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Kerala Plus One Chemistry Notes Chapter 12 Organic Chemistry: Some Basic Principles and Techniques

Introduction
The element carbon organic chemistry the element carbon has the unique property called catenation due to which it forms covalent bonds with other carbon atoms. It also forms covalent bonds with atoms of other elements like hydrogen, oxygen, nitrogen, sulphur, phosphorus and halogens. The resulting compounds are studied under a separate branch of chemistry called organic chemistry.

General Introduction
In early years of chemistry, compounds were classified into two types. Compounds derived from non-living sources such as rocks, minerals etc. were called ‘inorganic compounds’ and those derived from plants and animals were regarded as ‘organic compounds’. On account of the special nature of organic compounds and their occurrence in living world alone, it was believed that they were produced by a vital force existing in living organisms. This led to the belief that such compounds could not be synthesised in the laboratory. However in 1828, F. Wohler succeeded in preparing urea (an organic compound) from an inorganic material, ammonium cyanate.

Tetravalance Of Carbon

Shapes of Organic Compounds
Carbon (atomic number 6) has the ground state electronic configuration 1 s² 2s²p¹_x2p¹_y2p°_z. Carbon attains noble gas configuration only by sharing electrons with other atoms. Carbon atom, therefore, forms four covalent bonds in all its compounds.

During the formation of bonds (which is an energy releasing process) the two electrons in the 2s orbital get unpaired and one is promoted to the empty 2p_z orbital. This corresponds to the excited state of carbon which has four half-filled orbitals (four valence electrons). The shapes of molecules such as methane (CH₄), ethene (C₂H₄) and ethyne (C₂H₂) are explained in terms of the use of sp³, sp² and sp hybridised orbitals by the carbon atoms in the respective molecules.

Structural Representations Of Organic Compounds

Complete, Condensed and Bond-line Structural Formulas
Structures of organic compounds are represented in several ways. The Lewis structure or dot structure, dash structure, condensed structure and bond line structural formulas are some of the specific types. In Lewis structures bonds are represented by lines and lone pair electrones by dots. For example,

In bond-line structural representation of organic compounds, carbon and hydrogen atoms are not shown and the lines representing carbon-carbon bonds are drawn in a zig-zag fashion. The only atoms specifically written are oxygen, chlorine, nitrogen etc. The terminals denote methyl (-CH₃) groups.

Classification Of Organic Compounds

I. Acyclic or open chain compounds
These compounds are also called as aliphatic compounds and consist of straight or branched chain compounds, for example:

II. Alicyclic or closed chain or ring compounds
Alicyclic (aliphatic cyclic) compounds contain carbon atoms joined in the form of a ring (homocyclic). Sometimes atoms other than carbon are also present in the ring (heterocyclic). Some examples of this type of compounds are:
CycloiTexane Tetrahydrofuran These exhibit some of the properties similar to those of aliphatic compounds ane:

Aromatic compounds
Aromatic compounds are special types of compounds. You will learn about these compounds in detail in Unit 13. These include benzene and other related ring compounds (benzenoid). Like alicyclic compounds, aromatic comounds may also have hetero atom in the ring. Such compounds are called heterocyclic aromatic compounds. Some of the examples of various types of aromatic compounds are:

Benzenoid aromatic compounds

Non-benzenoid compound

Heterocyclic aromatic compounds

Organic compounds can also be classified on the basis of functional groups, into families or homolo-gous series.

Homologous series
Homologous series may be defined as a series of similarly constituted organic compound in which the members possess the same functional group and have similar chemical properties and the neighbouring (or consecutive) members differ by – CH₂ unit in their molecular formula,
eg: Alkane family (Cn H_2n+2), alkenes (CnH_2n) alcohols (Cn H_2n+1OH).

Nomenclature Of Organic Compounds
IUPAC names of unbrached saturated hydrocarbons

Nomenclature of branched chain alkanes:
We encounter a number of branched chain alkanes. The rules for naming them are given below.
1. First of all, the longest carbon chain in the molecule is identified.For example,

2. The numbering is done in such a way that the branched carbon atoms get the lowest possible numbers.

and the numbering is not from right to left.

3. The names of alkyl groups attached as a branch are then prefixed to the name of the parent alkane and position of the substituents is indicated by the appropriate numbers. If different alkyl groups are present, they are listed in alphabetical order. Thus, name for the compound shown above is: 6- Ethyl-2- methylnonane.
[Note: the numbers are separated from the groups by hyphens and there is no break between methyl and nonane.]

4. If two or more identical substituent groups are present then the numbers are separated by commas. The names of identical substituents are not repeated, instead prefixes such as di (for 2), tri (for 3), tetra (for 4), Penta (for 5), Hexa (for 6) etc. are used. While writing the name of the substituents in alphabetical order, these prefixes, however, are not considered. Thus, the following compounds are named as:

5. If the two substituents are found in equivalent positions, the lower number is given to the one coming first in the alphabetical listing. Thus, the following compound is 3-Ethyl-6-methyl octane and ‘ not6-Ethyl-3-methyl octane.

6. The branched alkyl groups can be named by following the above mentioned procedures. However, the carbon atom of the branch that attaches to the root alkane is numbered 1 as exemplified below.

Cyclic Compounds
A saturated monocyclic compound is named by prefixing ‘cyclo’ to the corresponding straight chain alkane. If side chains are present, then the rules given above are applied. Names of some cyclic compounds are given below.

Nomenclature Of Organic Compounds Having Functional Group (S)
When a functional group (other than C=C and -C ≡ C) is present in the molecule, it is indicated by adding secondary suffix after the primary suffix. The terminal ‘e’ of the primary suffix is removed before adding secondary suffix (whose name begins with a, i, o, u or y). It is to be noted that some functional group such as alkoxy (-OR), nirto (-NO₂), halogeno etc. are indicated by the prefixes.

For example, CH₃-CH₂-OH Ethane -e+ol = Ethanol (using secondary suffix) CH₃-CH₂-CHO Propane -e+al = Propanal (using secondary suffix) CH₃-CH₂-NO₂ Nitroethane (using prefix)

The systematic name of an organic compound containing functional group can be derived using the following sequence of steps.

The longest carbon chain (parent chain) containing the functional groups is identified. This gives the word root. The name of the compound is then obtained as follows:
Prefixes – word root – primary suffix – secondary suffix The following examples will illustrate the rules

In case of compounds containing more than one similar functional group, the world di, tri etc is added before the secondary suffix which indicates the functional group. In doing so, the last letter ‘e’ of the parent alkane has to be retained. However, the endingne of the parent alkane is dropped in case of compounds having more than one double or triple bond. For example,

If the molecule contains two or more different functional groups, the parent chain must contain maximum possible number of functional groups. The carbon atoms in the parent chain are numbered in such a way that the functional group of higher priority gets the lower number. The priority of various functional groups follows the order.

COOH>-CO-O-CO->-COOR>-COCI>-CONH₂>-CN>- HC=O>CO>-OH>-NH₂>>C=C<-C≡C->-X>-NO₂>R-
The functional group with higher priority is indicated by suitable secondary suffix and the other functional groups are treated as substituents which are specified by suitable prefixes.
For example,

Nomenclature of Substituted Benzene Compounds
For IUPAC nomenclature of substituted benzene compounds, the substituent is placed as prefix to the word benzene as shown in the following examples.

If benzene ring is disubstituted, the position of substituents is defined by numbering the carbon atoms of the ring such that the substituents are located at the lowest numbers possible. For example, the compound(b) is named as 1, 3-Dibromobenzene and not as 1,5 dibromobenzene. Substituent of the base compound is assigned number 1 and then the direction of numbering is chosen such that the next substituent gets the lowest number. The substituents appear in the name in alphabetical order. Some examples are given below.

In the trivial system of nomenclature the terms ortho (o), meta (m) and para (p) are used as prefixes to indicate the relative positions 1,2-;1,3- and Ir-respectively. Thus, 1,3-dibromobenzene (b) is named as m-dibromobenzene (meta is abbreviated as m-) and the other isomers of dibromobenzenel, 2-(a) and 1,4-(c), are named as ortho (or just o-) and para (or just p-)-dibromobenzene, respectively. The substituents appear in the name in alphabetical order.

Isomerism
The phenomenon of existence of two or more compounds possessing the same molecular formula but different properties is known as isomerism. Such compounds are called as isomers.

Structural Isomerism
Compounds having the same molecular formula but different structures (manners in which atoms are linked) are classified as structural isomers. Some typical examples of different types of structural isomerism are given below:

(i) Chain isomerism:
When two or more compounds have similar molecular formula but different carbon skeletons, these are referred to as chain isomers and the phenomenon is termed as chain isomerism.

ii) Position isomerism:
When two or more compounds differ in the position of substituent atom or functional group on the carbon skeleton, they are called position isomers and this phenomenon is termed as position isomerism.

iii) Functional group isomerism :
Two or more compounds having the same molecularformula but different functional groups are called functional isomers and this phenomenon is termed as functional group isomerism.

iv) Metamerism :
It arises due to different alkyl chains on either side of the functional group in the molecule. For example, C_aH₁₀O represents Methoxypropane (CH₃OC₃H₇) and Ethoxyethane (C₂H₅OC₂H₅).

Fundamental Concepts In Organic Reaction Mechanism
An organic reaction takes place by the attack of a reagent on an organic compound which is designated as a substrate. The steps of an organic reaction showing the breaking and formation of bonds in such substrate leading to the formation of the final product are referred to as its mechanism.

Fission Of A Covalent Bond
A covalent bond can be broken in two different ways,

(i) Homolytic fission :
If a covalent bond breaks in such a way that each atom takes away one electron of the shared pair. It is called homolytic fission or homolysis. The fragments with odd or unpaired electrons formed by homolysis are known as free radicals. For example,

Usually homolysis occurs at high temperature or in presence of high energy radiations. Reactions occurring through homolytic fission are known as free radical reactions (non-polar reactions).

(ii) Heterolytic fission :
When a covalent bond breaks in such a way that both the electrons of the covalent bond are taken away by one of the bonded atoms, the mode of cleavage is called heterolytic cleavage or heterolysis. The products of heterolysis of a covalent bond are positive and negative ions, eg:

Inductive Effect (I Effect)
It is the permanent polarisation of a sigma bond in a molecule by the influence of an adjacent polar bond or group.

For illustration, let us consider a carbon chain in which one terminal carbon atom is joined to a chlorine atom. Since the chlorine atom is more electronegative than carbon, the sigma electrons of the C-Cl bond are displaced towards the chlorine atom. As a result, the chlorine atom acquires a small negative charge and C acquires a small positive charge as shown below. The magnitude of+charge is of the order C₁ > C₂ >C₃.

This type of electron displacement of sigma electrons along a saturated carbon chain due to the presence of an electron-withdrawing group (or electron-donating group) is called Inductive effect:

This effect decreases sharply with increasing dis-tance from the substituent and becomes negligible afterthe third carbon in a chain. Atoms orgroups of atoms that attract or withdraw electrons from a chain are said to have electron withdrawing inductive effect or-l effect, eg:,
-NO₂ >- CN >- COOH >- F >- Cl >- Br >-l

Atoms or groups which push or donate electrons to a carbon chain are said to have electron releasing inductive effect or + l effect. Alkyl groups have +l effect and the order of + l effect of alkyl groups is

Resonance Structure
There are many organic molecules whose behaviour cannot be explained by a single Lewis structure. An example is that of benzene. Its cyclic structure containing alternating C-C single and C=C double bonds shown is inadequate for explaining its characteristic properties.

As per the above representation, benzene should exhibit two different bond lengths, due to C-C single and C=C double bonds. Thus, the structure of benzene cannot be represented adequately by the above structure. Further, benzene can be represented equally well by the energetically identical structures I and ll. The actual structure of benzene cannot be adequately represented by any of these structures, rather it is a hybrid of the two structures (I and II) called resonance structures.

The resonance structures are hypothetical and individually do not represent any real molecule. They contribute to the actual structure in proportion to their stability. The energy of actual structure of the molecule (the resonance hybrid) is lower than that of any of the canonical structures. The difference in energy is called the resonance stabilisation energy or simply the resonance energy. The more the number of important contributing structures, the more is the resonance energy. Resonance is particularly important when the contributing structures are equivalent in energy.

Resonance Effect
The resonance effect is defined as ‘the polarity produced in the molecule by the interaction of two π- bonds or between a π-bond and lone pair of electrons present on an adjacent atom’.The effect is transmitted through the chain. There are two types of resonance or mesomeric effect designated as R or M effect.
(i) Positive Resonance Effect (+R effect)
In this effect, the transfer of electrons is away from an atom or substituent group attached to the conjugated system. This electron displacement makes certain positions in the molecule of high electron densities.

(ii) Negative Resonance Effect (- R effect)
This effect is observed when the transfer of electrons is towards the atom or substituent group attached to the conjugated system.The atoms or substituent groups, which represent +R or-R electron displacement effects are as follows:
+R effect: – halogen, -OH, -OR, -OCOR, -NH₂, – NHR,-NR₂,-NHCOR,
– R effect: – COOH, -CHO, >C=O, – CN, -NO₂

Electromeric Effect (E – Effect)
This is temporary effect which involves the complete transfer of n electrons of a multiple bond to one of the bonded atoms in presence of an attacking reagent. However, when the attacking reagent is removed, the polarised molecule shifts back to its original electronic condition.

With transfer of π electrons takes place towards the attacking reagent the effect is called +E effect and when the transfer of π electrons occurs-S away from the attacking reagent the effect is called -E effect.

Hyperconjugation
Hyperconjugation is a general stabilising interaction. It involves delocalisation of σ electrons of C—H bond of an alkyl group directly attached to an atom of unsaturated system or to an atom with an unshared p orbital. The σ electrons of C—H bond of the alkyl group enter into partial conjugation with the attached unsaturated system or with the unshared p orbital. Hyperconjugation is a permanent effect.

Methods Of Purification Of Organic Compounds
Sublimation:
On heating, some solid substances change from solid to vapour state without passing through liquid state. The purification technique based on the above principle is known as sublimation and is used to separate sublimable compounds from nonsublimable impurities.

Crystallisation:
It is based on the difference in the solubilities of the compound and the impurities in a suitable solvent. The impure compound is dissolved . in a solvent in which it is sparingly soluble at room temperature but appreciably soluble at higher temperature. The solution is concentrated to get a nearly saturated solution. On cooling the solution, pure compound crystallises out and is removed by filtration. The filtrate (mother liquor) contains impurities and small quantity of the compound.

Distillation:
This important method is used to separate (i) volatile liquids from non-volatile impurities and (ii) the liquids having sufficient difference in their boiling points. Liquids having different boiling points vaporise at different temperatures. The vapours are cooled and the liquids so formed are collected separately. Chloroform (b.p 334 K) and aniline (b.p. 457 K) are easily separated by the technique of distillation. On boiling, the vapours of lower boiling component are formed first. The vapours are condensed by using a condenser and the liquid is collected in a receiver. The vapours of higher boiling component form later and the liquid can be collected separately.

Fractional Distillation:
If the difference in boiling points of two liquids is not much, simple distillation cannot be used to separate them. The vapours of such liquids are formed within the same temperature range and are condensed simultaneously. The technique of fractional distillation is used in such cases. In this technique, vapours of a liquid mixture are passed through a fractionating column before condensation. The fractionating column is fitted over the mouth of the round bottom flask. Vapours of the liquid with higher boiling point condense before the vapours of the liquid with lower boiling point. The vapours rising up in the fractionating column become richer in more volatile component. By the time the vapours reach to the top of the fractionating column, these are rich in the more volatile component. The vapours become richer in low boiling component. On reaching the top, the vapours become pure in low boiling component and pass through the condenser and the pure liquid is collected in a receiver.

After a series of successive distillations, the remaining liquid in the distillation flask gets enriched in high boiling component. Each successive condensation and vaporisation unit in the fractionating column is called a theoretical plate. One of the technological applications of fractional distillation is to separate different fractions of crude oil in petroleum industry. Distillation under reduced pressure: This method is used to purify liquids having very high boiling points and those, which decompose at or below their boiling points. Such liquids are made to boil at a temperature lower than their normal boiling points by reducing the pressure on their surface. A liquid boils at a temperature at which its vapour pressure is equal to the external pressure. The pressure is reduced with the help of a water pump or vacuum pump. Glycerol can be separated from spent-lye in soap industry by using this technique.

Steam Distillation:
This technique is applied to separate substances which are steam volatile and are immiscible with water. In steam distillation, steam from a steam generator is passed through a heated flask containing the liquid to be distilled. The mixture of steam and the volatile organic compound is condensed and collected. The compound is later separated from water using a separating funnelAniline is separated by this technique from aniline-water mixture.

Differential Extraction:
When an organic compound is present in an aqueous medium, it is separated by shaking it with an organic solvent in which it is more soluble than in water. The organic solvent and the aqueous solution should be immiscible with each other so that they form two distinct layers which can be separated by separatory funnel. The organic solvent is later removed by distillation or by evaporation to get back the compound.

Chromatography:
The name chromatography is based on the Greek word chroma, for colour since the method was first used for the separation of coloured substances found in plants. In this technique, the mixture of substances is applied onto a stationary phase, which may be a solid ora liquid.

A pure solvent, a mixture of solvents, or a gas is allowed to move slowly over the stationary phase. The components of the mixture get gradually separated from one another. The moving phase is called the mobile phase. Based on the principle involved, chromatography is classified into different categories. Two of these are:

1. Adsorption chromatography, and
2. Partition chromatography.

1. Adsorption Chromatography:
Adsorption chromatography is based on the fact that different compounds are adsorbed on an adsorbent to different degrees. Commonly used adsorbents are silica gel and alumina. When a mobile phase is allowed to move over a stationary phase (adsorbent), the components of the mixture move by varying distances over the stationary phase. Following are two main types of chromatographic techniques based on the principle of differential dsorption.

1. Column chromatography, and
2. Thin layer chromatography.

Column Chromatography
Column chromatography involves separation of a mixture overa column of adsorbent (stationary phase) packed in a glass tube. The column is fitted with a stopcock at its lower soluble in the organic solvent. The mixture adsorbed on adsorbent is placed on the top of the adsorbent column packed in a glass tube. An appropriate eluant which is a liquid or a mixture of liquids is allowed to flow down the column slowly. Depending upon the degree to which the compounds are adsorbed, complete separation takes place. The most readily adsorbed substances are retained near the top and others come down to various distances in the column.

Thin layer chromatography (TLC) is another type of adsorption chromatography, which involves separation of substances of a mixture over a thin layer of an adsorbent coated on glass plate. A thin layer of an adsorbent (silica gel or alumina) is spread overa glass plate of suitable size. The plate is known as thin layer chromatography plate or chrome plate. The solution of the mixture to be separated is applied as a small spot about 2 cm above one end of the TLC plate. The glass plate is then placed in a closed jar containing the eluant. As the eluant rises up the plate, the components of the mixture move up along with the eluant to different distances depending on their degree of adsorption and separation takes place. The relative adsorption of each component of the mixture is expressed in terms of its retardation factor i.e. R_f value
R_f = x/y

Where distance moved by the substance from base line is x and distance moved by the solvent from base line is y. The spots of coloured compounds are visible on TLC plate due to their original colour. Another detection technique is to place the plate in a covered jar containing a few crystals of iodine. Spots of compounds, which adsorb iodine, will show up as brown spots. Sometimes an appropriate reagent may also be sprayed on the plate. For example, amino acids may be detected by spraying the plate with ninhydrin solution.

Partition Chromatography:
Partition chromatography is based on continuous differential partitioning of components of a mixture between stationary and mobile phases. Paper chromatography is a type of partition chromatography. In paper chromatography, a special quality paper known as chromatography paper is used. Chromatography paper contains water trapped in it, which acts as the stationary phase. A strip of chromatography paper spotted at the base with the solution of the mixture is suspended in a suitable solvent ora mixture of solvents. This solvent acts as the mobile phase. The solvent rises up the paper by capillary action and flows over the spot. The paper selectively retains different components according to their differing partition in the two phases. The paper strip so developed is known as a chromatogram. The spots of the separated coloured compounds are visible at different heights from the position of initial spot on the chromatogram. The spots of the separated colourless compounds may be observed either under ultraviolet light or by the use of an appropriate spray reagent as discussed under thin layer chromatography.

Qualitative Analysis Of Organic Compounds
Thin layer chromatography (TLC) is another type of adsorption chromatography, which involves separation of substances of a mixture over a thin layer of an adsorbent coated on glass plate. A thin
The elements present in organic compounds are carbon and hydrogen. In addition.to these, they may also contain oxygen, nitrogen, sulphur, halogens and phosphorus.

Detection of Carbon and Hydrogen
Carbon and hydrogen are detected by heating the compound with copper(ll) oxide. Carbon present in the compound is oxidised to carbon dioxide (tested with lime-water, which develops turbidity) and hydrogen to water (tested with anhydrous copper sulphate, which turns blue).

Detection of Other Elements
Nitrogen, sulphur, halogens and phosphorus present in an organic compound are detected by “Lassaigne’s test”. The elements present in the compound are converted from covalent form into the ionic form by fusing the compound with sodium metal. Following reactions take place:

C, N, Sand X come from organic compound. Cyanide, sulphide and halide of sodium so formed on sodium fusion are extracted from the fused mass by boiling it with distilled water. This extract is known as sodium fusion extract.

1. Test for Nitrogen
The sodium fusion extract is boiled with iron(ll) sulphate and then acidified with concentrated sulphuric acid. The formation of Prussian blue colour confirms the presence of nitrogen. Sodium cyanide first reacts with iron(ll) sulphate and forms sodium hexacyanoferrate(ll). On heating, with concentrated sulphuric acid some iron(ll) ions are oxidised to iron(lll) ions which react with sodium hexacyanoferrate(ll) to produce iron(lll) hexacyanoferrate(ll) (ferric ferrocyanide) which is Prussian blue in colour.

2. Test for Sulphur
1. The sodium fusion extract is acidified with acetic acid and lead acetate is added to it. A black precipitate of lead sulphide indicates the presence of sulphur.

2. On treating sodium fusion extract with sodium nitroprusside, appearance of a violet colour further indicates the presence of sulphur.

In case, nitrogen and sulphur both are present in an organic compound, sodium thiocyanate is formed. It gives blood-red colour and no Prussian blue since there are no free cyanide ions.

If sodium fusion is carried out with excess of sodium, the thiocyanate decomposes to yield cyanide and sulphide. These ions give their usual tests.
NaSCN + 2Na → NaCN + Na₂S

3. Test for Halogens
The sodium fusion extract is acidified with nitric acid and then treated with silver nitrate. A white precipitate, soluble in ammonium hydroxide shows the presence of chlorine, a yellowish precipitate, sparingly soluble in ammonium hydroxide shows the presence of bromine and a yellow precipitate, insoluble in ammonium hydroxide shows the presence of iodine.
X^– + Ag⁺ → AgX

X represents a halogen – Cl, Br or I. If nitrogen or sulphur is also present in the compound, the sodium fusion extract is first boiled with concentrated nitric acid to decompose cyanide or sulphide of sodium formed during Lassaigne’s test. These ions would otherwise interfere with silver nitrate test for halogens.

4. Test for Phosphorus
The compound is heated with an oxidising agent (sodium peroxide). The phosphorus present in the compound is oxidised to phosphate. The solution is boiled with nitric acid and then treated with ammonium molybdate. A yellow colouration or precipitate indicates the presence of phosphorus.

Quantitative Analysis
The percentage composition of elements present in an organic compound is determined by the methods based on the following principles:

Carbon and Hydrogen
Both carbon and hydrogen are estimated in one experiment. A known mass of an organic compound is burnt in the presence of excess of oxygen and copper(ll) oxide. Carbon and hydrogen in the compound are oxidised to carbon dioxide and water respectively.
C_xH_y + (x + y /4)O₂ → xCO₂ +(y/2)H₂O

The mass of water produced is determined by passing the mixture through a weighed U-tube containing anhydrous calcium chloride. Carbon dioxide is absorbed in another U-tube containing concentrated solution of potassium hydroxide. These tubes are connected in series. The increase in masses of calcium chloride and potassium hydroxide gives the amounts of water and carbon dioxide from which the percentages of carbon and hydrogen are calculated. Let the mass of organic compound be m g, mass of water and carbon dioxide produced be m₁ and m₂g respectively;

Nitrogen
There are two methods for estimation of nitrogen: (i) Dumas method and (ii) Kjeldahl’s method.
(i) Dumas method:
The nitrogen containing organic compound, when heated with copper oxide in an atmosphere of carbon dioxide, yields free nitrogen in addition to carbon dioxide and water.
C_xH_yN_z+(2x + y/2)CuO → x CO₂ + y / 2H₂O +Z / 2N₂ + (2x + y / 2)CU

Traces of nitrogen oxides formed, if any, are reduced to nitrogen by passing the gaseous mixture over a heated copper gauze. The mixture of gases so produced is collected over an aqueous solution of potassium hydroxide which absorbs carbon dioxide. Nitrogen is collected in the upper part of the graduated tube. Let the mass of organic compound = m g Volume of nitrogen collected = V₁ mL
Room temperature = T₁K
Volume of nitrogen at STP = 7\(\frac{p_{1} v_{1} \times 273}{760 \times T_{1}}\) (LetitbeVmL)

Where p₁ and V₁ are the pressure and volume of nitrogen, p,is different from the atmospheric pressure at which nitrogen gas is collected. The value of pt is obtained by the relation;
p₁ Atmospheric pressure-Aqueous tension 22400 mL N₂ at STP weighs 28 g.
\frac{p_{1} v_{1} \times 273}{760 \times T_{1}}

(ii) Kjeldahl’s method:
The compound containing nitrogen is heated with concentrated sulphuric acid. Nitrogen in the compound gets converted to ammonium sulphate.The resulting acid mixture is then heated with excess of sodium hydroxide. The liberated ammonia gas is absorbed in an excess of standard solution of sulphuric acid. The amount of ammonia produced is determined by estimating the amount of sulphuric acid consumed in the reaction. It is done by estimating unreacted sulphuric acid left after the absorption of ammonia by titrating it with standard alkali solution. The difference between the initial amount of acid taken and that left after the reaction gives the amount of acid reacted with ammonia.
taken = V mL
Volume of NaOH of molarity, M. used for titration of excess of H₂SO₄ = V₁ mL
V₁L of NaOH of molarity M= V1 /2 mL of H₂SO₄ 0f molarity M
Volume of H₂SO₄ of molarity M unused= (V-1/1/2) mL (V-1/1/2) mL of H2S04 of molarity M = 2(V-V1/2) mL of NH₃ solution of molarity M.
1000 mL of 1 M NH₃ solution contains 17g NH₃ or 14 g of N
2(V-V1/2) mL of NH₃ solution of molarity M contains:

Kjeldahl method is not applicable to compounds containing nitrogen in nitro and azo groups and nitrogen present in the ring (e.g. pyridine) as nitrogen of these compounds does not change to ammonium sulphate under these conditions.

Halogens
Carius method: A known mass of an organic compound is heated with fuming nitric acid in the presence of silver nitrate contained in a hard glass tube known as Carius tube, in a furnace. Carbon and hydrogen present in the compound are oxidised to carbon dioxide and water. The halogen present forms the corresponding silver halide (AgX). It is filtered, washed, dried and weighed.
Let the mass of organic compound taken = m g
Mass of AgX formed = m, g
1 mol of AgX contains 1 mol of X
Mass of halogen in m1g of AgX

Sulphur
A known mass of an organic compound is heated in a Carius tube with sodium peroxide or fuming nitric acid. Sulphur present in the compound is oxidised to sulphuric acid. It is precipitated as barium sulphate by adding excess of barium chloride solution in water. The precipitate is filtered, washed, dried and weighed. The percentage of sulphur can be calculated from the mass of barium sulphate. Let the mass of organic compound taken = m g and the mass of barium sulphate formed = m₁ g
1 mol of BaSO₄ = 233 g BaSO₄ = 32 g sulphur

Phosphorus
A known mass of an organic compound is heated with fuming nitric acid whereupon phosphorus present in the compound is oxidised to phosphoric acid. It is precipitated as ammonium phosphomolybdate, (NH₄)₃PO₄.12MoO₃, by adding ammonia and ammonium molybdate. Alternatively, phosphoric acid may be precipitated as MgNH₄PO₄ by adding magnesia mixture which on ignition yields Mg₂P₂O₇. Let the mass of organic compound taken = m g and mass of ammonium phospho molydate = m1g.
Molar mass of (NH₄)₃PO₄.12MoO₃ = 1877 g.

where, 222 u is the molar mass of Mg₂P₂O₇, m, the mass of organic compound taken, mv the mass of Mg₂P₂O₇ formed and 62, the mass of two phosphorus atoms present in the compound Mg₂P₂O₇.

Oxygen
The percentage of oxygen in an organic compound is usually found by difference between the total percentage composition (100) and the sum of the percentages of all other elements. However, oxygen can also be estimated directly as follows:

A definite mass of an organic compound is decomposed by heating in a stream of nitrogen gas. The mixture of gaseous products containing oxygen is passed over red-hot coke when all the oxygen is converted to carbon monoxide. This mixture is passed through warm iodine pentoxide (IJOJ) when carbon monoxide is oxidised to carbon dioxide producing iodine.

The percentage of oxygen can be derived from the amount of carbon dioxide or iodine produced.
Let the mass of organic = mg compund taken
Mass of carbon dioxide = m₁g
44g carbon dioxide = 32 g oxygen

Presently, the estimation of elements in an organic compound is carried out by using microquantities of substances and automatic experimental techniques. The elements, carbon, hydrogen and nitrogen present in a compound are determined by an apparatus known as CHN elemental analyser. The analyser requires only a very small amount of the substance (1 -3 mg) and displays the values on a screen within a short time.

Students can Download Chapter 11 The p Block Elements 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 11 The p Block Elements

Introduction
There are six groups of p-block elements in the periodic table numbering from 13to 18. Boron, carbon, nitrogen, oxygen, fluorine and helium head the groups. Their valence shell electronic configuration is ns² np^1-6(except for He). The inner core of the electronic configuration may, however, differ. The difference in inner core of elements greatly influences their physical properties (such as atomic and ionic radii, ionisation enthalpy, etc.) as well as chemical properties.

In groups 13, 14 and 15, the group oxidation state is the most stable state for lighter elements of the group. However, the oxidation state two units less than the group oxidation state becomes progressively more stable down a group. This is due to the reluctance of ns² electrons to participate in bond formation in the case of heavier elements. This phenomenon is known as inert pair effect. Since p-block contains non-metals (and metalloids), these elements have higher electronegativities and higher ionisation enthalpies. In contrast to metals which form cations, non-metals readily form anions.

The combined effect of size and availability of cf orbitals considerably influences the ability of these elements to form π bonds. The first member of a group differs from the heavier members in its ability to form pπ -pπ multiple bonds to itself ( e.g., C=C, C° C, N° N) and to other second row elements e.g., C=0, C=N, C° N, N=0). This type of π – bonding is not particularly strong for the heavier p-block elements. The heavier elements do form π bonds but this involves d orbitals.

Group 13 Elements: The Boron Family

Electronic Configuration
The outer electronic configuration of these elements is ns² np¹. This difference in electronic structures affects the other properties and consequently the chemistry of all the elements of this group.

Atomic Radii
On moving down the group, atomic radius is expected to increase. However, a deviation can be seen. Atomic radius of Ga is less than that of Al. This can be understood from the variation in the inner core of the electronic configuration. The presence of additional 10 d-electrons offer only poor screening effect for the outer electrons from the increased nuclear charge in gallium. Consequently, the atomic radius of gallium (135 pm) is less than that of aluminium (143 pm).

Ionization Enthalpy
The ionisation enthalpy values as expected from the general trends do not decrease down the group. The decrease from B to Al is associated with increase in size. The observed discontinuity in the ionisation enthalpy values between Al and Ga, and between In and Tl are due to inability of d- and f-electrons, which have low screening effect, to compensate the increase in nuclear charge.

Electronegativity
Down the group, electronegativity first decreases from B to Al and then increases marginally.

Physical Properties
Boron is non-metallic in nature. It is extremely hard and black coloured solid. It exists in many allotropic forms.

Chemical Properties
Oxidation state and trends in chemical reactivity The sum of its first three ionization enthalpies of boron is very high due to its small size. This prevents it to form +3 ions and forces it to form only covalent compounds. But as we move from BtoAl.the sum of the first three ionisation enthalpies of Al considerably decreases, and is, therefore, able to form Al³⁺ ions. The tendency to behave as Lewis acid decreases with the increase in the size down the group. BCl₃ easily accepts a lone pair of electrons from ammonia to form BCl₃.NH₃.

i) Reactivity towards air
Boron has crystalline form which is unreactive. Alu-minium forms a very thin oxide layer on the surface which protects the metal from further attack.

ii) Reactivity towards acids and alkalies
Boron does not react with acids and alkalies even at moderate temperature, but aluminium has amphoteric character.

iii) Reactivity towards halogens
2E(s) + 3X₂(g) → 2EX₃(S) (X = F, Cl, Br, I)

Important Trends And Anomalous Properties Of Boron
The tri-chlorides, bromides and iodides of all these elements being covalent in nature are hydrolysed in water. Species like tetrahedral [M(OH)₄]^– and octahedral [M(H₂O)6]³⁺, except in boron, exist in aqueous medium. The monomeric trihalides, being electron deficient, are strong Lewis acids. Boron trifluoride easily reacts with Lewis bases such as NH₃ to complete octet around boron.
F₃B+: NH₃ → F₃B ← NH₃

It is due to the absence of d orbitals that the maximum covalence of B is 4. Since the d orbitals are available with Al and other elements, the maximum covalence can be expected beyond 4. Most of the other metal halides (e.g., AlCl₃ are dimerised through halogen bridging (e.g., Al₂Cl₆). The metal species ‘ completes its octet by accepting electrons from halogen in these halogen bridged molecules.

Some Important Compounds Of Boron

Borax
It is the most important compound of boron. Formula of the compound is Na₂B₄O₇.10H₂O . In fact it contains the tetranuclear units [B₄O₅(OH)₄]^2- and correct formula; therefore, is Na₂[B₄O₅(OH)₄].8H₂O.

On heating, borax first loses water molecules. On further heating it turns into a transparent liquid, which solidifies into glass like material known as borax bead.

Orthoboric acid
Orthoboric acid, H₃B0₃ is a white crystalline solid, with soapy touch. It is sparingly soluble in water but highly soluble in hot water.
Na₂B₄O₇ + 2HCl + 5H₂O → 2NaCl + 4B(OH)₃

Boric acid is a weak monobasic acid. It is not a protonic acid but acts as a Lewis acid by accepting electrons from a hydroxyl ion:
B(OH)₃ +2HOH → [B(OH)₄]^– + H₃O⁺

Structure of boric acid is given below.

Diborane (B₂H₆)
The simplest boron hydride is diborane (B₂H₆). Diborane can be prepared by treating BF₃ with lithium aluminium hydride in ether. A convenient laboratory method is oxidation of sodium borohydride with iodine.
2NaBH₄ + l₂ → B₂H₆ + 2Nal +H₂

On a commercial scale, diborane is produced by the action of BF₃ on sodium hydride.

Diborane is a colourless toxic gas. It catches fire on exposure to air releasing large amount of energy.
B₂H₆+ 6H₂O → 2B(OH)₃ + 6H₂

Reaction of diborane with NH₃ gives an addition product B₂H₆.2NH₃ which on heating gives borazine (B₃N₃H₃), commonly known as inorganic benzene due to its structural similarity with benzene. Boron forms a series of hydridoborates, the most important being (BH₄)^–.NaBH₄ (sodium borohydride) is a good reducing agent.

Each boron atom in B₂H₆ is sp³ hybridised. The structure contains two types of H- atoms the four-terminal hydrogen atoms and two bridged hydrogen atoms. The four-terminal H atoms and two B atoms lie in the same plane. Above and below this plane lie the bridged H atoms. B-H bonds formed by the terminal hydrogen atoms are normal covalent bonds while the bridge B-H bonds are three centre two-electron bonds. Each B atom forms four bonds even though boron has only three valence electrons. Hence B2H6 is an electron deficient compound.

Group 14 Elements: The Carbon Family
Carbon, silicon, germanium, tin, and lead form the carbon family.
Occurrence:
Carbon is widely distributed in nature in the free and combined states. Graphite, diamond, coal, etc are elemental forms of carbon while in the combined state it occurs as metal carbonates, hydrocarbons and CO₂ in air. Silicon is present in nature as silica and silicates. Ge is found only in traces. Tin occurs as cassiterite (SnO₂) and lead as galena (PbS)

Electronic Configuration
The valence shell electronic configuration of these elements is ns²np². The inner core of the electronic configuration of elements in this group also differs.

Covalent Radius
There is a considerable increase in covalent radius from C to Si, thereafter from Si to Pb a small increase in radius is observed. This is due to the presence of completely filled d and f orbitals in heavier members.

Ionization Enthalpy
The first ionization enthalpy of group 14 members is higher than the corresponding members of group 13. The influence of inner core electrons is visible here also. In general, the ionisation enthalpy decreases down the group.

Electronegativity
Due to small size, the elements of this group are slightly more electronegative than group 13 elements. The electronegativity values for elements from Si to Pb are almost the same.

Physical Properties
All group 14 members are solids. Carbon and silicon are non-metals, germanium is a metalloid, whereas tin and lead are soft metal.

Chemical Properties Oxidation states and trends in chemical reactivity
The group 14 elements have four electrons in outermost shell. The common oxidation states exhibited by these elements are +4 and +2.
Carbon also exhibits negative oxidation states. Since the sum of the first four ionization enthalpies is very high, compounds in +4 oxidation state are generally covalent in nature. In heavier members the tendency to show +2 oxidation state increases in the sequence Ge<Sn (i) Reactivity towards oxygen
All members when heated in oxygen form oxides. There are mainly two types of oxides, monoxide, and dioxide of formula MO and MOs respectively.

(ii) Reactivity towards water

(iii) Reactivity towards halogen
These elements can form halides of formula MX₂, and MX₄ (where X = F, Cl, Br, I). Except carbon, all other members react directly with halogen under suitable condition to make halides.

Hydrolysis can be understood by taking the example of SiCl₄. It undergoes hydrolysis by initially accepting lone pair of electrons from water molecule in d orbitals of Si, finally leading to the formation of Si(OH)₄ as shown below:

Important Trends And Anomalous Behaviour Of Carbon
Carbon differs from rest of the members of its group. It is due to its smaller size, higher electronegativity, higher ionisation enthalpy and unavailability of d orbitals. In carbon, only s and p orbitals are available for bonding and, therefore, it can accommodate only four pairs of electrons around it. This would limit the maximum covalence to four whereas other members can expand their covalence due to the presence of d orbitals.
Carbon has the ability to form pπ – pπ multiple bonds with itself and with other atoms of small size and high electronegativity.

Few examples are: C=C, C° C, C=0, C=S, and C° N. Carbon atoms have the tendency to link with one another through covalent bonds to form chains and rings. This property is called catenation.

Allotropes Of Carbon

Diamond
It has a crystalline lattice. In diamond, each carbon atom undergoes sp³ hybridisation and linked to four other carbon atoms by using hybridised orbitals in tetrahedral fashion. The C-C bond length is 154 pm. In this structure, directional covalent bonds are present throughout the lattice. It is very difficult to break extended covalent bonding and, therefore, diamond is a hardest substance on the earth. It is used as an abrasive for sharpening hard tools.

Graphite
Graphite has layered structure. Layers are held by van der Waals forces and distance between two layers is 340 pm. Each layer is composed of planar hexagonal rings of carbon atoms. C—C bond length within the layer is 141.5 pm. Each carbon atom in hexagonal ring undergoes sp² hybridisation and makes three sigma bonds with three neighbouring carbon atoms. Fourth electron forms a π bond. The electrons are delocalised over the whole sheet. Electrons are mobile and, therefore, graphite conducts electricity along the sheet. Graphite cleaves easily between the layers and, therefore, it is very soft and slippery. For this reason graphite is used as a dry lubricant in machines running at high temperature, where oil cannot be used as a lubricant.

Fullerenes
Fullerenes are prepared by heating graphite in an electric arc in the presence of helium or argon. The sooty material formed by condensation of the vapours consists of C₆₀ with smaller amounts of C₇₀ and other fullerenes. C₆₀ is named as Buckminster fullerence. The general name fullerence refers to the family of spheroidal carbon-cage molecules. The shape of C₆₀ resembles that of a soccer ball. It contains twelve five-membered rings and twenty 6-membered rings of carbon. The 6-membered rings are fused both to other five and six membered rings. However, the 5-membered rings are fused only to six-membered rings. Both carbon-carbon single (1.435 Å) and double (1.383 Å) bonds are present in this structure. Carbon black, coke and charcoal are impure amorphous forms of graphite or fullerenes. Carbon black is formed by burning hydrocarbon in limited supply of air. Charcoal and coke are obtained by heating wood and coal respectively in the absence of air.

Uses of Carbon
Being good conductor, graphite is used for electrodes in batteries and industrial electrolysis. Crucibles made from graphite are inert to dilute acids and alkalies. Being highly porous, activated charcoal is used in adsorbing poisonous gases. Diamond is a precious stone and used in jewellery.

Some Important Compounds Of Carbon And Silicon
Oxides of Carbon
Two important oxides of carbon are carbon monoxide, CO and carbon dioxide, CO₂.

Carbon Monoxide
Direct oxidation of C in limited supply of oxygen or air yields carbon monoxide.

On commercial scale it is prepared by the passage of steam over hot coke. The mixture of CO and H₂ thus produced is known as water gas or synthesis gas.

When air is used instead of steam, a mixture of CO and N₂ is produced, which is called producer gas.

Water gas and producer gas are very important industrial fuels. Carbon monoxide in water gas or producer gas can undergo further combustion forming carbon dioxide with the liberation of heat. CO arises has the ability to form a complex with haemoglobin, which is about 300 times more stable than the oxygen-haemoglobin complex. This prevents haemoglobin in the red blood corpuscles from carrying oxygen round the body and ultimately resulting in death.

Carbon Dioxide
It is prepared by complete combustion of carbon and carbon-containing fuels in excess of air.

On commercial scale it is obtained by heating limestone. Carbon dioxide, which is normally present to the extent of ~0.03 % by volume in the atmosphere, is removed from it by the process known as photosynthesis. It is the process by which green plants convert atmospheric CO₂ into carbohydrates such as glucose. The overall chemical change can be expressed as:

The increase in combustion of fossil fuels and decomposition of limestone for cement manufacture in recent years seem to increase the CO₂ content of the atmosphere. This may lead to increase in green house effect and thus, raise the temperature of the atmosphere which might have serious consequences. Carbon dioxide can be obtained as a solid in the form of dry ice by allowing the liquified CO₂ to expand rapidly. Dry ice is used as a refrigerant for ice-cream and frozen food.

Resonance structures of carbon dioxide

Silicon Dioxide, SiO₂
Quartz, cristobatite and tridymite are some of the crystalline forms of silica, and they are interconvertible at suitable temperature. In Silicon dioxide, each silicon atom is covalently bonded in a tetrahedral manner to four oxygen atoms. Each oxygen atom in turn covalently bonded to another silicon atoms.

Silicones
They are a group of organosilicon polymers, which have (R₂SiO) as a repeating unit. The starting materials for the manufacture of silicones are alkyl or aryl substituted silicon chlorides, RnSiCl_(4-n), where R is alkyl or aryl group.

Silicates
The basic structural unit of silicates if SiO₄^4- tertrahedra. Feldspar, zerolites, mica, asbestose, etc. are examples of silicates. In silicates, either the SiO₄^4- will be present as discrete units or several such units are joined togetherth rough sharing of corner of the tetrahedra using one to four oxygen atoms per silicate unit. Like this, different silicates assume different forms such as chain, ring, sheet or three-dimensional structures. Glass and cement are examples of man-made silicates.

Zeolites
Zeolites are alumino silicates. If a few Si atoms of the three-dimensional network structure of SiO₂ are replaced by Al atoms, the resulting structure is called alumino silicate structure. This structure evidently has negative charge and Na⁺.K⁺ pr Ca²⁺ ions balance the negative charge. Zeolites are used as catalysts in petrochemical industry for cracking of hydrocarbons. ZSM-5 is a type of zeolite used in the conversion of alcohol to gasoline. Zeolites are also used in softening hard water.

Students can Download Chapter 10 The s Block Elements 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 10 The s Block Elements

Introduction
Group 1 of the periodic table consists of the elements: Lithium, Sodium, Potassium, Rubidium, Caesium and Francium. They are collectively known as alkali metals.

Group 2 consists of Beryllium, Mgnesium, Calcium, Strontium, Barium and Radium. These elements except of beryllium are known as the alkaline earth metals. The general electronic configuration of s-block elements is [noble gasjns1 for alkali metals and [noble gas] ns² for alkaline earth metals. The first elements of Group 1 and Group 2 respectively exhibit diagonal similarity, which is commonly referred to as diagonal relationship in the periodic table. The diagonal relationship is due to the similarity in ionic sizes and /or charge/radius ratio of the elements.

Group 1 Elements: Alkali Metals

1) Electronic Configuration:
All the alkali metals have one valence electron, ns¹ outside the noble gas core. The loosely held s-electron readily lose electron to give monovalent M⁺ ions.

2) Atomic And Ionic Radii:
The atomic and ionic radii of alkali metals increase on moving down the group. Hence, ionization enthalpies of the alkali metals are considerably low and decrease down the group.

3) Hydration Enthalpy:
The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes. Li⁺ > Na⁺ > K⁺ > Rb⁺ > Cs⁺ Li⁺ has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., LiCl- 2H₂O

Physical Properties
When heat is supplied to alkali metal or its salt the electrons are excited to higher energy levels. As these electrons return to their original level; radiations are emitted which fall in the visible region of electromagnetic spectrum. Thus they appear coloured. Li imparts crimson red colour, K gives violet colour and Na gives golden yellow colour to the flame.

Chemical Properties
The reactivity of these metals increases with their size. They burn vigorously in oxygen forming oxides. Lithium forms monoxide, sodium forms peroxide, the other metals form superoxides. The superoxide O₂ ion is stable only in the presence of large cations such as K, Rb, Cs.
4Li + O₂ → 2LizO(oxide)
2Na + O₂ → Na₂O₂ (peroxide)
M + O₂ → MO₂(superoxide)
(M=K, Rb, Cs)
Because of their high reactivity towards air and water, they are normally kept in kerosene oil.lt may be noted that although lithium has most negative E° value.

They also react with proton donors such as alcohol, gaseous ammonia and alkynes.AII the alkali metal hydrides are ionic solids with high melting points.
2M + H₂ → 2M⁺H^–.

The alkali metals readily react vigorously with halogens to form ionic halides, M⁺X^–. However, lithium halides are somewhat covalent. It is because of the high polarisation capability of lithium-ion. The alkali metals are strong reducing agents, lithium being the most and sodium the least powerful. The alkali metals dissolve in liquid ammonia giving deep blue solutions. The solutions are paramagnetic and on standing slowly liberate hydrogen.

General Characteristics Of The Compounds Of The Alkali Metals

Oxides And Hydroxides
Reactivity of alkali metals with oxygen increases down the group. Lithium, when heated in air, forms the normal oxide (Li₂O) while sodium forms the per-oxide (Na₂O₂). Potassium, Rubidium and caesium form superoxides (MO₂).
4Li + O₂ → 2Li₂O; 2Na+ O₂ → Na₂O₂; K + O₂ → KO₂

The normal oxides dissolve in water to form hydroxides (MOH) which are strong bases. However, LiOH is only slightly soluble in water and it decomposes on heating. The peroxides and superoxides also dis-solve in water to form basic hydroxides. The basic character of alkali metal hydroxides increases down the group.

Halides
Alkali metals react vigorously with halogens to form metal halides of the general formula MX. 2M+X₂ → 2MX X=F, Cl, Br or l and M= alkali metal Reactivity of alkali metal towards halogen increases from Li to Cs. Halides of alkali metals are ionic compounds readily soluble in water. But LiF is almost insoluble due to high lattice energy.

Anomalous Properties Of Lithium
The anomalous behaviour of lithium is due to the :

1. exceptionally small size of its atom and ion, and
2. high polarising power (i.e., charge/ radius ratio).

As a result, there is increased covalent character of lithium compounds which is responsible for their solubility in organic solvents.

Points Of Similarities Between Lithium And Magnesium
The similarity between lithium and magnesium is particularly striking and arises because of their similar sizes: atomic radii, Li = 152 pm, Mg= 160 pm; ionic radii: Li⁺ = 76 pm, Mg²⁺ = 72 pm. The main points of similarity are:

1. Both lithium and magnesium are hander and lighter than other elements in the respective groups.
2. Lithium and magnesium react slowly with water. Their oxides and hydroxides are much less soluble and their hydroxides decompose on heating. Both form a nitride, Li₃N and Mg₃N₂, by direct combination with nitrogen.
3. The oxides, Li₂O and MgO do not combine with excess oxygen to give any superoxide.
4. The carbonates of lithium and magnesium decompose easily on heating to form the oxides and CO₂.

Some Important Compounds Of Sodium Sodium Carbonate (Washing Soda), Na₂CO₃.10H₂O
Sodium carbonate is generally prepared by Solvay Process.
The equations for the complete process may be written as:
2NH₃ + H₂O + CO₂ → (NH₄)₂CO₃
(NH₄)₂CO₃ + H₂O + CO₂ → 2NH₄HCO₃
NH₄HCO₃ +NaCl → NH₄Cl + NaHCO₃
2NaHCO₃ → Na₂CO₃ +CO₂ +H₂O

In this process, NH₃ is recovered when the solution containing NH₄Cl is treated with Ca(OH)₂. On heating washing soda becomes monohydrate and then completely anhydrous i.e., soda ash.

Sodium Chloride, NaCl
The most abundant source of sodium chloride is seawater. Common salt is generally obtained by evaporation of seawater. Crude sodium chloride, generally obtained by crystallisation of brine solution, contains sodium sulphate, calcium sulphate, calcium chloride and magnesium chloride as impurities. Calcium chloride, CaCl₂, and magnesium chloride, MgCl₂ are impurities because they are deliquescent (absorb moisture easily from the atmosphere). To obtain pure sodium chloride, the crude salt is dissolved in minimum amount of water and filtered to remove insoluble impurities. The solution is then saturated with hydrogen chloride gas. Crystals of pure sodium chloride separate out. Calcium and magnesium chloride, being more soluble than sodium chloride, remain in solution.

Uses:

• It is used as a common salt or table salt for domestic purpose.
• It is used for the preparation of Na₂O₂, Na0H and Na₂CO₃.

Sodium Hydroxide (Caustic Soda), NaOH
Sodium hydroxide is generally prepared commercially by the electrolysis of sodium chloride in Castner-Kellner cell. A brine solution is electrolysed using a mercury cathode and a carbon anode.

The amalgam is treated with water to give sodium hydroxide and hydrogen gas.
2 Na – amalgam + 2H₂O → 2NaOH + 2Hg +H₂
The sodium hydroxide solution at the surface reacts with the C02 in the atmosphere to form Na₂CO₃.

Uses:
It is used in (i)the manufacture of soap, paper, artificial silk and a number of chemicals,(ii) in petroleum refining, (iii) in the purification of bauxite, (iv) in the textile industries for mercerising cotton fabrics and (v) for the preparation of pure fats and oils.

Biological Importance Of Sodium And Potassium
Sodium ions participate in the transmission of nerve signals. The concentration gradient of Na₊ and K₊ demonstrates that-a discriminatory mechanism called sodium-potassium pump, operates across the cell membranes.

Group 2 Elements: Alkaline Earth Metals
The group 2 elements (except beryllium) are known as alkaline earth metals. The first element beryllium differs from the rest of the members and shows diagonal relationship to aluminium.
1) Electronic Configuration:
These elements have two electrons in the s-orbital of the valence shell. Their general electronic configuration may be represented as [noble gas] ns².

2) Atomic And Ionic Radii:
Within the group, the atomic and ionic radii increase with increase in atomic number due to the increased nuclear charge in these elements. They have low ionisation enthalpy and it decreases down the group with increase in size.

3) Hydration Enthalpy:
Hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group. Be²⁺ > Mg²⁺ > Ca²⁺ > Sr²⁺ > Ba²⁺ The hydration enthalpies of alkaline earth metal ions are larger than those of alkali metal ions.

Physical Properties
Calcium, Strontium and Barium impart characteristic brick red, crimson and apple green colours respectively to the flame. Inflame the electrons are excited to higher energy levels and when they drop back to the ground state, energy is emitted in the form of visible light. The electrons in Be and Mg are too strongly bound to get excited by flame. Hence, these elements do not impart any colour to the flame.

Chemical Properties
The alkaline earth metals are less reactive than the alkali metals. The reactivity of these elements increases on going down the group.
Reactivity towards air and water Beryllium and Magnesium are kinetically inert to oxygen and water because of the formation of an oxide film on their surface. However, powdered beryllium burns brilliantly on ignition in air to give BeO and Be₃N₂.
Reactivity towards halogen
M + X₂ → MX₂ (X = F, Cl, Br, I)

Reactivity towards hydrogen
All the elements except beryllium form their hydrides, MH₂.BeH₂, however, can be prepared by the reaction of BeCl₂ with LiAlH₄.
2BeCl₂ +LiAlH₄ → 2BeH₂ +LiCl + AlCl₃

Reactivity towards acids:
The alkaline earth metals readily react with acids liberating dihydrogen.

General Characteristics Of Compounds Of The Alkaline Earth Metals
i) Oxides and Hydroxides: Alkaline earth metals burn in air or oxygen to form their oxides. (Oxides are also prepared by the thermal decomposition of their carbonates). Be, Mg and Ca form monoxides (MO). The tendency to form peroxide increases as the size of the metal ion increases. Strontium and barium form peroxides (MO₂)
2M + O₂ → MO (M = Be, Mg or Ca)
M+O₂ → MO₂ (M = Sr or Ba)

BeO is amphoteric in character, while the oxides of the rest of the elements in group 2 are basic. The oxides of Ca, Sr and Ba react with water to form their corresponding hydroxides.

The hydroxides of alkaline earth metals are bases except Li(OH)₂ which is amphoteric. The basic strength increases from Mg(OH)₂ to Ba(OH)₂. The solubility and thermal stability of hydroxides increase downward in the group. Be(OH)₂ and Mg(OH)₂ are almost insoluble. Ca(OH)₂ is sparingly soluble, while Sr(OH)₂ and Ba(OH)₂ are increasingly more soluble.

ii) Halides: Group 2 metals directly combine with halogen to form divalent halides of the formula

The s-Block Elements
MX₂ where X is the halogen. The metal halides are also formed by the action of halogen acids on metals, their oxides, carbonates and hydroxides. BeCl₂ is, however, prepared by passing Cl₂ over a hot mixture of BeO and coke.
In contrast to the halides of other alkaline earth metals, beryllium halides are covalent. In the solid-state BeCl₂ has a polymeric chain structure involving Be-CI-Be bridges. The anhydrous halides are hygroscopic and form hydrates such as MgCl₂.6H₂O, CaCl₂.6H₂O etc. Due to this reason, anhydrous calcium chloride is widely used as a dehydrating agent. Fluorides are relatively less soluble due to high lattice energies,

iii) Salts of Oxoacids:
The alkaline earth metals also form salts of oxoacids. Some of these are : Carbonates, Sulphates and Nitrates.

Anomalous Behaviour Of Beryllium
Beryllium differs from the rest element in many of its properties. These are

1. Beryllium has high ionisation enthalpy.
2. Small size of Be atom
3. Be does not exhibit coordination number more than four.
4. The oxides and hydroxides of Be are amphoteric in nature.

Diagonal Relationship Between Beryllium And Aluminium
The ionic radius of Be²⁺ is estimated to be same as that of the Al³⁺ ion. Hence Be resembles Al in some ways. Some of the similarities are:

1. Like AI, Be is not readily attacked by acids because of the presence of an oxide film on the surface of the metal.
2. Beryllium hydroxide dissolves in excess of alkali to give a beryllate ion just as aluminium hydroxide gives aluminate ion.
3. The chlorides of both Be and Al have Ch bridged chloride structure in vapour phase. Both the chlorides are soluble in organic solvents and are strong Lewis acids. They are used as Friedel Craft catalysts.
4. Be and Al ions have strong tendency to form complexes, BeF₄^2-, AlF₆^3-.

Some Important Compounds Of Calcium
Important compounds of calcium and their preparations are given below.

Calcium Oxide Or Quick Lime, CaO
It is prepared by the following reaction.
CaCO₃ \(\rightleftharpoons \) Ca0 + CO₂
CO₂ is removed as soon as it is produced to enable the reaction to proceed to completion.
CaO + H₂O → Ca(OH)₂
This process is called slaking of lime. CaO is a basis oxide.

Uses:

• Primary material for manufacturing cement
• It is used in the manufacturing of caustic soda
• Used to purify sugar

Calcium Hydroxide (Slaked Lime), Ca(OH)₂
It is prepared by adding water to CaO. The aqueous solution of Ca(OH)₂ is known as lime water and the suspension of slaked lime is known as milk of lime. When CO₂ is passed through lime water it turns milky due to the formation of CaCO₃
Ca(OH)₂ + CO₂ → CaCO₃ +H₂O

Uses:

• It is used in whitewash due to its disinfectant nature.
• Used in the preparation of bleaching powder.
• Used to purify sugar.

Calcium Carbonate, CaCO₂
It occurs in limestone, chalk, marble etc.
It can be prepared by the following reactions.
Ca(OH)₂ + CO₂ → CaCO₃ + H₂O
CaCl₂ + Na₂CO₃ → CaCO₃ + 2NaCl
CaCO₃ reacts with dilute acids to liberate carbon dioxide.

Uses:

• It is used as a flux in the extraction of metals.
• It is used as the building material of quick lime.

Calcium Sulphate (Plaster Of Paris), CaSO₄.½H₂O
It is obtained by heating gypsum (CaSO₂.2H₂O)

Above 393K anhydrous calcium sulphate is formed. This is known as ‘dead burnt plaster’

Used:

• It is used in building industry as well as plasters.
• Used to make casts of statues.

Cement
Cement is prepared by combining CaO with other materials such as clay with silica, SiO₂ along with Oxides of Al, iron and magnesium. The average composition of portland cement is:
CaO, 50-60%;
SiO₂, 20-25%;
Al₂O₃, 5-10%;
MgO, 2-3%;
Fe₂O₃, 1-2% and
SO₃, 1-2%.
When limestone and clay are heated we get cement clinker. This clinker is mixed with gypsum to form cement.

Setting of Cement:
When mixed with water the setting of cement takes place to give a hard mass. It is due to the rearrangement and hydration of molecules of constituents. Gypsum is added to slow down the setting process so it gets sufficiently hardened.

Uses:

• Used in construction of building.

Biological Importance Of Magnesium And Calcium
Human body contains about 25g of Mg and 1200g of Ca. Mg is a cofactor in enzymes which use ATP in phosphate transfer process in our body. Photosynthesis in plants takes place in presence of chlorophyll which contains Mg. About 99% of body calcium is found in teeth and bones. Calcium concentration in plasma is regulated at 100mg/litre in presence of hormones such as calcitonin and parathyroid hormone.

Students can Download Chapter 9 Hydrogen 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 9 Hydrogen

Introduction
Hydrogen has the simplest atomic structure all the elements around us in nature. It consists of only one proton and one electron.

Position Of Hydrogen In The Periodic Table
Hydrogen is the first element in the periodic table. Hydrogen has electronic configuration 1 s1. On one hand, its electronic configuration is similar to the outer electronic configuration (ns¹) of alkali metals. On the other hand, it is short by one electron to the corresponding noble gas configuration, helium (1s²). It has resemblace to both alkali metals and halogens.

Dihydrogen, H₂

Isotopes Of Hydrogen
There are three isotopes of hydrogen with mass numbers 1,2 and 3. They are called protium, deuterium and tritium respectively. Their natural abundances . are in the ratio l:1.56 × 10^-2: 1 × 10^-17 respectively.

1. Protium (ordinary hydrogen)(\(_{ 1 }^{ 1 }{ H }\)): It is the most abundant isotope of hydrogen. Its nucleus contains one proton and no neutron.
2. Deuterium (heavy hydrogen, \(_{ 1 }^{ 2 }{ H }\) or D): Heavy hydrogen is prepared from heavy water (D₂O) which is obtained by electrolysis of ordinary water.
3. Tritium has 2 neutrons in the nucleus.

Preparation Of Dihydrogen, H₂
It is usually prepared by the following reactions:

3. Reaction of steam on hydrocarbons or coke at high temperatures in the presence of catalyst yields hydrogen.

The mixture of CO and H₂ is called water gas. As this mixture of CO and H₂ is used for the synthesis of methanol and a number of hydrocarbons, it is also called synthesis gas or ‘syngas’. Nowadays ‘syngas’ is produced from sewage, sawdust, scrap wood, newspapers etc. The process of producing ‘syngas’ from coal is called ‘coal gasification’.

This reaction is called water-gas shift reaction.

Properties Of Dihydrogen

Physical Properties
Dihydrogen is a colourless, odourless, tasteless, combustible gas. It is lighter than air and insoluble in water.

Chemical Properties
Dihydrogen is not particularly reactive because of its high bond dissociation enthalpy. However, hydrogen forms compounds with almost all elements at high temperature or in presence of catalysts.
Reaction with halogens:
H₂ (g) + X₂(g) → 2HX(g) (X = F, Cl, Br, l)
Reaction with dioxygen:

Uses Of Hydrogen

1. Hydrogen is used in the manufacture of ammonia by Haber process, water gas, fertilisers etc.
2. It is used in the hydrogenation of vegetable oils and as a reducing agent.
3. It is used in the production of methanol and synthetic petrol.
4. Liquid hydrogen is used in as rocket fuel along with liquid oxygen.
5. It is used in oxy-hydrogen torch used for welding.

Hydrides
Hydrogen can form binary compounds with almost all elements. These are known as hydrides.
The hydrides are classified into three categories:

1. Ionic or saline or salt like hydrides
2. Covalent or molecular hydrides
3. Metallic or non-stoichiometric hydrides

Ionic Or Saline Hydrides
These are stoichiometric compounds of dihydrogen formed with most of the s-block elements which are highly electropositive in character. However, significant covalent character is found in the lighter metal hydrides such as LiH, BeH₂ and MgH₂.

Covalent Or Molecular Hydride
Dihydrogen forms molecular compounds with most of the p-block elements. Most familiar examples are CH₄, NH₃, H₂O and HF. For convenience hydrogen compounds of nonmetals have also been considered as hydrides. Molecular hydrides are further classified according to the relative numbers of electrons and bonds in their Lewis structure into :

1. electron-deficient,
2. electron-precise,and
3. electron-rich hydrides.

Group13 elements form electron deficient compounds. They act as Lewis acids i.e., electron acceptors. eg.B₂H₆ Group 14 elements form electron precise compounds. They have required number of electrons. eg.CH₄. Electron-rich hydrides have excess electrons which are present as lone pairs. Elements of group 15-17 form such compounds. (NH₃ has 1 – lone pair, H₂O – 2 and HF -3 lone pairs).They will behave as Lewis bases.

Metallic Or Non-Stoichiometric (Or Interstitial) Hydrides
These are formed by many d-block and f-block elements. However, the metals of group 7, 8 and 9 do not form hydride. Even from group 6, only chromium forms CrH. These hydrides conduct heat and electricity though not as efficiently as their parent metals do. Unlike saline hydrides, they are almost always nonstoichiometric, being deficient in hydrogen. For example, LaH_2.87 & YbH_2.55

Water
Water is a colourless tasteless liquid. A major part of all living organisms is made up of water.The unusual properties of water is due to the presence of extensive hydrogen bonding between water molecules.

Structure Of Water
In the gas phase water is a bent molecule with a bond angle of 104.5°, and O-H bond length of 95.7 pm

(a) The bent structure of water;
(b) the water molecule as a dipole

Structure Of Ice
The crystalline form of water is ice. At atmospheric pressure, ice crystallises in the hexagonal form, but at very low temperatures it condenses to cubic form. Hydrogen bonding gives ice a rather open type structure with wide holes. These holes can hold some other molecules of appropriate size interstitially. Density of ice is less than that of water. Therefore, an ice cube floats on water. In winter season ice formed on the surface of a lake provides thermal insulation which ensures the survival of the aquatic life.

Chemical Properties of Water
1) Amphoteric Nature:
It has the ability to act as an acid as well as a base i.e., it behaves as an amphoteric substance. In the Bronsted sense it acts as an acid with NH₃ and a base with H₂S.

2) Redox Reactions Involving Water
Water can be reduced and oxidised:
2H₂O(l) + 2Na(s) → 2NaOH(aq) + H₂(g): reduction Water is oxidised to O₂ during photosynthesis
6CO₂(g) +12H₂O(I) → C₆H₁₂O₆ (aq) + 6H₂O(I) + 6O₂(g)

3) Hydrolysis Reaction:
Due to high dielectric constant, it has a very strong hydrating tendency.
P₄O₁₀(s) + 6H₂O(l) → 4H₃PO₄(aq)

4) Hydrates Formation:
From aqueous solutions, many salts can be crystallised as hydrated salts. Such an association of water is of different types viz.,
i) Coordinated water e.g.,
[Cr(H₂O)₆]₃+3Cl^–
ii) Interstitial water.g., BaCl₂.2H₂O
iii) hydrogen-bonded water.g.,
[Cu(H₂O)₄]²⁺ SO₄^2-.H₂O in CuSO₄.5H₂O

Hard And Soft Water
Water which produces lather with soap solution readily is called soft water. For example, rainwater, distilled water etc. Water which does not produce lather with soap solution readily is called hard water, eg: Sea water, water from certain rivers.

Hardness of water is due to the presence of bicarbonates, chlorides and sulphates of calcium and magnesium. The calcium and magnesium ions present in hard water form insoluble salts with soap and prevent the formation of lather.

Temporary Hardness
Temporary hardness is due to the presence of mag-nesium and calcium hydrogencarbonates.
It can be removed by boiling.

Permanent Hardness
It is due to the presence of soluble salts of magnesium and calcium in the form of chlorides and sulphates in water. Permanent hardness is not removed by boiling. It can be removed by the following methods:
i) Treatment with washing soda (sodium carbonate):
Washing soda reacts with soluble calcium and magnesium chlorides and sulphates in hard water to form insoluble carbonates.
MCl₂ → MCO₃ ↓ 2NaCl (M=Mg, Ca)
MSO₄ + Na₂CO₃ → MCO₃ ↓ +NaSO₄

ii) Calgon’s method:
Sodium hexametaphosphate (Na₆P₆O₁₈), commercially called ‘calgon’, when added to hard water, the following reactions take place.
Na₆P₆O₁₈ → Na^{+ + Na₄P₆O₁₈2- (M=Mg, Ca)
M2+ + Na₄P₆O₁₈2- → [Na₂MP₆O₁₈]2- + 2Na+}

iii) Ion-exchange method:
This method is also called zeolite/perm utit process. Hydrated sodium aluminium silicate iszeolite/permutit.Forthe sake of simplicity, sodium aluminium silicate (NaAlSiO₄) can be written as NaZ.
2NaZ(s) + M²⁺(aq) → MZ₂(s) + 2Na⁺(aq) (M=Mg, Ca)
MZ₂ (S) + 2NaCl(aq) → 2NaZ(s) + MCl₂(aq)

iv) Synthetic resins method:
Nowadays hard . water is softened by using synthetic cation exchangers. This method is more efficient than zeolite process.Ion exchange resin (RSO₃H) is changed to RNa by treating it with NaCI. Here R is resin anion.
2RNa(s) + M²⁺(aq) → R₂M(s) + 2Na⁺(aq)

The resin exchanges Na+ ions with Ca²⁺ and Mg²⁺ ions present in hard waterto make the water soft.

HYDROGEN PEROXIDE (H₂O₂)
It can be prepared by the following methods.

Structure
Hydrogen peroxide has a non-planar structure.

Chemical Properties
i) Oxidising action in acidic medium
PbS(s) + 4H₂O₂(aq) → PbSO₄(s) + 4H₂O(l)

ii) Reducing action in acidic medium
HOCl + H₂O → H₃O⁺ +Cl^– + O₂

iii) Oxidising action in basic medium
Mn²⁺ +H₂O₂ → Mn⁴⁺ + 2OH^–

iv) Reducing action in basic medium
2MnO₄^– + 3H₂O₂ → MnO₂ + O₂ + 2H₂O + OH^–

Uses

1. As a bleaching agent for textiles, wood and paper pulp
2. In the manufacture of chemicals such as sodium perborate, epoxides etc.
3. A dilute solution of H₂O₂ is used as a disinfectant. This solution is used as an antiseptic for wounds, teeth and ears under the name perhydrol.
4. iv) It is used in pollution control treatment of domestic and industrial effluents.

Heavy water. D₂O
It is extensively used as a moderator in nuclear reactors and in exchange reactions for the study of reaction mechanisms. It can be prepared by exhaustive electrolysis of water or as a by-product in some fertilizer industries.lt is used for the preparation of other deuterium compounds.

Dihydrogen As A Fuel
Due to extensive use, our reserves of fossil fuels are fast depleting. A prospective alternative in this regard is what is known as hydrogen economy. The major idea behind hydrogen economy is the storage and transportation of energy in the form of gaseous and liquid hydrogen. Hydrogen can replace fossil fuels in automobiles, and coal or coke in industrial processes involving reduction. Hydrogen fuel can release more energy per unit weight of the fuel than our conventional fuels. Hydrogen oxygen fuel cells can be used for generating power in automobiles. Liquid hydrogen has already been used as rocket fuel along with liquid oxygen.

The technology involves the production of bulk quantities of hydrogen and its storage in liquid form in vacuum insulated cryogenic tanks. Transport of liquid hydrogen by road or rail, or through pipelines is feasible. Certain metal alloys can be used as smaller storage units for hydrogen.

Students can Download Chapter 8 Redox Reactions 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 8 Redox Reactions

Introduction
The reaction which involve both oxidation and reduction reactions is called Redox reaction.

Classical Idea Of Redox Reactions Oxidation And Reduction Reactions
“Oxidation” is defined as the addition of oxygen/electronegative element to a substance or removal of hydrogen/electropositive element from a substance. Examples of oxidation:

1. Addition of oxygen 2Mg + O2 → 2MgO
2. Removal of hydrogen 2H₂S + O₂ → 2S + 2H₂O
3. Addition of electronegative element Mg + Cl₂ → MgCl₂

The term reduction been broadened these days to include removal of oxygen/electronegative element from a substance or addition of hydrogen /electropositive element.

1. Removal of electronegative element FeCl₃ + H₂ → 2FeCl₂ + 2HCl
2. Removal of Oxygen (2H₂O → 2Hg + O₂)
3. Addition of Hydrogen (H₂ + Cl₂ → 2HCl)

Redox Reactions In Terms Of Electron Transfer Reactions
According to electronic concept, the processes which involves loss of electrons are called oxidation reactions. Similarly, processes which involve gain of electrons are called reduction reactions.
The atom which reduced, act as oxidising agent and the atom which oxidised act as reducing agent.For example;
2Na(s) + Cl₂(g) → 2Na⁺Cl^–(s) or 2NaCl(s)
Here Na is oxidised and Cl is redused.

Competitive Electron Transfer Reaction
Place a strip of metallic zinc in an aqueous solution of copper nitrate. You may notice that the strip becomes coated with reddish metallic copper and the blue colour of the solution disappears. Formation of Zn²⁺ ions among the products can easily be judged when the blue colour of the solution due to Cu2+ has disappeared. The reaction is,
Zn(s) + Cu²⁺(aq) → Zn²⁺(aq) + Cu(s) zinc is oxidised, releasing electrons, something must be reduced, accepting the electrons lost by zinc. Copper ion is reduced by gaining electrons from the zinc.

Oxidation Number
Oxidation number of an element may be defined as the charge which an atom of the element has or appears so have when present in the combined state in a compound.

1. Electrons shared between two like atoms are divided equally between the sharing atoms.
2. Electrons shared between two unlike atoms are counted with the more electronegative atom. Atoms can assume positive, zero or negative values of oxidation numbers depending on their state of combination. Oxidation number can be a fraction in some cases.

The rules for calculation of oxidation number are:
1. In elements, in the free or the uncombined state, each atom bears an oxidation number of zero. Evidently each atom in H2 has the oxidation number zero.

2. For ions composed of only one atom, the oxidation number is equal to the charge on the ion. Thus Na+ ion has an oxidation number of +1, Mg²⁺ion, +2, Fe³⁺ ion, +3, Cl^– ion, -1, O^2- ion, -2; and so on. In their compounds all alkali metals have oxidation number of +1, and all alkaline earth metals have an oxidation number of +2. Aluminium is regarded to have an oxidation number of +3 in all its compounds.

3. The oxidation number of oxygen in most compounds is-2. However, we come across two kinds of exceptions here.in peroxides (e.g., H₂O₂, Na₂O₂), each oxygen atom is assigned an oxidation number of—1, in superoxides (e.g., KO₂, RbO₂) each oxygen atom is assigned an oxidation number of -(½). The second exception appears rarely, i.e. when oxygen is bonded to fluorine. In such compounds e.g., oxygen difluoride (OF₂) and dioxygen difluoride (O₂F₂), the oxygen is assigned an oxidation number of +2 and +1, respectively. The number assigned to oxygen will depend upon the bonding state of oxygen but this number would now be a positive figure only.

4. The oxidation number of hydrogen is +1, except when it is bonded to metals in binary compounds (that is compounds containing two elements). For example, in LiH, NaH, and CaH₂, its oxidation number is —1.

5. In all its compounds, fluorine has an oxidation number of-1. Other halogens (Cl, Br, and I) also have an oxidation number of-1, when they occur as halide ions in their compounds. Chlorine, bromine and iodine when combined with oxygen, for example in oxoacids and oxoanions, have positive oxidation numbers.

6. The algebraic sum of the oxidation number of all the atoms in a compound must be zero. In polyatomic ion, the algebraic sum of all the oxidation numbers of atoms of the ion must equal the charge on the ion. Thus, the sum of oxidation number of three oxygen atoms and one carbon atom in the carbonate ion, (CO₃)^2- must equal -2. A term that is often used interchangeably with the oxidation number is the oxidation state. Oxidation state of a metal is a compound is sometimes represented by Stock notation. According to this, the oxidation number is written as Roman numeral in parenthesis after the symbol of the metal in the molecular formula. e.g.,Fe(ll)0, Sn(IV), Cl₄,Mn(IV)O₂.

Problem
Using Stock notation, represent the following compounds HAUCl₄, Ti₂O, FeO, Fe₂O₃, Cul, CuO, MnO and MnO₂.

Solution
By applying various rules of calculating the oxidation number of the desired element in a compound, the oxidation number of each metallic element in its compound is as follows:
HAuCl₄ → Au has 3
Tl₂O → Tl has 1
FeO → Fe has 2
Fe₂O₃ → Fe has 3
Cul → Cu has 1
CuO → Cu has 2
MnO → Mn has 2
MnO₂ → Mn has 4

Therefore, these compounds may be represented as
HAU(III)Cl₄, Tl₂(I)O, Fe(II)O, Fe₂(III)O₃, Cu(I)l, Cu(II)O, Mn(II)O, Mn(IV)O₂.

In terms of oxidation number, oxidation may be defined as a chemical change in which there occurs an increase in the oxidation number of an atom or atoms. Reduction may be defined as a chemical change in which there occurs a decrease in the oxidation number of an atom or atoms. Thus, a redox reaction may be defined as a reaction in which the oxidation number of atoms undergoes a change.

Types Of Redox Reactions
1. Combination Reactions:
A combination reaction may be denoted in the manner
A + B → C

2. Decomposition Reaction:
Decomposition reactions are the opposite of combination reactions.
For example, 2H₂O → 2H₂ + O₂

3. Displacement Reaction:
In a displacement reaction, an ion (or an atom) in a compound is replaced by an ion (or an atom) of another element. It may be denoted as:
X +YZ → XZ + Y
Displacement reactions fit into two categories:
metal displacement and non-metal displacement.

a) Metal displacement:
A metal in a compound can be displaced by another metal in the uncombined state.
CuSO₄(aq) + Zn(s) → Cu(s) + ZnSO₄(aq)

b) Non-metal displacement:
The non-metal displacement redox reactions include hydrogen displacement and a rarely occurring reaction involving oxygen displacement.
2Na(s) + 2H₂O(I) → 2NaOH(aq) + H₂(g)

The power of these elements as oxidising agents decreases as we move down from fluorine to iodine in group 17 of the periodic table.

Note:
fluorine is the strongest oxidising agent; there is no way to convert F^– ions to F₂ by chemical means. The only way to achieve F₂ from F^– is to oxidise electrolytically,

4. Disproportionation Reactions:
In a disproportionation reaction an element in one oxidation state is simultaneously oxidised and reduced.

Balancing Of Redox Reactions
There are two ways to balance a redox equation.
They are oxidation number method and Half Reaction Method.

a) Oxidation Number Method
The various steps involved in this method are:

1. Write the skeletal equation and assign oxidation numbers to each element. Identify the elements undergoing change in oxidation number.
2. Find out the increase or decrease of oxidation number per atom. Multiply the increase or decrease of oxidation number with number of atoms undergoing the change.
3. Multiply the formulae of the oxidising agent and the reducing agent by suitable integers so as to equalize the total increase or decrease in oxidation number as determined in the above step.
4. Balance the equation with respect to all atoms other the term reduction has than oxygen and hydrogen.
5. Balance oxygen atoms by adding equal number of H₂O molecules to the side deficient in oxygen atoms.
6. For reaction taking place in acidic medium, add H⁺ ions to the side of deficient in hydrogen atoms.
7. For reaction taking place in basic medium, add H₂O molecules to the side deficient in hydrogen atoms and simultaneously add equal number of OH ions on the other side of the equation.

Problem
Permanganate ion reacts with bromide ion in basic medium to give manganese dioxide and bromate ion. Write the balanced ionic equation forthe reaction.
Solution:
The skeletal ionic equation is:
MnO₄^–(aq) + Br^–(aq) → MnO₂(s) + BrO₃^–(aq)

Assign oxidation numbers for Mn and Br

this indicates that permanganate ion is the oxidant and bromide ion is the reductant.

Calculate the increase and decrease of oxidation number, and make the increase equal to the decrease.

As the reaction occurs in the basic medium, and the ionic charges are not equal on both sides, add 2 OH^– ions on the right to make ionic charges equal.
2MnO₄^–(aq) + Br^–(aq) → 2MnO₂(s) + BrO₃^–(aq) + 2OH^–(aq)

Finally, count the hydrogen atoms and add appropri- ‘ ate number of water molecules (i.e. one H20 molecule) on the left side to achieve balanced redox change.
2MnO₄^–(aq) + Br^–(aq) → 2MnO₂(s) + Br0₃^–(aq) + 2OH^–(aq)

b) Half Reaction Method
This method involves identifying the oxidation and reduction reactions in the given skeletal equation and then splitting the reaction accordingly as two half reactions. Each half reaction is then balanced systematically in various steps as outlined below.

Step 1.
Write the skeletal equation and identify the oxidant and reductant.

Step 2.
Write the half reactions for oxidation and reduction separately.

Step 3.
Balance the half reaction with respect to atoms that undergo change in oxidation number. Add electron to whichever side is necessary, to make up for difference in ON.

Step 4.
Balance O-atoms by adding proper number of H₂O molecules to the side deficient in oxygen atoms.

Step 5.
For ionic equations in acid medium, add sufficient H⁺ ions to the side deficient in hydrogen. If the reaction occurs in basic medium, add sufficient H₂O molecules to the side deficient in H atoms to balance H atoms and equal number of hydroxyl ions on the opposite side.

Step 6.
Equalise the number of electrons lost or gained by multiplying the half reaction with suitable integer and add the half reactions to get the final balanced equation.

Problem
Permanganate (VII) ion, MnO₄^– in basic solution oxidises iodide ion, l^– to produce molecular iodine (l₂) and manganese (IV) oxide (MnO₂). Write a balanced ionic equation to represent this redox reaction.
Solution:

Redox Reactions As The Basis For Titrations
In redox systems, the titration method can be adopted to determine the strength of a reductant/ oxidant using a redox sensitive indicator. The usage of indicators in redox titration is illustrated below:
1. In one situation, the reagent itself is intensely coloured, e.g., permanganate ion, MnO₄^–. Here MnO₄^– – acts as the self indicator. The visible endpoint, in this case, is achieved after the last of the reductant (Fe²⁺ or C₂O₄^2-) is oxidised and the first lasting tinge of pink colour appears at MnO₄^– concentration as low as 10^-6 mol dm^-3 (10^-6 mol L^-1), This ensures a minimal ‘overshoot’ in colour beyond the equivalence point, the point where the reductant and the oxidant are equal in terms of their mole stoichiometry.

2. If there is no dramatic auto-colour change (as with Mn04 – titration), there are indicators which are oxidised immediately after the last bit of the reactant is consumed, producing a dramatic colour change. The best example is afforded by Cr2072-, which is not a self-indicator, but oxidises the indicator substance diphenylamine just after the equivalence point to produce an intense blue colour, thus signalling the endpoint.

Redox Reactions And Electrode Pro-Cesses
When zinc rod is dipped in copper sulphate solution, zinc gets oxidised to Zn²⁺ while Cu²⁺ ions are reduced to Cu due to direct transfer of electrons. However, if a zinc rod dipped in ZnSO₄ solution taken in a breaker is connected externally by a conducting wire to a copper rod placed in CuSO₄ solution in another beaker, electrons are transferred indirectly from Zn to Cu. Now, each beaker contains both the oxidised and reduced form of the same substance ‘ called a redox coupe. In this experiment the redox couples developed are Zn²⁺/Zn and Cu²⁺/Cu When the solutions in the two beakers (called electrodes) are joined by a salt bridge (a U-tube containing a solution of KCl, solidified in presence of agar-agar), electrons flow from Zn to Cu while current flows in the reverse direction. The salt bridge provides electrical continuity between the solutions without allowing them to mix with each other. The flow of current is due to a potential difference between Cu and Zn electrodes (or half cells). This experimental set up gives an electrochemical cell or galvanic cell.

The potential of an electrode is a measure of its ability to lose (oxidation) or gain (reduction) electrons. When the concentrations of solutions in the half cells are unity and the temperature is 298 K, the potential of each electrode is known as Standard Electrode Potential (E°). By convention, E° of hydrogen electrode is zero volts and the potential of other electrodes will be a measure of the relative tendency of the active species to be in oxidised/reduced form. A negative E°shows that the redox couple is a stronger reducing agent than H⁺/H₂ couple.

A positive E° shows that the redox couple is a weaker reducing agent than H⁺/H₂ couples. The values of standard reduction potentials of various electrodes are given in the increasing order in an electrochemical series (electromotive series)

Students can Download Chapter 7 Chemical Equilibrium 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 7 Chemical Equilibrium

Introduction
Chemical equilibria are very important in numerous biological and environmental processes. At equilibrium state, the rate of product formed is equal to the rate of reactants formed. The mixture of reactants and products at equilibrium state is called an equilibrium mixture. A equilibrium mixture involving ions in aqueous solutions which is called as ionic equilibrium

Equilibrium In Physical Processes
Phase transformation processes are the familiar example for equilibrium in Physical process.
They are,
Solid \(\rightleftharpoons \) liquid
Liquid \(\rightleftharpoons \) gas
Solid \(\rightleftharpoons \) gas

Solid Liquid Equilibrium
Consider a perfectly insulated thermos flask containing some ice and water at 273 K and normal atmospheric pressure. Since the flask is insulated, there will be no exchange of heat between its contents and the surroundings. It is seen that as long as the temperature remains constant, there is no change in the mass of ice and water. This represents an equilibrium state between ice and water and maybe represented as

We observe there is no change in mass of both ice and water. Since the rate of both reactions are equal.
rate of melting = rate of freezing For any pure substance at 1 atmospheric pressure the temperature at which the solid and liquid phases are at equilibrium is called the normal melting point or normal freezing point of the substance.

Liquid – Vapour Equilibrium
In order to understand the liquid-vapour equilibrium, let us consider evaporation of water in a closed vessel. Consider a closed vessel connected to a manometer. The water vapour present in the vessel is first removed by placing some drying agent such as anhydrous calcium chloride in it for some time. The drying agent is then removed. Now the level of mercury in both the limbs of the manometer will be same. Introduce some water into the vessel and allow to stay at room temperature. Now water starts evaporating. A Pressure will gradually develop within the vessel due to the formation of water vapours. The change of pressure can be easily measured from the manometer. As evaporation continues, the pressure goes on increasing and the level of mercury in the right limb of the manometer starts rising. After some time it is observed that pressure becomes constant. This shows that the quantity of water vapour is not increasing any more, although liquid water is still present in the vessel. This indicates that a state of dynamic equilibrium has been attained between liquid water and water vapours.

At equilibrium, both reaction take place at the same rate. Thus at equilibrium,
rate of evaporation = rate of condensation

The pressure exerted by the vapours in equilibrium with the liquid at a particular temperature is called
vapour pressure of the liquid.

It may be noted that the equilibrium between the vapours and the liquid is attained only in a closed vessel. If the vessel is open, the vapours leave the vessel and get dispersed. Hence the rate of conden-sation will never become equal to the rate of evapo-ration.

Solid – Vapour Equilibrium
Consider systems where solids sublime to vapour phase, For example,

Equilibrium involving Dissolution of Solid or Gas in Liquids
Solids in liquids: In a saturated solution, a dynamic equilibrium exits between the solute molecules in the solid state and in the solution: the rate of dissolution of sugar = rate of crystallisation of sugar. Gases in liquids: This equilibrium is governed by Henry’s law, which states that the mass of a gas dissolved in a given mass of a solvent at any temperature is proportional to the pressure of the gas above the solvent

General Characteristics of Equilibria involving Physical Processes
For the physical processes discussed above, following characteristics are common to the system at equilibrium:

1. Equilibrium is possible only in a closed system at a given temperature.
2. Both the opposing processes occur at the same rate and there is a dynamic but stable condition.
3. All measurable properties of the system remain constant.
4. When equilibrium is attained for a physical process, it is characterised by constant value of one of its parameters at a given temperature.
5. The magnitude of such quantities at any stage indicates the extent to which the reaction has proceeded before reaching equilibrium.

Equilibrium In Chemical Processes – Dynamic Equilibrium
Consider a general reversible reaction
A+B \(\rightleftharpoons \) C+D

Suppose the reaction is carried out in a closed container. In the beginning, the concentrations of A and B are maximum and the concentrations of C and D are minimum (equal to zero). As the reaction proceeds, the concentrations of A and B will decrease whereas the concentrations of C and D will increase. Hence the rate of the forward reaction will be high in the beginning and it will decrease gradually because of the fall in concentrations of A and B. On the other hand the velocity of the reverse reaction will be minimum at the beginning and it will increase gradually due to the increase in concentrations of C and D. Finally a stage will be reached when the rate of the forward reaction becomes equal to the rate of the reverse reaction. This state of the system is known as the state of chemical equilibrium. At this state the concentrations of the reactants and the products remain constant.

We can also start with C and D and make the reaction to proceed in the reverse direction. The concentration of C and D decreases and A and B increases. Finally, equilibrium is attained. One such example is given.
H_2(g) +l_2(g) \(\rightleftharpoons \) 2Hl_(g)

Law Of Chemical Equilibrium And Equilibrium Constant
The relation between rates of reaction and concentrations was given by Guldberg and Wage in 1864. This relation is known as law of mass action.
The relation is,
\(K_{c}=\frac{[C][D]}{[A][B]}\)
For a general reversible reaction of the type,
aA + bB \(\rightleftharpoons \) cC + dD
the equilibrium constant maybe represented as
\(K_{ c }=\frac { [c]^{ c }[D]^{ d } }{ [A]^{ a }{ \left[ B \right] }^{ b } } \)
The equation is known as the expression for the law of chemical equilibrium.

The law of chemical equilibrium or equilibrium law may thus be stated as :
At a given temperature, the product of concentrations of the reaction products raised to the respective stoichiometric coefficient in the balanced chemical equation divided by the prod-uct of concentrations of the reactants raised to their individual stoichiometric coefficients has a constant value. This is known as the Equilibrium Law or Law of Chemical Equilibrium.
If equilibrium constant for the backward reaction is
K’_c then K’_c = \(\frac{1}{K_{e}}\)

Homogeneous Equilibria
In a homogeneous system, all the reactants and products are in the same phase. For example, in the gaseous reaction,
N_2(g) + 3H_2(g) \(\rightleftharpoons \) 2NH_3(g)

Heterogeneous Equilibria
Equilibrium in a system having more than one phase
is called heterogeneous equilibrium.
For example, H₂O_(l) \(\rightleftharpoons \) H₂O_(g)

Applications Of Equilibrium Constants

Predicting the Extent of a Reaction

• If K_c >10³, products predominate over reactants, i.e., if K_c is very large, the reaction proceeds nearly to completion.
• If K_c < 10^-3, reactants predominate over products, i.e., if K_c is very small, the reaction proceeds rarely.
• If K_c is in the range of 10^-3 to 10³, appreciable concentrations of both reactants and products are present.

Predicting The Direction Of The Reaction
The equilibrium constant is also used to find in which direction the reaction mixture of reactants and products will proceed. For this, we have to calculate the reaction quotient (Q_c) and compare with the equilibrium constant (K_c).

The concentrations of the species in Q_c are not necessarily equilibrium values.
For a general reaction aA + bB → cC + dD
\(Q_{ c }=\frac { [c]^{ c }[D]^{ d } }{ [A]^{ a }{ \left[ B \right] }^{ b } } \)
If Q_c > K_c, the reaction will proceed in the direction of the reactants (i.e., reverse reaction).
If Q_c < K_c, the reaction will proceed in the direction of the products (i.e., forward reaction).
If Q_c = K_c, the reaction mixture is already at equilibrium.

Calculating Equilibrium Concentrations
Step 1.
Write the balanced equation forthe reaction.

Step 2.
Under the balanced equation, make a table that lists foreach substance involved in the reaction:
a) the initial concentration,
b) the change in concentration on going to equilibrium, and
c) the equilibrium concentration.

In constructing the table, define x as the concentration ’ (mol/L) of one of the substances that reacts on going to equilibrium, then use the stoichiometry of the reaction to determine the concentrations of the other substances in terms of x.

Step 3.
Substitute the equilibrium concentrations into the equilibrium equation forthe reaction and solve for x. If you are to solve a quadratic equation choose the mathematical solution that makes chemical sense.

Step 4.
Calculate the equilibrium concentrations from the calculated value of x.

Step 5.
Check your results by substituting them into the equilibrium equation.

Problem
3.00 mol of PCl₅ kept in 1L closed reaction vessel was allowed to attain equilibrium at 380 K. Calculate composition of the mixture at equilibrium. K_c = 1.80

Solution
Let x mol of PCl₅ dissociated, At equilibrium:
(3 – x) x x
K_c = [PCl₃][Cl₂][PCl₅]
1.8 = x²/(3 – x)
x² + 1.8x – 5.4 = 0
x = [-1.8 ± √(1.8)² – 4(-5.4)]/2
x = [-1.8 ± √3.24 + 21.6]/2
x = [-1.8 ± 4.98]/2
x = [-1.8 + 4.98]/2
x = 1.59
[PCl₅] = 3.0 -x = 3 – 1.59 = 1.41 M
[PCl₃] = [Cl₂] = x = 1.59 M

Relationship Between Equilibrium Constant K, Reaction Quotient Q And Gibbs Energy G

• ∆G is negative, then the reaction is spontaneous and proceeds in the forward direction.
• ∆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.
• ∆G is O, reaction has achieved equilibrium; at this point, there is no longer any free energy left to drive the reaction.

A mathematical expression of this thermodynamic view of equilibrium can be described by the following equation:

∆G = ∆G° + RT InQ
where, G° is standard Gibbs energy.
At equilibrium, when ∆G = 0 and Q=K_c the equation becomes,
∆G = ∆G° +RTIn K = 0
∆G° = -RTInK
InK = -∆G° / RT
Therefore, K = e^∆Gv/RT

Factors Affecting Equilibria
In order to decide what course the reaction adopts and make a qualitative prediction about the effect of a change in conditions on equilibrium we use Le Chatelier’sprinciple. It 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. This is applicable to both physical and chemical equilibria.

Effect of Concentration Change
When the concentration of any of the reactants or products in a reaction at equilibrium is changed, the composition of the equilibrium mixture changes so as to minimize the effect of concentration changes.

Effect of Pressure Change
A pressure change obtained by changing the volume can affect the yield of products in case of a gaseous reaction where the total number of moles of gaseous reactants and total number of moles of gaseous products are different.

Effect of Inert Gas Addition
If the volume is kept constant and an inert gas such as argon is added which does not take part in the reaction, the equilibrium remains undisturbed. It is because the addition of an inert gas at constant volume does not change the partial pressures orthe molar concentrations of the substance involved in the reaction. So the reaction quotient does not change.

Effect of Temperature Change
Whenever an equilibrium is disturbed by a change in the concentration, pressure or volume, the composition of the equilibrium mixture changes because the reaction quotient, Q_c no longer equals the equilibrium constant, K_c However, when a change in temperature occurs, the value of equilibrium constant, K_c is changed. In general, the temperature dependence of the equilibrium constant depends on the sign of ∆H for the reaction.

• The equilibrium constant for an exothermic reaction (negative ∆H) decreases as the temperature increases.
• The equilibrium constant for an endothermic reaction (positive ∆H) increases as the temperature increases.

Temperature changes affect the equilibrium constant and rates of reactions.

Effect of a Catalyst
A catalyst increases the rate of the chemical reaction by making available a new low energy pathway for the conversion of reactants to products. It increases the rate of forward and reverse reactions that pass through the same transition state and does not affect equilibrium. Catalyst lowers the activation energy for the forward and reverse reactions by exactly ‘ the same amount.

Ionic Equilibrium In Solution
Michael Faraday classified the substances into two categories based on their ability to conduct electricity. One category of substances conduct electricity in their aqueous solutions and are called electrolytes while the other do not and are thus, referred to as non-electrolytes.

Faraday further classified electrolytes into strong and weak electrolytes.

Strong electrolytes on dissolution in water are ionized almost completely, while the weak electrolytes are only partially dissociated.

Acids. Bases And Salts

Arrhenius Concept of Acids and Bases
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). The ionization of an acid HX _(aq) can be represented by the following equations:
HX_(aq) → H⁺(aq) + X^–(aq)
or
HX(aq) + H₂O(l) -> H₃O⁺(aq) + X^–(aq)

The Bronsted-Lowry Acids and Bases
The Danish chemist, Johannes Bronsted and the English chemist, Thomas M. Lowry gave a more general definition of acids and bases. According to Bronsted-Lowry theory, acid is a substance that is capable of donating a hydrogen ion l-T and bases are substances capable of accepting a hydrogen ion, H⁺. In short, acids are proton donors and bases are proton acceptors.

The acid-base pair that differs only by one proton is called a conjugate acid-base pair. Therefore, OH^– is called the conjugate base of an acid H₂O and NH₄⁺ is called conjugate acid of the base NH₃. If Bronsted acid is a strong acid then its conjugate base is a weak base and vice versa.
Consider the example of ionization of hydrochloric acid in water.

Ionization Of Acids And Bases

The Ionization constant of water and its ionic product
Water undergoes self ionisation to a small extent as follows.

Since [H₂O] is constant, K[H₂O]² may be taken as a new constant K_w. Thus,
K_w= [H₃O⁺][OH^–]

Where K_w is called ionic product of water. Its value is 1 x10‘14 mol2 L2 at 298 K. In pure water, the concen-tration of hydronium ions and hydroxyl ions are equal. Therefore in pure water,
[H₃O⁺] = [OH^–] = 1 × 10^-7 mol L^-1

Since the ionisation of water increases with increase of temperature, K_w increases with rise of temperature.

The pH Scale
Hydronium ion concentration in molarity is more conveniently expressed on a logarithmic scale known as the pH scale.

The pH of a solution is defined as the negative logarithm to base 10 of the activity (a_H+) of hydrogen ion.
i.e., pH = – log aH^H+ = – log {[H+]/,mol L^-1}
Acidic solution has pH < 7 Basic solution has pH > 7
Neutral solution has pH = 7

Ionization Constants of Weak Acids

Here, c= initial concentration of the undissociated acid, HXat time, t = 0. α = extent up to which HX is ionized into ions.
K_a = c²a² / c(1 – α) = cα²/1 – A
K_a is called the dissociation or ionization constant.

Ionization of Weak Bases
The equilibrium constant for base ionization is called base ionization constant and is represented by K_b.

When equilibrium is reached, the equilibrium constant can be written as:
K_b = (cα)² / c(1 – α) = cα² / (1 – α)
considering the base-dissociation equilibrium reaction:
K_b = [BH⁺][OH^–]/[B]
Then multiplying and dividing the above expression by [H⁺], we get:
K_b = [BH⁺][OH^–][H⁺]/[B][H⁺]
= {[OH^–][H⁺]}{[BH⁺]/[B][H⁺]}
= K_w/K_a
Then we get the following relation;
pK_a + PK_b = pK_q = 14 (at 298 K)

Common ion effect in the ionization of Acids and Bases.
Common ion effect my be defined as the suppression of the dissociation of a weak electrolyte (weak acid or weak base) by the addition of some strong electrolyte containing a common ion.

Factors Affecting Acid Strength
Dissociation of an acid depends on the strength and polarity of the H-A bond.
Electronegativity of A increases CH₄ < NH₃ < H₂O < HF Acid strength increases

Common Ion Effect in the Ionization of Acids and Bases
Ka = [H⁺] [Ac^–] / [HAc] acetate ions to an acetic acid solution results in decreasing the concentration of hydrogen ions, [H⁺], Also, if H⁺ ions are added from an external source then the equilibrium moves in the direction of undissociated acetic acid. This phenomenon is an example of common ion effect.

Hydrolysis of Salts and the pH of their Solutions
Salts formed by the reactions between acids and bases in definite proportions, undergo ionization in water. The cations/anions formed on ionization of salts either exist as hydrated ions in aqueous solutions or interact with water to reform corresponding acids/bases depending upon the nature of salts. The later process of interaction between water and cations/anions or both of salts is called hydrolysis.

Buffer Solutions
The solutions which resist change in pH on dilution or with the addition of small amounts of acid or alkali are called Buffer Solutions.

Solubilityequilibriaof Sparingly Soluble Salts

Solubility Product Constant
The equilibrium between the undisolved solid and the ions in a saturated solution can be represented by the equation:

We call K_sp the solubility product constant or simply solubility product.

Thus, solubility product of a salt is the product of concentration of ions in its saturated solution, raised to a power equal to the number of times the ions occur in the equation representing the dissociation of the salt.

The term K_sp in equation is given by Q_sp when the concentration of one or more species is not the concentration under equilibrium. Obviously under equilibrium conditions K_sp = Q_sp but otherwise it gives the direction of the processes of precipitation or dissolution.

Common Ion Effect on Solubility of Ionic Salts
The solubility of salts of weak acids like phosphates increases at lower pH. This is because at lower pH the concentration of the anion decreases due to its proto-nation. This, in turn, increases the solubility of the salt so that K_sp = Q_sp.

Ncert Supplementary Syllabus

Designing Buffer Solution
Knowledge of pK_a, pK_b and equilibrium constant help us to prepare the buffer solution of known pH. Let us see how we can do this.

Preparation of Acidic Buffer
To prepare a buffer of acidic pH we use weak acid and its salt formed with strong base. We develop the equation relating the pH, the equilibrium constant, K_a of weak acid and ratio of concentration of weak acid and its conjugate
base. For the general case where the weak acid HA ionises in water,

ratio of concentration of conjugate base (anion) of the acid and the acid present in the mixture. Since acid is a weak acid, it ionises to a very little extent ‘and concentration of [HA] is negligibly different from concentration of acid taken to form buffer. Also, most of the conjugate base, [A^–], comes from the ionisation of salt of the acid. Therefore, the concentration of conjugate base will be negligibly different from the concentration of salt. Thus, equation (A-2) takes the form: pH-pK_a + log\(\frac{[\mathrm{Salt}]}{[\mathrm{Acid}]}\)

In the equation (A-1), if the concentration of [A^–] is equal to the concentration of [HA], then pH = pK_a because value of log 1 is zero. Thus if we take molar concentration of acid and salt (conjugate base) same, the pH of the buffer solution will be equal to the pK_a of the acid. So for preparing the buffer solution of the required pH we select that acid whose pK_a is close to the required pH. For acetic acid pK_a value is 4.76, therefore pH of the buffer solution formed by acetic acid and sodium acetate taken in equal molar concentration will be around 4.76.

A similar analysis of a buffer made with a weak base and its conjugate acid leads to the result,

pH of the buffer solution can be calculated by using the equation pH + pOH =14.

We know that pH + pOH = pK_w and pK_a + pK_b = pK_w On putting these values in equation (A-3) it takes the form as follows:

If molar concentration of base and its conjugate acid (cation) is same then pH of the buffer solution will be same as pK_a for the base. pK value for ammonia is 9.25; therefore a buffer of pH close to 9.25 can be obtained by taking ammonia solution and ammonium chloride solution of equal molar concentration. For a buffer solution formed by ammonium chloride and ammonium hydroxide, equation (A-4) becomes:

pH of the buffer solution is not affected by dilution because ratio under the logarithmic term remains unchanged.

Students can Download Chapter 5 States of Matter 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 5 States of Matter

Introduction
The observable characteristics of chemical systems represent bulk properties of matter. Chemical properties do not depend the physical state of matter but chemical reactions do.

Inter Molecular Forces
Intermolecular forces are the forces of attraction and repulsion between interacting particles (atoms and molecules). This term does not include the electrostatic forces that exist between the two oppositely charged ions and the forces that hold atoms of a molecule together i.e., covalent bonds.

Dispersion Forces Or London Forces
A nonpolar atom or molecule has a positive centre surrounded by a symmetrical negative electron cloud. The displacement of electron cloud creates an instantaneous dipole temporarily. This instantaneous dipole distorts the electron distribution of other atoms or molecules which are close to it and induces dipole in them also. In this way, a large number of nonpolar molecules become temporarily polar and they mutually attracted by weak attractive forces. These forces are very weak and are known to operate in all types of molecules.

Dipole-Dipole Forces
These type of interactions occur in polar molecules having permanent dipoles such as HCl, HBr, H₂S, etc. Such molecules possess partial charges of opposite sign at their ends. The positive end of one molecule attracts the negative end of the other molecule and vice versa. A simple example is the of H-Cl in which chlorine being more electronegative acquires slight negative charge whereas hydrogen becomes slightly positively charged. The dipole-dipole inter-action then takes place in H-Cl as follows.

Dipole-Induced Dipole Forces
These type of interactions are found in a mixture, containing polar and nonpolar molecules. When a nonpolar molecule is brought neara polar molecule, the positive end of the polar molecule attracts the electron cloud of the nonpolar molecule. Thus a polarity is induced in the nonpolar molecule. Then there will be attractive interacting between the polar molecule and the induced dipole of the nonpolar molecule.

Hydrogen Bond
Hydrogen bond can be defined as the attractive force which binds hydrogen atom of one molecule with the electronegative atom (F, O or N) 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: H^δ+ – X^δ-

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:

Thermal Energy
Thermal energy is the energy of a body arising from motion of its atoms or molecules. It is directly proportional to the temperature of the substance. The movement of particles using thermal energy is called thermal motion.

Intermolecular Forces Vs Thermal Interactions
Intermolecular forces tend to keep the molecules together but thermal energy of the molecules tends to keep them apart. Three states of matter are the result of balance between intermolecular forces and the thermal energy of the molecules.

The Gaseous State

• Gases are highly compressible.
• Gases exert pressure equally in all directions.
• Gases have much lower density than the solids and liquids.
• The volume and the shape of gases are not fixed. These assume volume and shape of the container.
• Gases mix evenly and completely in all proportions without any mechanical aid.

The Gas Laws

Boyle’s Law (Pressure – Volume Relationship)
On the basis of his experiments, Robert Boyle reached to the conclusion that at constant temperature, the pressure of a fixed amount of gas
varies inversely with its volume. This is known as Boyle’s law. It can be written as p ∝ \(\frac{1}{V}\)
where temperature(T) and number of moles(n)are constant.

If a fixed amount of gas at constant temperature T occupying volume V₁ at pressure p₁ undergoes expansion, so that volume becomes V₂ and pressure becomes p₂, then according to Boyle’s law :
p₁V₁ = p₂V₂ = constant
\(\Rightarrow \frac{p_{1}}{p_{2}}=\frac{V_{2}}{V_{1}}\)

Here T is constant and the graph is called isotherm

Charles’ Law (Temperature – Volume Relationship)
Charles’ law, which states that pressure remaining constant, the volume of a fixed mass of a gas is directly proportional to its absolute temperature. According to this law

Here we use new temperature scale called the kelvin temperature scale or Absolute temperature scale, t °C in Celsius scaleis equal to (273.15+t) kelvin in kelvin scale.

Each line of the volume vs temperature graph is called isobar. The lowest hypothetical or imaginary temperature at which gases are supposed to occupy zero volume is called Absolute zero.
We can see that the volume of the gas at – 273.15 °C will be zero.

Gay Lussac’s Law (Pressure-Temperature Relationship)
The relationship between pressure and temperature was given by Joseph Gay Lussac and is known as Gay Lussac’s law. It states that at constant volume, pressure of a fixed amount of a gas varies directly with the temperature. Mathematically,
P ∝ T
⇒ \(\frac{p}{T}\) = constant
This relationship can be derived from Boyle’s law and Charles’ Law.Each line of Pressure vs temperature (kelvin) graph at constant molar volume is called isochore.

Avogadro Law (Volume -Amount Relationship)
In 1811 Italian scientist Amedeo Avogadro tried to combine conclusions of Dalton’s atomic theory and Gay Lussac’s law of combining volumes which is now known as Avogadro law. It states that equal volumes of all gases under the same conditions of temperature and pressure contain equal number of molecules.

Mathematically we can write v α n where n is the number of moles.

The number of molecules in one mole of a gas has been determined to be 6.022 *1023and is known as Avogadro constant. A gas that follows Boyle’s law, Charles’ law and Avogadro law strictly is called an ideal gas.

Ideal Gas Equation
The combination of Boyle’s law, Charles’ law, and Avagadro’s law leads to an equation which gives the combined effect of change of temperature and pressure on the volume of a gas.
According to Boyle’s law, V α \(\frac{1}{P}\) ——- (i) (at constant T and n)
According to Charles’ Law, V α T ——- (ii) (at constant P and n)
According to Avogardro’s Law, V α n ——- (iii) (at constant T and P

Where R is a constant known as the universal gas constant. The equation is known as ideal gas equation.

Density and Molar Mass of a Gaseous Substance
Ideal gas equation can be rearranged as follows:

we get \(\frac{d}{M}=\frac{p}{R T}\)
(where d is the density)
On rearranging equation we get the relationship for calculating molar mass of a gas.
\(M=\frac{d R T}{p}\)

Dalton’s Law of Partial Pressures
The law was formulated by John Dalton in 1801. It states that the total pressure exerted by the mix-ture of non-reactive gases is equal to the sum of the partial pressures of individual gases.
P_Total = P₁ + P₂ + P₃ + ………. (at constant T, V)

where p_total is the total pressure exerted by the mixture of gases and p₁, p₂, p₃ etc. are partial pressures of gases.

Partial pressure in terms of mole fraction

Kinetic Molecular Theory Of Gases
Maxwell, Boltzmann, and others put forward a theoretical model of the gas. The theory is known as
Kinetic molecular theory of gases or microscopic
model of gases.
Postulates of kinetic molecular theory.

1.  All gases are made up of a large number of extremely small particles called molecules.
2. The molecules are separated from one another by large distances so that the actual volume of the molecules is negligible as compared to the total volume of gas.
3. The molecules are in a state of continuous rapid motion in all directions. During their motion, they keep on colliding with one another and also with the walls of the container.
4. Molecular collisions are perfectly elastic i.e. there is no net loss or gain of energy in their collisions. However, there may be redistribution of energy during such collisions.
5. There are no attractive forces between the molecules. They move completely independent of each other.
6. The pressure exerted by the gas is due to the bombardment of its molecules on the walls of the container.
7. At any instant, different molecules possess different velocities and hence different energies. However, the average kinetic energy of the molecules is directly proportional to its absolute temperature.

Behaviour Of Real Gases:
Deviation From Ideal Gas Behaviour
There are two types of curves are seen in the graph. In the curves for dihydrogen and helium, as the pressure increases the value of pV also increases. The second type of plot is seen in the case of other gases like carbon monoxide and methane. In these plots first, there is a negative deviation from ideal behaviour, the pV value decreases with increase in pressure and reaches to a minimum value characteristic of a gas. After that pV value starts increasing. The curve then crosses the line for ideal gas and after that shows positive deviation continuously. It is thus, found that real gases do not follow ideal gas equation perfectly under all conditions.

We find that two assumptions of the kinetic theory do not hold good. These are

1. There is no force of attraction between the molecules of a gas.
2. Volume of the molecules of a gas is negligibly small in comparison to the space occupied by the gas.

If assumption (a) is correct, the gas will never liquify. This means that forces of repulsion are powerful enough and prevent squashing of molecules in tiny volume. If assumption (b) is correct, the pressure vs volume graph of experimental data (real gas) and that theoritically calculated from Boyles law (ideal gas) should coincide.

The volume occupied by the molecules also becomes significant because instead of moving in volume V, these are now restricted to volume (V-nb) where nb is approximately the total volume occupied by the molecules themselves. Here, b is a constant. Having taken into account the corrections for pressure and volume, we can rewrite equation as This equation is known as van der Waals’ equation.

Value of ‘a’ is measure of magnitude of intermolecular attractive forces within the gas and is independent of temperature and pressure.
Real gases show ideal behaviour when conditions of temperature and pressure are such that the intermolecular forces are practically negligible. The real gases show ideal behaviour when pressure approaches zero.

The deviation from ideal behaviour can be measured in terms of compressibility factor Z, which is the ratio of product pV and nRT. Mathematically
\(z=\frac{p V}{n R T}\)

For ideal gas Z = 1 at all temperatures and pressures because pV = nRT.
At high pressure, all the gases have Z > 1. These are more difficult to compress. At intermediate pressures, most gases have Z < 1. The temperature at which a real gas obeys ideal gas law over an appreciable. range of pressure is called Boyle temperature or Boyle point.

Liquefaction Of Gases
The highest temperature at which liquefaction of the gas first occurs is called Critical temperature (T_c). Volume of one mole of the gas at critical temperature is called critical volume (V_c) and pressure at this temperature is called critical pressure (P_c).
The critical temperature, pressure, and volume are called critical constants.

Liquid State
Intermolecular forces are stronger in liquid state than in gaseous state.

Vapour Pressure
The pressure exerted by the vapour on the walls of the container is known as vapour pressure.

Surface Tension
Liquids tend to minimize their surface area. The molecules on the surface experience a net downward force and have more energy than the molecules in the bulk, which do not experience any net force. This characteristic property of liquids is known as surface tention. Liquids tend to have minimum number of molecules at their surface due to surface tention.

If surface of the liquid is increased by pulling a molecule from the bulk, attractive forces will have to be overcome. This will require expenditure of energy. The energy required to increase the surface area of the liquid by one unit is defined as surface energy.

Viscosity
It is a common observation that certain liquids flow faster than others. For example, liquid like water, ether, etc. flow rapidly while liquids like glycerine, castor oil, honey, etc. flow slowly. These differences in flow rates result from a property known as viscosity. Every liquid has some internal resistance to flow. This internal resistance to flow possessed by a liquid is called its viscosity. Liquids which flow slowly have high internal resistance and are said to have high viscosity. On the other hand, liquids which flow rapidly have low internal resistance and are said to have low viscosity.

Viscosity is also related to intermolecular forces in liquids. If the intermolecular forces are large, the vis-cosity will be high. Viscosity of a liquid decreases with rise in temperature. This is because at higher temperature the attractive forces between molecules are overcome by the increased kinetic energies of the molecules.

Ncert Supplementary Syllabus

Kinetic Energy And Molecular Speeds
Molecules of gases remain in continuous motion. While moving they collide with each other and with the walls of the container. This results in change of their speed and redistribution of energy. So the speed and energy of all the molecules of the gas at any instant are not the same. Thus, we can obtain only average value of speed of molecules. If there are n number of molecules in a sample
and their individual speeds are u₁, u₂, ……… u_n, then average speed of molecules u_av can be calculated as follows:
\(u_{a v}=\frac{u_{1}+u_{2}+\ldots u_{n}}{n}\)

Maxwell and Boltzmann have shown that actual distribution of molecular speeds depends on temperature and molecular mass of a gas. Maxwell derived a formula for calculating the number of molecules possessing a particular speed. Fig. A(1) shows schematic plot of number of molecules vs. molecular speed at two different temperatures T₁ and T₂ (T₂ is higher than T₁). The distribution of speeds shown in the plot is called Maxwell-Boltzmann distribution of speeds.

The graph shows that number of molecules possessing very high and very low speed is very small. The maximum in the curve represents speed possessed by maximum number of molecules. This speed is called most probable speed, u_mp. This is very close to the average speed of the molecules. On increasing the temperature most probable speed increases. Also, speed distribution curve broadens at higher temperature. Broadening of the curve shows that number of molecules moving at higher speed increases. Speed distribution also depends upon mass of molecules. At the same temperature, gas molecules with heavier mass have slower speed than lighter gas molecules. For example, at the same temperature, lighter nitrogen molecules move faster than heavier chlorine molecules. Hence, at any given temperature, nitrogen molecules have higher value of most probable speed than the chlorine molecules. Though at a particular temperature the individual speed of molecules keeps changing, the distribution of speeds remains same.

The kinetic energy of a particle is given by the expression:
Kinetic Energy = \(\frac{1}{2}\) mu²
Therefore, if we want to know average translational kinetic energy, \(\frac{1}{2} m \overline{u^{2}}\) , for the movement of a gas particle in a straight line, we require the value of mean of square of speeds, \(\overline{u^{2}}\), of all molecules. This is represented as follows:
\(u^{2}=\frac{u_{1}^{2}+u_{2}^{2}+\ldots . u_{n}^{2}}{n}\)

The mean square speed is the direct measure of the average kinetic energy of gas molecules. If we take the square root of the mean of the square of speeds then we get a value of speed which is different from most probable speed and average speed. This speed is called root mean square speed and is given by the expression as follows:
\(u_{m s}=\sqrt{u_{2}}\)

Root mean square speed, average speed and the most probable speed have following relationship:
u_rms u_av u_mp

The ratio between the three speeds is given below:
u_mp : u_av : u_rms :: 1 : 1.128 : 1.224

Students can Download Chapter 3 Classification of Elements and Periodicity in Properties 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 3 Classification of Elements and Periodicity in Properties

Introduction
The systematic classification of elements made the study of elements easy. In this unit, we will study the historical development of the periodic table and also learn how elements are classified.

Genesis Of Periodic Classification
While Dobereiner initiated the study of periodic relationship, it was Mendeleev who was responsible for publishing the Periodic Law for the first time. It states as follows:

The properties of the elements are a periodic function of their atomic weights.
Mendeleev arranged elements in horizontal rows and vertical columns of a table in order of their increasing ‘ atomic weights. Elements with similar properties occupied the same vertical column or group. He realized that some of the elements did not fit in with his scheme of classification if the order of atomic weight was strictly followed. He ignored the order of atomic weights, thinking that the atomic measurements might be incorrect, and placed the elements with similar properties together.

At the same time, keeping his primary aim of arranging the elements of similar properties in the same group, he proposed that some of the elements were still undiscovered and, therefore, left several gaps in the table. He left the gap under aluminium and a gap under silicon, and called these elements Eka-Aluminium and Eka-Silicon. Mendeleev predicted the existence of gallium and germanium, and their general physical properties. These elements were discovered later.

Modern Periodic Law And The Present Form Of The Periodic Table
Modem periodic law states that “The physical and chemical properties of the elements are periodic functions of their atomic numbers”. Atomic number is equal to the nuclear charge and the elements are arranged in the increasing order of atomic number.

The period number correspond to the highest principal quantum number (n) of the elements.

Nomenclature Of Elements With Atomic Number Greater Than 100
The names (IUPAC) are derived directly form the atomic number using numerical roots for zero and numbers 1 to 9. The roots are linked together in the order of digits and ‘ium’ is added at the end. The roots for 0,1, 2 9 are nil, un, bi, tri, quad, pent, hex, sept, oct and enn respectively. For example, the element with atomic number 110 will have the name Ununnilium (Un+ un+nil + ium), The element with atomic number 114 has the name Ununquadium (un + un + quad + ium) and the element with atomic number 120 will be Unbinilium (un + bi + nil + ium).

Electronic Configurations And Types Of Elements: s, p, d, f- Blocks

The s-Block Elements
The elements of Group 1 (alkali metals) and Group 2 (alkaline earth metals) which have ns1 and ns2 outermost electronic configuration belong to the s-Block Elements. They are all reactive metals with low ionization enthalpies.

They lose the outermost electron(s) readily to form 1+ ion (in the case of alkali metals) or 2+ ion (in the case of alkaline earth metals). The metallic character and the reactivity increase as we go down the group. Because of high reactivity they are never found pure in nature.

The compounds of the s-block elements, with the exception of those of lithium and beryllium are predominantly ionic.

The p-Block Elements
The p-Block Elements comprise those belonging to group 13 to 18 and these together with the s-Btock Elements are called the Representative Elements or Main Group Elements. The outermost electronic configuration varies from ns2np1 to ns2np6in each period. At the end of each period is a noble gas element with a closed valence shell ns2np6 configuration. All the orbitals in the valence shell of the noble gases are completely filled by electrons and it is very difficult to alter this stable arrangement by the addition or removal of electrons. The noble gases thus exhibit very low chemical reactivity. Preceding the noble gas family are two chemically important groups of non-metals. They are the halogens (Group 17) and the chalcogens (Group 16). These two groups of elements have high negative electron gain enthalpies and readily add one or two electrons respectively to attain the stable noble gas configuration. The non-metallic character increases as we move from left to right across a period and metallic character increases as we go down the group.

The d-Block Elements (Transition Elements)
These are the elements of group 3 to 12 in the centre of the Periodic Table. These are characterised by the filling of inner d orbitals by electrons and are therefore referred to as d-Block Elements. These elements have the general outer electronic configuration (n-1) d1-10ns^2. They are all metals. They mostly form coloured ions, exhibit variable valence (oxidation states), paramagnetism and oftenly used as catalysts. However, Zn, Cd and Hg which have the electronic configuration, (n-1) d10ns2 do not show most of the properties of transition elements. In a way, transition metals form a bridge between the chemically active metals of s-block elements and the less active elements of groups 13 and 14 and thus take their familiar name “Transition Elements”.

The f-Block Elements (Inner-Transition Elements)
The two rows of elements at the bottom of the Periodic Table, called the Lanthanoids, Ce(Z = 58) -Lu(Z = 71) and actinoids, Th(Z = 90)-Lr(Z= 103) are characterised by the outer electronic configuration (n-2)f1-14 (n-1 )d°-1ns2. The last electron added to each element is filled in f- orbital. These two series of ‘ elements are hence called the Inner Transition Elements (f-Block Elements). They are all metals. Within each series, the properties of the elements are quite similar. The elements after Uranium are called Transuranium Elements.

Metals, Non-metals and Metalloids. In addition to displaying the classification of elements into s, p, d and f-blocks, they can be divided into Metals and Non-Metals. Metals usually have high melting and boiling points. They are good conductors of heat and electricity. They are malleable (can be flattened into thin sheets by hammering) and ductile (can be drawn into wires). In contrast, non-metals are located at the top right hand side of the Periodic Table.

In fact, in a horizontal row, the property of elements change from metallic on the left to non-metallic on the right. Non-metals are usually solids or gases at room temperature with low melting and boiling points (boron and carbon are exceptions). They are poor conductors of heat and electricity. Most nonmetallic solids are brittle and are neither malleable nor ductile. The elements become more metallic as we go down a group; the nonmetallic character increases as one goes from left to right across the Periodic Table. The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) running diagonally across the Periodic Table show properties that are characteristic of both metals and nonmetals. These elements are called Semi-metals or Metalloids.

Periodic Trends In Properties Of Elements
Most of the properties such as atomic radius, ionic radius, Ionisation enthalpy, electron gain enthalpy and electron negativity are directly related to electronic configuration of their atoms. They show variation with change in atomic number within a period or a group.

Trends In Physical Properties

1. Atomic Radius :
lt is defined as the distance from the centre of the nucleus of an atom to the outermost shell of electrons. Electron cloud surrounding the atom does not have a sharp boundary since, the determination of the atomic size cannot be precise. Hence it is expressed in terms of different types of radii. Some of these are covalent radius and metallic radius. Covalent radius is defined as one half of the distance between the centres of nuclei of two similar atoms bonded by a single covalent bond. Metallic radius may be defined as half of the internuclear distance between two adjacent atoms in the metallic crystal.

2. Ionic Radius:
The removal of an electron from an atom results in the formation of a cation, whereas gain of an electron leads to an anion. The ionic radii can be estimated by measuring the distances between cations and anions in ionic crystals. When we find some atoms and ions which contain the same number of electrons, we call them isoelectronic species. For example, O2-, F~, Na+ and Mg2+ have the same number of electrons (10). Their radii would be different because of their different nuclear charges.

3. Ionization Enthalpy:
A quantitative measure of the tendency of an element to lose electron is given by its Ionization Enthalpy. It represents the energy required to remove an electron from an isolated gaseous atom (X) in its ground state. To understand the trends in ionization enthalpy, we have to consider two factors: (i) the attraction of electrons towards the nucleus, and (ii) the repulsion of electrons from each other. The effective nuclear charge experienced by a valence electron in an atom will be less than the actual charge on the nucleus because of “shielding” or “screening” of the valence electron from the nucleus by the intervening core electrons.

The first ionization enthalpy of boron (Z = 5) is slightly less than that of beryllium (Z = 4) even though the former has a greater nuclear charge. It is because, s-electron is attracted to the nucleus more than a p-electron. In beryllium, the electron removed during the ionization is an s-electron whereas the electron removed during ionization of boron is a p-electron. The penetration of a 2s-electron to the nucleus is more than that of a 2p-electron; hence the 2p electron of boron is more shielded from the nucleus by the inner core of electrons than the 2s electrons of beryllium. Therefore, it is easier to remove the 2p-electron from boron Compared to the removal of a 2s-electron from beryllium.

Thus, boron has a smaller first ionization enthalpy than beryllium. Another “anomaly” is the smaller first ionization enthalpy of oxygen compared to nitrogen. This arises because in the nitrogen atom, three 2p-electrons reside in different atomic orbitals (Hund’s rule) whereas, in the oxygen atom, two of the four 2p-electrons must occupy the same 2p-orbital resulting in an increased electron-electron repulsion. Consequently, it is easier to remove the fourth 2p-electron from oxygen than it is, to remove one of the three 2p-electrons from nitrogen.

4. Electron Gain Enthalpy :
When an electron is added to a neutral gaseous atom (X) to convert it into a negative ion, the enthalpy change accom-panying the process is defined as the Electron Gain Enthalpy (∆_eg H).

5. Electronegativity:
A qualitative measure of the ability of an atom in a chemical compound to attract shared electrons to itself is called electronegativity. Unlike ionization enthalpy and electron gain enthalpy, it is not a measurable quantity. However, a number of numerical scales of electronegativity of elements viz., Pauling scale, Mulliken-Jaffe scale, Allred-Rochow scale have been developed.

Trends In Chemical Properties
1. Oxidation State :
The atomic property, valency is better explained in terms of oxidation state. It is the charge which an atom of element has or appears to have when present in the combined state. Electronegative elements generally acquire negative oxidation states while electropositive elements acquire positive oxidation states.

2. Anomalous properties of second-period elements:
The first element of each group in s and p block differs in many respects from the remaining members of the respective groups. This is due to their small size, high charge/ radius ratio, high electronegativity and availability of less valence orbitals. The first member has only 4 valence orbitals (2s, 2p) whereas the second member of the same group will have nine valence orbitals (3s, 3p, 3d) for bonding. B can form only (BF₄)^– while Al forms (AlF₆)^3-

In group 1 only Li forms covalent compounds and in many respects, Li resembles Mg of group 2. Similarly, Be resembles Al of group 13. This type of similarity in properties is known as diagonal relationship.

Chemical Reactivity
Across a period ionisation enthalpy increases and electron gain enthalpy becomes more negative. Thus elements at the extreme left show lower ionisation enthalpies (more electropositive nature) and those at the right (excluding nobel gases) show larger negative electron gain enthalpies (more electronegative). Therefore high chemical reactivity is found with elements at the two extremes compared to those at the centre. Electropositivity leads to metallic behaviour and electronegativity leads to non-metallic behaviour.