chemistry higher secondary first year guide and higher secondary second year chemistry practical book | pdf free download
HIGHER SECONDARY - FIRST YEAR
VOLUME - II
REVISED BASED ON THE RECOMMENDATIONS OF THE
TEXT BOOK DEVELOPMENT COMMITTEE
A Publication Under
Government of Tamilnadu
Distribution of Free Textbook Programme
(NOT FOR SALE)
Untouchability is a sin
Untouchability is a crime
Untouchability is inhuman
College Road, Chennai - 600 006Syllabus : Higher Secondary - First Year Chemistry
Unit I - Chemical Calculations
Significant figures - SI units - Dimensions - Writing number in scientific
notation - Conversion of scientific notation to decimal notation - Factor label
method - Calculations using densities and specific gravities - Calculation of formula
weight - Understanding Avogadro’s number - Mole concept-mole fraction of the
solvent and solute - Conversion of grams into moles and moles into grams -
Calculation of empirical formula from quantitative analysis and percentage
composition - Calculation of molecular formula from empirical formula - Laws of
chemical combination and Dalton’s atomic theory - Laws of multiple proportion
and law of reciprocal proportion - Postulates of Dalton’s atomic theory and
limitations - Stoichiometric equations - Balancing chemical equation in its molecular
form - Oxidation reduction-Oxidation number - Balancing Redox equation using
oxidation number - Calculations based on equations. - Mass/Mass relationship -
Methods of expressing concentration of solution - Calculations on principle of
volumetric analysis - Determination of equivalent mass of an element -
Determination of equivalent mass by oxide, chloride and hydrogen displacement
method - Calculation of equivalent mass of an element and compounds -
Determination of molar mass of a volatile solute using Avogadro’s hypothesis.
Unit 2 - General Introduction to Metallurgy
Ores and minerals - Sources from earth, living system and in sea -
Purification of ores-Oxide ores sulphide ores magnetic and non magnetic ores -
Metallurgical process - Roasting-oxidation - Smelting-reduction - Bessemerisation
- Purification of metals-electrolytic and vapour phase refining - Mineral wealth of
Unit 3 - Atomic Structure - I
Brief introduction of history of structure of atom - Defects of Rutherford’s
model and Niels Bohr’s model of an atom - Sommerfeld’s extension of atomic
structure - Electronic configuration and quantum numbers - Orbitals-shapes of s,
(v)p and d orbitals. - Quantum designation of electron - Pauli’s exclusion principle
- Hund’s rule of maximum multiplicity - Aufbau principle - Stability of orbitals -
Classification of elements based on electronic configuration.
Unit 4 - Periodic Classification - I
Brief history of periodic classification - IUPAC periodic table and IUPAC
nomenclature of elements with atomic number greater than 100 - Electronic
configuration and periodic table - Periodicity of properties Anomalous periodic
properties of elements.
Unit 5 - Group-1s Block elements
Isotopes of hydrogen - Nature and application - Ortho and para hydrogen
- Heavy water - Hydrogen peroxide - Liquid hydrogen as a fuel - Alkali metals
- General characteristics - Chemical properties - Basic nature of oxides and
hydroxides - Extraction of lithium and sodium - Properties and uses.
Unit 6 - Group - 2s - Block elements
General characteristics - Magnesium - Compounds of alkaline earth metals.
Unit 7 -p- Block elements
General characteristics of p-block elements - Group-13. Boron Group -
Important ores of Boron - Isolation of Born-Properties - Compounds of Boron-
Borax, Boranes, diboranes, Borazole-preparation. properties - Uses of Boron
and its compounds - Carbon group - Group -14 - Allotropes of carbon -
Structural difference of graphite and diamond - General physical and chemical
properties of oxides, carbides, halides and sulphides of carbon group - Nitrogen
- Group-15 - Fixation of nitrogen - natural and industrial - HNO -Ostwald process
- Uses of nitrogen and its compounds - Oxygen - Group-16 - Importance of
molecular oxygen-cell fuel - Difference between nascent oxygen and molecular
oxygen - Oxides classification, acidic basic, amphoteric, neutral and peroxide -
Ozone preparation, property and structure - Factors affecting ozone layer.
Unit 8 - Solid State - I
Classification of solids-amorphous, crystalline - Unit cell - Miller indices -
Types of lattices belong to cubic system.
Unit 9 - Gaseous State
Four important measurable properties of gases - Gas laws and ideal gas
equation - Calculation of gas constant ‘‘R” - Dalton’s law of partial pressure -
Graham’s law of diffusion - Causes for deviation of real gases from ideal behaviour
- Vanderwaal’s equation of state - Critical phenomena - Joule-Thomson effect
and inversion temperature - Liquefaction of gases - Methods of Liquefaction of
Unit 10 - Chemical Bonding
Elementary theories on chemical bonding - Kossel-Lewis approach - Octet
rule - Types of bonds - Ionic bond - Lattice energy and calculation of lattice
energy using Born-Haber cycle - Properties of electrovalent compounds -
Covalent bond - Lewis structure of Covalent bond - Properties of covalent
compounds - Fajan’s rules - Polarity of Covalent bonds - VSEPR Model -
Covalent bond through valence bond approach - Concept of resonance -
Coordinate covalent bond.
Unit 11 - Colligative Properties
Concept of colligative properties and its scope - Lowering of vapour
pressure - Raoul’s law - Ostwald - Walker method - Depression of freezing
point of dilute solution - Beckmann method - Elevation of boiling point of dilute
solution - Cotrell’s method - Osmotic pressure - Laws of Osmotic pressure -
Berkley-Hartley’s method - Abnormal colligative properties Van’t Hoff factor
and degree of dissociation.
Unit 12 - Thermodynamics - I
Thermodynamics - Scope - Terminology used in thermodynamics -
Thermodynamic properties - nature - Zeroth law of thermodynamics - Internal
energy - Enthalpy - Relation between ‘‘H and “E - Mathematical form of First
law - Enthalpy of transition - Enthalpy of formation - Enthalpy of combustion -
(vii)Enthalpy of neutralisation - Various sources of energy-Non-conventional energy
Unit 13 - Chemical Equilibrium - I
Scope of chemical equilibrium - Reversible and irreversible reactions -
Nature of chemical equilibrium - Equilibrium in physical process - Equilibrium in
chemical process - Law of chemical equilibrium and equilibrium constant -
Homogeneous equilibria - Heterogeneous equilibria.
Unit 14 - Chemical Kinetics - I
Scope - Rate of chemical reactions - Rate law and rate determining step -
Calculation of reaction rate from the rate law - Order and molecularity of the
reactions - Calculation of exponents of a rate law - Classification of rates based
on order of the reactions.
Unit 15 - Basic Concepts of Organic Chemistry
Catenation - Classification of organic compounds - Functional groups -
Nomenclature - Isomerism - Types of organic reactions - Fission of bonds -
Electrophiles and nucleophiles - Carbonium ion Carbanion - Free radicals -
Electron displacement in covalent bond.
Unit 16 - Purification of Organic compounds
Characteristics of organic compounds - Crystallisation - Fractional
Crystallisation - Sublimation - Distillation - Fractional distillation - Steam distillation
Unit 17 - Detection and Estimation of Elements
Detection of carbon and hydrogen - Detection of Nitrogen - Detection of
halogens - Detection of sulphur - Estimation of carbon and hydrogen - Estimation
of Nitrogen - Estimation of sulphur - Estimation of halogens.
Unit 18 - Hydrocarbons
Classification of Hydrocarbons - IUPAC nomenclature - Sources of
alkanes - General methods of preparation of alkanes - Physical properties -
(viii)Chemical properties - Conformations of alkanes - Alkenes - IUPAC nomenclature
of alkenes - General methods of preparation - Physical properties - Chemical
properties - Uses - Alkynes - IUPAC Nomenclature of alkynes - General
methods of preparation - Physical properties - Chemical properties - Uses.
Unit 19 - Aromatic Hydrocarbons
Aromatic Hydrocarbons - IUPAC nomenclature of aromatic hydrocarbons
- Structure of Benzene - Orientation of substituents on the benzene ring -
Commercial preparation of benzene - General methods of preparation of Benzene
and its homologues - Physical properties - Chemical properties - Uses -
Carcinogenic and toxic nature.
Unit 20 - Organic Halogen Compounds
Classification of organic hydrogen compounds - IUPAC nomenclature of
alkyl halides - General methods of preparation - Properties - Nucleophilic
substitution reactions - Elimination reactions - Uses - Aryl halide - General
methods of preparation - Properties - Uses - Aralkyl halides - Comparison
arylhalides and aralkyl halides - Grignard reagents - Preparation - Synthetic uses.
(ix)CHEMISTRY PRACTICALS FOR STD XI
I. Knowledge of using Burette, Pipette and use of logarithms is to be
II. Preparation of Compounds.
1. Copper Sulphate Crystals from amorphous copper sulphate solutions
2. Preparation of Mohr’s Salt
3. Preparation of Aspirin
4. Preparation of Iodoform
5. Preparation of tetrammine copper (II) sulphate
III. Identification of one cation and one anion from the following. (Insoluble
salt should not be given)
++ ++ ++ 2+ ++ ++ ++ ++ +
Cation : Pb , Cu , Al , Mn , Zn , Ca , Ba , Mg , NH .
Anions : Borate, Sulphide, Sulphate, Carbonate, Nitrate, Chloride, Bromide.
IV. Determination of Melting point of a low melting solid.
V. Acidimetry Vs Alkalimetry
1. Preparation of Standard solution of Oxalic acid and Sodium
2. Titration of HCl Vs NaOH
3. Titration of HCl Vs Na CO
4. Titration of Oxalic acid Vs NaOH
10. CHEMICAL BONDING
To know about bonding as binding forces between atoms to form
To learn about Kossel-Lewis approach to chemical bonding, the octet
rule, its limitations and Lewis representations of simple molecules.
To know about ionic bond, lattice energy and Born-Haber cycle.
To understand covalent bond, directional character.
To learn about VSEPR model and predict the geometry of simple
To understand the concepts of hybridisation, 1DQGERQGVUHVRQDQFH
and coordinate covalent bonds.
10.1 Elementary theories on Chemical Bonding
The study on the "nature of forces that hold or bind atoms together to
form a molecule" is required to gain knowledge of the following-
i) to know about how atoms of same element form different
compounds combining with different elements.
ii) to know why particular shapes are adopted by molecules.
iii) to understand the specific properties of molecules or ions and the
relation between the specific type of bonding in the molecules.
Existence of a strong force of binding between two or many atoms is
referred to as a Chemical Bond and it results in the formation of a stable
compound with properties of its own. The bonding is permanent until it is
acted upon by external factors like chemicals, temperature, energy etc. It is
known that, a molecule is made up of two or many atoms having its own
characteristic properties which depend on the types of bonding present.
Classification of molecules
Molecules having two identical atoms like H , O , Cl , N etc. are
2 2 2 2
called as homonuclear diatomic molecules. Molecules containing two
different atoms like CO, HCl, NO, HBr etc., are called as heteronuclear
diatomic molecules. Molecules containing identical but many atoms
bonded together such as P, S etc., are called as homonuclear
polyatomics. In most of the molecules, more than two atoms of different
kinds are bonded such as in molecules like NH , CH COOH, SO , HCHO
3 3 2
and they are called as heteronuclear polyatomics.
Chemical bonds are basically classified into three types consisting of
(i) ionic or electrovalent bond (ii) covalent bond and (iii) coordinate-
covalent bond. Mostly, valence electrons in the outer energy level of an
atom take part in the chemical bonding.
In 1916, W.Kossel and G.N.Lewis, separately developed theories of
chemical bonding inorder to understand why atoms combined to form
molecules. According to the electronic theory of valence, a chemical bond
is said to be formed when atoms interact by losing, gaining or sharing of
valence electrons and in doing so, a stable noble gas electronic
configuration is achieved by the atoms.
Except Helium, each noble gas has a stable valence shell of eight
electrons. The tendency for atoms to have eight electrons in their outershell
by interacting with other atoms through electron sharing or electron-transfer
is known as the octet rule of chemical bonding.
10.1.1 Kossel-Lewis approach to Chemical Bonding
W.Kossel laid down the following postulates to the understanding of
In the periodic table, the highly electronegative halogens and the
highly electropositive alkali metals are separated by the noble gases.
Therefore one or small number of electrons are easily gained and
transferred to attain the stable noble gas configuration.
The formation of a negative ion from a halogen atom and a positive
ion from an alkali metal atom is associated with the gain and loss of
an electron by the respective atoms.
The negative and positive ions so formed attains stable noble gas
electronic configurations. The noble gases (with the exception of
helium which has two electrons in the outermost shell) have filled
outer shell electronic configuration of eight electrons (octet of
electrons) with a general representation ns np .
The negative and positive ions are bonded and stabilised by force of
Kossel's postulates provide the basis for the modern concepts on
electron transfer between atoms which results in ionic or electrovalent
For example, formation of NaCl molecule from sodium and chlorine
atoms can be considered to take place according to Kossel's theory by an
electron transfer as:
──→ ─ ─
(i) Na Na + e
Ne 3s Ne
where Ne = electronic configuration of Neon
= 2s 2p
─ ─→ ─ ─
(ii) Cl + e Cl
Ne3s 3p Ar
Ar = electronic configuration of
+ - + -
(iii) Na+Cl NaCl(or)Na Cl
───→ ─ ─
NaCl is an electrovalent or ionic compound made up of sodium ions
and chloride ions. The bonding in NaCl is termed as electrovalent or ionic
bonding. Sodium atom loses an electron to attain Neon configuration and
also attains a positive charge. Chlorine atom receives the electron to attain
the Argon configuration and also becomes a negatively charged ion. The
coulombic or electrostatic attraction between Na and Cl ions result in
Similarly formation of MgO may be shown to occur by the transfer of
two electrons as:
──→ ─ ─
(i) Mg Mg + 2e
───→ ─ ─
(ii) O + 2e O
2 4 2 6
He2s 2p He2s 2p (or) Ne
2+ 2- 2+ 2-
(iii) Mg +O─ ──→ ─ ─ MgO(or)Mg O
The bonding in MgO is also electrovalent or ionic and the electrostatic
forces of attraction binds Mg ions with O ions. Thus, "the binding forces
existing as a result of electrostatic attraction between the positive and
negative ions", is termed as electrovalent or ionic bond. The
electrovalency is considered as equal to the number of charges on an ion.
Thus magnesium has positive electrovalency of two while chlorine has
negative electrovalency of one.
The valence electron transfer theory could not explain the bonding in
molecules like H , O , Cl etc., and in other organic molecules that have
2 2 2
G.N.Lewis, proposed the octet rule to explain the valence electron
sharing between atoms that resulted in a bonding type with the atoms
attaining noble gas electronic configuration. The statement is : "a bond is
formed between two atoms by mutual sharing of pairs of electrons to attain
a stable outer-octet of electrons for each atom involved in bonding". This
type of valence electron sharing between atoms is termed as covalent
bonding. Generally homonuclear diatomics possess covalent bonds.
It is assumed that the atom consists of a `Kernel' which is made up of a
nucleus plus the inner shell electrons. The Kernel is enveloped by the outer
shells that could accommodate a maximum of eight electrons. The eight
outershell electrons are termed as octet of electrons and represents a stable
electronic configuration. Atoms achieve the stable outer octet when they are
involved in chemical bonding.
In case of molecules like F , Cl , H etc., the bond is formed by the
2 2 2
sharing of a pair of electrons between the atoms. For example, consider the
formation of a fluorine molecule (F). The atom has electronic
2 2 5
configuration. He2s 3s 3p which is having one electron less than the
electronic configuration of Neon. In the fluorine molecule, each atom
contributes one electron to the shared pair of the bond of the F molecule.
In this process, both the fluorine atoms attain the outershell octet of a noble
gas (Argon) (Fig. 10.1(a)). Dots ( ) represent electrons. Such structures are
called as Lewis dot structures.
Lewis dot structures can be written for combining of like or different
atoms following the conditions mentioned below :
Each bond is the result of sharing of an electron pair between the
atoms comprising the bond.
Each combining atom contributes one electron to the shared pair.
The combining atoms attain the outer filled shells of the noble gas
If the two atoms share a pair of electrons, a single bond is said to be
formed and if two pairs of electrons are shared a double bond is said to be
formed etc. All the bonds formed from sharing of electrons are called as
Fig. 10.1(a) F molecule
In carbon dioxide (CO ) two double bonds are seen at the centre carbon
atom which is linked to each oxygen atom by a double bond. The carbon
and the two oxygen atoms attain the Neon electronic configuration.
Fig. 10.1 (b) CO molecule
When the two combining atoms share three electron pairs as in N
molecule, a triple bond is said to be formed. Each of the Nitrogen atom
shares 3 pairs of electrons to attain neon gas electronic configuration.
Fig. 10.1 (c) N molecule
10.2 Types of Bond
There are more than one type of chemical bonding possible between
atoms which makes the molecules to show different characteristic
properties. The different types of chemical bonding that are considered to
exist in molecules are (i) ionic or electrovalent bond which is formed as a
result of complete electron transfer from one atom to the other that
constitutes the bond; (ii) covalent bond which is formed as a result of
mutual electron pair sharing with an electron being contributed by each
atom of the bond and (iii) coordinate - covalent bond which is formed as a
result of electron pair sharing with the pair of electrons being donated by
only one atom of the bond. The formation and properties of these types of
bonds are discussed in detail in the following sections.
10.3 Ionic (or) Electrovalent bond
The electrostatic attraction force existing between the cation and the
anion produced by the electron transfer from one atom to the other is
known as the ionic (or) electrovalent bond. The compounds containing such
a bond are referred to as ionic (or) electrovalent compounds.
Ionic bond is non directional and extends in all directions. Therefore, in
solid state single ionic molecules do not exist as such. Only a network of
cations and anions which are tightly held together by electro-static forces
exist in the ionic solids. To form a stable ionic compound there must be a
net lowering of energy. That is, energy is released as a result of electovalent
bond formation between positive and negative ions.
When the electronegativity difference between the interacting atoms are
greatly different they will form an ionic bond. In fact, a difference of 2 or
more is necessary for the formation of an ionic bond. Na has
electronegativity 0.9 while Cl has 3.0, thus Na and Cl atoms when brought
together will form an ionic bond.
For example, NaCl is formed by the electron ionisation of sodium atom
to Na ion due to its low ionisation potential value and chlorine atom to
chloride ion by capturing the odd electron due to high electron affinity.
Thus, NaCl (ionic compound) is formed. In NaCl, both the atoms possess
──→ ─ ─
i) Na(g) Na + e
2 6 1 2 6
2s2p3s 2s sp
─ ─→ ─ ─
ii) Cl(g) + e Cl
2 5 2 6
3s3p 3s , 3p
iii) Na + Cl ───→ ─ ─ NaCl
Sodium Chloride ionic/crystalline
ion ion compound is formed
Fig. 10.2 Electron transfer between Na and Cl atoms during ionic bond
formation in NaCl
In CaO, which is an ionic compound, the formation of the ionic bond
involves two electron transfers from Ca to O atoms. Thus, doubly charged
positive and negative ions are formed.
──→ ─ ─
Ca Ca +2e (Calcium Cation)
+ O 2e─ ─→ ─ ─ O
2 4 2 6
2s 2p 2s 2p
Ca + O───→ ─ ─ CaO
attraction ionic compound
Ionic bond may be also formed between a doubly charged positive ion
with single negatively charged ion and vice versa. The molecule as a whole
remains electrically neutral. For example in MgF , Mg has two positive
charges and each fluorine atom has a single negative charge. Hence, Mg
binds with two fluoride (F ) ions to form MgF which is electrically neutral.
Mg Mg + 2e
2 6 2 2 6
(2s 2p 3s (2s 2p )
2e + 2F 2F
2 5 2 6
(2s 2p) (2s 2p )
i.e:- Mg + 2F MgF
Magnesium - fluoride
(an ionic compound).
Similarly in Aluminium bromide (AlBr ), Aluminium ion has three
positive charges and therefore it bonds with three Bromide ions to form
AlBr which is a neutral ionic molecule.
Al Al + 3e
6 2 1 2 6
2p 3s 3p (2s , 2p )
3 Br + 3e 3 Br
2 5 2 6
(4s 4p) (4s 4p )
Al + 3Br AlBr (ionic bond)
10.3.1 Lattice energy and Born - Haber's cycle
Ionic compounds in the crystalline state exist as three dimensionally
ordered arrangement of cations and anions which are held together by
columbic interaction energies. The three dimensional network of points that
represents the basic repetitive arrangement of atoms in a crystal is known as
lattice or a space lattice. Thus a qualitative measure of the stability of an
ionic compound is provided by its enthalpy of lattice formation.
Lattice enthalpy of an ionic solid is defined as the energy required to
completely separate one mole of a solid ionic compound into gaseous
constituent ions. That is, the enthalpy change of dissociation of MX ionic
solid into its respective ions at infinity separation is taken the lattice
MX M + X
(s) → (g) (g)
H = L.E
Lattice enthalpy is a positive value.
For example, the lattice enthaply of NaCl is 788 kJ.mol . This means
that 788 kJ of energy is required to separate 1 mole of solid NaCl into 1
mole of Na and 1 mole of Cl (g) to an infinite distance.
In ionic solids, the sum of the electron gain enthaply and the
ionisation enthalpy may be positive but due to the high energy released in
the formation of crystal lattice, the crystal structure gets stabilised.
Born Haber's Cycle
Determination of Lattice enthalpy
It is not possible to calculate the lattice enthalpy directly from the forces
of attraction and repulsion between ions but factors associated with crystal
geometry must also be included. The solid crystal is a three-dimensional
entity. The lattice enthalpy is indirectly determined by the use of Born -
Haber Cycle. The procedure is based on Hess's law, which states that the
enthalpy change of a reaction is the same at constant volume and pressure
whether it takes place in a single or multiple steps long as the initial
reactants and the final products remain the same. Also it is assumed that the
formation of an ionic compound may occur either by direct combination of
elements (or) by a step wise process involving vaporisation of elements,
conversion of gaseous atoms into ions and the combination of the gaseous
ions to form the ionic solid.
For example consider the formation of a simple ionic solid such as an
alkali metal halide MX, the following steps are considered.
M M M
1/2X X X
û+ = enthalpy change for sublimation of M to M
1 (s) (g)
û+ = enthalpy change for dissociation of 1/2 X to X
2 2(g) (g)
û+ = ionization energy of M to M
3 (g) (g)
û+ = electronic affinity or electron gain energy for conversion of X
û+ = the lattice enthalpy for formation of solid MX (1 mole).
û H = enthalpy change for formation of MX solid directly from the
respective elements such as 1 mole of solid M and 0.5 moles of
According to Hess's law,
o o o o o o
û+ = û+ û+ û+ û+ û+
f 1 2 3 4 5
Some important features of lattice enthalpy are:
i. The greater the lattice enthalpy the more stabler the ionic bond
ii. The lattice enthalpy is greater for ions of higher charge and smaller
iii. The lattice enthalpies affect the solubilities of ionic compounds.
Calculation of lattice enthalpy of NaCl
Let us use the Born - Haber cycle for determining the lattice enthalpy of
NaCl as follows :
7KHVWDQGDUGHQWKDOS\FKDQJHû H overall for the reaction,
Na + 1/2 Cl NaCl is - 411.3 kJmol
(s) 2(g) → (s)
Na + ½ Cl NaCl
(s) 2(g) (s)
7KHYDOXHRIû+ calculated using the equation of Born - Haber cycle should be
reversed in sign
10 Sublimation ½ Dissociation
Ionization Energy Electron Affinity
Δ 3 Δ4
Fig. 10.3 Born-Haber cycle for Lattice enthalpy determination
involving various stepwise enthalpic processes for NaCl
Since the reaction is carried out with reactants in elemental forms and
products in their standard states, at 1 bar, the overall enthalpy change of the
reaction is also the enthalpy of formation for NaCl. Also, the formation of
NaCl can be considered in 5 steps.The sum of the enthalpy changes of these
steps is considered equal to the enthalpy change for the overall reaction
from which the lattice enthalpy of NaCl is calculated.
û+ for Na(s) Na(g) is + 108.70 (kJ mol )
û+ for ½ Cl (g) Cl(g) is + 122.0
°2 2 →
û+ for Na(g) Na (g) + e is + 495.0
Electron affinity :
û+ for e + Cl(g) Cl (g) is - 349.0
Lattice enthalpy :
û+ for Na (g) + Cl (g) NaCl(g) is ?
11 û H û+ û+ û+ û+ û+
f ° °1 °2 °3 °4 °5
-411.3 = 108.70 + 122.0 + 495 -
û+ = -788.0 kJ mol
But the lattice enthalpy of NaCl is defined by the reaction
NaCl(g) Na (g) + Cl (g) only.
/DWWLFHHQWKDOS\YDOXHIURPû+ is written with a reversed sign.
Lattice enthalpy of NaCl = +788.0 kJ mol .
Calculation of lattice enthalpy of MgBr from the given data.
The enthalpy of formation of MgBr according to the reaction
Mg(s) + Br (l) MgBr Vû H = -524 kJ/mol
2 → 2 f °
û+ for Mg(s) Mg(g) = + 148 kJ mol
2+ - -1
û+ for Mg(g) Mg (g) +2e = +2187kJ mol
û+ for Br (l) Br (g) = 31 KJ mol
3 2 2
û+ for Br (g) 2Br(g) = 193 KJ mol
- - -1
û+ for Br(g) + e (g) Br = -331 KJ mol
for Mg (g)+2Br (g) Mg Br(s) = ?
°6 → 2
ûH û+ û+ û+ û+ û+ û+
f ° °1 °2 °3 °4 °5 °6
-524 kJ mol = (+148 + 2187 + 31
+ 193 - 2(331) + H ) kJ mol
= -2421 KJ mol= H
Hence, lattice enthalpy of Mg Br û+ = 2421 kJ mol
10.3.2 Properties of electrovalent (or) ionic compounds
Ionic compounds possess characteristic properties of their own like
physical state, solubility, melting point, boiling point and conductivity. The
nature of these properties are discussed as follows.
i. Due to strong coulombic forces of attraction between the oppositely
charged ions, electrovalent compounds exist mostly as hard
12 crystalline solids. Due to the hardness and high lattice enthalpy, low
volatility, high melting and boiling points are seen.
ii. Because of the strong electrostatic forces, the ions in the solid are
not free to move and act as poor conductor of electricity in the solid
state. However, in the molten state, or in solution, due to the
mobility of the ions electrovalent compounds become good
conductor of electricity.
iii. Ionic compounds possess characteristic lattice enthalpies since they
exist only as ions packed in a definite three dimensional manner.
They do not exist as single neutral molecule or ion.
iv. Ionic compounds are considered as polar and are therefore, soluble
in high dielectric constant solvents like water. In solution, due to
solvation of ions by the solvent molecules, the strong interionic
attractions are weakened and exist as separated ions.
v. Electrovalent compounds having the same electronic configuration
10.4 Covalent bond
A covalent bond is a chemical bond formed when two atoms mutually
share a pair of electron. By doing so, the atoms attain stable octet electronic
configuration. In covalent bonding, overlapping of the atomic orbitals
having an electron from each of the two atoms of the bond takes place
resulting in equal sharing of the pair of electrons. Also the interatomic bond
thus formed due to the overlap of atomic orbitals of electrons is known as a
covalent bond. Generally the orbitals of the electrons in the valency shell of
the atoms are used for electron sharing. The shared pair of electrons lie in
the middle of the covalent bond. Including the shared pair of electrons the
atoms of the covalent bond attain the stable octet configuration. Thus in
hydrogen molecule (H ) a covalent bond results by the overlap of the two s
orbitals each containing an electron from each of the two H atoms of the
molecule. Each H atom attains '1s filled K shell.