OXIDATION

Oxidation is defined as the addition of an oxygen/electronegative element to a substance or removal of a hydrogen/ electropositive element from a susbtance. For example,

REDUCTION

Reduction is defined as the removal of oxygen/electronegative elements from a substance or addition of hydrogen or electropositive elements to a substance. For example,

REDOX REACTION IN TERMS OF ELECTRON TRANSFER REACTION

A few examples of redox reaction on the basis of electronic concept are given below:
According to the electronic concept every redox reaction consists of two steps known as half reactions.
(i) Oxidation reaction: Half reactions that involve loss of electrons are called oxidation reactions.
(ii) Reduction reaction: Half reactions that involve gain of electrons are called reduction reactions.
Oxidizing agent: Acceptor of electrons.
Reducing agent: Donar of electrons.

OXIDATION NUMBER

It is the oxidation state of an element in a compound which is the charge assigned to an atom of a compound is equal to the number of electrons in the valence shell of an atom that are gained or lost completely or to a large extent by that atom while forming a bond in a compound.

Rules for Assigning Oxidation Numbers

(i) The oxidation number of an element in its elementary form is zero.
For example, H2, 02, N2 etc. have oxidation number equal to zero.
(ii) In a single monatomic ion, the oxidation number is equal to the charge on the ion. For example, Na+ ions have an oxidation number of +1 and Mg2+ ions have +2.
(iii) Oxygen has oxidation number -2 in its compounds.
(iv) In non-metallic compounds of hydrogen like HCl, H2S, H2O oxidation number of hydrogen = + 1 but in metal hydrides oxidation number of hydrogen = -1 [LiH, NaH, CaH2 etc.]
(v) In compounds of metals and non-metals metals have positive oxidation number while non-metals have negative oxidation number. For example, In NaCl. Na has +1 oxidation number while chlorine has -1.
(vi) If in a compound there are two non-metallic atoms the atoms with high electronegativity are assigned negative oxidation number while other atoms have positive oxidation number.
(vii) The algebraic sum of the oxidation number of all atoms in a compound is equal to zero.
(viii) In polyatomic ion the sum of the oxidation no. of all the atoms in the ion is equal to the net charge on the ion.
For example, in (C03)2—Sum of carbon atoms and three oxygen atoms is equal to -2.
Fluorine (F2) is so highly reactive non-metal that it displaces oxygen from water.

DISPROPORTIONATION REACTION

In a disproportionation reaction an element in one oxidation state is simultaneously oxidized and reduced. For example,

Hence, the oxygen of peroxide, which is present in -1 oxidation state is connected to zero oxidation state and in 02 and in H2O decreases to -2 oxidation state.

FRACTIONAL OXIDATION NUMBERS

Elements as such do not have any fractional oxidation numbers.
When the same elements are involved in different bonding in a species, their actual oxidation states are whole numbers but an average of these is fractional. For example, In C302

Fractional O.N. of a particular element can be claculated only if we know about the structure of the compound or in which it is present.

BALANCING OF REDOX REACTIONS

Oxidation Number Method:

Following steps are involved:
(a) Write the correct formula for each reactant and product.
(b) By assigning the oxidation change in oxidation number can be identified.
(c) Calculate the increase and decrease in oxidation number per atom with respect to the reactants. If more than one atom is present then multiply by suitable coefficient.
(d) Balance the equation with respect to all atoms. Balance hydrogen and oxygen atoms also.
(e) If the reaction is carried out in an acidic medium, use H+ ions in the equation. If it is in basic medium use OH– ions.
(f) Hydrogen atoms in the expression can be balanced by adding (H20) molecules to the reactants or products.
If there are the same number of oxygen atoms on both sides of the equation then it represents the balanced redox reaction.

Half Reaction Method:

In this method two half equations are balanced separately and then added together to give a balanced equation.

REDOX REACTIONS AS THE BASIS FOR TITRATION

Potassium Permanganate Titration:

In these titrations potassium permanganate (pink in colour) acts as an oxidising agent in the acidic medium while oxalic acid or some ferrous salts acts as a reducing agent.
The ionic equation can be written as:

These are the examples of redox titration.
On both these titrations, potassium permanganate itself acts as indicator.
It is commonly known as a self indicator.
The appearance of pink colour in the solution represents the end points.

Potassium Dichromate Titration :

In place of potassium permanganate, potassium dichromate can also be used in the presence of dil. H2S04.
The ionic equation for the redox reaction with FeS04 (Fe2+ ions) is given.

LIMITATION OF CONCEPT OF OXIDATION NUMBER

According to the concept of oxidation number, oxidation means increase in oxidation number – by loss of electrons and reduction means decrease in oxidation number by the gain of electrons. However, during oxidation there is decrease in electron density while increase in electron density around the atom undergoing reduction.

REDOX REACTIONS AND ELECTRODE

Processes-Electrochemical Cells

A device in which the redox reaction is carried indirectly and the decrease in energy appears as the electrical energy are called electrochemical cell.

Electrolytic Cell

The cell in which electrical energy is converted into chemical energy. Example, when a lead storage battery is recharged, it acts as electrolytic cell.

IMPORTANT TERMS AND DEFINITIONS

SYSTEM

Refers to the portion of the universe which is under observation.

SURROUNDINGS

Everything else in the universe except the system is called surroundings.
The Universe = The System + The Surroundings.

OPEN SYSTEM

In a system, when there is exchange of energy and matter taking place with the surroundings, then it is called an open system.
For Example: Presence of reactants in an open beaker is an example of an open system.

CLOSED SYSTEM

A system is said to be a closed system when there is no exchange of matter ‘ but exchange of energy is possible. For example: The presence of reactants in a closed vessel made of conducting material.

ISOLATED SYSTEM

In a system, when no exchange of energy or matter takes place with the surroundings, it is called an isolated system.
For example: The presence of reactants in a thermoflask, or substance in an insulated closed vessel is an example of an isolated system.

HOMOGENEOUS SYSTEM

A system is said to be homogeneous when all the constituents present are in the same phase and is uniform throughout the system.
For example: A- mixture of two miscible liquids.

HETEROGENEOUS SYSTEM

A mixture is said to be heterogeneous when it consists of two or more phases and the composition is not uniform.
For example: A mixture of insoluble solid in water. ’

THE STATE OF THE SYSTEM

The state of a thermodynamic system means its macroscopic or bulk properties which can be described by state variables:
Pressure (P), volume (V), temperature (T) and amount (n) etc.
They are also known as state functions.

ISOTHERMAL PROCESS

When the operation is carried out at constant temperature, the process is said to be isothermal. For isothermal processes, dT = 0 Where dT is the change in temperature.

ADIABATIC PROCESS

It is a process in which no transfer of heat between system and surroundings takes place.

ISOBARIC PROCESS

When the process is carried out at constant pressure, it is said to be isobaric. i.e. dP = 0

ISOCHORIC PROCESS

A process when carried out at constant volume, it is known as isochoric in nature.

CYCLIC PROCESS

If a system undergoes a series of changes and finally returns to its initial state, it is said to be a cyclic process.

REVERSIBLE PROCESS

When in a process, a change is brought in such a way that the process could, at any moment, be reversed by an infinitesimal change. The change r is called reversible.

INTERNAL ENERGY

It is the sum of all the forms of energies that a system can possess.
In thermodynamics, it is denoted by AM which may change, when
— Heat passes into or out of the system
— Work is done on or by the system
— Matter enters or leaves the system.

CHANGE IN INTERNAL ENERGY BY DOING WORK

Let us bring the change in the internal energy by doing work.
Let the initial state of the system be state A and Temp.
TA Internal energy = uA
On doing’some mechanical work the new state is called state B and the temp. TB.
It is found to be :TB > TA
uB is the internal energy after change.
∴ Δu = uB – uA

CHANGE IN INTERNAL ENERGY BY TRANSFER OF HEAT

Internal energy of a system can be changed by the transfer of heat from the surroundings to the system without doing work.
Δu = q, Where q is the heat absorbed by the system.
It can be measured in terms of temperature difference.
q is +ve when heat is transferred from the surroundings to the system. q is -ve when heat is transferred from the system to surroundings.
When change of state is done both by doing work and transfer of heat.
Δu = q + w

FIRST LAW OF THERMODYNAMICS (LAW OF CONSERVATION OF ENERGY):

It states that energy can neither be created nor be destroyed. The energy of an isolated system is constant.
Δu = q + w.

WORK (PRESSURE-VOLUME WORK)

Let us consider a cylinder which contains one mole of an ideal gas in which a frictionless piston is fitted.

WORK DONE IN ISOTHERMAL AND REVERSIBLE EXPANSION OF IDEAL GAS

ISOTHERMAL AND FREE EXPANSION OF AN IDEAL GAS

For isothermal expansion of an ideal gas into vacuum W = 0

ENTHALPY (H):

It is defined as the total heat content of the system.
It is equal to the sum of internal energy and pressure-volume work.
Mathematically, H = U + PV

CHANGE IN ENTHALPY

Change in enthalpy is the heat absorbed or evolved by the system at constant pressure.
ΔH = qp
For exothermic reaction (System loses energy to Surroundings),
ΔH and qp both are -Ve.
For endothermic reaction (System absorbs energy from the Surroundings).
ΔH and qp both are +Ve.

EXTENSIVE PROPERTY

An extensive property is a property whose value depends on the quantity or size of matter present in the system.
For example: Mass, volume, enthalpy etc. are known as extensive property.

INTENSIVE PROPERTY

Intensive properties do not depend upon the size of the matter or quantity of the matter present in the system.
For example: temperature, density, pressure etc. are called intensive properties.

HEAT CAPACITY

The increase in temperature is proportional to the heat transferred.
q = coeff. x ΔT
q = CΔT, Where, coefficient C is called the heat capacity.
C is directly proportional to the amount of substance.
Cm = C/n
It is the heat capacity for 1 mole of the substance.

MOLAR HEAT CAPACITY

It is defined as the quantity of heat required to raise the temperature of a substance by 1° (kelvin or Celsius).

SPECIFIC HEAT CAPACITY

It is defined as the heat required to raise the temperature of one unit mass of a substance by 1° (kelvin or Celsius).
q = C x m x ΔT where, m = mass of the substance and ΔT = rise in temperature.

RELATION BETWEEN Cp AND Cv FOR AN IDEAL GAS

At constant volume heat capacity = Cv
At constant pressure heat capacity = Cp
At constant volume qv= CvΔT = ΔU
At constant pressure qp = Cp ΔT = ΔH
For one mole of an ideal gas
ΔH = ΔU + Δ (PV) = ΔU + Δ (RT)
ΔH = ΔU + RΔT
On substituting the values of ΔH and Δu, the equation is modified as :Cp ΔT = CvΔT + RΔT
or Cp-Cv = R

MEASUREMENT OF ΔU and ΔH-CALORIMETRY

DETERMINATION OF ΔU:

ΔU is measured in a special type of calorimeter, called bomb calorimeter.

WORKING OF CALORIMETER:

The calorimeter consists of a strong vessel called (bomb) which can withstand very high pressure.
It is surrounded by a water bath to ensure that no heat is lost to the surroundings.

PROCEDURE:

A known mass of the combustible substance is burnt in the pressure of pure dioxygen in the steel bomb.
Heat evolved during the reaction is transferred to the water and its temperature is monitored.

ENTHALPY CHANGES DURING PHASE TRANSFORMATION

ENTHALPY OF FUSION

Enthalpy of fusion is the heat energy or change in enthalpy when one mole of a solid at its melting point is converted into liquid state.

ENTHALPY OF EVAPORATION

It is defined as the heat energy or change in enthalpy when one mole of a liquid at its boiling point changes to gaseous state.

ENTHALPY OF SUBLIMATION

Enthalpy of sublimation is defined as the change in heat energy or change in enthalpy when one mole of solid directly changes into gaseous state at a temperature below its melting point.

STANDARD ENTHALPY OF FORMATION

Enthalpy of formation is defined as the change in enthalpy in the formation of 1 mole of a substance from its constituting elements under standard conditions of temperature at 298K and 1 atm pressure.

ENTHALPY OF COMBUSTION

It is defined as the heat energy or change in enthalpy that accompanies the combustion of 1 mole of a substance in excess of air or oxygen.

THERMOCHEMICAL EQUATION

A balanced chemical equation together with the value of ΔrH and the physical state of reactants and products is known as thermochemical equation.

CONVENTIONS REGARDING THERMOCHEMICAL EQUATIONS

The coefficients in a balanced thermochemical equation refer to the number of moles of reactants and products involved in the reaction.

HESS'S LAW OF CONSTANT HEAT SUMMATION

The total amount of heat evolved or absorbed in a reaction is the same whether the reaction takes place in one step or in number of steps.

SPONTANEITY

SPONTANEOUS PROCESS

A process which can take place by itself or has a tendency to take place is called spontaneous process.
Spontaneous processes need not be instantaneous.
Its actual speed can vary from very slow to quite fast.
A few examples of spontaneous process are:
(i) Common salt dissolves in water of its own.
(ii) Carbon monoxide is oxidised to carbon dioxide of its own.

ENTROPY (S)

The entropy is a measure of the degree of randomness or disorder of a system.
Entropy of a substance is minimum in solid state while it is maximum in gaseous state.
The change in entropy in a spontaneous process is expressed as ΔS

GIBBS ENERGY AND SPONTANEITY

A new thermodynamic function, the Gibbs energy or Gibbs function G, can be defined as G = H-TS
ΔG = ΔH – TΔS
Gibbs energy change = enthalpy change – temperature x entropy
Change ΔG gives a criteria for spontaneity at constant pressure and temperature,
(i) If ΔG is negative (< 0) the process is spontaneous. (ii) If ΔG is positive (> 0) the process is non-spontaneous.

FREE ENERGY CHANGE IN REVERSIBLE REACTION

INTERMOLECULAR FORCES

Intermolecular forces are the forces of attraction and repulsion between interacting particles that have permanent dipole moments.
This interaction is stronger than the London forces but is weaker than ion-ion interaction because only partial charges are involved.
The attractive forces decrease with the increase of distance between dipoles.
The interaction energy is proportional to 1/r6 where r is the distance between polar molecules.

ION-DIPOLE INTERACTION

This is the force of attraction which exists between the ions (cations or anions) and polar molecules.
The ion is attracted towards the oppositely charged end of dipolar molecules.
The strength of attraction depends upon the charge and size of the ion and the dipole moment and the size of the polar molecule. For example: Solubility of common salt (NaCl) in water.

ION-INDUCED DIPOLAR INTERACTIONS

In this type of interaction permanent dipole of the polar molecule induced dipole on the electrically neutral molecule by deforming its electron cloud.
Interaction energy is proportional to 1/r6 where r is the distance between two molecules.

LONDON FORCES OR DISPERSION FORCES

As we know that in non-polar molecules, there is no dipole moment because of their electronics .
Charge cloud is symmetrically distributed.
But, it is believed that at any instant of time, the electron cloud of the molecule may be distorted so that an instantaneous dipole or momentary dipole is produced in which one part of the molecule is slightly more negative than the other part.
This momentary dipole induced dipoles in the neighbouring molecules.
Thus, the force of attraction exists between them and is exactly the same as between permanent dipoles.
This force of attraction is known as London forces or Dispersion forces.
These forces are always attractive and the interaction energy is inversely proportional to the sixth power of the distance between two interacting particles, (i.e. 1/r6 where r is the distance between two particles). This can be shown by fig. given below.

HYDROGEN BONDING

When a hydrogen atom is attached to a highly electronegative element by covalent bond, electrons are shifted towards the more electronegative atom.
Thus a partial positive charge develops on the hydrogen atom. Now, the positively charged hydrogen atom of one molecule may attract the negatively charged atom of some other molecule and the two molecules can be linked together through a weak force of attraction.

THERMAL ENERGY

The energy arising due to molecular motion of the body is known as thermal energy. Since motion of the molecules is directly related to kinetic energy and kinetic energy is directly proportional to the temperature.

THE GASEOUS STATE

Physical Properties of Gaseous State:
(i) gases have no definite volume and they do not have specific shape,
(ii) Gases mix evenly and completely in all proportions without any mechanical aid.
(iii) Their density is much lower than solids and liquids.
(iv) They are highly compressible and exert pressure equally in all directions.

BOYLE'S LAW (PRESSURE-VOLUME RELATIONSHIP)

At constant temperature, the volume of a given mass of gas is inversely proportional to its pressure.

CHARLE'S LAW

At constant pressure, the volume of a given mass of a gas is directly proportional to its absolute temperature.

GAY-LUSSAC'S LAW (PRESSURE-TEMPERATURE RELATIONSHIP)

At constant volume, pressure of a given mass of a gas is directly proportional to the temperature.

AVOGADRO LAW (VOLUME-AMOUNT RELATIONSHIP)

Avogadro’s law states that equal volumes of all gases under the same conditions of temperature and pressure contain equal number of molecules.
V α n, Where n is the number of moles of the gas.
Avogadro constant: The number of molecules in one mole of a gas
= 6.022 x 1023

IDEAL GAS

A gas that strictly follows Boyle’s law, Charles’ law and Avogadro law, is called an ideal gas.
Real gases follow these laws only under certain specific conditions. When forces of interaction are practically negligible.

IDEAL GAS EQUATION

This is the combined gas equation of three laws and is known as the ideal gas equation.

DALTON'S LAW OF PARTIAL PRESSURE

When two or more non-reactive gases are enclosed in a vessel, the total pressure exerted by the gaseous mixture is equal to the sum of the partial pressure of individual gases.
Let P1 ,P2, and P3 be the pressure of three non reactive gases A, B, and C.
When enclosed separately in the same volume and under the same condition.
PTotal = P1+ P2 + P3
Where, PTotal = P is the total pressure exerted by the mixture of gases.

AQUEOUS TENSION

Pressure of non reacting gases are generally collected over water and therefore are moist.
Pressure of dry gas can be calculated by subtracting vapour pressure of water from total pressure of moist gas.
P2Dry gas = PTotal – Aqueous Tension

PARTIAL PRESSURE IN TERMS OF MOLE FRACTION

Let at the temperature T, three gases enclosed in the volume V, exert partial pressure P1 , P2 and P3 respectively, then

KINETIC MOLECULAR THEORY OF GASES

(i) Gases consist of large number of very small identical particles (atoms or molecules),
(ii) Actual volume occupied by the gas molecule is negligible in comparison to empty space between them.
(iii) Gases can occupy all the space available to them. This means they do not have any force of attraction between their particles.
(iv) Particles of a gas are always in constant random motion.
(v) When the particles of a gas are in random motion, pressure is exerted by the gas due to collision of the particles with the walls of the container.
(vi) Collisions of the gas molecules are perfectly elastic. This means there is no loss of energy after collision. There may be only an exchange of energy between colliding molecules.
(vii) At a particular temperature distribution of speed between gaseous particles remains constant.
(viii) Average kinetic energy of the gaseous molecule is directly proportional to the absolute temperature.

DEVIATION FROM IDEAL GAS BEHAVIOUR

REAL GAS

A gas which does not follow ideal gas behaviour under all conditions of temperature and pressure, is called real gas. Deviation with respect to pressure can be studied by plotting pressure Vs volume curve at a given temperature. (Boyle’s law)

COMPRESSIBILITY FACTOR (Z)

Deviation from ideal behaviour can be measured in terms of compressibility factor, Z.

VAN DER WAALS EQUATION

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Where V is a constant for molecular attraction while ‘V is a constant for molecular volume.
(a) There is no force of attraction between the molecules of a gas.
(b) Volume occupied by the gas molecule is negligible in comparison to the total volume of the gas.
Above two assumptions of the kinetic theory of gas was found to be wrong at very high pressure and low temperature.

LIQUIFACTION OF GASES

Liquifaction of gases can be achieved either by lowering the temperature or increasing the pressure of the gas simultaneously.
Thomas Andrews plotted isotherms of C02 at various temperatures shown in figure.

CRITICAL TEMPERATURE (Tc)

It is defined as that temperature above which a gas cannot be liquified however high pressure may be applied on the gas.
Tc = 8a/27bR, (Where a and b are van der Waals constants)

CRITICAL PRESSURE (Pc)

It is the pressure required to Liquify the gas at the critical temperature.
Pc = a/27b2
The volume occupied by one mole of the gas at the critical temperature and the critical pressure is called the critical volume (Vc).
For Example. For C02 to Liquify.
Tc = 30.98°C
Pc = 73,9 atm.
Vc = 95-6 cm3/mole
All the three are collectively called critical constants.

LIQUID STATE

Characteristics of Liquid State:
(i) In liquid, intermolecular forces are strong in comparison to gas.
(ii) They have definite volume but irregular shapes or we can say that they can take the shape of the container.
(iii) Molecules of liquids are held together by attractive intermolecular forces.

VAPOUR STATE

The pressure exerted by the vapour of a liquid, at a particular temperature in a state of dynamic equilibrium, is called the vapour pressure of that liquid at that temperature.
Vapour Pressure depends upon two factors:
(i) Nature of Liquid
(ii) Temperature

SURFACE TENSION

It is defined as the force acting per unit length perpendicular to the line drawn on the surface of liquid.
S.I. unit of Surface Tension = Nm-1
Surface Tension decreases with increase in temperature, because force acting per unit length decreases due to increase in kinetic energy of molecules.

VISCOSITY

It is defined as the internal resistance to flow possessed by a liquid.
The liquids which flow slowly have very high internal resistance, which is due to strong intermolecular forces and hence are said to be more viscous.

When liquid flows, the layer immediately below it tries to retard its flow while the one above tries to accelerate.
Thus, force is required to maintain the flow of layers.

EFFECT OF TEMP, ON VISCOSITY

Viscosity of liquids decreases as the temperature rises because at high temperature, molecules have high kinetic energy and can overcome the intermolecular forces to slip past one another.

BOYLE'S LAW

It states that, under isothermal conditions pressure of a given mass of a gas is inversely proportional to its volume.

CHEMICAL BOND

The force that holds different atoms in a molecule is called a chemical bond.

OCTET RULE

Atoms of different elements take part in chemical combinations in order to complete their octet or to attain the noble gas configuration.

VALENCE ELECTRONS

It is the outermost shell electron which takes part in a chemical combination.

FACTS STATED BY KOSSEL IN RELATION TO CHEMICAL BONDING

In the periodic table, the highly electronegative halogens and the highly electro-positive alkali metals are separated by noble gases.
Formation of an anion and cation by the halogens and alkali metals are formed by gain of electron and loss of electron respectively.
Both the negative and positive ions acquire the noble gas configuration.
The negative and positive ions are stabilized by electrostatic attraction Example,

MODES OF CHEMICAL COMBINATION

By the transfer of electrons: The chemical bond which is formed by the complete transfer of one or more electrons from one atom to another is termed as electrovalent bond or ionic bond.
By sharing of electrons: The bond which is formed by the equal sharing of electrons between one or two atoms is called covalent bond. In these bonds electrons are contributed by both.
Co-ordinate bond: When the electrons are contributed by one atom and shared by both, the bond is formed and it is known as dative bond or co-ordinate bond.

IONIC OR ELECTROVALENT BOND

Ionic or Electrovalent bond is formed by the complete transfer of electrons from one atom to another.
Generally, it is formed between metals and non-metals. We can say that it is the electrostatic force of attraction which holds the oppositely charged ions together.
The compound which is formed by ionic or electrovalent bonds is known as electrovalent compounds.
For Example, ,
(i) NaCl is an electrovalent compound. Formation of NaCl is given below:

Na+ ion has the configuration of Ne while Cl– ion represents the configuration of Ar.
(ii) Formation of magnesium oxide from magnesium and oxygen.

Electrovalency: Electrovalency is the number of electrons lost or gained during the formation of an ionic bond or electrovalent bond.

FACTORS AFFECTING THE FORMATION OF IONIC BOND

(i) Ionization enthalpy: As we know that ionization enthalpy of any element is the amount of energy required to remove an electron from the outermost shell of an isolated gaseous atom to convert it into cation. Hence, lesser the ionization enthalpy, easier will be the formation of a cation and have greater chance to form an ionic bond. Due to this reason alkali metals have more tendency to form an ionic bond. For example, in formation of Na+ ion I.E = 496 kJ/mol. While in case of magnesium, it is 743 kJ/mole. That’s why the formation of a positive ion for sodium is easier than that of magnesium.
Therefore, we can conclude that lower the ionization enthalpy, greater the chances of ionic bond formation.
(ii) Electron gain enthalpy (Electron affinities): It is defined as the energy released when an isolated gaseous atom takes up an electron to form anion. Greater the negative electron gain enthalpy, easier will be the formation of anion. Consequently, the probability of formation of ionic bonds increases.
For example, Halogens possess high electron affinity. So, the formation of anion is very common in halogens.

(iii) Lattice energy or enthalpy: It is defined as the amount of energy required to separate 1 mole of ionic compound into separate oppositely charged ions.

LATTICE ENERGY OF AN IONIC COMPOUND DEPENDS UPON FOLLOWING FACTORS:

(i) Size of the ions: Smaller the size, greater will be the lattice energy.
(ii) Charge on the ions: Greater the magnitude of charge, greater the interionic attraction and hence higher the lattice energy.

GENERAL CHARATERISTICS OF IONIC COMPOUNDS

(i) Physical State: They generally exist as crystalline solids, known as crystal lattice. Ionic compounds do not exist as single molecules like other gaseous molecules e.g., H2 , N2 , 02 , Cl2 etc.
(ii) Melting and boiling points: Since ionic compounds contain high interionic force between them, they generally have high melting and boiling points.
(iii) Solubility: They are soluble in polar solvents such as water but do not dissolve in organic solvents like benzene, CCl4etc.
(iv) Electrical conductivity: In solid state they are poor conductors of electricity but in molten state or when dissolved in water, they conduct electricity.
(v) Ionic reactions: Ionic compounds produce ions in the solution which gives a very fast reaction with oppositely charged ions. For example,

COVALENT BOND-LEWIS LANGMUIR CONCEPT

When the bond is formed between two or more atoms by mutual contribution and sharing of electrons, it is known as covalent bond.
If the combining atoms are the same , the covalent molecule is known as homoatomic.
If they are different, they are known as heteroatomic molecules.
For Example,

LEWIS REPRESENTATION OF SIMPLE MOLECULE (THE LEWIS STRUCTURES)

The Lewis dot Structure can be written through the following steps:
(i) Calculate the total number of valence electrons of the combining atoms.
(ii) Each anion means addition of one electron and each cation means removal of one electron. This gives the total number of electrons to be distributed.
(iii) By knowing the chemical symbols of the combining atoms.
(iv) After placing shared pairs of electrons for a single bond, the remaining electrons may account for either multiple bonds or as lone pairs. It is to be noted that the octet of each atom should be completed.

FORMAL CHARGE

In polyatomic ions, the net charge is the charge on the ion as a whole and not by particular atom. However, charges can be assigned to individual atoms or ions. These are called formal charges. It can be expressed as

LIMITATIONS OF THE OCTET RULE

(i) The incomplete octet of the central atoms: In some covalent compounds the central atom has less than eight electrons, i.e., it has an incomplete octet. For example,

Li, Be and B have 1, 2, and 3 valence electrons only.

(ii) Odd-electron molecules: There are certain molecules which have an odd number of electrons; the octet rule is not applied for all the atoms.

(iii) The expanded Octet: In many compounds there are more than eight valence electrons around the central atom. It is termed as an expanded octet. For Example,

OTHER DRAWBACKS OF OCTET THEORY

(i) Some noble gases also combine with oxygen and fluorine to form a number of compounds like XeF2 , XeOF2 etc.
(ii) This theory does not account for the shape of the molecule.
(iii) It does not give any idea about the energy of The molecule and relative stability.

BOND LENGTH

It is defined as the equilibrium distance between the centres of the nuclei of the two bonded atoms. It is expressed in terms of A. Experimentally, it can be defined by X-ray diffraction or electron diffraction method.

BOND ANGLE

It is defined as -the angle between the lines representing the orbitals containing the bonding – electrons.
It helps us in determining the shape. It can be expressed in degree. Bond angle can be experimentally determined by spectroscopic methods.

BOND ENTHALPY

It is defined as the amount of energy required to break one mole of bonds of a particular type to separate them into gaseous atoms.
Bond Enthalpy is also known as bond dissociation enthalpy or simple bond enthalpy. Unit of bond enthalpy = kJ mol-1
Greater the bond enthalpy, stronger is the bond. For e.g., the H—H bond enthalpy in hydrogen is 435.8 kJ mol-1.
The magnitude of bond enthalpy is also related to bond multiplicity.
Greater the bond multiplicity, more will be the bond enthalpy. For e.g., bond enthalpy of C —C bond is 347 kJ mol-1 while that of C = C bond is 610 kJ mol-1.
In polyatomic molecules, the term mean or average bond enthalpy is used.

BOND ORDER

According to Lewis, in a covalent bond, the bond order is given by the number of bonds between two atoms in a molecule. For example,
Bond order of H2 (H —H) =1
Bond order of 02 (O = O) =2
Bond order of N2 (N = N) =3
Isoelectronic molecules and ions have identical bond orders. For example, F2 and O22- have bond order = 1. N2, CO and NO+ have bond order = 3.
With the increase in bond order, bond enthalpy increases and bond length decreases. For example,

RESONANCE STRUCTURES

There are many molecules whose behaviour cannot be explained by a single-Lewis structure, Tor example, Lewis structure of Ozone represented as follows:

Thus, according to the concept of resonance, whenever a single Lewis structure cannot explain all the properties of the molecule, the molecule is then supposed to have many structures with similar energy.
Positions of nuclei, bonding and nonbonding pairs of electrons are taken as the canonical structure of the hybrid which describes the molecule accurately.
For 03, the two structures shown above are canonical structures and the III structure represents the structure of 03 more accurately. This is also called resonance hybrid.
Some resonating structures of some more molecules and ions are shown as follows:

POLARITY OF BONDS

Polar Covalent bonds
Non-Polar Covalent bonds

NON-POLAR COVALENT BONDS:

When the atoms joined by covalent bonds are the same like; H2, 02, Cl2, the shared pair of electrons is equally attracted by two atoms and thus the shared electron pair is equidistant to both of them.
Alternatively, we can say that it lies exactly in the centre of the bonding atoms. As a result, no poles are developed and the bond is called a non-polar covalent bond. The corresponding molecules are known as non-polar molecules.
For Example,

POLAR BONDS:

When covalent bonds formed between different atoms of different electronegativity, the shared electron pair between two atoms gets displaced towards highly electronegative atoms.
For Example, in HCl molecule, since electronegativity of chlorine is high as compared to hydrogen thus, electron pair is displaced more towards chlorine atom, thus chlorine will acquire a partial negative charge (δ–) and hydrogen atom have a partial positive charge (δ+) with the magnitude of charge same as on chlorination. Such a covalent bond is called polar covalent bond.

DIPOLE MOMENT

Due to polarity, polar molecules are also known as dipole molecules and they possess dipole moment. Dipole moment is defined as the product of magnitude of the positive or negative charge and the distance between the charges.

APPLICATIONS OF DIPOLE MOMENT

(i) For determining the polarity of the molecules.
(ii) In finding the shapes of the molecules. For example, the molecules with zero dipole moment will be linear or symmetrical. Those molecules which have asymmetrical shapes will be either bent or angular.
(e.g., NH3with μ = 1.47 D).
(iii) In calculating the percentage ionic character of polar bonds.

THE VALENCE SHELL ELECTRON PAIR REPULSION (VSEPR) THEORY

Sidgwick and Powell in 1940, proposed a simple theory based on the repulsive character of electron pairs in the valence shell of the atoms. It was further developed by Nyholm and Gillespie (1957).

MAIN POSTULATES ARE THE FOLLOWING:

(i) The exact shape of a molecule depends upon the number of electron pairs (bonded or non bonded) around the central atoms.
(ii) The electron pairs have a tendency to repel each other since they exist around the central atom and the electron clouds are negatively charged.
(iii) Electron pairs try to take such a position which can minimize the rupulsion between them.
(iv) The valence shell is taken as a sphere with the electron pairs placed at maximum distance.
(v) A multiple bond is treated as if it is a single electron pair and the electron pairs which constitute the bond as single pairs.

VALENCE BOND THEORY

Valence bond theory was introduced by Heitler and London (1927) and developed by Pauling and others. It is based on the concept of atomic orbitals and the electronic configuration of the atoms.
Let us consider the formation of hydrogen molecules based on valence-bond theory.
Let two hydrogen atoms A and B having their nuclei NA and NB and electrons present in them are eA and eB .
As these two atoms come closer new attractive and repulsive forces begin to operate.
(i) The nucleus of one atom is attracted towards its own electron and the electron of the other and vice versa.
(ii) Repulsive forces arise between the electrons of two atoms and nuclei of two atoms. Attractive forces tend to bring the two atoms closer whereas repulsive forces tend to push them apart.

ORBITAL OVERLAP CONCEPT

According to the orbital overlap concept, covalent bond formed between atoms results in the overlap of orbitals belonging to the atoms having opposite spins of electrons.
Formation of hydrogen molecule as a result of overlap of the two atomic orbitals of hydrogen atoms is shown in the figures that follows:

Stability of a Molecular orbital depends upon the extent of the overlap of the atomic orbitals.

TYPES OF ORBITAL OVERLAP

Depending upon the type of overlapping, the covalent bonds are of two types, known as sigma (σ ) and pi (π) bonds.

(i) Sigma (σ bond):

Sigma bond is formed by the end to end (head-on) overlap of bonding orbitals along the internuclear axis.
The axial overlap involving these orbitals is of three types:
s-s overlapping: In this case, there is overlap of two half-filled s-orbitals along the internuclear axis as shown below:

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

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

(ii) pi (π bond):

π bond is formed by the atomic orbitals when they overlap in such a way that their axes remain parallel to each other and perpendicular to the internuclear axis.The orbital formed is due to lateral overlapping or side wise overlapping.

STRENGTH OF SIGMA AND PI BONDS

Sigma bond (σ bond) is formed by the axial overlapping of the atomic orbitals while the π-bond is formed by side-wise overlapping. Since axial overlapping is greater as compared to side wise. Thus, the sigma bond is said to be a stronger bond in comparison to a π-bond.

DISTINCTION BETWEEN SIGMA AND PI BOND

HYBRIDIZATION

Hybridisation is the process of intermixing of the orbitals of slightly different energies so as to redistribute their energies resulting in the formation of new sets of orbitals of equivalent energies and shape.

SALIENT FEATURES OF HYBRIDIZATION

(i) Orbitals with almost equal energy take part in the hybridisation.
(ii) Number of hybrid orbitals produced is equal to the number of atomic orbitals mixed,
(iii) Geometry of a covalent molecule can be indicated by the type of hybridisation.
(iv) The hybrid orbitals are more effective in forming stable bonds than the pure atomic orbitals.

CONDITIONS NECESSARY FOR HYBRIDIZATION

(i) Orbitals of valence shells take part in the hybridisation.
(ii) Orbitals involved in hybridisation should have almost equal energy.
(iii) Promotion of electrons is not a necessary condition prior to hybridisation.
(iv) In some cases filled orbitals of valence shells also take part in hybridisation.

TYPES OF HYBRIDIZATION

(i) sp hybridisation: When one s and one p-orbital hybridise to form two equivalent orbitals, the orbital is known as sp hybrid orbital, and the type of hybridisation is called sp hybridisation.
Each of the hybrid orbitals formed has 50% s-characer and 50%, p-character. This type of hybridisation is also known as diagonal hybridisation.

(ii) sp2 hybridisation: In this type, one s and two p-orbitals hybridise to form three equivalent sp2 hybridised orbitals.
All the three hybrid orbitals remain in the same plane making an angle of 120°. Example. A few compounds in which sp2 hybridisation takes place are BF3, BH3, BCl3 carbon compounds containing double bonds etc.

(iii) sp3 hybridisation: In this type, one s and three p-orbitals in the valence shell of an atom get hybridised to form four equivalent hybrid orbitals.
There is 25% s-character and 75% p-character in each sp3 hybrid orbital.
The four sp3 orbitals are directed towards four corners of the tetrahedron.
The structures of NH2 and H20 molecules can also be explained with the help of sp3 hybridisation.

The angle between sp3 hybrid orbitals is 109.5°. A compound in which sp3 hybridisation occurs is, (CH4).

FORMATION OF MOLECULAR ORBITALS

Linear Combination of Atomic Orbitals (LCAO)
The formation of molecular orbitals can be explained by the linear combination of atomic orbitals.
Combination takes place either by addition or by subtraction of wave function as shown below.

The molecular orbital formed by addition of atomic orbitals is called bonding molecular orbital while molecular orbital formed by subtraction of atomic orbitals is called antibonding molecular orbital.
Conditions for the combination of atomic orbitals:
(1) The combining atomic orbitals must have almost equal energy.
(2) The combining atomic orbitals must have the same symmetry about the molecular axis.
(3) The combining atomic orbitals must overlap to the maximum extent.

TYPES OF MOLECULAR ORBITALS

Sigma (σ) Molecular Orbitals: They are symmetrical around the bond-axis.
pi (π) Molecular Orbitals: They are not symmetrical, because of the presence of positive lobes above and negative lobes below the molecular plane.

ELECTRONIC CONFIGURATION AND MOLECULAR BEHAVIOUR

The distribution of electrons among various molecular orbitals is called electronic configuration of the molecule.

BOND ORDER

Bond order is defined as half of the difference between the number of electrons present in bonding and antibonding molecular orbitals.
Bond order (B.O.) = 1/2 [Nb-Na]
The bond order may be a whole number, a fraction or even zero.
It may also be positive or negative.
Nature of the bond: Integral bond order value for single double and triple bond will be 1, 2 and 3 respectively.

BOND LENGTH

Bond order is inversely proportional to bond-length. Thus, greater the bond order, smaller will be the bond-length.
Magnetic Nature: If all the molecular orbitals have paired electrons, the substance is diamagnetic. If one or more molecular orbitals have unpaired electrons, it is paramagnetic e.g., 02 molecule.

BONDING IN SOME HOMONUCLEAR (DIATOMIC) MOLECULES

(1) Hydrogen molecule (H2): It is formed by the combination of two hydrogen atoms. Each hydrogen atom has one electron in Is orbital, so, the electronic configuration of hydrogen molecule is

This indicates that two hydrogen atoms are bonded by a single covalent bond. Bond dissociation energy of hydrogen has been found = 438 kJ/mole. Bond-Length = 74 pm No unpaired electron is present therefore, it is diamagnetic.

(2) Helium molecule (He2): Each helium atom contains 2 electrons, thus in the He2 molecule there would be 4 electrons. The electrons will be accommodated in σ1s and σ*1s molecular orbitals:

HYDROGEN BONDING

When highly electronegative elements like nitrogen, oxygen, flourine are attached to hydrogen to form a covalent bond, the electrons of the covalent bond are shifted towards the more electronegative atom.
Thus, a partial positive charge develops on the hydrogen atom which forms a bond with the other electronegative atom. This bond is known as the hydrogen bond and it is weaker than the covalent bond. For example, in HF molecules, hydrogen bonds exist between the hydrogen atom of one molecule and fluorine atom of another molecule. It can be depicted as

TYPES OF H-BONDS:

Intermolecular hydrogen bond
Intramolecular hydrogen bond.

(i)Intermolecular hydrogen bond:

It is formed between two different molecules of the same or different compounds. For Example, in HF molecules, water molecules etc.

(ii) Intramolecular hydrogen bond:

In this type, hydrogen atom is in between the two highly electronegative F, N, O atoms present within the same molecule.
For example, in o-nitrophenol, the hydrogen is in between the two oxygen atoms.

GENESIS OF PERIODIC CLASSIFICATION

DOBEREINER'S TRIADS

In 1829, Dobereiner arranged certain elements with similar properties in groups of three in such a way that the atomic mass of the middle element was nearly the same as the average atomic masses of the first and the third elements.
A few triads proposed by him are listed.

LIMIATIONS OF DOBEREINER'S TRIADS

The triads given by Dobereiner were helpful in grouping some elements with similar characteristics together, but he could not arrange all the elements known at that time into triads.

NEWLAND'S LAW OF OCTAVES

John Newlands proposed the law of octaves by stating that when elements are arranged in order of increasing atomic masses, every eighth element has properties similar to the first.
Newlands called it the law of octaves because a similar relationship exists in the musical notes also.
This can be illustrated as:

LIMITATION OF NEWLAND'S LAW OF OCTAVES

(i) This classification was successful only up to the element calcium. After that, every eighth element did not possess the same properties as the element lying above it in the same group.
(ii) When noble gas elements were discovered at a later stage, their inclusion in the table disturbed the entire arrangement.

MENDELEEV'S PERIODIC TABLE

Mendeleev’s Periodic Law: The physical and chemical properties of the elements are a periodic function of their atomic masses.
Mendeleev arranged the elements known at that time in order of increasing atomic masses and this arrangement was called the periodic table.
Elements with similar characteristics were present in vertical rows called groups. The horizontal rows were known as periods.

DESCRIPTION OF MENDELEEV'S PERIODIC TABLE

(i) In the periodic table, the elements are arranged in vertical rows called groups and horizontal rows known as periods.
(ii) There are nine groups indicated by Roman Numerals as I, II, III, IV, V, VI, VII, VIII and zero. Group VIII consists of nine elements which are arranged in three triads. The zero group contains elements belonging to inert gases or noble gases and elements present have zero valency.
(iii) There are seven periods (numbered from 1 to 7) or, horizontal rows in the Mendeleev’s periodic table.

IMPORTANCE OF MENDELEEV'S PERIODIC TABLE

(i) This made the study of the elements quite systematic in the sense that if the properties of one element in a particular group are known, those of others can be predicted.
(ii) This helped to a great extent in the discovery of these elements at a later stage.
(iii) Mendeleev corrected the atomic masses of certain elements with the help of their expected positions and properties.

DEFECTS IN MENDELEEV'S PERIODIC TABLE

(i) Hydrogen has been placed in group IA along with alkali metals. But it also resembles halogens of group VII A in many properties.Thus, its position in Mendeleev’s periodic table is controversial.
(ii) Although the elements in Mendeleev’s periodic table have been arranged in order of their atomic masses, in some cases the element with higher atomic mass precedes the element with lower atomic mass.
(iii) We know that the isotopes of an element have different atomic masses but the same atomic number.
Since, periodic table has been framed on the basis of increasing atomic masses of the elements, different positions must have been allotted to all the isotopes of a particular element.
(iv) According to Mendeleev, the elements placed in the same group must resemble their properties. But there is no similarity among the elements in the two sub-groups of a particular group. (v) In some cases, elements with similar properties have been placed in different groups. (vi) Lanthanoids and actinoids were placed in two separate rows at the bottom of the periodic table without assigning a proper reason. (vii) No proper explanation has been offered for the fact that why the elements placed in group show resemblance in their properties.

MENDELEEV'S PERIODIC TABLE

Physical and chemical properties of the elements are the periodic function of their atomic numbers.
Present Form of the Periodic Table (Long form of Periodic Table)
The long form of periodic table, also called Modem Periodic Table, is based on Modern periodic law. In this table, the elements have been arranged in order of increasing atomic numbers.
Nomenclature of Elements with Atomic No. more than 100

STRUCTURAL FEATURES OF THE PERIODIC GROUPS

The long form of the periodic table also consists of the vertical rows called groups.
There are in all 18 groups in the periodic table.
Unlike the Mendeleev periodic table, each group is an independent group.

CHARACTERISTICS OF GROUPS

(i) All the elements present in a group have the same general electronic configuration of the atoms.
(ii) The elements in a group are separated by definite gaps of atomic numbers (2, 8, 8,18, 18,32).
(iii) The atomic sizes of the elements in the group increase down the group due to an increase in the number of shells.
(iv) The physical properties of the elements such as m.p., b.p. density, solubility etc., follow a systematic pattern.
(v) The elements in each group have generally similar chemical properties.

PERIODS

Horizontal rows in a periodic table are known as periods.
There are seven periods in the long form of the periodic table.
Characteristics of periods:
(i) In all the elements present in a period, the electrons are filled in the same valence shell.
(ii) The atomic sizes generally decrease from left to right.

s-BLOCK ELEMENTS

General electronic configuration: ns1-2 Characteristics of s-block elements:
(i) All the elements are soft metals.
(ii) They have low melting and boiling points.
(iii) They are highly reactive.
(iv) Most of them impart colours to the flame.
(v) They generally form ionic compounds.
(vi) They are good conductors of heat and electricity.

p-BLOCK ELEMENTS

General electronic configuration: ns2np1-6
Characteristics of p-block elements:
(i) The compounds of these elements are mostly covalent in nature.
(ii) They show variable oxidation states.
(iii) In moving from left to right in a period, the non-metallic character of the elements increases.
(iv) The reactivity of elements in a group generally decreases downwards.
(v) At the end of each period is a noble gas element with a closed valence shell ns2 np6 configuration.
(vi) Metallic character increases as we go down the group.

d-BLOCK ELEMENTS

General electronic configuration: (n -1) d1-10 ns0-2
The d-block elements are known as transition elements because they have incompletely filled d-orbitals in their ground state or in any of the oxidation states.
(i) They are all metals with high melting and boiling points. (ii) The compounds of the elements are generally paramagnetic in nature. (iii) They mostly form coloured ions, exhibit variable valence (oxidation states). (iv) They are oftenly used as catalysts.

f-BLOCK ELEMENTS

General electronic configuration: (n – 2) f1-14 (n -1) d0-1 ns2
They are known as inner transition elements because in the transition elements of d-block, the electrons are filled in (n – 1) d sub-shell while in the inner transition elements of f-block the filling of electrons takes place in (n – 2) f subshell, which happens to be one inner subshell.
(i) 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). (ii) These two series of elements are called Inner Transition Elements (f-Block Elements). (iii) They are all metals. Within each series, the properties of the elements are quite similar. (iv) Most of the elements pf the actinoid series are radio-active in nature.

METALS

(i) Metals comprise more than 78% of all known elements and appear on the left side of the Periodic Table.
(ii) Metals are solids at room temperature.
(iii) Metal usually has high melting and boiling points.
(iv) They are good conductors of heat and electricity.
(v) They are malleable and ductile.

NON-METALS

(i) Non-metals are located at the top right hand side of the Periodic Table.
(ii) Non-metals are usually solids or gases at low temperature with low melting and boiling points.
(iii) They are poor conductors of heat and electricity.
(iv) The non-metallic character increases as one goes from left to right across the Periodic Table.
(v) Most non-metallic solids are brittle and are neither malleable nor ductile.

METALLOIDS

The elements (e.g., silicon, germanium, arsenic, antimony and tellurium) show the characteristic of both metals and non-metals.
These elements are also called semimetals.

NOBLE GASES

These are the elements present in group 18.
Each period ends with a noble gas element.
All the members are of gaseous nature and because of the presence of all the occupied filled orbitals, they have very little tendency to take part in chemical combinations.
These are also called inert gases.