonds have some partial ionic character, the ionic bonds also have partial covalent character. The partial covalent character of ionic bonds was discussed by Fajans in terms of the following rules:

• The smaller the size of the cation and the larger the size of the anion, the greater the covalent character of an ionic bond.
• The greater the charge on the cation, the greater the covalent character of the ionic bond.
• For cations of the same size and charge, the one, with electronic configuration (n-1)d”ns°, typical of transition metals, is more polarising than the one with a noble gas configuration, ns2 np6, typical of alkali and alkaline earth metal cations.

The cation polarises the anion, pulling the electronic charge toward itself and thereby increasing the electronic charge between the two. This is precisely what happens in a covalent bond, i.e., buildup of electron charge density between the nuclei. The polarising power of the cation, the polarisability of the anion and the extent of distortion (polarisation) of anion are the factors, which determine the per cent covalent character of the ionic bond.

The Valence Shell Electron Pair Repul-Sion (Vspert) Theory
The main postulates of VSEPR theory are as follows:

• The shape of a molecule depends upon the number of valence shell electron pairs (bonded or nonbonded) around the central atom.
• Pairs of electrons in the valence shell repel one another since their electron clouds are negatively charged.
• These pairs of electrons tend to occupy such positions in space that minimise repulsion and thus maximise distance between them.
• The valence shell is taken as a sphere with the electron pairs localising on the spherical surface at maximum distance from one another.
• A multiple bond is treated as if it is a single electron pair and the two or three electron pairs of a multiple bond are treated as a single super pair.
• Where two or more resonance structures can represent a molecule, the VSEPR model is applicable to any such structure.

The repulsive interaction of electron pairs de-crease in the order:
Lone pair (lp) – Lone pair (lp) > Lone pair (lp) – Bond pair (bp) > Bond pair (bp) – Bond pair (bp)

Valence Bond Theory
Valence bond theory was introduced by Heitlerand London (1927) and developed further by Pauling and others. A discussion of the valence bond theory is based on the knowledge of atomic orbitals, electronic configurations of elements, the overlap criteria of atomic orbitals, the hybridization of atomic orbitals and the principles of variation and superposition. First, we consider the formation of H2. When the attractive forces become greater than the repulsive forces, the molecule is formed and the system gets minimum energy. Because energy is released when a bond is formed. The energy so released is called bond enthalpy.

Orbital Overlap Concept
When two atoms approach each other, their atomic orbitals undergo partial interpenetration. This partial interpenetration of atomic orbitals is called overlapping of atomic orbitals. The electrons belonging to these orbitals are said to be shared and this results in the formation of a covalent bond. The main ideas of orbital of overlap concept of formation of covalent bonds are

• Covalent bonds are formed by the overlapping of half filled atomic orbitals present in the valence shell of the atoms taking part in bonding.
• The orbitals undergoing overlapping must have electrons with opposite spins.
Overlapping of atomic orbitals results in decrease of energy and formation of covalent bond.
• The strength of a covalent bond depends upon the extent of overlapping. The greater the overlapping, the stronger is the bond formed.

The above treatment of formation of covalent bond involving the overlap of half-filled atomic orbitals is called valence bond theory.

Types of Overlapping and Nature of Covalent Bonds
The covalent bond may be classified into two types depending upon the types of overlapping:
(i) Sigma(σ) bond, and
(ii) pi(π) bond

(i) Sigma( σ) bond :
This type of covalent bond is formed by the end to end (hand-on) overlap of bonding orbitals along the internuclear axis. This is called as head on overlap or axial overlap. This can be formed by any one of the following types of combinations of atomic orbitals.
s-s overlapping:
In this case, there is overlap of two half filled s-orbitals along the internuclear axis as shown below:

s-p overlapping:
This type of overlap occurs between half filled s-orbitals of one atom and half filled p-orbitals of another atom.

p-p overlapping :
This type of overlap takes place between half filled p-orbitals of the two approaching atoms.

(ii) pi(π) bond :
In the formation of n bond the atomic orbitals overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis. The orbitals formed due to side wise overlapping consists of two saucer type charged clouds above and below the plane of the participating atoms.

Strength of Sigma and pi Bonds
Basically the strength of a bond depends upon the extent of overlapping. In case of sigma bond, the overlapping of orbitals takes place to a larger extent. Hence, it is stronger as compared to the pi bond where the extent of overlapping occurs to a smaller extent. Further, it is important to note that pi bond between two atoms is formed in addition to a sigma bond. It is always present in the molecules containing multiple bond (double ortriple bonds).

Hybridisation
Hybridisationis defined as the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies, resulting in the formation of new set of orbitals of equivalent energies and shape.

The number of hybrid orbitals is equal to the number of the atomic orbitals that get hybridised.

These hybrid orbitals are stable due to their arrangement which provides minimum repulsion between electron pairs. Therefore, the type of hybridisation indicates the geometry of the molecules.

It is not necessary that only half filled orbitals participate in hybridisation. In some cases, even filled orbitals of valence shell take part in hybridisation.

Types of Hybridisation
There are various types of hybridisation involving s, p and d orbitals. The different types of hybridisation are as under:

(I) sp hybridisation:
This type of hybridisation involves the mixing of one s and one p orbital resulting in the formation of two equivalent sp hybrid orbitals. Each sp hybrid orbitals has 50% s-character and 50% p-character. Such a molecule in which the central atom is sp- hybridised and linked directly to two other central atoms possesses linear geometry.The two sp hybrids point in the opposite direction which provides more effective overlapping resulting in the formation of stronger bonds.

Example of molecule having sp hybridisation BeCl2:
The ground state electronic.configuration of Be is 1s²2s². In the exited state one of the 2s-electrons is promoted to vacant 2p orbital to account for its divalency. One 2s and one 2p-orbitalsget hybridised to form two sp hybridised orbitals. These two sp hybrid orbitals are oriented in opposite direction forming an angle of 180°. Each of the sp hybridised orbital overlaps with the 2p-orbital of chlorine axially and form two Be-Cl sigma bonds.

II) sp² hybridisation :
In this hybridisation there is involvement of one s and two p-orbitals in orderto form three equivalent sp² hybridised orbitals. For example, in BCl2 molecule, the ground state electronic configuration of central boron atom is 1s²2s²2p¹. In the excited state, one of the 2s
electrons is promoted to vacant 2p orbital as a result boron has three unpaired electrons.

These three orbitals (one 2s and two 2p) hybridise to form three sp2 hybrid orbitals. The three hybrid orbitals so formed are oriented in a trigonal planar arrangement and overlap with 2p orbitals of chlorine to form three B-Cl bonds. Therefore, in BCl3 the geometry is trigonal planar with ClBCl bond angle of 120°

III) sp³ hybridisation:
This type of hybridisation can be explained by taking the example of CH4 molecule in which there is mixing of one s-orbital and three p-orbitals of the valence shell to form four sp³ hybrid orbital of equivalent energies and shape.

There is 25% s-character and 75% p-character in each sp³ hybrid orbital. The four sp3 hybrid orbitals so formed are directed towards the four corners of the tetrahedron.

The angle between sp³ hybrid orbital is 109.5°

The structure of NH3 and H2O molecules can also be explained with the help of sp3hybridisation. In NH3, the valence shell (outer) electronic configuration of nitrogen in the ground state is 2s²$$p_{x}^{1} 2 p_{y}^{1} 2 p_{z}^{1}$$ having three unpaired electrons in the sp³ hybrid orbitals and a lone pair of electrons is present in the fourth one. These three hybrid orbitals overlap with 1s orbitals of hydrogen atoms to form three N-H sigma bonds. Due to the force of repulsion, the molecule gets distorted and the bond angle is reduced to 107° from 109.5°. The geometry of such a molecule will be pyramidal.

Other Examples of sp³, sp² and sp Hybridisation:
sp³ Hybridisation in C2H6 molecule:
In ethane molecule both the carbon atoms assume sp3 hybrid state. One of the four sp³ hybrid orbitals of carbon atom overlaps axially with similar orbitals of other atom to form sp³-sp³ sigma bond while the other three hybrid orbitals of each carbon atom are used in forming sp³-s sigma bonds with hydrogen atoms Therefore in ethane C-C bond length is 154 pm and each C-H bond length is 109 pm.

sp² Hybridisation in C2H4:
In the formation of ethene molecule, one of the sp² hybrid orbitals of carbon atom overlaps axially with sp² hybridised orbital of another carbon atom to form C-C sigma bond. While the other two sp² hybrid orbitals of each carbon atom are used for making sp²-s sigma bond with two hydrogen atoms. The unhybridised orbital (2px or 2py) of one carbon atom overlaps sidewise with the similar orbital of the other carbon atom to form weak π bond, which consists of two equal electron clouds distributed above and below the plane of carbon and hydrogen atoms.

Thus, in ethene molecule, the carbon-carbon bond consists of one sp²-sp² sigma bond and one pi (π) bond between p orbitals which are not used in the hybridisation and are perpendicular to the plane of molecule; the bond length is 134 pm. The C-H bond is sp^s sigma with bond length 108 pm. The H-C-H bond angle is 117.6° while the H-C-C angle is 121°.

sp Hybridisation in C2H2:
In the formation of ethyne molecule, both the carbon atoms undergo sp- hybridisation having two unhybridised orbital i.e., 2py and 2px. One sp hybrid orbital of one carbon atom overlaps axially with sp hybrid orbital of the other carbon atom to form C-C sigma bond, while the other hybridised orbital of each carbon atom overlaps axially with the half filled s orbital of hydrogen atoms forming σ bonds. Each of the two unhybridised p orbitals of both the carbon atoms overlaps sidewise to form two π bonds between the carbon atoms. So the triple bond between the two carbon atoms is made up of one sigma and two pi bonds

Hybridisation of Elements involving d-Orbitals
The elements present in the third period contain d orbitals in addition to s and p orbitals. The energy of the 3d orbitals are comparable to the energy of the 3s and 3p orbitals. The energy of 3d orbitals are also comparable to those of 4s and 4p orbitals. As a consequence the hybridisation involving either 3s, 3p, and 3d or 3d, 4s and 4p is possible. However, since the difference in energies of 3p and 4s orbitals is significant, no hybridisation involving 3p, 3d and 4s orbitals is possible.

1. Formation of PCl5 (sp³d hybridisation):
The ground state and the excited state outer electronic configurations of phosphorus (Z=15) are represented below.

Now the five orbitals (i.eone s, three p, and one d orbitals) are available for hybridisation to yield a set of five sp3d hybrid orbitals which are directed towards the five comers of a trigonal bipyramidal.

Three sigma bond known as equatorial bonds lie in one plane and make an angle of 120° with each other.
The remaining two P-Cl bonds(called axial bonds)-one lying above and the other lying below the equatorial plane, make an angle of 90° with the plane.

As the axial bond pairs suffer more repulsive interaction from the equatorial bond pairs, therefore axial bonds have been found to be slightly longer and hence slightly weaker than the equatorial bonds; which makes PCl5 molecule more reactive.

2. Formation of SF6 (sp³d² hybridisation):
In SF6 the central sulphur atom has the ground state outer electronic configuration 3s²3p4. In the exited state the available six orbitals i.e., one s, three p and two d are singly occupied by electrons. These orbitals hybridise to form six new sp³d² hybrid orbitals, which are projected towards the six corners of a regular octahedron in SF6. These six sp³d² hybrid orbitals overlap with singly occupied orbitals of fluorine atoms to form six S-F sigma bonds. Thus SF6 molecule has a regular octahedral geometry.

The structure of SF6 is given below.

Molecular Orbital Theory
Molecular orbital (MO) theory was developed by F. Hund and R.S. Mulliken in 1932. The salient features of this theory are :

• The electrons in a molecule are present in the various molecular orbitals as the electrons of atoms are present in the various atomic orbitals.
• The atomic orbitals of comparable energies and proper symmetry combine to form molecular orbitals.
• While an electron in an atomic orbital is influenced by one nucleus, in a molecular orbital it is influenced by two or more nuclei depending upon the number of atoms in the molecule. Thus, an atomic orbital is monocentric while a molecular orbital is polycentric.
• The number of molecular orbital formed is equal to the number of combining atomic orbitals. When two atomic orbitals combine, two molecular orbitals are formed. One is known as bonding molecular orbital while the other is called antibonding molecular orbital.
• The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital.
• Just as the electron probability distribution around a nucleus in an atom is given by an atomic orbital, the electron probability distribution around a group of nuclei in a molecule is given by a molecular orbital.
• The molecular orbitals like atomic orbitals are filled in accordance with the aufbau principle obeying the Pauli’s exclusion principle and the Hund’srule.

Formation of Molecular Orbitals
Linear Combination of Atomic Orbitals (LCAO)
The atomic orbitals of these atoms may be represented by the wave functions ψA and ψB. The formation of molecular orbitals is the linear combination of atomic orbitals that can take place by addition and by subtraction of wave functions of individual atomic orbitals as shown below.

Energy Level Diagram for Molecular orbitals

The increasing order of energies of various molecular orbitals for O2 and F2 is given below:
σ1s < σ*1s < σ2s < σ*2s < σ2px <(π2px = π2py) <(π*2px = π*2pz) < σ*2px

This sequence of energy levels of molecular orbitals is not correct for the remaining molecules Li2, Be2, B2, C2, N2. For molecules such as B2, C2, N2 etc. the increasing order of energies of various molecular orbitals is
σ1s < σ*1s < σ2s < σ*2s <(π2px = π2py) < σ2px <(π*2px = π*2pz) < σ*2pz

The important characteristic feature of this order is that the energy of σ 2pz molecular orbital is higher than that of π2px and π2py molecular orbitals. If the bonding influence is stronger a stable molecule results and if the antibonding influence is stronger,the molecule is unstable.

Bonding In Some Homonuclear Diatomic Molecules
Bond Order:
Bond order is defined as half of the difference between the number of electrons in the bonding molecular orbitals and that in the antibonding molecular orbitals.
Nh – N
i.e. Bond Order= $$\frac{N_{b}-N_{a}}{2}$$. Where Nb is the number of electrons in the bonding molecular orbitals and Na is the number of electrons in the antibonding mo-lecular orbitals.

Significance of bond order:
Bond order conveys the following important informations about a molecule.
i) If the value of bond order is positive, it indicates a stable molecule and if the value of bond order is negative or zero, the molecule is unstable and is not formed.
ii) Bond dissociation energy of a diatomic molecule is directly proportional to the bond order of the molecule. The greater the bond order, the higher is the bond dissociation energy.
iii) Bond order is inversely proportional to the bond length. The higherthe bond ondervalue, smaller is the bond length. For example, the bond length in N2 molecule (having bond order 3) is less than that in O2 molecule (having bond order 2).

Magnetic character:
If all the electrons in the mol-ecules of a substance are paired, the substance will be diamagnetic. On the other hand, if there are un-paired electrons in the molecule, the substance will be paramagnetic.

Hydrogen Bonding
Hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O orN) of another molecule. When hydrogen is bonded to strongly electronegative element ‘X’, the electron pair shared between the two atoms moves far away from hydrogen atom. As a result the hydrogen atom becomes, highly electropositive with respect to the other atom ‘X’. Since there is displacement of electrons towards X, the hydrogen acquires fractional positive charge (δ+) while ‘X’ attain fractional negative charge (δ). This results in the formation of a polar molecule having electrostatic force of attraction which can be represented as: Hδ+ – Xδ-

The magnitude of H-bonding depends on the physical state of the compound. It is maximum in the solid state and minimum in the gaseous state. Thus, the hydrogen bonds have strong influence on the structure and properties of the compounds.

Types of Hydrogen Bonds
There are two types of hydrogen bonds

1. Intermolecular hydrogen bond
2. Intramolecular hydrogen bond

1. Intermolecular hydrogen bond:
It is formed between two different molecules of the same or different compounds. For example, H-bond in case of HF molecule, alcohol or water molecules, etc.

2. Intramolecular hydrogen bond:
It is formed when hydrogen atom is in between the two highly electronegative (F, O, N) atoms present within the same molecule. For example, in o-Nitrophenol the hydrogen is in between the two oxygen atoms as shown below:

Plus One Physics Chapter Wise Questions and Answers Chapter 5 Law of Motion

Students can Download Chapter 5 Law of Motion Questions and Answers, Plus One Physics Chapter Wise Questions and Answers helps you to revise the complete Kerala State Syllabus and score more marks in your examinations.

Kerala Plus One Physics Chapter Wise Questions and Answers Chapter 5 Law of Motion

Plus One Physics Law of Motion One Mark Questions and Answers

Plus One Physics Laws Of Motion Questions Chapter 5 Question 1.
Which one of the following is not a force?
(a) Impulse
(b) Tension
(c) Thrust
(d) Weight
(a) Impulse
Tension, thrust, weight are all common forces in mechanics whereas impulse is not a force.
Impulse = Force × Time duration.

Plus One Physics Laws Of Motion Questions And Answers Chapter 5 Question 2.
A passenger getting down from a moving bus, falls in the direction of the motion of the bus. This is an example for
(a) Inertia of motion
(b) Second law of motion
(c) Third law of motion
(d) Inertia of rest
(a) Inertia of motion
A passenger getting down from a moving bus, falls in the direction of the motion of the bus. This is because his feet come to rest on touching the ground and the remaining body continues to move due to inertia of motion.

Plus One Physics Important Questions And Answers Pdf Chapter 5  Question 3.
Which one of the following is not a contact force?
(a) Viscous force
(b) Magnetic force
(c) Friction
(d) Buoyant force
(b) Magnetic force

Plus One Physics Chapter Wise Questions And Answers Chapter 5 Question 4.
A jet engine works on the principle of
(a) Conservation of linear momentum
(b) Conservation of mass
(c) Conservation of energy
(d) Conservation of angular momentum
(a) Conservation of linear momentum
A jet engine works on the principle of linear momentum.

State And Prove Impulse Momentum Theorem Chapter 5 Question 5.
Newton’s second and third laws of motion lead to the conservation of
(a) linear momentum
(b) angular momentum
(c) potential energy
(d) kinetic energy
(a) linear momentum
Newton’s second and third laws lead to the conservation of linear momentum.

Hsslive Plus One Physics Chapter Wise Questions And Answers Chapter 5 Question 6.
A large force is acting on a body for a short time. The impulse imparted is equal to the change in
(a) acceleration
(b) momentum
(c) energy
(d) velocity
(b) momentum
If a large force F acts for a short time dt, the impulse imparted is
I = F.dt, = $$\frac{d p}{d t}$$.dt
I = dp = change in momentum.

Laws Of Motion Class 11 Questions With Solutions Pdf Chapter 5 Question 7.
When a shell explodes, the fragments fly apart though no external force is acting on it. Does this violate Newton’s first law of motion?
No. The explosion takes place due to the internal force. The internal force does not change the position of centre of mass.

Prove Impulse Momentum Theorem Chapter 5 Question 8.
In taking a catch, a cricket player moves his hands backward on holding the ball. Why?
We know F = $$\frac{\Delta P}{\Delta t}$$
When ∆t increases, the force acting on hand decreases.

350 degrees f to c … T C T F 32 x 59 is the method for converting degrees Fahrenheit to degrees Celsius.

Plus One Physics Important Questions And Answers Chapter 5 Question 9.
Name the factor on which inertia depends.
Mass

Laws Of Motion Class 11 Test Paper Chapter 5 Question 10.
Why does a swimmer push the water backwards?
A swimmer pushes the water backward in order to be pushed forward (Newton’s third law).

Laws Of Motion Previous Year Questions Chapter 5 Question 11.
Rocket works on the principle of conservation of_______.
Momentum

Motion Questions And Answers Pdf Chapter 5 Question 12.
A man experience a backward jerk, while firing bullet from gun. Which law is applicable here? Answer:
Conservation of momentum.

Plus One Physics Laws Of Motion Notes Chapter 5 Question 13.
If you jerk a piece of paper under a book quick enough, the book will not move. Why?
This is due to inertia of rest.

Class 11 Physics Chapter 5 Important Questions Chapter 5 Question 14.
Why it is difficult to walk on a slipper road?
We will not get required reaction from slippery road.

Laws Of Motion Class 11 Important Questions Chapter 5 Question 15.
A stone, when thrown on a glass window, smashes the window pan to pieces. But a bullet fired from the gun passes through it making a hole why?
This is due to inertia of rest of glass window.

Question 16.
Why an athlete runs some distance before taking a jump?
An athletic runs some distance before taking a jump to gain some initial momentum. It helps the athlete to jump more.

Question 17.
Why a horse can not pull a cart and run in empty space?
The horse-cart system moves forward due to reaction of ground on the feet of horse. In free space, there is no reaction. So it can not pull cart.

Question 18.
Why parachute descends slowly?
Parachute has large surface area. This increases fluid friction and slows down the motion of parachute.

Question 19.
Sand is thrown on tracks with snow. Why?
The presence of snow on tracks reduces friction and driving is not safe. If sand is thrown, friction will be increased and driving becomes safe.

Question 20.
It is difficult to move a cycle along a road with its brakes on. Explain.
When the cycle is moved with its brakes on, wheels can only skid. There will be sliding friction. The sliding friction is more compared to rolling friction. Hence it is difficult to move a cycle with its brakes on.

Plus One Physics Law of Motion Two Mark Questions and Answers

Question 1.
Two masses are in the ratio 1:5

1. What is inertia.
2. What is the ratio of inertia of above case?

1. The inability of a body to change it’s state of rest or uniform motion is called inertia.
2. Mass is a measure of inertia. Hence ratio of inertia is 1:5.

Question 2.
More force is required to push a body than pull to get same speed on a ground with some friction. Why?
When we push, the action on the surface and normal reaction on the body increases. (Friction is directly proportional to normal reaction).

As a result more force is required to push the body. When we pull, normal reaction decreases. Hence friction decreases. Hence less force is required to pull the body.

Question 3.
A lift in a multistoried building is moving from ground floor to third floor. What will happen to weight of a person sitting in side of the lift.

1. A When starts to move up from ground floor.
2. When the lift moves with constant speed.

1. A weight increases weight w = mg + ma
2. weight is constant ie. w = mg

Question 4.
Why it is advisable to hold a gun tight to one’s shoulder when it is being fired?
The recoiling gun can hurt the shoulder. If gun is held tightly against the shoulder, the body and gun act a system. This will reduce recoil velocity as it is inversly proportional to mass of system.

Question 5.
Why shockers are used in vehicles?
When there is a jerk or jump, the time for which force acts (∆t) increases. As the product of force and time for which force acts (F∆t) remains constant, increase in At will reduce the force. This provide smooth motion.

Plus One Physics Law of Motion Three Mark Questions and Answers

Question 1.
Give the magnitude and direction of net force on

1. a drop of rain falling down with a constant velocity.
2. a stone of mass 0.1 kg just after it dropped from the window of a tram accelerating at 1 ms-2.

1. Net force is zero

2. When stone is dropped, gravitational force will act on the stone.
Gravitational force F = mg
= 0.1 × 10
= 1 N downward.

Question 2.
An external force is always required to break the inertia of a body which is either in the state of rest or state of uniform motion.

1. Which law governs this statement?
2. Can all forces produce acceleration? Why?
3. A boy holding a spring balance in his hand suspend a mass 2kg from it. If the balance slips from his hand and falls down, find the reading of the balance while it is in the air.

1. Newtons first law of motion.
2. No. If resultant force acting on the body is zero, the body will move with constant velocity or remain at rest.
3. Zero

Question 3.
A man weighs 70 kg. He stands on a weighing scale in a lift which is moving.

1. upward with a uniform speed of 10 m/s.
2. downward with an uniform acceleration of 5 m/s2.
3. upward with an uniform accelerate of 5 m/s2. (Take g = 10m/s2). Find weight in each case.

1. Weight W = mg
= 70 × 10 = 700 N.

2. W = mg – ma
= 70 × 10 – 70 × 5
= 700 – 350
= 350 N

3. W = mg + ma
= 70 × 10 + 70 × 5
= 700 + 350
= 1050N.

Question 4.
A body of mass ‘m’ is placed on a rough inclined plane having coefficient of friction µs. The inclination of plane is given as ‘θ’.

1. Which component of weight brings the body towards the bottom along the plane.
2. Find how much force is required to pull the body along the plane.

1. mg Sinθ brings the body downwards
2. When the body moves upwards the frictional force (Fs) acts downwards
Total pulling Force = mg Sinθ + Frictional force (Fs) (u, mgCosθ).

Question 5.
Four person sitting in the back seat of a car at rest, is pushing on the front seat.

1. Does the car move. Why?
3. Long jumpers take a long run before the jump. Why?

1. No. Action and reaction cancel each other.
2. Newtons third law of motion.
3. To get large inertia of motion.

Question 6.
A Cricket player lowers his hands while catching a Cricket ball to avoid injury.

1. What do you mean by impulsive force?
2. Prove impulse – momentum theorem.

1. The forces which acton bodies for short time are called impulsive forces.
Example:

• In hitting a ball with a bat
• In firing a gun

2. F = $$\frac{d p}{d t}$$
F∆t = dp
impulse = change in momentum.

Plus One Physics Law of Motion Four Mark Questions and Answers

Question 1.
A bead sliding on a wire A moves to C through B as shown in the figure. The bead at A has a speed of200cms

1. what is speed at B?
2. To what height will it rise before it returns?
3. Why the ball moves up even after reaching the bottom most point B?

1. mgh = 1/2 mv2
m × 10 × 0.8 = 1/2 mv2
V2 = 2 × 10 × 0.8
V = $$\sqrt{2 \times 10 \times 0.8}$$
V = 4 m/s.

2. 80 cm (if friction is neglected).

3. when the ball reaches at B, the potential energy is converted into kinetic energy. Due to this kinetic energy the ball raises to the point c.

Question 2.
Figure shows a block (mass m1) on a smooth horizontal surface, connected by a thin cord that passes over a pulley to a second block (m2), which hangs vertically.

1. Obtain formula for the acceleration of the system and tension in the cord.
2. If m1 and m2 interchanges its position, will it affect the tension of the string?
3. What is the acceleration of the system if m1 = 5 kg and m2 = 2kg?

1. When the body m2 moves in down ward direction.
m2g – T = m1 a
T = m2g – m1a.

2. New tension can be found from the relation
m1g – T = m2a
T = m1g – m2 a.

3. Acceleration of system, a

Question 3.
The collision of two ice hockey players are shown in figure.

Analyse the data given in the figure and answer the following questions.

1. Which conservation law is applicable in this case.
2. In which direction and at what speed do they travel after they stick together.
[Hint – towards right can be taken us +ve direc¬tion and vice versa]
3. If we assume the friction of playing ground is zero, predict the nature of motion and the point at which they come to rest.

1. Conservation of linear momentum.

2. Total momentum before collision = Total momentum after collision.
110 × 4 + 90 × -6 = (110 + 90)v
v = 0.5 m/s
-ve direction, (in the direction of man mass 90 kg).

3. Uniform motion They will not stop.

Question 4.
A circular track of radius 300m is kept with outside of track raised to make 5 degree with the horizontal.

1. Name the process in which outside of the road is raised little above the inner.
2. Obtain an expression for the optimum speed to avoid skidding (considering to friction)

1. Banking of roqd

2.

Consider a vehicle along a curved road with angle of banking θ. Then the normal reaction on the ground will be inclined at an angle θ with the vertical.
The vertical component can be divided into N Cosθ (vertical component) and N sinθ (horizontal component). The frictional force can be divided into two components. Fcosθ (horizontal component) and F sinθ (vertical component).
From the figure
N cos θ = F sinθ + mg
N cosθ – F sinθ = mg ______(1)
The component Nsin0 and Fsinθ provide centripetal force. Hence
N sinθ + F cos θ = $$\frac{\mathrm{mv}^{2}}{\mathrm{R}}$$ ______(2)
eq (1) by eq (2)

Dividing both numerator and denominator of L.H.S by N cosθ. We get

This is the maximum speed at which vehicle can move over a banked curved road.
Optimum speed:
Optimum speed is the speed at which a vehicle can move over a curved banked road without using unnecessary friction. Putting µ = 0 in the above equation we get
v0 = $$\sqrt{\mathrm{Rg} \tan \theta}$$.

Question 5.
A circular track of radius 400m is kept with outer side of track raised to make 5° with the horizontal (coefficient of friction 0.2)
(a) Name such track?
(b) What is optimum speed to avoid wear and tear of type?
(c) What is the maximum permissible speed to avoid skidding?
(a) Banking.

Plus One Physics Law of Motion Five Mark Questions and Answers

Question 1.
A horse pulls a cart with constant force so that the cart moves with a constant speed.

1. Does it violate Newtons second law of motion?
2. If not, how will you account for the non acceleration of the cart?
3. Will the speed of the cart increase, decrease or remain the same if the horse applied more force?
4. A body of mass 5kg is acted upon by two perpendicular forces 8N and 6N. Give the magnitude and direction of the acceleration of the body.

1. No.

2. The force applied by the car is balanced by the frictional force. Hence the cart moves with constant velocity.

3. If the horse is applied more force, the speed of the cart increases.

4. The resultant force,

F = 10N
We know, F = ma
10 = 5 × a
acceleration, a = $$\frac{5}{10}$$ = 2 m / sec2

The angle of resultant force,

θ = tan-1 6/8
θ = 36°521
The angle of acceleration θ = 36°521.

Question 2.

1. Friction is the force which opposes the relative motion between two surfaces in contact with each other. What is a limiting static friction? State the laws related to this.
2. Show that the coefficient of friction is equal to the tan of the angle between the resultant and normal reactions.
3. For a body of mass 5kg on a plane at a limiting static friction of 30 degrees. What is the force of friction?

1. The maximum value of static friction is called limiting static friction.

• The magnitude of the limiting friction is independent of the area of contact between the surfaces.
• The limiting static friction is directly proportional to the normal reaction R.

ie f α R
fs = µsR.

2. Angle of friction is the angle whose tangent gives the coefficient of friction.

Consider a body placed on a surface. Let N be the normal reaction and limit is the limiting friction. Let ‘θ’ be the angle between Resultant vector and normal reaction. From the triangle OBC,

∴ tanθ = µ.

3. Tangent of the angle cient of friction.
µs = tanθ
µs = tan 30
µs = $$1 / \sqrt{3}$$
Friction F = µsmg
= $$1 / \sqrt{3}$$ × 5 × 10
F = $$\frac{50}{\sqrt{3}}$$N.

Question 3.
The rate of change of linear momentum of a body is directly proportional to the external force applied on it, and takes place always in the direction of force applied.

1. Name this law.
2. Using this law obtain the expression for force.
3. The motion of a particle of mass m is described by y = ut + $$\frac{1}{2}$$ gt2. Find the force acting on it.

1. Newton’s Second Law.

2. Consider a body of mass ‘m’ moving with a momentum $$\vec{p}$$. Let $$\vec{F}$$ be the force acting on it for time internal ∆t. Due to this force the momentum is changed from $$\vec{p}$$ to p + ∆p. Then according to Newtons second law, we can write

Where K is a constant pf proportionality. When we take the limit ∆t → 0, we can write

3.

Hence force F = mg.

Question 4.
Recoil of gun is based on the principle of conservation of momentum.

1. State the principle of conservation of momentum.
2. Explain the reoil velocity of gun.
3. A bullet of mas 100g is fired from a rife of mass 200 kg with a spped of 50 m/s. Calculate the recoil velocity of the rife.

1. According to law of conservation of linear momentum, if the external force acting on a body is zero, total linear momentum remains constant. According to Newton’s second law.
F = $$\frac{d p}{d t}$$
If F = 0, $$\frac{d p}{d t}$$ = 0 i.e; P is constant.

2. Let M, m be the mass of gun and bullet respectively. Let V and ν be the velocities of gun and bullet after firing.
According to consevation of momentum
Total momentum before firing = Total momentum after firing
∴ O = MV + m ν
-MV = mν
The above equation shows that when bullet moves in forward direction, the gun moves in back direction. This motion of gun is called recoil of gun.

3. M = 200kg, m = 100g = 0.1kg
ν = 50 m/s, V = ?
MV = mν
200 × V = 0.1 × 50
V = $$\frac{0.1 \times 50}{200}$$m/s.

Question 5.
While firing a bullet, the gun must be held tight to the shoulder.

1. Which conservation law helps you to explain this
2. “In the firing process, the speed of the gun is very low compared to the speed of the bullet.” Substantiate the above statement using mathematical expressions.
3. A shell of 20kg moving at 50m/s bursts in to two parts of masses 15kg and 5kg. If the larger part continues to move in the same direction at 70 m/s. What is the velocity and direction of motion of the other piece.

1. Conservation of momentum.

2. Total momentum is conserved
∴ mu + MV = 0
V = $$\frac{-m u}{M}$$ M is very large. Hence v is small

3. MV = m1 u1 + m2 u2
20 × 50 = 5u1 + 15 × 70
5u1 = 50
u1 = 10m/s.

Question 6.
While firing a bullet, the gun must be held tight to the shoulder.

1. This is a consequence of______
2. Show that recoil velocity is opposite to the muzzle velocity of the bullet.
3. A gun of mass 5 kg fire a bullet of mass 5g, vertically upwards to a height of 100m. Calculate the recoil velocity of gun.

1. Conservation of linear momentum.

2. Let M be the mass of gun and m be the mass of bullet. When gun fires, the gun and bullet acquire velocities V and v respectively.
According to conservation of momentum.
Total momentum before firing = Total momentum afterfiring
m × o + M × o = mu + MV
O = mv + MV
ie. – MV = mv
V = $$\frac{-m v}{M}$$

3. M = 5kg, m = 5 × 10-3 kg, h = 100m
v2 = u2 + 2as
0 = u2 + 2 × 10 × 100
Velocity of bullet,

Question 7.
A standing passenger falls backwards when the bus starts suddenly.

1. Explain why this happens?
2. Which Newtons law gives the above concept. State the law.
3. Obtain an expression for force using Newtons law.

1. Due to inertia of rest, the body continues in the state of rest.

2. Newtons first law:
Everybody continues in its state of rest or of uniform motion along a straight line unless it is compelled by an external unbalanced force to change that state:

3. Consider a body of mass ‘m’ moving with a momentum $$\vec{p}$$. Let $$\vec{F}$$ be the force acting on it for time internal ∆t. Due to this force the momentum is changed from $$\vec{p}$$ to p + ∆p. Then according to Newtons second law, we can write

Where K is a constant pf proportionality. When we take the limit ∆t → 0, we can write

Question 8.
According to Newton’s law of motion rate of change of momentum is directly proportional to applied force.
a. Impulse has the unit similarto that of

1. Momentum
2. force
3. time
4. Energy

b. A man falling from certain height receives more injuries when he falls on a marble floor than when he falls on a heap of sand. Explain. Why?
c.

Force – time graph for a body starting from rest is shown in the figure. What is the velocity of the body at the end of 12 second? (Mass of the body is 5 kg)
a. 1. Momentum.

b. When a man falls on a marble floor, the momentum is reduced to zero in lesser time. Due to this, the rate of change of momentum is large. So greater force acts on a man falls on marble floor.

c. The area of force – time graph gives change in momentum.
ie. change in momentum,
mv = 1/2 × (12 – 4) × (20 -10)
mv = 40

Plus One Physics Law of Motion NCERT Questions and Answers

Question 1.
Give the magnitude and direction of the net force acting on
(a) a drop of rain falling down with a constant speed.
(b) a cork of mass 10g floating on water
(c) a kite skillfully held stationary in the sky
(d) a car moving with a constant velocity of 30km h-1 on a rough road
(e) a high – speed electron in space far from all material objects, and free of electric and magnetic fields.
Applying Newton’s first law of motion, we find that no net force acts in any of the situations, (a) to (d). Again, no force in situation (e). This is because electron is far away from all material agencies producing electromagnetic and gravitational forces.

Question 2.
A constant retarding force of 50 N is applied to a body of mass 20kg moving initially with a speed of 15ms-1. How long does the body take to stop?
Acceleration, a = –$$\frac{50 \mathrm{N}}{20 \mathrm{kg}}$$ = -2.5ms-2
u = 15ms-1, v = 0, t = ?
v = u + at
0 = 15 – 2.5t or 2.5t = 15 or
t = $$\frac{15}{2.5}$$s = 6.0s.

Question 3.
A constant force acting on a body of mass 3.0kg changes its speed from 2.0ms-1 to 3.5 ms-1 in 25s. The direction of motion of the body remains unchanged. What is the magnitude and direction of the force?
m = 3kg; u = 2ms-1; v = 3.5 ms-1;
t = 25s ; F = ?
v = u + at
3.5 = 2 + 25a or a = 0.06 ms-2
F = ma = 3kg × 0.06 ms-2 = 0.18N.
The direction of force is along the direction of motion.

Question 4.
A bob of mass 0.1 kg hung from the ceiling of a room by a string 2 m long is set into oscillation. The speed of the bob at its mean position is 1 ms-1. What is the trajectory of the bob if the string is cut when the bob is

1. at one of its extreme positions.
2. at its mean position.

1. At the extreme position, the speed of the bob is zero. If the string is cut, it will fall vertically down wards.
2. At the mean position, the bob has a horizontal velocity. If the string is cut, it will fall along a parabolic path.

Question 5.
A man of mass 70kg stands on a weighing scale in a lift which is moving

1. upwards with a uniform speed of 10ms-1
2. downwards with a uniform acceleration of 5ms-2
3. upwards with a uniform acceleration of 5ms-2 What would be the readings on the scale in each case?
4. What would be the reading if the lift mechanism failed and it hurtled down freely under gravity?

1. a = 0, R = mg = 70 × 10 N = 700N
2. mg – R = , ma ; R – mg – ma = (g – a)
= 70(10 – 5) N = 350N
3. R – mg = ma or R = m(g + a)
= 70(10 + 5)N = 1056 N
4. In the event of free fall, it is a condition of weight lessness.

Question 6.
A nucleus is at rest in the laboratory frame of reference. Show that if it dist integrates into two smaller nuclei, the products must move in opposite directions.
Applying principle of conservation of momentum,

Question 7.
A shell of mass 0.020 kg is fired by a gun of mass 100kg. If the muzzle speed of the shell is 80ms-1, what is the recoil speed of the gun?
m = 0.02kg, M = 100kg, v = 80ms-1, V = ?

= -0.016ms-1 = -1.6cm s-1
Negative sign indicates that gun moves in a direction opposite to the direction of motion of the bullet.

Kerala Plus One Physics Model Question Paper 4

Time: 2 Hours
Cool off time: 15 Minutes
Maximum: 60 Scores

General Instructions to candidates

• There is a ‘cool off time’ of 15 minutes in addition to the writing time.
• Use the ‘cool off time’ to get familiar with the questions and to plan your answers.
• Calculations, figures, and graphs should be shown in the answer sheet itself.
• Malayalam version of the questions is also provided.
• Give equations wherever necessary.
• Electronic devices except non-programmable calculators are not allowed in the Examination Hall.

Answer any four questions from question numbers 1 to 5. Each carries one score.

Engineering Physics MCQ with answers in PDF format.

Question 1.
Name the weakest force among the fundamental forces.

Question 2.
The work done during an isochoric process is …………….

The harmonic sequence formula is a sort of average calculator that is estimated by dividing the number of utilities.

Question 3.
Highway police detect over speeding vehicles by using ……………….
a. Magnus effect
b. Pascals law
c. Doppler effect
d. Bernoulli’s theorem

Question 4.
Two forces 3N and 4N are acting perpendicular to each other. The magnitude of the resultant force is

Question 5.
Say true/false: “Trade winds are produced due to conduction.”

Answer any five questions from question numbers 6 to 11. Each carries two scores.

Question 6.
The displacement (S) of a body in time ‘t’ is given by S = at2 + bt. Find the dimensions of a and b.

Question 7.
Give the magnitude and direction of the net force on a stone of mass 0.1 kg.
a. Just after it is dropped from the window of a train accelerating 1 ms2.
b. Lying on the floor of a train which is accelerating with 1 ms-2, the stone being at rest relative to the train.

Question 8.
A body is rolling on a horizontal surface. Derive an equation for its kinetic energy,

Question 9.
The stress-strain graphs for two materials A and B are shown below (the graphs are drawn using the same scale) Which one is more elastic? Why?

Question 10.
“A heavy and a light body have the same kinetic energy.” Which one has greater momentum? Why?

Question 11.
The following figures refer to the steady flow of a nonviscous liquid. Which of the two figures is correct? Why?

Answer any five questions from question num 1.5 m numbers 12 to 17. Each carries three scores.

Question 12.
The side of a cube is measured as 3.405 cm.
a. How many significant figures are there in the measurement?
b. If the percentage error in the measurement of the side of the cube is 3%, find; the percentage error in its volume.

Question 13.
According to the conservation of energy “energy can neither be created nor be destroyed”
a. Prove law of conservation of mechanical energy in the case of a freely falling body.
b. The bob of a pendulum of length 1.5 m is released from the position A shown in the figure. What is the speed with which the bob arrives at the lowermost point B, given that 5% of its initial energy is dissipated against air resistance?

Question 14.
Acceleration due to gravity on earth changes with depth and height.
a. What is the weight of a body placed at the center of the earth? Why?
b. Find the height at which the acceleration due to gravity is 1/4th that at the surface of the earth.

Question 15.
A metal sphere of density ‘p’ and radius ‘a is falling through an infinite column of liquid of density ‘o’ and coefficient of viscosity Ty
a. Name any two forces acting on the; sphere.
b. With the help of Stokes theorem, derive an equation for the terminal velocity of I the sphere.

Question 16.
Conduction is the mode of transfer of heat in solids. of Write the unit of thermal conductivity.
b. “Burns produced by steam is severe than that produced by boiling water”Why?

Question 17.
A gas has ‘f’ degrees of freedom.
a. Calculate its Cp, Cv, and γ.
b. Define the mean free path.

Answer any five questions from question numbers 18 to 22. Each carries two scores.

Question 18.
A satellite moves in a circular orbit of radius ‘r’ with an orbital velocity.
a. Derive an equation for the orbital velocity of a satellite.
b. The time taken by Saturn to complete one orbit around the Sun is 29.5 times the earth year. If the distance of the earth from the Sun is 1.5 × 108km, then what will be the distance of the Saturn from the Sun?

Question 19.
In the simple harmonic motion, force is directly proportional to the displacement from the mean position.
a. Give an example of a harmonic oscillator.
b. Derive equations for the kinetic and potential energies of a harmonic oscillator.
c. Show graphically the variation of kinetic energy’ and potential energy of a harmonic oscillator.

Question 20.
A stretched string can be used as a musical instrument.
a. What is the fundamental frequency of a stretched string?
b. With neat diagrams, derive equations for the second and third harmonics of a stretched string.

Question 21.
A body having an initial velocity ‘v0’ has an acceleration ‘a’.
a. Using the velocity-time graph, derive an equation for displacement of the above body.
b. Draw the velocity Time graph and speed Time graph of a body thrown vertically in the air.

Question 22.
A javelin is thrown with an initial velocity ‘ V0‘ at an angle ” with the horizontal.
a. What are the horizontal and vertical velocities of the body
i. At the point of projection
ii. At maximum height
b. Find the angle of projection at which the maximum height attained by the javelin is equal to the horizontal range.

Answer any three questions from question numbers 23 to 26. Each carries five scores.

Question 23.
a. What is meant by ‘banking of roads’?
b. With a neat diagram, derive an equation for the maximum velocity of a car on a banked road.
c. What is the optimum speed of the car along the banked road?

Question 24.
The moment of inertia of a thin rod of mass M and length 1 about an axis perpendicular to the rod at its midpoint is $$\frac { { Ml }^{ 2 } }{ 12 }$$.
a. What is the radius of gyration in the above case?
b. A student has to find the moment of inertia of the above rod about an axis (AB) perpendicular to the rod and passing through one end of the rod. Name and state the law used for this case.
c. Using the theorem, find the moment of inertia of the rod about AB.

Question 25.
Small drops of water assume spherical shape due to surface tension.
a. Define surface tension.
b. Derive an equation for the excess pressure inside a liquid drop of radius ‘R’ having surface tension σ.
c. Why do farmers plow the fields before summer?

Question 26,
Carnot engine is considered as an ideal heat engine.
a. Draw the PV graph of Carnot’s cycle.
b. Derive an equation to find the work done during an adiabatic process.
c. Calculate the efficiency of a heat engine working between ice point and steam point.

Doppler effect

Zero

Doppler effect

7N

False

[S] = [L]
[at2] = [L]
a = [LT-2]
[bt] = [L]
[b] = [LT-1]

a. Only force is gravitational. F = mg = 0.1 × 9.8 = 9.8 N downward j
b. Gravitational force is cancelled by normal I reaction.
∴ F2 = ma = 0.1 × 1 = 0.1 N, direction of motion of train.

In the two graphs, the slope of a graph of material A is greater than the slope of a graph of material B. So material A is more elastic than B. For material A the break-even point (D) is higher.

Momentum is greater for a heavy body.

Figure b is correct. According to an equation of continuity, the speed of liquid is larger at a smaller area. From Bernoulli’s theorem due to larger speed, the pressure will be lower at a smaller area and therefore the height of liquid column will also be at lesser height, while in Fig(a) height of liquid column at the narrow area is higher.

a. Law of conservation of energy. Energy can neither be created nor be destroyed, but it can be transformed from one form into another. Consider a body of mass’s’ placed at

b. Changing in PE after dissipation.

a. i. Weight, F, = mg acting downward
ii. Viscous force, F2 acting upward,
b. By strokes, formula F = 6πrηV Viscous force = Apparent weight of sphere in the solid

a. W m-1K-1
b. Boiling water contains only a specific amount of heat energy required for it to boil. However, as steam is formed from boiling water, it contains the heat energy of boiling water, along with the latent heat of vaporization.i.e., 1kg of steam at 100°C contains 22.6 × 105 J more heat than 1 kg of water at 100°C. Hence, as steam has more heat energy, it can cause more severe burns than boiling water.

b. Mean free path is an average distance between two successive collisions.

a. It is the velocity required to put the satellite into its orbit around the earth.

The gravitational force on the satellite

The centripetal force required by the satellite to stay in this orbit is

in this orbit is In equilibrium the centripetal force is given by the gravitational force

a. Oscillation of simple pendulum Oscillation of loaded spring
b. Let m be the mass of the particle executing SHM. Let v be the velocity at any instant,

Potential energy is the work required to take a particle against the restoring., force. Let a particle be displaced through a distance x from the mean position. Then restoring force, F = – kx, where k is the force constant. Now if we displace the particle further through a distance dx, Small work done, dw = – Fdx = kx dx Total work done from 0 to x

a. Fundamental mode (or) First harmonic: If the string is plucked in the middle and released, then it vibrates in one segment with nodes at its ends and an antinode in the middle.

This is the lowest frequency with which string vibrates.
b. Second harmonic If the string is pressed in the middle and plucked at one-fourth of its length, then the string vibrates in two segments.

Third harmonic If the striping is pressed at one-third of its length from one end and plucked at one-sixth its length, it will vibrate in three segments.

Thus a collection of all possible mode is called harmonic series and n is called harmonic number.

The area under the velocity-time graph gives the displacement of the body. Displacement, x = area OABD x = area of triangle ABC+ area of rectangle OACD.

a.
i. Horizontal Vx = V0 cosθ Vertical Vy = V0sinθ
ii. Horizontal V’x = VO cosθ

a. To avoid skidding and damage to tires of vehicles, the outer part of a road is slightly raised than the inner part. This is known as banking of roads.

The forces on the car are:
1. The weight of the car vertically downwards.
2. Normal reaction Racing normal to the road.
3. Frictional force acting parallel to the road.
Since there is no vertical acceleration,
R cosθ = mg + F sinθ
or R cosθ – F sinθ = mg …(1)
Now for maximum speed, F = μ, R
The centripetal force is provided by horizontal components of Rand Fas shown in the figure.

a. The radius of gyration (k). It is the defined as the distance from an axis of rotation at which, if the whole mass of the body was concentrated, then its moment of inertia about that point would be the same as the moment of inertia of actual distribution of mass. l = Mk2
The radius of gyration (k) of a body is the square root of a ratio of the moment of inertia and a total mass of the body.
ie., a radius of gyration, k= $$k=\sqrt { \frac { l }{ M } }$$

b. Theorem of parallel axes: This theorem is good for any shape. The moment of inertia of the body about any axis is equal to the sum of a moment of inertia of a.parallel axis passing through the center of mass and product of its mass of the body and square of the distance between the two parallel axes.

where I am the moment
c. Using parallel axes theorem, the moment of inertia about AB,

a. Surface tension (a) is the property due to which the free surface of a liquid at rest behaves like an elastic stretched membrane tending to contract so as to occupy a minimum surface area.

Thus it is measured as the force acting per unit length of an imaginary line drawn on the liquid surface, the direction of force being perpendicular to this line and tangential to the liquid surface.

b. Consider a liquid drop of radius R and surface tension o. Let P be the excess pressure inside the drop. The work done by the force due to excess pressure is

c. On plowing, the gap between sand particles act as a capillary tube, so that groundwater reaches the surface easily due to capillary rise.

b. Work was done in the adiabatic process: We have a small amount of work done when volume changes through at pressure P.

(volume changes from v1 to v2 diabolically)

Plus One Maths Chapter Wise Previous Questions Chapter 15 Statistics

Kerala State Board New Syllabus Plus One Maths Chapter Wise Previous Questions and Answers Chapter 15 Statistics.

Kerala Plus One Maths Chapter Wise Previous Questions Chapter 15 Statistics

Plus One Maths Statistics 3 Marks Important Questions

Question 1.
Consider the numbers 4,7,8,9,10,12,13,17 (MARCH-2010)
i) Find the mean of the numbers.
ii) Find the mean deviation about the mean.
iii) Find the standard deviation.

Question 2.
Consider the following data; 35,49,30,32,50,41,34,45,36 (MARCH-2013)
i) Find its median.
ii) Find its mean deviation about median
i) 30, 32, 34, 35, 36, 41, 45, 49, 50
Median is the 5th observation when the data is arranged in ascending order. Hence median = 36
ii)

Question 3.
The mean and standard deviation of marks obtained by 50 students of 50 students in a class in two subjects mathematics are given below:  (IMP-2014)

Which one of the subject shows highest variability in marks? Why?

Thus Accountancy with highest CV shows highest variability and Mathematics with lowest CV shows lowest variability.

Plus One Maths Statistics 4 Marks Important Questions

Free Online Polynomial in Ascending Order Calculator helps people to rearrange the given polynomial expression in ascending order in a fraction of seconds.

Question 1.
A public Opinion polling agency surveyed 200 government employees. The following table shows the ages of the employees interviewed: (MARCH-2011)

i) Calculate the mean age of the employees interview.
ii) Compute the mean deviation of the ages about the mean age.

The frequency distribution calculator helps you find the distribution frequency of the numbers in the data set.

Question 2.
Consider the following frequency table. (IMP-2011)

i) Find the mean.
ii) Find the mean deviation about mean.

Question 3.
Consider the following data in respect of marks of 50 students in Mathematics and Physics. (IMP-2011)

i) Find coefficient of variation of Mathematics and Physics.
ii) Which subject shows more variability?
iii) Which subject shows more consistent?

ii) Greater CV more variability; therefore Mathematics is more variable than Physics.
iii) Less CV more consistent, therefore Physics is more consistent.

Question 4.
Find the Standard deviation for the following data: (IMP-2012)

Question 5.
Consider the frequency distribution. (MARCH-2013)

i) Find the mean.
ii) Calculate the variance and the standard deviation.
i)

ii)

Question 6.
Consider the frequency distribution. (MARCH-2013)

i) Find the mean.
ii) Calculate the variance and standard deviation.
i)

ii)

Question 7.
Consider the following frequency table (MARCH-2014)

i) Find the mean.
ii) Find the mean deviation about the mean.

Question 8.
Find the standard deviation of the data: (IMP-2014)

Plus One Maths Statistics 6 Marks Important Questions

Question 1.
The scores of two batsmen A and B in 5 innings during a certain match are as follows: (IMP-2010)

Find:
i) Mean score of each batsman.
ii) Standard deviation of the scores of each batsman.
iii) Which of the batsman is more consistent?

Question 2.
Calculate mean, variance and standard deviation for the following distribution. (IMP-2012)

Standard Deviation = √20100 = 141.8

Question 3.
Calculate the median and Mean deviation about median for the following data. (IMP-2012)

Median class is the class in which the $$\left(\frac{50}{2}=25\right)^{t h}$$ observation lies. Therefore median class is 20 – 30.

Question 4.
Consider the following distribution; (MARCH-2012)

i) Calculate the mean of the distribution.
ii) Calculate the standard deviation of the distribution.
i)

ii)

Question 5.
Consider the following distribution. (IMP-2012)

i) Find the mean.
ii) Find the standard deviation.
iii) Find the coefficient of variation of marks.
i)

ii)

iii)

Question 6.
Consider the frequency distribution (MARCH-2014)

i) Find the mean.
ii) Calculate the variance and the standard deviation.
i)

ii)

Question 7.
i) If $$overline { x }$$ is the mean and a is the standard deviation of a distribution, then the coefficient of variation is ……… (MARCH-2015)
a) $$\frac{\bar{x}}{\sigma} \times 100$$
b) $$\frac{\sigma}{\bar{x}}$$
c) $$\frac{\sigma}{\bar{x}} \times 100$$
d) $$\frac{\bar{x}}{\sigma} \times 50$$
ii) Find the standard deviation for the following data:

i)
$$\frac{\bar{x}}{\sigma} \times 100$$
ii)

Question 8.
i) The sum of all the deviations of the observations of a data from its A.M. is …………. (IMP-2013)
a) Zero
b) Maximum
c) Minimum
d) Negative number
ii) Calculate the Mean. Variance and Standard deviations of the following frequency distribution.

i) a) zero
ii)

Question 9.
i) Suppose the mean of certain number of observation is 50 and the sum of all the observations is 450. Write down the number of observations (MARCH-2016)
ii) Find the mean deviation about mean for the following data:

Question 10.
i) If the variance of a certain distribution is 8, write its standard deviation. (MAY-2017)
ii) Find the mean, standard deviation and coefficient of variation for the following frequency distribution.

i) √8
ii)

Question 11.
i) Find the variance for the observations 2,4,6,8 and 10. (MARCH-2017)
ii) Consider the frequency distribution

i)

ii)

Plus One Maths Notes Chapter 6 Linear Inequalities

Kerala State Board New Syllabus Plus One Maths Notes Chapter 6 Linear Inequalities.

Kerala Plus One Maths Notes Chapter 6 Linear Inequalities

Two real numbers or two algebraic expressions related by the symbols <, >, ≤ or ≥ form an inequality. In this unit we study linear inequalities in one and two variables, their formation and solution graphically.

Solve inequality calculator or quadratic inequalities with our free step-by-step algebra calculator.

I. Linear Inequalities in One Variable
The solution of an inequality in one variable is a value of the variable ‘x’ which makes it a true statement.

Equal numbers can be added or subtracted from both sides of the inequation.

If we multiply or divide both sides of an inequation by a positive number, the inequality sign will not be changed.

If we multiply or divide both sides of an inequation by a negative number, the inequality sign will be reversed.

To represent x < a (or x > a) on a number line, put a circle on the number ‘a’ and a dark line to the left (or right) of the number ‘a’.

To represent x ≤ a (or x ≥ a) on a number line, put a dark circle on the number ‘a’ and a dark line to the left (or right) of the number ‘a’.

II. Linear Inequalities in two Variables
The region containing all the solutions of an inequality is called the solution region.

In order to identify the half-plane represented by inequality, it is just sufficient to take any point (a, b) [say point (0, 0)] not on the line and check whether it satisfies the inequality or not. If it satisfies, then the inequality represents the half-plane and shade the region which contains the point, otherwise, the inequality represents that half-plane which does not contain the point within it.

If the inequality is of the type ax + by ≥ c or ax + by ≤ c, then the point on the line ax + by = c is also included in the solution. So draw a dark line in the solution region.

If the inequality is of the type ax + by > c or ax + by < c, then the point on the line ax + by = c are not to be included in the solution. So draw a broken or dotted line in the solution region.

Plus One Maths Notes Chapter 9 Sequences and Series

Kerala State Board New Syllabus Plus One Maths Notes Chapter 9 Sequences and Series.

Kerala Plus One Maths Notes Chapter 9 Sequences and Series

I. Sequence and Series
A sequence can be regarded as a function whose domain is the set of natural numbers or some subset of it of the type {1, 2, 3, ….., k}.
Generally denoted by a1, a2, …….., an, ………
Let a1, a2, …….., an, …….. be a sequence. Then the expression a1 + a2 + ……. + an + …….. is called the series associated with the given sequence.

II. Arithmetic Progression (AP)
A sequence a1, a2, ……, an, …….. is called an arithmetic sequence or arithmetic progression if an+1 = a2 + d, n ∈ N, where a1 is called the first term and the constant term d is called the common difference of the AP.

Standard form of an AP:
a, a + d, a + 2d, …….. where a is the first term and d is a common difference.

If a constant is added to each term of an AP, the resulting sequence is also an AP.

If a constant is subtracted to each term of an AP, the resulting sequence is also an AP.

If each term of an AP is multiplied by a constant k, the resulting sequence is also an AP. But the resulting AP will have a common difference kd.

If each term of an AP is divided by a constant k, the resulting sequence is also an AP. But the resulting AP will have a common difference $$\frac{d}{k}$$.

nth term, an = a + (n – 1)d

Sum of n terms, Sn = $$\frac{n}{2}$$ [2a + (n – 1)d]

Sn = $$\frac{n}{2}$$ [t1 + tn]

Arithmetic mean between a and b is $$\frac{a+b}{2}$$

III. Geometric Progression (GP):
A sequence a1 + a2 + ……… + an + …….. is called Geometric sequence or Geometric progression if $$\frac{a_{k+1}}{a_{k}}=r$$, k ≥ 1, where a1 is called the first term and the constant term r is called the common ratio of the AP.

Standard form of a GP:
a, ar, ar2,…… where a is the first term and r is a common difference.

nth term, tn = arn-1

Sum of n terms,

Geometric mean between a and b is √ab

Arithmetic mean ≥ Geometric mean.

Infinite G.P, and its Sum G. P. of the form a + ar + ar2 + ar3 + …… ∞ is called infinite G.P.
S = $$\frac{a}{1-r}$$; |r| < 1

Infinite Series Calculator‘ is an online tool that helps to calculate the summation of infinite series for a given function.

IV. Special Series

Plus One Maths Notes Chapter 8 Binomial Theorem

Kerala State Board New Syllabus Plus One Maths Notes Chapter 8 Binomial Theorem.

Kerala Plus One Maths Notes Chapter 8 Binomial Theorem

Binomial theorem gives the expansion of (a + b)n for a rational number ‘n’. In this Unit, we study the binomial theorem for positive integral indices only.

Expanding Binomial Calculator is a free online tool.

I. Binomial Theorem
The expansion of a binomial for any positive integral ‘n’ is given by the binomial theorem.

There are (n + 1) terms in the expansion of (a + b)n.

The sum of the indices of ‘a’ and ‘b’ in every term of the expansion is ‘n’.

The general term in the expansion is tr+1 = nCr an-r br

Middle term in the expansion:
1. If ‘n’ is even, $$\left(\frac{n}{2}+1\right)^{t h}$$ term.

2. If ‘n’ is odd, $$\left(\frac{n+1}{2}\right)^{t h}$$ and $$\left(\frac{n+1}{2}+1\right)^{b_{t}}$$ term.

Plus One Malayalam Textbook Answers Unit 4 Chapter 6 Shasthrakriya

Kerala State Board New Syllabus Plus One Malayalam Textbook Answers Unit 4 Chapter 6 Shasthrakriya Text Book Questions and Answers, Summary, Notes.

Plus One Accountancy Chapter Wise Questions and Answers Kerala

HSE Kerala Board Syllabus HSSLive Plus One Accountancy Chapter Wise Questions and Answers Pdf Free Download in both English Medium and Malayalam Medium are part of SCERT Kerala Plus One Chapter Wise Questions and Answers. Here HSSLive.Guru has given Higher Secondary Kerala Plus One Accountancy Chapter Wise Questions and Answers based on CBSE NCERT syllabus.

 Board SCERT, Kerala Text Book NCERT Based Class Plus One Subject Accountancy Chapter All Chapters Category Kerala Plus One

Kerala Plus One Accountancy Chapter Wise Questions and Answers

We hope the given HSE Kerala Board Syllabus HSSLive Plus One Accountancy Chapter Wise Questions and Answers Pdf Free Download in both English Medium and Malayalam Medium will help you. If you have any query regarding Higher Secondary Kerala Plus One Accountancy Chapter Wise Questions and Answers based on CBSE NCERT syllabus, drop a comment below and we will get back to you at the earliest.

Plus One Economics Chapter Wise Questions and Answers Chapter 4 Presentation of Data

Students can Download Chapter 4 Presentation of Data Questions and Answers, Plus One Economics Chapter Wise Questions and Answers helps you to revise the complete Kerala State Syllabus and score more marks in your examinations

Kerala Plus One Economics Chapter Wise Questions and Answers Chapter 4 Presentation of Data

Plus One Economics Presentation of Data One Mark Questions and Answers

Plus One Economics Chapter Wise Questions And Answers Pdf Question 1.
Which of the following comes under geometric diagram?
(a) Histogram
(b) Bar diagram
(c) Ogives
(d) Frequency polygon
(b) Bar diagram

Plus One Economics Chapter Wise Questions And Answers Question 2.
Which of the following comes under frequency diagrams?
(a) Bar diagram
(b) Histogram
(c) Pie diagram
(d) All the above
(b) Histogram

Plus One Economics Chapter Wise Questions And Answers Malayalam Question 3.
To draw time-series graph, time is presented on:
(a) X-axis
(b) Y-axis
(c) any of two
(a) X-axis

Plus One Statistics Chapter Wise Questions And Answers Question 4.
Name the types of graphs.

1. One dimensional graph
2. Two-dimensional graph
3. Three-dimensional graph
4. Pictograms

State whether true or false.

1. The width of bars in a bar diagram need not be equal.
2. The width of rectangles in a histogram should essentially be equal.
3. Histograms can only be formed with continuous classification of data.
4. Histogram and column diagram are the same method of presentation of data.
5. Mode of a frequency distribution can be drawn graphically with the help of histogram,
6. The median of a frequency distribution cannot be drawn from the Ogive.

1. true
2. false
3. true
4. true
5. true
6. true

Plus One Economics Presentation of Data Two Mark Questions and Answers

Hsslive Economics Plus One Chapter Wise Questions And Answers Question 1.
Which of the following is a cumulative frequency curve?
(a) Bar diagram
(b) Histogram
(c) Ogive
(d) Pie diagram
(c) Ogive

Plus One Economics Questions And Answers Question 2.
Distinguish between captions and stubs.
Captions refers to the column headings and stubs refers to the row heading.

Histogram And Frequency Polygon Questions With Answers Question 3.
Match the following.

 A B Source note Row headings Captions Gives origin of data Stubs Explains the specific feature Footnote Column Headings

 A B Source note Gives origin of data Captions Column Headings Stubs Row Headings Footnote Explains the specific feature

Plus One Economics Presentation of Data Three Mark Questions and Answers

Presentation Of Data Class 11 Question 1.
What kind of diagrams are more effective in representing the following?

1. Monthly rainfall in a year
2. Composition of the population of Delhi by religion
3. Components of cost in a factory

1. Simple bar diagram
2. Sub-divided or component bar diagram
3. Pie diagram

Pie Chart Class 11 Economics Question 2.
Name different types of diagrams.
The different types of diagrams are:
1. Geometric diagram

• Bar diagrams
• Pie diagram

2. Frequency diagram

• Histogram
• Frequency polygon
• Frequency curve -Ogive

3. Arithmetic line graph

Question 3.
“Diagrams and graphs help us visualize the whole meaning of numerical complex data at a single glance”. Comment.
One of the most convincing and appealing ways in which statistical results may be presented is through diagrams and graphs. The special feature of graphs and diagrams is that they do away with figures altogether. Diagrams and graph is a statistical method which can be used for simplifying the complexity of quantitative data and t make them easily intelligible.

It presents dry and uninteresting statistical facts in the shape of attractive and appealing pictures and charts. They are important methods of visual aids and are appealing t the eye and mind of the observer.

Question 4.
“There are generally three forms of diagrammatic presentation of data” explain.
There are various methods to present data. But generally, three forms of presentation of data are there
which are noted below:

1. Geometric diagram
2. Frequency diagram
3. Arithmetic line graph

1. Geometric Diagram:
Bar diagram and pie diagram come in the category of geometric diagram for presentation of data. The bar diagrams are of three types-simple, multiple and component bar diagrams.

2. Frequency Diagram:
Data in the form of grouped frequency distributions are generally represented by frequency diagrams like histogram, frequency polygon, frequency curve, and ogive

3. Arithmetic Line Graph:
An arithmetic line graph is also called time-series graph and is a method of diagrammatic presentation of data. A line graph by joining these plotted points, thus, obtained is called arithmetic line graph or time-series graph.

Question 5.
Explain Ogive?
Cumulative frequency of any class is equal to the sum of the frequencies of all the classes preceding that class and its own frequency e.g., frequencies are 10, 7, 12, 17 and 22. Cumulative frequencies are 10, 10 + 7 = 17, 17 + 12 = 29, 29 + 17 = 46 and 46 + 22 = 68.
Cumulative frequency of the last class = Total frequency.

For drawing an Ogive, cumulative frequency (i.e. number of values) is taken on the Y-axis and limits of class intervals on the X-axis.
Ogive is of two types:

1. less than
2. more than

In a “less than” type Ogive, we plot the upper limit of each class along the X-axis and in a “more than” type Ogive, we plot the lower limit of each class along the X-axis. Along the Y-axis, we plot the cumulative frequencies at the end of each class. Ogive can be drawn even if the class interval are unequal or open end. Ogives are performed over frequency curves for comparative study.

Question 6.
Illustrate how classes can be formed while presenting the data?
Classes can be formed in two ways:

1. Exclusive type
2. Inclusive type

1. Exclusive Type:
When the class intervals are so fixed that the upper limit of one class is the lower limit of the new class, it is known as exclusive method of classification.

 Marks (Percentage) No. of students 0-10 15 10-20 17 20-30 22 30-40 30 40-50 39 50-60 45

In this method, higher value of the variable in the class is not included in that class i.e.,

 Marks (Percentage) No. of students 0 and more but less than 10 15 10 and more but less than 20 17 20 and more but less than 30 22 30 and more but less than 40 30 40 and more but less than 50 39 50 and more but less than 60 45

2. Inclusive Type:
In this method, the students getting say 39% marks will be included in class 30 – 39 itself i.e.,

 Marks (Percentage) No. of students 0-9 5 10-19 8 20-29 7 30-39 13 40-49 25

Plus One Economics Presentation of Data Four Mark Questions and Answers

Question 1.
a. Bar diagram is a

1. one-dimensional diagram
2. two-dimensional diagram
3. diagram with no dimension
4. none of the above

b. Data represented through a histogram can help in finding graphically the

1. mean
2. mode
3. median
4. all the above

c. Ogives can be helpful in locating graphically the

1. mode
2. mean
3. median
4. none of the above

d. Data represented through arithmetic line graph help in understanding

1. long term trend
2. cyclicity in data
3. seasonality in data
4. all the above

a. 1. one-dimensional diagram
b. 3. mode
c. 3. median
d. 1. long term trend

Question 2.
Point out major parts of a statistical table.

1. Table number
2. Title
4. Stub
6. Body or field
7. Footnote
8. Source note

Question 3.
Give the rules for constructing tables.
The rules of constructing diagrams are:

• Every diagram should be titled.
• It should suit the size of the paper
• It should be neat and attractive
• It should be neatly indexed
• It should contain footnotes
• The details in diagram should be self-explanatory

Plus One Economics Presentation of Data Five Mark Questions and Answers

Question 1.
Explain the advantages of diagrammatic presentation.
The advantages of diagrammatic presentation are given below.

1. Diagram give a clear picture of data
2. Comparison can be made easy
3. Diagrams can be used university at any place
4. It saves time and energy
5. The data can be remembered easily

Question 2.
Show how pie diagram is drawn for the following data?

 Items Production in K.G. Tea 3260 Coffee 1850 Cocoa 900 Total 6010

Question 3.
Give steps in the preparation of pie diagram.
A pie diagram is also a component diagram, but unlike a component bar diagram, a circle whose area is proportionally divided among the components it represents. It is also called a pie chart. The circle is divided into as many parts as there are components by drawing straight lines from the centre to the circumference.

The following steps in the preparation of pie diagram are given below:

• Convert each component as percentage of the total.
• Multiply the percentage by 360/100 = 3.6 to convert into degree.
• Starting with the twelve o’clock position on the circle draw the largest component circle
• Draw other components in clockwise succession in descending order of magnitude except for each all components

Like all others and miscellaneous which are shown last:

• Use different columns or shades to distinguish between different components
• Explain briefly the different components either within the components in the figure or outside by arrow.

Plus One Economics Presentation of Data Eight Mark Questions and Answers

Question 1.
Write short notes on the following

1. pie diagrams
2. frequency curves
3. frequency polygon
4. ogive
5. arithmetic line graph

1. Pie Diagram:
A pie diagram is also a component diagram, but unlike a component bar diagram, a circle whose area is proportionally divided among the components it represents. It is also called a pie chart. The circle is divided into as many parts as there are components by drawing straight lines from the centre to the circumference. Pie charts usually are not drawn with absolute values of a category.

The values of each category are first expressed as percentage of the total value of all the categories. A circle in a pie chart, irrespective of its value of radius, is thought of having 100 equal parts of 3.6° (3607100) each. To find out the angle, the component shall subtend at the centre of the circle, each percentage figure of every component is multiplied by 3.6°.

2. Frequency Polygon:
A frequency polygon is a plane bounded by straight lines, usually four or more lines. Frequency polygon is an alternative to histogram and is also derived from histogram itself. A frequency polygon can be fitted to a histogram for studying the shape of the curve. The simplest method of drawing a frequency polygon is to join the midpoints of the topside of the consecutive rectangles of the histogram.

3. Frequency Curve:
The frequency curve is obtained by drawing a smooth freehand curve passing through the points of the frequency polygon as closely as possible. It may not necessarily pass through all the points of the frequency polygon but it passes through them as closely as possible

4. Ogive:
Ogive is also called cumulative frequency curve. As there are two types of cumulative frequencies, for example, less than type and more than type, accordingly there are two ogives for any grouped frequency distribution data. Here in place of simple frequencies as in the case of frequency polygon, cumulative frequencies are plotted along y-axis against class limits of the frequency distribution.

For less than give the cumulative frequencies are plotted against the respective upper limits of the class intervals whereas for more than ogives the cumulative frequencies are plotted against the respective lower limits of the class interval. An interesting feature of the two ogives together is that their intersection point gives the median

5. Arithmetic Line Graph:
An arithmetic line graph is also called time-series graph and is a method of diagrammatic presentation of data. Init, time (hour, day/date, week, month, year, etc.) is plotted along x-axis and the value of the variable (time series data) along y-axis. A line graph by joining these plotted points, thus, obtained is called arithmetic line graph (time series graph). It helps in understanding the trend, periodicity, etc. in a long term time series data.

Question 2.
3 Forms of presentation of data

1. Textual
2. Tabular
3. Diagrams & graphs Prepare a flow chart.