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INORGANIC CHEMISTRY
DIWAKAR EDUCATION HUB Page 1
INORGANIC CHEMISTRY
CHEPTER – 1 - CHEMICAL PERIODICITY AND
CHEMISTRY OF MAIN GROUP ELEMENTS
CHEPTER – 2 - CHEMICAL BONDING AND STRUCTURE
OF MOLECULES
CHEPTER – 3 - ACID BASE CHEMISTRY
CHEPTER – 4 - s – BLOCK ELEMENTS

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Chart
S. N. TOPIC
CHEMICAL SCIENCE
PAGE
NO.
CHEMICAL PERIODICITY AND CHEMISTRY OF MAIN GROUP ELEMENTS
1. Periodicity- Periodic Properties Like Atomic And Iconic Ionic Radii 4-25
2. Ionization Energy 26-27
3. Electronegativity 27-38
4. Electron Affinity 38-40
5. Atomic Volume 40-44
6. Metallic Character 44-50
7. Melting And Boiling Points 50-61
8. Polarizing Power And Polarizability 61-69
9. Oxidation State 69-70
10. Oxidizing And Reducing Nature 70-74
11. The Inert-Pair Effect 74-78
12. Compounds Like Oxides, Hydrides, Halides, Oxoacids etc. 78-101
13. GENERAL CHARACTERISTICS OF p-BLOCK ELEMENTS 101-106
14. 𝐝𝛑 − 𝐩𝛑 𝐛𝐨𝐧𝐝𝐢𝐧𝐠 106-112
15. Diagonal Relationship 112-113
16. Chemical Periodicity And Chemistry Of Main Group Elements MCQs 113-137
CHEMICAL BONDING AND STRUCTURE OF MOLECULES
17. Structure And Bonding In Homo And Heteronuclear Molecules Bond,
Ionic Bond, Covalent Bond
138-152
18. Valence Bond Theory 152-157
19. LCAO Theory 157-158
20. Octet Theory 158-164
21. Lewis Electron Dot Structures 164-170
22. VSEPR Theory 170-179
23. Resonance 179-185
24. Hybridization 185-204
25. Fluxionality TBP Molecules 204-211

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26. Dipole Moment 211-216
27. Bond Angles, Bond Distance And Bond Energies Including Shapes Of
Molecules
216-219
28. Chemical Bonding And Structure Of Molecules 219-248
ACID BASE CHEMISTRY
29. Concept Of Acids And Bases 249-269
30. Lewis Acid And Base 269-273
31. Hard and Soft Acids and Bases (HSAB) Principle 273-280
32. Relative Strength Of Acids And Bases 280-283
33. Amines 283-293
34. Oxides 293-296
35. Oxoacids 296-298
36. Acid Strength 298-299
37. Inductive effect 299-304
38. Resonance Effects 304-306
39. Factors Affecting Strength Of Hydroacids 306-311
40. Usanovich Acid-Base Concept 311-311
41. Non-Aqueous Solvents 311-312
42. Acid Base Chemistry MCQs 312-337
s-BLOCK ELEMENTS
43. Main Group Elements And Their Compounds Like Oxides, Peroxides,
Superoxides, Carbonates, Hydroxides etc., Diagonal Relationship,
Coordination Chemistry Of s-Block Elements, Metallic Crystal, Bonding
In Metals.
338-396
44. s-Block Elements 396-420

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CHEPTER – 1 - CHEMICAL PERIODICITY
PERIODICITY
In the context of chemistry and the periodic table, periodicity refers to trends or recurring variations
in element properties with increasing atomic number. Periodicity is caused by regular and predictable
variations in element atomic structure.
Mendeleev organized elements according to recurring properties to make a periodic table of elements.
Elements within a group (column) display similar characteristics. The rows in the periodic table (the
periods) reflect the filling of electrons shells around the nucleus, so when a new row begins, the
elements stack on top of each other with similar properties. For example, helium and neon are both
fairly unreactive gases that glow when an electric current is passed through them. Lithium and sodium
both have a +1 oxidation state and are reactive, shiny metals.
Uses of Periodicity
Periodicity was helpful to Mendeleev because it showed him gaps in his periodic table where elements
should be. This helped scientists find new elements because they could be expected to display certain
characteristics based on the location they would take in the periodic table. Now that the elements have
been discovered, scientists and students used periodicity to make predictions about how elements will
behave in chemical reactions and their physical properties. Periodicity helps chemists predict how the
new, superheavy elements might look and behave.
Properties That Display Periodicity
Periodicity can include many different properties, but the key recurring trends are:
 Ionization Energy - This is the energy needed to completely remove an electron from an
atom or ion. Ionization energy increases moving left to right across the table and decreases
moving down a group.
 Electronegativity - A measure of how readily an atom forms a chemical
bond. Electronegativity increases moving left to right across a period and decrease moving
down a group.
 Atomic Radius - This is half the distance between the middle of two atoms just touching
each other. Atomic radius decreases moving left to right across a period and increases
moving down a group. Ionic radius is the distance for ions of the atoms and follows the
same trend. Although it might seem like increasing the number of protons and electrons in
an atom would always increase its size, the atom size doesn't increase until a new electron
shell is added. Atom and ion sizes shrink moving across a period because the increasing
positive charge of the nucleus pulls in the electron shell.

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 Electron Affinity - This is a measure of readily an atom accepts an electron. Electron affinity
increases moving across a period and decreases moving down a group. Nonmetals usually
have higher electron affinities than metals. The noble gases are an exception to the trend
since these elements have filled electron valence shells and electron affinity values
approaching zero. However, the behavior of the noble gases is periodic. In other words,
even though an element group might break a trend, the elements within the group display
periodic properties.
Modern Periodic Law
 Properties of elements are the periodic function to their atomic numbers.
 The periodicity in properties is due to repetition of similar outer shell electronic
configuration at a certain regular intervals.
 In modern periodic table is based on modern periodic law in which elements are arranged in
increasing order of their atomic numbers.
 In the modern periodic table, the elements are arranged in rows and columns. These rows
and columns are known as periods and groups respectively.
 The table consists of 7 periods and 18 groups
 Period indicates the value of ‘n’ (principal quantum number) for the outermost or valence
shell.
 Same number of electrons is present in the outer orbitals (that is, similar valence shell
electronic configuration
Periodic Properties of Elements
The basic law governing modern periodic table states that the properties of elements are periodic
functions of their atomic number. These properties reappear at regular intervals or follow a particular
trend at regular intervals. This phenomenon is known as the periodicity of elements.
The periodic properties of elements occur due to the recurrence of similar electronic configuration that
is having the same number of electrons in the outermost orbit. In a particular group, the number of
valence electrons remains the same. On the other hand, the number of valence electrons increases, as
we move from left to right across a period. The chemical property of an element depends on the
number of electrons in the valence shell.

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Explanation for periodic properties
The periodic properties of an element depend on valency and number of shells in an atom. As we move
down a group the number of shell increases successively such that the number of the shell of an
element is equal to the number of periods to which it belongs. As we move across a period, the number
of shell remains the same. For example, elements of the second period have two shells.
The combining capacity of an atom is known as its valency. It is equal to the number of electrons that an
atom can accept or donate in order to complete its octet. As we move down a group, the number of
electrons in the valence shell remains the same. Hence, the valency of a group is constant. Valency
depends on the number of electrons in the outermost shell of an atom. If the number of electrons is 1,
2, 3, 4 then the respective valences will be 1, 2, 3, 4. If the number of electrons in the outermost shell
will be 5, 6, 7 then the valency will be 8 – 5 = 3, 8 – 6 = 2 and 8 – 7 = 1. Valency is the combining capacity
of an atom hence will always have a positive value and largely affects the periodic properties.
In a period, the number of electrons increases from left to right. As a result, the number of electrons
needed to complete the octet also changes. Hence, the valency successively increases to four in group
14 and then subsequently decreases to 1 in group 17.
IUPAC Nomenclature for Elements with Atomic Number ? 100
Digit Name Abbreviation

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0 nil n
1 un u
2 bi b
3 tri t
4 quad q
5 pent p
6 hex h
7 sept s
8 oct o
9 enn e
Classification of Elements
Elements Valence Shell Electronic
Configuration
Nature Position in Modern
Periodic Table
s-block
elements
ns1-2
( n = 1 to 7). Metals 1 and 2 group elements
p – block
elements
ns2
np1-6
( n = 2 to 7). Metalloids & non metals but
some of them are metals
also.
groups 13 to 18
d-Block
Elements
(n-1)d1-10
ns1-2
(n = 4 to 7). Metals 3 to 12 groups
3d series – Sc(21) to Zn
(30)
4d series – Y (39) to Cd

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(48)
5d series – La (57), Hf
(72) to Hg (80
f-Block
Elements
(n-2)f1-14
(n-1)s2
(n-1)p6
(n-
1)d0-1
ns2
(n = 6 and 7).
Radioactive group 3
4f series – Lanthanides –
14 Elements
Ce (58) to Lu (71)
5f series – Actinides –
14 Elements
Th (90) to Lw (103)
Need for the Periodic Classification of Elements
All existing matter in our surroundings is made up of basic units known as elements. Initially, in 1800,
only 31 chemical elements were discovered. After some advancement in technology in 1865, about 63
more elements were discovered. This created the need for the periodic classification of elements.
Presently, there are 118 elements known to us. Out of these 118 chemical elements, some elements are
man-made.

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CHEPTER – 2 - BONDING AND STRUCTURE OF MOLECULES
CHEMICAL BONDING
Chemical Bonding refers to the formation of a chemical bond between two or more atoms, molecules,
or ions to give rise to a chemical compound. These chemical bonds are what keep the atoms together in
the resulting compound.
The attractive force which holds various constituents (atom, ions, etc.) together and stabilizes them
by the overall loss of energy is known as chemical bonding. Therefore, it can be understood
that chemical compounds are reliant on the strength of the chemical bonds between its constituents;
The stronger the bonding between the constituents, the more stable the resulting compound would be.
The opposite also holds true; if the chemical bonding between the constituents is weak, the resulting
compound would lack stability and would easily undergo another reaction to give a more stable
chemical compound (containing stronger bonds). To find stability, the atoms try to lose their energy.
Whenever matter interacts with another form of matter, a force is exerted on one by the other. When
the forces are attractive in nature, the energy decreases. When the forces are repulsive in nature, the
energy increases. The attractive force that binds two atoms together is known as the chemical bond.
ELECTROVALENT OR IONIC BOND
An ionic bond is formed by the complete transference of one or more electrons from the outer energy
shell (valency shell) of one atom to the outer energy shell of the other atom. In this way, both the atoms
acquire the stable electronic configurations of the nearest noble atom. The atom from which the
electrons are transferred i.e., the atom which loses the electrons, acquires a positive charge and be-
comes cation.
The atom which gains the electrons acquires a negative charge and becomes anion. The electrostatic
attraction between the oppositely charged ions results in the formation of an ionic bond or
electrovalent bond between the two atoms and the compounds are called ionic compounds or
electrovalent compounds, e.g.,
The electropositive elements (like Na, K, Ca, Ba, Sr, Mg, etc.) donate electron(s) and changed into
cations, while the electronegative elements (like oxygen, sulphur and halogens) accept electron(s) and
changed into anions. For example,
Cation formation is favoured by low ionization energy.

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Anion formation is favoured by high electron affinity. The cations and anions are held together by
electrostatic attraction forces or coulombic forces.
Combination of cations and anions is an exergonic process. The energy released in the reaction of 1 gm.
mole of a crystal from the gaseous ions is called lattice energy of that crystal. Higher the lattice energy
of a crystal, the greater is the ease of its formation.
An ionic bond is formed only if sum of electron affinity of the anion-forming element and the lattice
energy exceeds the ionization energy of the cation-forming element. The number of electron(s) lost or
gained by the atom to>form ionic compound is known as valency of the respective atom in that
compound. For example, in the above example valency of sodium is + 1 while that of chlorine is -1.
There are certain factors that favour the formation of ionic bond as follows:
1. One of the atoms (the metal) must have a low ionization energy so that it easily loses its
electron(s).
2. The other atom (non-metal) must have the capacity to hold the extra electrons, i.e., it must
have high electron affinity.
3. One of the atoms (metal) should be larger in size, while the other (non-metal) should be
smaller in size.
4. The combining elements should differ by at least 1.9 in electronegativity.
5. The electrostatic attraction between charged ions in the crystal, i.e., lattice energy, should
be high.
6. The cation and anion should have inert gas electronic configuration.
In general, the metals (elements on the left of the periodic table) have low ionization energies and non-
metals (elements on the right of the periodic table) have high electron affinity. Ionic bonds are favoured
between these elements.
The ionic compounds are made up of ions and are highly soluble in solvents with high dielectric
constant, such as water (dielectric constant = 80) and other polar solvents, but insoluble in non-polar
solvents like benzene, ether, etc. Ionic compounds dissociate into water. The ions have a tendency to
become hydrated and release hydration energy.

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Normally hydration energy is sufficient to overcome the lattice energy and thus the ions are separated
from each other forming a solution. Some ionic compounds, on the other hand, are sparingly soluble
(almost insoluble) as the lattice energy is higher than the hydration energy, e.g., BaSO4, PbSO4, AgCl,
AgBr, Agl, Ag2CrO4, etc.
Ionic compounds have high melting and boiling points due to strong electrostatic attraction between
ions of their crystals. They conduct electricity in solution as well as in molten form (fused state).
In solution, ionic compounds show ionic reactions, which are quite fast and instantaneous. In crystals of
the ionic compounds the constituent units are ions and not molecules, arranged in a regular pattern to
form the crystal lattice.
Major Factors Influencing the Formation of Ionic Bond
1. Ionization energy
It is defined as the amount of energy required to remove the most loosely bound electron from an
isolated gaseous atom of an element. The lesser the ionization energy, the greater is the ease of the
formation of a cation.
2. Electron affinity
It is defined as the amount of energy released when an electron is added to an isolated gaseous atom of
an element. The higher the energy released during this process, the easier will be the formation of an
anion.
Thus, low ionization energy of a metal atom and high electron affinity of a non-metal atom facilitate the
formation of an ionic bond between them.
3. Lattice energy
It is defined as the amount of energy released when cations and anions are brought from infinity to their
respective equilibrium sites in the crystal lattice to form one mole of the ionic compound. The higher the
lattice energy, the greater is the tendency of the formation of an ionic bond. The higher the charges on
the ions and smaller the distance between them, the greater is the force of attraction between them.
COVALENT BOND
Atoms may combine with one another by sharing of electrons in their valency shells so that the
combining atoms attain the nearest noble gas configurations. The shared electrons contribute towards
the stability of both the atoms. This type of linkage is called covalent linkage or covalent bond, and the
compounds are called covalent compounds.

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Covalent bond is formed between two atoms of comparable electronegativity. In the formation of a
covalent bond, equal number of electron(s) is (are) shared between the two concerned atoms. These
shared electrons become the common property of both atoms. Sharing of electrons may occur in three
ways.
(i) When each atom contributes one electron and hence the contributed pair of electrons is shared by
each atom. The bond formed is known as single bond and is represented by ‘-’.
Other common examples are NH3, BF3, alkanes, halogen acids, etc.
(ii) In some cases each atom contributes two electrons and thus two pairs of electrons are shared by
each atom. The bond is known as double bond and represented by ‘=’. Common examples are CO2, O2,
olefins, etc.
(iii) Similarly, triple bond (=) is formed by the sharing of three pairs of electrons (three electrons are
contributed by each atom). Common examples are N2, alkynes, etc.
Compounds having covalent bond(s) is (are) known as covalent compounds.
For example, when two hydrogen atoms approach each other each atom contributes one electron and
the pair of electrons is shared by both the atoms to form a molecule of hydrogen.
H + H → H: H or H – H
Similarly, hydrogen and chlorine atoms share one electron each to form HCI
H + CI → H: CI: or H – CI

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The bonding atoms may share more than one pair of electrons depending upon their requirement of
completing the octet. For example, in the formation of oxygen molecule, each oxygen atom has six
electrons in the valence shall, and therefore, they contribute two electrons each for sharing. Thus two
electrons pairs are shared and there is a double bond between the two oxygen atoms.
Similarly, in the formation of a nitrogen molecule, three electron pairs are shared and there is a triple
bond between the two nitrogen atoms The number of electrons which an atom contributes for sharing
in a covalent bond is called the covalency. Thus, covalency of hydrogen, chlorine, oxygen and nitrogen is
1, 1, 2 and 3, respectively.
Elements having very high ionization energies are incapable of transferring electrons and elements
having very low electron affinity cannot take up electrons. The atoms of such elements tend to share
their electrons with the atoms of other elements or with other atoms of the same element in a way that
both the atoms obtain octet configuration in their respective valence shell and thus achieve stability.
Such association through sharing of electron pairs among different or same kinds is known as Covalent
Bond.
Formation of Covalent Bond
Some Important Characteristics of Covalent Bond
1. Bond length
It is defined as the average distance between the nuclei of two bonded atoms in a molecule. In the
formation of hydrogen molecule when two hydrogen atoms approach each other, a stage is reached
where the attractive forces balance the repulsive forces.
At this point, the potential energy of the system becomes minimum and the atoms get bonded together.
The distance between the nuclei of two hydrogen atoms is called bo7id length of H-H bond and has been
found to be 0.74 A°.

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It may be noted that bond length decreases with multiplicity of the bond between the two atoms. Thus,
C = C bond is shorter than C = C bond which, in turn, is shorter than C-C bond.
2. Bond angle
Bond angle may be defined as the internal angle between the orbitals containing electron pairs in the
valency shell of the central atom in a covalent molecule. For example, the bond angles in H2O, NH3 and
CH4 molecules are 104.5°, 107° and 109.5°, respectively.
The bond angles give an idea of the distribution of the orbitals in three-dimensional space around the
central atom in the molecule and thus give an idea of the shape of the molecule.
3. Bond strength or bond energy
Energy is invariably required to break a chemical bond. For instance, in the breaking of 1 mole of
hydrogen gas into atoms, 458 kJ of energy is required. The bond strength in this case is said to be 458 kJ
per mole, i.e., per Avogadro’s number of bonds.
Bond strength or bond energy of a particular type of bond is defined as the energy required to break
one mole of bonds (i.e., Avogadro’s number of bonds) of that type in a substance in gaseous state.
The strength of the bond indicates the stability of the bond. Thus, N = N bond is more stable than O = O
bond. Hence nitrogen molecule is more stable than oxygen molecule. Consequently, nitrogen is much
less reactive than oxygen. The strength of F-F bond is lower than that of CI – CI bond. Hence, fluorine is
more reactive than chlorine.
Covalent Bonding can be Achieved in two Ways:
 Sharing of electrons between atoms of the same kind E.g. Formation of H2, Cl2, O2, etc.
 Sharing of electrons between atoms of different kind E.g. Formation of CH4, H2O, NH3, etc.
Covalent Bonding in Carbon Atom
As per the electronic configuration of Carbon, it needs to gain or lose 4 electrons to become stable,
which seems impossible as:
 Carbon cannot gain 4 electrons to become C4-, because it will be tough for 6 protons to hold
10 electrons and so the atom will become unstable.
 Carbon cannot lose 4 electrons to become C4+ because it would require a large amount of
energy to remove out 4 electrons and also the C4+ would have only 2 electrons held by
proton, which will again become unstable
Carbon cannot gain or donate electrons, so to complete its nearest noble gas configuration, it shares
electron to form a covalent bond.

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Properties of Covalent Bond
If the normal valence of an atom is not satisfied by sharing a single electron pair between atoms, the
atoms may share more than one electron pair between them. Some of the properties of covalent bonds
are:
 Covalent bonding does not result in the formation of new electrons. The bond only pairs
them.
 They are very powerful chemical bonds that exist between atoms.
 A covalent bond normally contains the energy of about ~80 kilocalories per mole (kcal/mol).
 Covalent bonds rarely break spontaneously after it is formed.
 Covalent bonds are directional where the atoms that are bonded showcase specific
orientations relative to one another.
 Most compounds having covalent bonds exhibit relatively low melting points and boiling
points.
 Compounds with covalent bonds usually have lower enthalpies of vaporization and fusion.
 Compounds formed by covalent bonding don’t conduct electricity due to the lack of free
electrons.
 Covalent compounds are not soluble in water.
Octet Rule
All atoms except noble gases have less than eight electrons in their valence shell. In other words, the
valence shells of these atoms do not have stable configurations. Therefore, they combine with each
other or with other atoms to attain stable electronic configurations.
Therefore,
“The tendency of atoms of various elements to attain stable configuration of eight electrons in their
valence shells is the cause of Chemical combination”
and
“The principle of attaining the maximum of eight electrons in the valence shell of atoms is called octet
rule.”
Lewis introduced simple symbols to denote the electrons present in the outer shell of atom known as
the valence electrons. These symbols are known as Electron Dot Symbols and the structure of the
compound is known as Lewis Dot Structure.

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Dot structure of methane
Conditions for writing the Lewis dot structures
 Sharing of an electron pair between the atoms results in the formation of covalent bonds.
 During bond formation, each bond consists of two electrons which are contributed by each
one of the combining atoms.
 By the mutual sharing of electrons, each atom attains octet configuration in its valence shell.
Electron dot structures of covalent molecules are written with respect to the octet rule. According to
this rule, all the atoms in the molecule will have eight electrons in their valence shell except the
Hydrogen atom. Hydrogen will have only two electrons because only two electrons complete its first
shell to attain helium configuration.
Thus the elements of group 17 such as Cl would share one electron to attain stable octet; the elements
of group 16 such as O and S would share two electrons; the elements of group 15 would share three
electrons and so on.
For Example, the oxygen atom which has six electrons in its valence shell completes its octet by sharing
its two electrons with two hydrogen atoms to form a water molecule.
Lewis Structure of Water Molecule
Types of Covalent Bonds
Depending upon the number of shared electron pairs, the covalent bond can be classified into:
 Single Covalent Bond
 Double Covalent Bond
 Triple Covalent Bond
Single Bonds
A single bond is formed when only one pair of the electron is shared between the two participating
atoms. It is represented by one dash (-). Although this form of covalent bond has a smaller density and is
weaker than a double and triple bond, it is the most stable.
For Example, HCL molecule has one Hydrogen atom with one valence electron and one Chlorine atom
with seven valence electrons. In this case, a single bond is formed between hydrogen and chlorine by
sharing one electron.

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CHEPTER – 3 - ACID BASE CHEMISTRY
The acid definition is given as any hydrogen that contains a substance capable of donating a proton (a
hydrogen ion) to the other substance. A base is an ion or molecule that is able to accept a hydrogen ion
from an acid.
Usually, the acidic substances are identified with their sour taste. Basically, an acid is a molecule that can
donate an H+
ion and also can remain energetically favorable after a loss of H+
ion. Acids are much
known to turn blue litmus into the red.
On the other side, bases are characterized by a slippery texture and a bitter taste. A base that is
dissolved in water is known as an alkali. When these substances react chemically with acids, they further
yield salts. Besides, the bases are much known to turn red litmus into blue.
Theories of Acids and Bases
There are 3 different theories that have been put forth to define these acids and bases. These 3 theories
include:
1. Arrhenius Theory
2. Bronsted-Lowry Theory
3. Lewis Theory of Acids and Bases.
A brief description of these theories is provided below. As discussed, acids and bases are defined via
three different theories.
1. Coming to the Arrhenius theory of acids and bases, it is stated that “an acid generates the H+
ions in a solution, and a base produces an OH–
ion in its solution.”
2. The theory of Bronsted-Lowry explains “an acid as a proton donor; a base as a proton
acceptor.”
3. And finally, the Lewis theory of acids and bases says “acids as electron-pair acceptors and
the bases as electron-pair donors.”
pH of Acids and Bases
To find the numeric value of the acidity or basicity level of a substance, the pH scale (pH stands for
‘potential of hydrogen’) can be used. Here, the pH scale is the most common and trusted procedure to
measure how acidic or basic a substance is. Also, a pH scale measure can differ from 0 to 14, where 14 is
the most basic, and 0 is the most acidic a substance can be.
The other way to check if a substance is acidic or basic is by using a litmus paper. There exist two types
of litmus paper available, used to identify the acids and bases. They are the red litmus paper and the
blue litmus paper. The blue litmus paper changes red under acidic conditions, whereas the red litmus
paper turns blue under alkaline or basic conditions.

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Properties of Acids and Bases
A molecule capable of creating a covalent bond with an electron pair is an acid. Acids are very commonly
found in some of the foods we consume like, for example, citrus fruits such as oranges and lemons,
which contain citric acid; vinegar contains acetic acid; in fact, our stomach uses hydrochloric acid for
digestion. Acids have a sour taste. It reacts with metals to form H2 and reacts with carbonates to form a
salt, carbon dioxide and water. Acids turn blue litmus paper red. Acid's strength can be measured on the
pH. Acids are sticky. Acids frequently cause nose burning.
Properties of Acids and Bases
Here, we will learn about the basic properties of acids and bases later, we will learn about the chemical
properties of acids and bases and also physical properties of acids and bases. Acids are a different group
of compounds because of the properties of their aqueous solutions. Properties of acids are as follows:
1. Acids can conduct electric currents because of the electrolyte nature, and some acids are
strong electrolytes because they can completely ionize in water producing many H+ ions.
2. Acids are sour. Oranges, lemons, and vinegar are few examples.
3. Acid changes the color of a few acid-base indicators. Two common indicators of acids are
litmus and phenolphthalein. Acid turns blue litmus into red, and phenolphthalein turns
colorless.
4. Acids react with metals to produce hydrogen gas.
Zn(s) + H2SO4(aq) → ZnSo4(aq) + H2(g)
Properties of Bases are:
1. Bases can be strong or weak. Base's aqueous solutions are also electrolytes.
2. Bases are often bitter. Soaps are less common as foods but are present in many household
products. Soaps are an example of the base, which is slippery.
3. Bases change the color of indicators too. Litmus paper turns blue, and phenolphthalein
turns pink.
4. Bases do not react with metals like acids.
How do Acids and Bases React with Metals?
Acids react with most metals to produce salt and hydrogen gas. As we know, metals that are more active
than acids can undergo a single displacement reaction.
Zinc metal reacts with hydrochloric acid and produces zinc chloride and hydrogen gas.
Zn(s) + 2HCl(aq) → ZnCl2(aq) + H2(g)
Bases too react with few metals like zinc or aluminum to produce hydrogen gas.

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Sodium hydroxide reacts with zinc and water to produce sodium zincate and hydrogen gas.
Zn(s) + 2NaOH(aq) + 2H2O(l) → Na2Zn(OH)4(aq) + H2(g)
Chemical Properties of Acids and Bases
Like we learned about the basic properties of acids and bases, now we will learn about the chemical
properties of acids and bases. We will learn what happens when bases meet metals and acids meet
metals.
Chemical Properties of Acids:
1. Acids react with Reactive Metals Like the Followings:
Acid + metal → salt + hydrogen
Copper and silver do not react with dilute acid.
Like:
2HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
2. Acids React with Bases
Acid + Base -> salt + water
6HNO3(aq) + Fe2O3(s) → 2Fe(NO3)3(aq) + 3H2O(l)
This is the reaction of acids and bases with metals.
3. Acids React with Carbonates
Acid + carbonate => any salt + water + carbon dioxide
H2SO4(aq) + CuCO3(s) →CuSO4(aq) + H2O(l) + CO2(g)
So these are the chemical properties of acids and bases.
Physical Properties of Acids
So, here are the physical properties of acids :
1. Acids are sour.
2. Acids are water-soluble.
3. Solutions of acids can turn blue litmus paper to red.
4. Acid solutions have pH values lesser than 7.

INORGANIC CHEMISTRY
5. React with metal carbonates to produce salt carbon dioxide and water.
Neutral Substances
These are the substances, which have no properties of either acid or base, which has a similar amount of
hydroxyl ions and hydrogen ions, and they do not modify the color of the litmus surface.
 Neutral substances do not display any acidic or basic characteristics.
 Their pH values approximately 7.
 Neutral substances have no effect on blue or red litmus paper.
 pH of pure water is exactly 7.
 Examples are Common salt (NaCl), Water, and more.
Uses of Acids and Bases
Various uses of acids and bases can be listed as follows:
Uses of Acids
 Vinegar, which is a diluted solution of acetic acid, has different household applications. It is
used primarily as a food preservative.
 Citric acid is an integral part of orange juice and lemon juice. It is also used as a food
preservative.
 Sulfuric acid is more widely used in batteries. Commonly, the batteries used to start
automobile engines contain this acid.
 The industrial production of dyes, explosives, paints, and fertilizers involves the use of nitric
acid and sulfuric acid.
 Phosphoric acid is a primary ingredient in various soft drinks.
Uses of Bases
 The manufacturing of paper and soap involves the use of sodium hydroxide. Also, NaOH is
used in the manufacturing of rayon.
 Ca(OH)2, which is also called calcium hydroxide or slaked lime, is used to manufacture the
bleaching powder.
 Dry mixes used in decoration or painting are made using a limited amount of calcium
hydroxide.
 Magnesium hydroxide, also called the milk of magnesia, is most commonly used as a
laxative. It also reduces if there is any excess acidity in the human stomach and is, thus, used
as an antacid substance.
 Ammonium hydroxide is an important reagent that is used in laboratories.
 Any excess acidity in soils is neutralized by employing the slaked lime.

INORGANIC CHEMISTRY
Experiments on Properties of Acids
Properties of acids and bases are studied with the help of their reactions with
1. Litmus solution
2. Zinc and
3. Solid sodium carbonate.
Materials Required For All Experiments
1. Test tubes,
2. Dropper
3. Test tube stand and holder
4. Matchbox
5. Burner
6. Flat bottom flask
7. Beaker
8. Litmus paper/ solution (red & blue)
9. Thistle funnel
10. Glass rods
11. Zinc granules
12. Freshly made lime water
13. Dilute HCl
14. Dilute NaOH
15. Solid sodium carbonate
Acids
 Acids are compounds which when dissolved in the water yield hydronium ions as the only
positively charged ions.
 Strength of acid is decided by the basicity of the acid (basicity is the number of hydrogen
ions present in the solution when acid is dissolved in water).
 Eg. Nitric acid gives hydrogen ion and nitric ion and hydrogen ion reacts with water ion
which gives hydronium ion.
An Experiment of Litmus Test
1. Take two test tubes and label them as A and B
2. Take 10 ml of blue litmus in test tube A and 10 ml of red litmus solution in test tube B
3. Take the dropper filled with nitric acid and add a few drops of it in the test tube A and test tube
B
4. Leave the solution for a few minutes

INORGANIC CHEMISTRY
5. It is observed that the blue colour of litmus in the test tube A changes its colour to red and is
observed that the red colour of litmus solution in test tube B remains the same.
6. We conclude that acid turns the colour of the litmus solution from blue to red
Experiment with Zinc
1. Take a dry and clean test tube
2. Place zinc granules in it
3. Fill it with liquid HCl submerging zinc granules and tilt it a little bit and fix it with properly in the
cork
4. And place the bunsen burner and lit it
5. Close the mouth of the test tube with cork so that no vapour can escape
6. Leave the solution for 2-3 minutes it burns with a robust sound releasing a colourless and
odourless gas
7. And if a burning match stick is brought near the mouth of the test tube it gives a pale blue flame
with a pop sound.
The reaction of the given experiment:
Zn + 2HCl → ZnCl2+H2
Experiment with Solid Sodium Carbonate
1. Take one gram of solid sodium carbonate with some distilled water in a flat bottomed flask
2. Take a dry double bore fork and with thistle funnel which has a delivery tube fitted with it
3. Close the mouth of the flat flask with the dry double bore fork
4. Add 2 ml of hydrochloric gas
5. Colourless and odourless gas is liberated which is passed through lime water through the
delivery tube
6. And it turns lime water milky
Na2CO3 + 2HCl →2NaCl+H20(l)+CO2(g)
Experiments on properties of Bases
Bases
 Bases are the compounds which react with the acid to give salt and water (neutralisation).
 Alkalis are bases which can be dissolved in water and yield hydroxyl ions as the only
negatively charged ions. eg calcium hydroxide gives calcium ion (positively charged) and
hydroxyl ion (negatively charged) thus when alkalis dissolved in water gives hydroxyl ion
 The strength of the base is given by the acidity of bases.
Note: all alkalis are bases but not all bases are alkalis

INORGANIC CHEMISTRY
 When sodium hydroxide reacts with hydrochloric acid it gives sodium chloride (salt) and water
An Experiment of Litmus Test
1. Take two test tubes and label them as A and B
2. Fill test tube A with blue litmus solution and fill test tube A with red litmus solution (each 10 ml).
3. Use the dropper to add a few drops of calcium hydroxide in the solution in the test tube
4. Leave the solution for a few minutes
5. Observation: The red litmus solution in test tube A turns blue and the blue litmus solution in test
tube B doesn't change its colour.
6. Interference: we conclude that calcium hydroxide is basic in nature and bases turn from red
litmus solution to blue litmus solution.
Experiment with Zinc
1. Take a dry and clean test tube
2. Place zinc granules in it
3. Fill it with liquid NaOH submerging zinc granules and tilt it a little bit and fix it with properly in
the cork
4. And place the bunsen burner and lit it
5. Close the mouth of the test tube with cork so that no vapour can escape
6. Leave the solution for 2-3 minutes it burns with a robust sound releasing a colourless and
odourless gas
7. And if a burning match stick is brought near the mouth of the test tube it gives a pale blue flame
with a pop sound.
The reaction of the given experiment:
2NaOH+Zn→Na2ZnO2+H2
Experiment with Solid Sodium Carbonate
1. Take one gram of Na2CO3 with some distilled water in a flat bottomed flask
2. Take a dry double bore fork and with thistle funnel which has a delivery tube fitted with it
3. Add dilute NaOH.
4. No reaction will take place.
Conclusions of Experiments
Acid Base
HCl acid turns blue litmus red but does not
affect red litmus
NaOH turns red litmus blue but does not affect blue
litmus

INORGANIC CHEMISTRY
Acids react with zinc and liberate hydrogen gas
and zinc chloride
Bases react with zinc and liberate hydrogen gas and
sodium zincate
Acids react with sodium carbonate and release
carbon dioxide.
Base does not react with sodium carbonate
IONIZATION OF ACIDS AND BASES
The ionization of a compound can be explained as a process where a neutral molecule splits into
charged ions when exposed to a solution.
Arrhenius theory says that acids are the compounds that are dissociated in an aqueous medium to
generate hydrogen ions, H+
(aq). On the other side, bases are the compounds that furnish hydroxyl ions,
OH−
(aq) in an aqueous solution or medium.
The difference between ionisation and dissociation.
The primary difference between dissociation and ionization or the is, dissociation is the process of
separating the charged particles which already exist in the compound, on the other side, ionization is
the formation of new charged particles, which are not present in the previous compound.
Arrhenius Theory
Arrhenius theory plays a major role in explaining the ionization of acids and bases because mostly
ionization occurs in an aqueous medium. Based on the degree of ionization of acids and bases, we can
define the strength of both acids and bases. Also, the degree of ionization differs for different
compounds of acidic and basic. A few acids, such as hydrochloric acid (HCl), perchloric acid (HClO4),
completely dissociate into their constituent ions in an aqueous medium.
All these acids are referred to as strong acids. Ionization of acids produces hydrogen ions, and therefore,
these compounds act as proton donors. In the same way, a few bases such as sodium hydroxide (NaOH),
lithium hydroxide (LiOH) too dissociate completely into their ions in an aqueous solution or medium.
These bases are referred to as strong bases. The ionization of these bases produces hydroxyl ions (OH−
).
Therefore, the ionization degree of acids and bases depends based on the degree of dissociation of
compounds into their constituent ions. The strong acids and bases have a high degree of ionization
when compared to the ionisation of weak acid and base. Also, a strong acid implies a good proton
donor, whereas a strong base implies a good proton acceptor - for example, dissociation of weak acid
HA.
HA(aq) + H2O(l) ⇌ H3O + (aq) + A-
(aq)
Explanation

INORGANIC CHEMISTRY
CHEPTER – 4 - s – BLOCK ELEMENTS
The elements of Group 1 and Group 2 of the modern periodic table are called S block elements. The two
types of s block elements are possible i.e. the elements with one electron (s1) or the elements with two
electrons (s2) in their s-subshell.
S block comprises 14 elements: hydrogen (H), lithium (Li), helium (He), sodium (Na), beryllium (Be),
potassium (K), magnesium (Mg), rubidium (Rb), calcium (Ca), cesium (Cs), strontium (Sr), francium (Fr),
barium (Ba), and radium (Ra).
What are S Block Elements?
The s block elements having only one electron in their s-orbital are called group one or alkali
metals whereas the s block elements having two electrons filling their s-orbital are called group two
or alkaline earth metals.
The electrons present in an atom occupy various sub-orbitals of available energy levels in the order of
increasing energy. The last electron of an atom may find itself in either of the s, p, d and f subshells.
Accordingly, the elements of the atom having their last valence electron present in the s-suborbital are
called the s block elements.
Electronic Configuration of S Block Elements
The electronic configuration of S block elements is explained below,
The alkali elements in s block consist of a single valence electron in their outermost shell. This outermost
electron is loosely held which makes these metals highly electropositive. Due to which they are not
available in the free state in nature. The general electronic configurations of s block elements – group 1
are as shown in the table below:

INORGANIC CHEMISTRY
Element Symbol Electronic configuration
Lithium Li 1s2
2s1
Sodium Na 1s2
2s2
2p6
3s1
Potassium K 1s2
2s2
2p6
3s2
3p6
4s1
Rubidium Rb 1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
5s1
Cesium Cs [Xe]6s1
Francium Fr [Rn]7s1
The electronic configurations of elements included in group 2 of S block elements are shown below:
Elements Symbols Electronic configuration
Beryllium Be [He]2s2
Magnesium Mg [Ne]3s2
Calcium Ca [Ar]4s2
Strontium Sr [Kr]5s2
Barium Ba [Xe]6s2
Radium Ra [Rn]7s2
Properties of S Block Elements
Both alkali and alkaline earth elements show a regular gradation in their properties among their
respective group elements. But the first member of both S block elements, namely, Lithium and
Beryllium differ much from the rest of their members but at the same time, they resemble more with
the diagonal element present in the next column.
The anomaly of these S block elements is due to;
1. Low atomic and ionic size
2. Greater charge density (charge/volume of the atom)
3. Greater polarization
4. Absence of d-orbitals.
Greater polarization of s block elements makes the first element more covalent and differentiates them
from the rest which are ionic.
The similarity in size and charge density makes them resemble the element diagonally placed in the next
group (diagonal relationship).

INORGANIC CHEMISTRY
It is observed that the physical and chemical properties of these s block elements change in a particular
trend as the atomic number of the elements increases. Changes in the various properties of the group
are as mentioned below:
Chemical Properties of S Block Elements
Atomic and Ionic Radii
When the s block elements of the modern periodic table are observed it is seen that the size of the alkali
metals is larger compared to other elements in a particular period. As the atomic number increases the
total number of electrons increases along with the addition of shells.
On moving down the group the atomic number increases. As a result, the atomic and ionic radius of the
alkali metals increases.
Ionization Enthalpy
As we go down the group the size of the atoms increases due to which the attraction between the
nucleus and the electrons in the outermost shell decreases. As a result, the ionization enthalpy
decreases. The ionization enthalpy of the alkali metals is comparatively lesser than other elements.
Hydration Enthalpy
As the ionic sizes of the elements increase, the hydration enthalpy decreases. Smaller the size of the ion
the hydration enthalpy is high as the atom has the capacity to accommodate a larger number of water
molecules around it due to high charge/radius ratio and hence gets hydrated.
Physical Properties of S block elements
 In the S block elements, the density of the alkali metals increases down the
group. Exception: the density of potassium is less than the density of sodium.

INORGANIC CHEMISTRY
 The alkali metals have a low melting and boiling point due to the weak metallic bonding.
 Alkali metals and its respective salts have the capability to impart colour to the oxidizing
flame due to the heat generated from the flame which excites the valence electrons from
one energy level to another energy level. This helps in the detection of alkali metals during
the flame test.
Diagonal Relationship within S Block Elements
A diagonal relationship in S block elements exists between adjacent elements which are located in the
second and third period of the periodic table. For example, Lithium of group 1A and second period
shows similarities with the properties of magnesium which are located in the 2nd group and 3rd period.
Similarly, properties of beryllium which are located in the 2nd group and 2nd period show a likeness
with properties of aluminium which is located in the third period and third group. The two elements
which show similarities in their properties can be called a diagonal pair or diagonal neighbours.
The properties of S block elements vary significantly when compared to the other elements of the sub-
group they belong to. The diagonal neighbours show a lot of similarities. Such a relationship is
exhibited as you move left to right and down the group; the periodic table has opposing factors.
For example, the electronegativity of the S block elements increases as we go across the period and
decreases as we go down the group. Therefore, when it is moved diagonally the opposite tendencies
cancel out and the value of electronegativity almost remains the same.
Similarities between Lithium and Magnesium
 The hardness of lithium and magnesium is higher than the other elements in their
respective groups.
 Chlorides of lithium and magnesium have the capability to be soluble in ethanol.
 They are lighter when compared to other elements in their groups.
 Lithium and magnesium react gently with water. The oxides and hydroxides are less soluble.
 In the presence of nitrogen, lithium and magnesium form their respective nitrides.
 Superoxides are not formed when lithium and magnesium react with excess oxygen.
 Carbon dioxide and their respective oxides are formed when carbonates of magnesium
and lithium are heated.

INORGANIC CHEMISTRY
Similarities between Beryllium and Aluminum
 Aluminium hydroxide and beryllium hydroxide react with excess alkali to form their
respective ions.
 Both these elements have the capacity to withstand the acid attack due to the presence of
an oxide film on the surface of the metal.
 Both these metals have the tendency to form complexes.
 Chlorides of both these metals possess the capacity to be soluble in organic solvents.
CHEMISTRY OF ALKALI METALS
Occurrence in Nature
All the discovered alkali metals occur in nature. Experiments have been conducted to attempt the
synthesis of ununennium (Uue), which is likely to be the next member of the group if the attempt is
successful. It is predicted that the next alkali metal after ununennium would be unhexpentium (Uhp), an
element that has not yet received even attempts at synthesis due to its extremely high atomic number.
Alkali metals
Alkali metals belong to the s-block elements occupying the leftmost side of the periodic table. Alkali
metals readily lose electrons, making them count among the most reactive elements on earth. In this
article, we will explain the electronic configurations, ionization enthalpy, hydration enthalpy and atomic,
ionic radii and other physical and chemical properties of the group one alkali metals.
In general ‘alkali’ refers to the basic or alkaline nature of their metal hydroxides. The compounds are
called alkali metals because when they react with water they usually form alkalies which are nothing but
strong bases that can easily neutralize acids.
Alkali metals have a corresponding [Noble gas] ns1 electronic configuration. They occupy the first
column of the periodic table. Alkali elements are Lithium(Li), Sodium(Na), Potassium (K), Rubidium
(Ru), Cesium (Cs) and Francium (Fr) occupying successive periods from first to seven. Francium is a
radioactive element with very low half-life.
However, the main reason why hydrogen (H) is not considered as an alkali metal is that it is mostly
found as a gas when the temperature and pressure are normal. Hydrogen can show properties or
transform into an alkali metal when it is exposed to extremely high pressure.

INORGANIC CHEMISTRY
Alkali Metals are very reactive and are present in the form of compounds only. They are electropositive
metals with unit valence.
The group I comprising Li, Na, K, Rb, Cs & Fr are commonly called alkali metals. These are called alkali
metals because hydroxides of these metals are strong alkali. For example NaOH and KOH Francium is
radioactive and has a very short life (half life of 21 minutes), therefore very little is known about it.
Overview of Alkali Metals
Metals Lithium Sodium Potassium Rubidium Cesium
Atomic Number 3 11 19 37 55
Configuration [He]2s1 [Ne]3s1 [Ar]4s1 [Kr]5s1 [Xe]6s1
Abundance (ppm) 65 28300 25900 310 7
Atomic size (pm) 152 186 227 248 265
Density g/cm3 0.53 0.97 0.86 1.53 1.9
Ionization energy kJ/mol 520 496 419 403 376
Hydration enthalpy kJ/mol -506 -406 -330 -310 -276
Reduction potential (v) -3.04 -2.714 -2.925 -2.930 -2.927
Flame colour Crimson red Yellow Violet Red Violet Blue

INORGANIC CHEMISTRY
Electronic Configuration of Alkali Metals
 Alkali metals have one electron in their valence shell.
 The electronic configuration is given by ns1. For example, the electronic configuration
of lithium is given by 1ns1 2ns1.
 They tend to lose the outer shell electron to form cations with charge +1 (monovalent ions).
 This makes them the most electropositive elements and due to the same reason, they are
not found in the pure state.
PHYSICAL PROPERTIES OF ALKALI METALS
The alkali metals have the high thermal and electrical conductivity, lustre, ductility, and malleability that
are characteristic of metals. Each alkali metal atom has a single electron in its outermost shell. This
valence electron is much more weakly bound than those in inner shells. As a result, the alkali metals
tend to form singly charged positive ions (cations) when they react with nonmetals. The compounds that
result have high melting points and are hard crystals that are held together by ionic bonds (resulting
from mutually attractive forces that exist between positive and negative electrical charges). In the
metallic state, either pure or in alloys with other alkali metals, the valence electrons become delocalized
and mobile as they interact to form a half-filled valence band. As with other metals, such a partially filled
valence band is a conduction band and is responsible for the valence properties typical of metals. In
passing from lithium to francium, the single electron tends to be less strongly held. Generally, the
energy necessary to remove the outermost electron from the atoms of an element, the ionization
energy, decreases in the periodic table toward the left and downward in each vertical file, with the
result that the most easily ionizable element in the entire table is francium, followed closely by cesium.
The alkali metals, which make up the extreme left-hand file, have ionization energies ranging from 124.3
kilocalories per mole (kcal/mole) in lithium to 89.7 kcal/mole in cesium (omitting the rare radioactive
element francium). The alkaline-earth metals, the next group to the right, have higher ionization
energies ranging from 214.9 in beryllium to 120.1 kcal/mole in barium.
The electronegativity scale of the elements compares the ability of the atoms of the various elements to
attract electrons to themselves. In the periodic table the electronegativities range from 0.7 for cesium,
the least electronegative of the elements, to 4.0 for fluorine, the most electronegative. Metals are
ordinarily considered to be those elements having values less than 2.0 on the electronegativity scale. As
a group the alkali metals are the least electronegative of the elements, ranging from 0.7 to 1.0 on the
scale, while the alkaline earths, the next group on the table, have electronegativities ranging from about
0.9 to 1.5.
The table summarizes the important physical and thermodynamic properties of the alkali metals. At
atmospheric pressure these metals are all characterized by a body-centred cubic crystallographic
arrangement (a standard pattern of atoms in their crystals), with eight nearest neighbours to each atom.
The closest distance between atoms, a characteristic property of crystals, increases with increasing
atomic weight of the alkali metal atoms. As a group, the alkali metals have a looser crystallographic
arrangement than any of the other metallic crystals, and cesium—because of its greater atomic
weight—has an interatomic distance that is greater than that of any other metal.
Vapour-pressure data for the alkali metals and for two alloys formed between elements of the group
show that the vapour pressures increase in regular fashion with increasing atomic weight. Cesium is the
most volatile of the alkali metals, with a boiling point of 671 °C (1,240 °F). The boiling points of the alkali

INORGANIC CHEMISTRY
metals decrease in regular fashion as the atomic numbers increase, with the highest, 1,317 °C (2,403 °F),
being that of lithium.
The melting points of the alkali metals as a group are lower than those of any other nongaseous group
of the periodic table, ranging between 179 °C (354 °F) for lithium and 28.5 °C (83.3 °F) for cesium.
Among the metallic elements, only mercury has a lower melting point (−38.9 °C, or −38.02 °F) than
cesium. The low melting points of the alkali metals are a direct result of the large interatomic distances
in their crystals and the weak bond energies associated with such loose arrays. These same factors are
responsible for the low densities, low heats of fusion, and small changes in volume upon fusion of the
metals. Lithium, sodium, and potassium are less dense than water.
The large size of an alkali metal atom (and the resulting low density of the metal) results from the
presence of only one, weakly bound electron in the large outer s-type orbital. Upon going from the
noble-gas configuration of argon (atomic number 18) to potassium (atomic number 19), the added
electron goes into the large 4s orbital rather than the smaller 3p orbital. When, however,
potassium, rubidium, or cesium metals are subjected to increasing pressure (up to one-half million
atmospheres or more), a number of phase transitions occur. Ultimately, occupation of a d-type orbital
becomes preferred over that of the s-orbital, with the result that these alkali metals resemble transition
metals. Under such circumstances, alloys with transition metals (such as iron) can form, a result that
does not occur at low pressures. It has been proposed that the lower-than-expected density
of Earth’s core may be the result of the formation of a potassium-iron alloy under the extreme pressures
that occur there.
The alkali metals have played an important role in quantum physics. Some alkali metal isotopes, such as
rubidium-87, are bosons. Dilute atomic gases of such alkali metal isotopes, confined by magnetic
fields or “laser mirrors” and cooled to temperatures near absolute zero, form Bose-Einstein
condensates. In this state, the cluster of atoms is in a single quantum state and exhibits macroscopic
behaviours normally seen only with atomic-sized particles. These include interference effects
and coherent motion of the entire “cloud” of atoms.
The general electronic configuration of alkali metals may be represented by [noble gas] where n = 2 to 7
Property
Elements
Li Na K Rb Cs Fr
(Radioactive)
Atomic Number 3 11 19 37 55 87
Electronic Configuration 2s1
3s1
4s1
5s1
6s1
7s1
Atomic Mass 6.94 22.99 39.10 85.47 13.91 223
Metallic radius (pm) 152 186 227 248 265 375
Ionic radius (M+
/pm) 76 102 138 152 167 180

CHEMICAL SCIENCE MCQs BOOK- 1
INORGANIC CHEMISTRY
CHEPTER – 1 - CHEMICAL PERIODICITY MCQs
1. Electronegativity is defined as the power
of an atom in a molecule to _____________
a) Repel electrons towards itself
b) Attract electrons towards itself
c) Expand itself
d) All of the mentioned
Answer: b
Explanation: Electronegativity is defined as
the power of an atom in a molecule to
attract electrons towards itself. Fluorine is
the most electronegative element.
2. The factors on which electronegativity
depends upon ____________
a) Valence state of atom
b) Hybridisation
c) Both valence state and hybridisation
d) None of the mentioned
Answer: c
Explanation: The factors on which
electronegativity depends upon is valence
state of atom and hybridisation.
3. How does the electronegativiy get
affected with the negative oxidation state?
a) It decreases
b) It increases
c) It remains constant
Answer: a
Explanation: Electronegativity decreases
with the negative oxidation state since the
tendency to attract an electron will
decrease with the negative charge of the
anion.
4. The electronegativity of sp2
hybridised
atom will be ____________
a) 3.29
b) 2.48
c) 3.69
d) 2.75
Answer: d
Explanation: The electronegativity of
sp2 hybridised atom will be 2.75. Fluorine is
the most electronegative element.
5. Which of the following is a permanent
electron displacement effect?
a) Inductomeric
b) Electromeric
c) Inductive
Answer: c
Explanation: Inductive effect is the
permanent electron displacement effect
and inductomeric and electromeric are
temporary electron displacement effects.
6. Arrange the following groups in the order
of decreasing (+I) effect.
a) C6H5O–
> COO–
> CR3 > CHR2 > H
b) C6H5O–
> H > CR3 > CHR2 > COO–
c) CR3 > C6H5O–
> H > COO–
> CHR2
d) C6H5O–
> COO–
> CHR2 > CR3 > H
Answer: a
Explanation: The correct order is- C6H5O–
>
COO–
> CR3 > CHR2 > H.
7. Arrange the following groups in the order
of decreasing (-I) effect.
a) CN > F > Br > Cl > COOH > I > H
b) COOH > CN > F > Br > Cl > I > H
c) H > COOH > CN > I > Cl > F > Cl
d) CN > COOH > F > Cl > Br > I > H
Answer: d
Explanation: The correct order is- CN >
COOH > F > Cl > Br > I > H.
8. Which of the following is an application
of inductive effect?
a) Bond length
b) Dipole moment
c) Strength of carboxylic acids

Answer: d
Explanation: Bond length, dipole moment
and strength of carboxylic acids are some of
the applications of inductive effect.
9. Relative basic strength of amines does
not depend upon ____________
a) Inductive effect
b) Mesomeric effect
c) Steric effect
d) Stabilisation of cation by hydration
Answer: b
Explanation: Relative basic strength of
amines does not depend upon mesomeric
effect. This effect is used in a qualitative
way and describes the electron withdrawing
or releasing properties of substituents
based on relevant resonance structures.
10. Due to presence of C – X polar bond in
alkyl halide, alkyl halides are ____________
a) More reactive than corresponding alkane
b) Less reactive than corresponding alkane
c) Equally reactive as corresponding alkane
Answer: a
Explanation: Due to the presence of C – X
polar bond in alkyl halide, alkyl halides are
more reactive than corresponding alkane.
11. Atomic radii ____________ along the
periods.
a) Increases
b) Decreases
c) Remains constant
d) Irregular
Answer: b
Explanation: Atomic radii decreases along
the periods because as the number of
electrons increases in the same shell of the
atom, the effective nuclear attraction
increases, thereby reducing the distance
between the outer shell and the nucleus i.e.
decreasing atomic radii.
12. O2-, F–, Na+ and Mg2+ are called as
__________
a) Isoelectronic species
b) Isoneutral species
c) Isotopes
d) Isobars
Answer: a
Explanation: O2-
, F–
, Na+
and Mg2+
are
known as isoelectronic species as they all
have the same no. of electrons(10).
Isotopes contain the same no. of protons
but a different number of neutrons. Isobars
have the same mass number but a different
atomic number.
13. X(g) → X+
(g) + e–
. What does this chemical
reaction need to occur?
a) Catalyst
b) Electron affinity
c) Electropositivity
d) Ionization energy
Answer: d
Explanation: The minimum amount of
energy that is required to remove an
electron from an atom is called Ionization
enthalpy. It is expressed in the units KJ/mol.
Electron affinity is the change in the energy
when we add an electron to a neutral atom.
Electropositivity is a metallic characteristic.
A catalyst speeds up the reaction.
14. What is the correct order of
electronegativity among the following
options?
a) Li<Na<K<Rb<Cs
b) Li<K<Na<Rb<Cs
c) Li>Na>K>Cs>Rb
d) Li>Na>K=Rb>Cs
Answer: d
Explanation: Electronegativity is the
measure of the ability to attract shared
electrons to itself of an atom in a chemical
compound. The values of electronegativities
of Li, Na, K, Rb and Cs are 1, 0.9, 0.8, 0.8
and 0.7 respectively.

15. Ionization energies are always positive.
a) True
b) False
Answer: a
Explanation: As the ionization energy is the
minimum amount of energy that is required
to remove an electron from an atom. As the
energy is always needed for the removal of
an electron from an atom, the values of
ionization energies are always positive.
6. What is the oxidation state of Mn in
KMnO4?
a) 5
b) 6
c) 7
d) 4
Answer: c
Explanation: The total charge of the
compound KMnO4 is zero as it is a neutral
and stable compound. As we know the
oxidation states of K and O are +1 and -2
respectively. So +1 + Mn charge + 4(-2) = 0;
Mn charge = 7.
17. The relationship between Li, Mg and Be,
Al is called the __________ relationship.
a) Diagonal
b) Periodic
c) Group
d) Triangle
Answer: a
Explanation: The Elements Li and Mg, Be
and Al are similar to each other in the case
of formation of different compounds in a
similar composition. They have a sort of
similar behaviour with each other, so they
are said to be in a diagonal relationship as
per their place in the periodic table.
18. N2O is a _________
a) Tear gas
b) Laughing gas
c) Acid
d) Base
Answer: c
Explanation: N2O is called laughing gas as it
has euphoric effects after being inhaled. It
is one of the World Health Organization’s
Essential Medicines. it is used for
recreational and anaesthetic purposes
mostly.
19. Which of the following is superoxide?
a) K2O
b) Na2O
c) MgO
d) KO2
Answer: d
Explanation: In the superoxide of the
element, the oxygen’s oxidation state is
given by -1/2. In peroxides, the oxidation
state of oxygen is -1 as in case of hydrogen
peroxide. The oxygen’s oxidation state in
KO2 is -1/2.
20. What can be tested using a litmus
paper?
a) Acidic nature only
b) Basic nature only
c) Both acidic nature and basic nature
d) Nothing
Answer: c
Explanation: A litmus paper is a dye that is
extracted from lichens. It used to test acidic
nature and basic nature of a substance. Red
colour indicates acidic nature and blue
colour indicated basic nature. Neutral
litmus paper is purple in colour.
21. __________ gave the idea for the first
time to classify elements as per their
properties.
a) Mendeleev
b) Dobereiner
c) Newland
d) John
Answer: b
Explanation: In the early 1800s, a German
chemist named Dobereiner came up with
the idea of classifying elements as per their

Explanation: According to Fajan's rules, the
greater the charge the greater is the
polarizing power of cation.
171. Maximum covalent character is
associated with the compound:
A) NaI
B) MgI2
C) AlCl3
D) AlI3
Answer: D
Explanation: Among cations, Al3+
has
smaller size and greater charge hence it has
greater polarizing power. Among anions,
iodide (I-
) is larger and hence can be
polarized easily. Hence the combination of
these two ions i.e. Al3+
and I-
results in more
covalent nature.
172. According to Fajan's rules, the covalent
bond is favored by
A) Large cation and small anion
B) Large cation and large anion
C) Small cation and large anion
D) Small cation and small anion
Answer: C
173. Polarisibility of halide ions increases in
the order:
A) F-,I-,Br-,Cl-
B) Cl-
,Br-
,I-
,F-
C) I-
,Br-
,Cl-
,F-
D) F-,Cl-,Br-,I-
Answer: D
174. Amongst LiCl, RbCl, BeCl2 and
MgCl2 the compounds with the greatest and
the least ionic character, respectively, are:
A) LiCl and RbCl
B) RbCl and BeCl2
C) RbCl and MgCl2
D) MgCl2 and BeCl2
Answer: B
175. Polarization power of cation increases
when:
A) size decrease
B) size increases
C) Anion has greater polarizing
power
D) covalent nature increases
Answer: A
176. Polarising power is directly
proportional to:
A) size of cation
B) charge on cation
C) electronegativity of cation
D) size of anion
Answer: B
CHEPTER – 2 – BONDING AND STRUCTURE OF MOLECULES MCQs
1. Atoms obtain octet configuration when
linked with other atoms. This is said by
_________
a) Lewis
b) Kossel
c) Langmuir
d) Sidgwick
Answer: a
Explanation: The above statement says that
the atoms achieve a stable octet
configuration when joined with other atoms
through chemical bonds as postulated by
Lewis. An example of this is the formation
of NaCl molecule where Na and Cl transfer
electrons to each other forming Na+
and Cl–
.
2. Find out the correct Lewis symbol for the
atom carbon among the following options.
a) .C:

b) :C.
c) :C:
d) .C.
Answer: c
Explanation: An American chemist G.N.
Lewis created Lewis symbols as a notation
to represent the valance electrons in an
atom. As the carbon atom has 4 electrons in
its outer shell, it is represented by 4 dots
around it.
3. What’s the group valance of atoms in the
halogen family?
a) 2
b) 1
c) 9
d) 7
Answer: b
Explanation: The group valance can be
calculated from Lewis symbols either by
subtracting it from eight (more than 4) or
having it equal (less than 4). The halogen
family has 7 electrons in their outer orbit.
So 8 – 7 = 1. Therefore the valency of the
halogen family is 1.
4. Highly electropositive Alkali metals are
separated from highly electronegative
halogens by _________
a) noble gases
b) oxygen family
c) f-block elements
d) 7th period
Answer: a
Explanation: According to Kossel, Highly
electropositive Alkali metals are separated
from highly electronegative halogens by
noble gases. This is because Alkali metals
are the 1st
group and halogens the
17th
group. Elements in 18th
group i.e.
nobles are preceded by group 17 elements
and succeeded by group 1 elements.
5. Sharing or transfer of electrons from one
atom to the other to attain stable octet
configuration follows _______
a) Duet rule
b) Triplet rule
c) Octet rule
d) Septet rule
Answer: c
Explanation: As per the electronic theory of
chemical bond that’s put forth by Lewis
&Kossel states that the atoms follow the
octet rule by sharing or transfer of electrons
from one atom to the other to attain stable
octet configuration.
6. In the covalent bond, atoms share
electrons to achieve octet configuration.
a) True
b) False
Answer: a
Explanation: In the year 1919, Langmuir
postulated the theory of covalent bond and
its formation by combining with Lewis
theory. An example of this is the formation
of Cl2, Two atoms of Cl combine by sharing
the 7th
electron in its outer shell.
7. Which of the following molecule doesn’t
involve covalent bond?
a) H2O
b) CCl4
c) NaCl
d) O2
Answer: c
Explanation: The formation of NaCl
molecule where Na and Cl transfer
electrons to each other forming Na+
and Cl–
.
There is no sharing of electrons i.e. no
covalent bond. Whereas the molecules H2O,
Cl2 and O2 involve sharing of electrons.
8. Calculate the formal charge of C in CH4.
a) 4
b) 1
c) -4
d) 0
Answer: d
Explanation: The formula for finding out the
formula charge of an in a molecule = total

number of valence electrons – total number
of non-bonding electrons – 1/2(total
number of bonding electrons). So here,
formal charge of C = 4 – 0 – 8/2 = 0.
9. Which of the following doesn’t follow
octet rule?
a) CH4
b) CCl4
c) HCl
d) NO2
Answer: d
Explanation: Though octet rule is widely
known, it does have a few limitations. The
compound nitrogen dioxide NO2 doesn’t
follow the octet rule. It’s a molecule with an
odd number of electrons. Even the nitric
oxide NO doesn’t follow.
10. Calculate the formal charge of the
middle atom in the ozone molecule.
a) 1
b) -1
c) 0
d) -2
Answer: a
Explanation: The formula for finding out the
formula charge of an in a molecule = total
number of valence electrons – total number
of non-bonding electrons – 1/2 (total
number of bonding electrons). So here, a
formal charge of central O is 6 – 2 – 6/2 = 1.
11. A chemical bond formation that involves
the complete transfer of electrons between
atoms is _______
a) ionic bond
b) covalent bond
c) metallic bond
d) partial covalent bond
Answer: a
Explanation: Ionic bond, which is otherwise
known as electrovalent bond forms
between two atoms by the transfer of
electrons between them. It generates
oppositely charged ions. Positively charged
ions are mostly metals and the vice-versa.
12. Formation of a compound through ionic
bond ______ the ionization energy of the
metal ion.
a) does not depends on
b) depend on
c) is independent regarding
d) may or may not depend on
Answer: b
Explanation: For the formation of the ionic
bond, the metal ion has to overcome to
energy for the removal of an electron from
its outer shell in order to become a cation,
that is ionization energy. Therefore
Formation of a compound through ionic
bond depends on the ionization energy of
the metal ion.
13. The enthalpy change that occurs when
an atom in the ground state gains an
electron, is electron gain enthalpy.
a) True
b) False
Answer: a
Explanation: Yes. electron gain enthalpy is
the enthalpy change for an atom in the
ground state to gain an electron. An atom
gains an electron, thus forming negatively
charged ion also known as a cation.
Symbolic representation is as follows: A(g) +
e–
→ A(g)
–
.
14. Electron gain enthalpy may be
________
a) exothermic
b) endothermic
c) both exothermic and endothermic
d) always zero

CHEPTER – 3 – ACID BASE CHEMISTRY MCQs
1. Which one is correctly matched?
a) Acids – pH range above7
b) Acids – pH range below 7
c) Acids – pH range 7(neutral)
d) Acids – pH range 8-9
Answers: b
Explanation: Acids have a pH range less
than 7, Water is the only solvent that has a
pH of 7 (neutral).
2. A Strong acid is same as concentrated
acid.
a) False
b) True
Answer: a
Explanation: Concentration of an acid
depends upon the water content whereas
the strength of an acid depends on
dissociation power.
3. When an acid reacts with a metal, which
one of the following gas is usually
liberated?
a) ammonia gas
b) chlorine
c) oxygen
d) Hydrogen gas
Answer: d
Explanation: When metal reacts with acid, a
soap bubble is formed and the bubble
contains Hydrogen gas (example: HCl,
H2SO4).
4. Which of the following is wrongly
mapped?
a) Sodium carbonate – Washing soda
b) Sodium chloride – common salt
c) Calcium carbonate – slaked lime
d) Sodium hydroxide – caustic soda
Answer: c
Explanation: calcium hydroxide is
commonly referred as slaked lime.
5. What will be the X in the following
equation?
MgO + 2HCl —-> X + H2O
a) Mg2Cl
b) 2MgCl
c) MgCl
d) MgCl2
Answer: d
Explanation: MgCl2 is the product formed
when magnesium oxide reacts with
hydrochloric acid and water is formed as a
by-product.
6. Which of the following is neither an acid
nor base?
a) CH3COOH
b) HCl
c) KCl
d) CH3OH
Answer: c
Explanation: CH3COOH and CH3OH are
organic acids, HCl is strong acid and KCl is a
salt.
7. Which one will change from red litmus to
blue?
a) NaCl
b) HCl
c) KOH
d) LiOH
Answer: b
Explanation: since HCl is a base it turns red
litmus to blue.
8. What is the pH of 0.0001 molar HCl
solution?
a) 1
b) 2
c) 3
d) 4
Answer: d
Explanation: It has a H+ concentration of 10-
4. The value of negative exponent(^-4) gives
pH value to be 4.
9. What will be the product when
HNO3 reacts with NH4OH?

a) NH4 NO3
b) 2NH4 NO3
c) NH4 (NO3)2
d) NH2NO3
Answer: a
Explanation: the reaction takes place as
follows:
HNO3 + NH4OH —-> NH4NO3 + H2O.
10. Find the odd one out.
a) Neutral salt : NaCl
b) Acid salt : CuSO4.5H2O
c) Basic salt: CuCO3.Cu(OH)2
d) Nonhydrated salt: KNO3
Answer: b
Explanation: CuSO4.5H20 is a hydrated salt.
An example of acid salt is NaHCO3.
11. CH3COOH ⇌ CH3COO– + H+ is in
__________________
a) ionic equilibrium
b) chemical equilibrium
c) dynamic equilibrium
d) physical equilibrium
Answer: a
Explanation: The equilibrium that is
attained between the ionized molecules
and the ions in the solution of weak
electrolyte is called Ionic Equilibrium.
CH3COOH ⇌ CH3COO–
+ H+
is an example;
CH3COO– and H+ are ions.
12. Electrolytes conduct electricity.
a) True
b) False
Answer: a
Explanation: Chemical substances which can
conduct electricity in their Aqua state or in
the molten state are called electrolytes. The
conduction of current through the
electrolyte is due to the movement of Ions,
hence the above statement is true.
13. Which of the following may not be a
strong electrolyte?
a) hydrochloric acid
b) sulfuric acid
c) nitric acid
d) ammonia
Answer: d
Explanation: Electrolytes which dissociate
almost completely into constituent ions in
aqueous solutions are known as strong
electrolytes. Therefore ammonia is not a
strong electrolyte because it can’t
dissociate completely.
14. All organic acids except sulfonic acid are
_____________ electrolytes.
a) weak
b) strong
c) not
d) neither strong nor weak
Answer: a
Explanation: Electrolytes which dissociate
into a lesser extent in aqua solution are
called weak electrolytes. All organic acids
except sulfonic acids and bases like
Ammonia, Ammonium hydroxide, amines,
etc are weak electrolytes.
15. Can nonelectrolytes conduct electricity?
a) yes
b) no
c) sometimes
d) cannot say
Answer: b
Explanation: Michael Faraday classified
substances into two categories; one is
electrolytes and nonelectrolytes,
nonelectrolytes do not dissociate into ions
in a solution. So they do not conduct
electricity.
16. Sugar solution __________ electricity.
a) do not conduct
b) conducts
c) depends on the type of sugar
d) cannot say
Answer: a
Explanation: Aqueous solution of sugar
does not conduct electricity, but Aqueous

solution of sugar conducts electricity. This is
because the aqueous solution of sugar is a
nonelectrolyte, whereas the salt solution is
an electrolyte.
17. Which of the following is in Ionic
Equilibrium?
a) 2AgI + Na2S ⇌Ag2S + 2NaI
b) 4 NH3 + 5 O2 ⇌4 NO + 6 H2O
c) TiCl4 + 2 H2O ⇌TiO2 + 4 HCl
d) H2O + H2O ⇌H3O+ + OH–
Answer: d
Explanation: Only H2O + H2O ⇌H3O+ + OH– is
in ionic equilibrium. As the equilibrium
established between the unionized
molecules and the ions in the solution of
weak electrolytes is known as Ionic
Equilibrium.
18. What is the degree of dissociation for
strong electrolytes?
a) 1
b) 0
c) less than 1
d) greater than 1
Answer: a
Explanation: Degree of dissociation is the
fraction of the total number of molecules
which dissociate into constituent ions, it is
represented by the symbol ɑ. As a strong
electrolyte dissociate completely, it values
is 1.
19. Degree of dissociation does not depend
on which of the following factors?
a) nature of the solute
b) nature of the solvent
c) sound
d) concentration
Answer: c
Explanation: Values of the degree of
dissociation or degree of ionization depends
upon the following factors: 1) the nature of
the solute, 2) the nature of the solvent 3)
concentration and 4) temperature of the
solution.
20. K in K = Cα2
/1 – ɑ represents
___________
a) dissociation constant
b) molar concentration
c) degree of dissociation degree of
ionization
d) degree of ionization
Answer: a
Explanation: The above equation represents
Ostwald’s dilution law, where K is the
dissociation constant, C is the molar
concentration of the solution and ɑ is a
degree of dissociation or degree of
ionization of the solution.
21. Which of the following is not a property
of an acid according to Robert Boyle?
a) turns blue Litmus red
b) sour in taste
c) neutralize bases
d) bitter in taste
Answer: d
Explanation: According to Robert Boyle,
acids are the substances which have a sour
taste, turns blue Litmus red, liberate
hydrogen with metals conduct electricity in
aqueous solution and neutralize bases. They
do not have a bitter taste.
22. Bases turn red litmus blue.
a) True
b) False
Answer: a
Explanation: Litmus is a mixture of different
dyes from lichens that is water soluble.
Acids change blue litmus red and bases
change red litmus blue. The original colour
of Litmus is purple. The pH of a base is in
between 7 and 14.
23. HCl is an Arrhenius ___________
a) acid
b) base
c) salt
d) water

Answer: b
Explanation: For any reversible reaction at
any stage other than equilibrium, the ratio
of the molar concentrations of the products
to that of the reactants, where is
concentration term is raised to the power
equal to the stoichiometric efficient to the
substance concerned is called the reaction
quotient, QC.
159. For a reaction aA + bB → cC + dD,
which is not in equilibrium the QC is given as
__________
a) [A]a[B]b/[C]c[D]d
b) [C]c
[D]d
/[A]a
[B]b
c) [A][B]/[C][D]
d) [C][D]/[A][B]
Answer: b
Explanation: A very basic reaction like aA +
bB → cC + dD, where the capital letters
represent the compounds or molecules and
the small letters are the coefficients of
them the reaction quotient QC, is given by
[C]c
[D]d
/[A]a
[B]b
.
160. What will happen If QC > KC?
a) QC decreases till equilibrium
b) QC increases till equilibrium
c) QC remains constant
d) cannot say
Answer: a
Explanation: If QC > KC, the value of QC will
tend to decrease to reach the value of
equilibrium constant (that is towards
equilibrium) and the reaction will continue
in the opposite direction, where QC is
reaction quotient and KC is the equilibrium
constant.
161. At equilibrium, KC is _______________
a) greater than reaction quotient
b) equal to the reaction quotient
c) less than the reaction question
d) independent of reaction question
Answer: b
Explanation: At equilibrium, the equilibrium
constant and the reaction quotient is equal.
The equilibrium constant is depicted by the
symbol KC and the reaction quotient is
represented by the symbol QC.
162. What do you think will happen if
reaction quotient is smaller than the
equilibrium constant?
a) equilibrium constant will change
b) reaction quotient remains constant
c) reaction quotient increases continuously
d) reaction quotient increases till KC
Answer: d
Explanation: If the reaction quotient is less
than the equilibrium constant KC, the
reaction quotient will tend to increase and
the reaction will proceed in the forward
direction, till it reaches the value of the
equilibrium constant.
163. What do you understand from the
reaction if reaction quotient is 2 and the
equilibrium constant is 3?
a) the equilibrium constant increases
b) the equilibrium constant decreases
c) the equilibrium constant remains the
same
d) reaction quotient increases
Answer: d
Explanation: In a reaction, if reaction
quotient is less than the equilibrium
constant, the reaction quotient will tend to
increase and the reaction will proceed in
the forward direction till it reaches
equilibrium.
CHEPTER – 4 – s – BLOCK ELEMENTS MCQs
1. Which of the following metal is not an
alkali metal?
a) magnesium
b) rubidium

c) sodium
d) caesium
Answer: a
Explanation: Alkali metals are the elements
of group 1. The outer shell configuration of
group 1 elements is ns1
, where n is the
number of it’s period. Magnesium is not an
alkali metal because it’s outer shell
configuration is ns2
.
2. Alkali metals have the biggest atomic
radii.
a) true
b) false
Answer: a
Explanation: The alkali metals have the
biggest atomic radii in their respective
periods, atomic radii increases as we go
down the group due to the addition of a
new shell in each subsequent step. So the
above statement is true.
3. The melting point of alkali metal is
_____________
a) depends on the atmosphere
b) low
c) high
d) zero
Answer: b
Explanation: The melting and boiling points
of alkali metals are quite low and they
decrease down the group due to weakening
of their metallic bonds. Francium is the only
element in this group which is a liquid at
room temperature.
4. Is there removal of second electron
difficult in alkali metals?
a) Yes
b) No
c) Maybe
d) Cannot say
Answer: a
Explanation: The first ionization enthalpy of
alkali metals is the lowest among the
elements in their respective periods and
increases on moving down the Group. The
second ionization enthalpies of the alkali
metals are very high because by releasing
an electron, ions require noble gas
configuration, so removal of the second
electron is difficult.
5. Alkali metals are strongly _____________
a) neutral
b) electropositive
c) electronegative
d) non-metallic
Answer: b
Explanation: Due to low ionization
enthalpies, alkali metals are strongly
electropositive or metallic in nature and
electropositive nature increases from
Lithium to caesium due to decrease in
ionization enthalpy.
6. Alkaline earth metals show +1 Oxidation
state and their atomic volume
_____________ down the group.
a) is irregular
b) increase
c) decrease
d) do not change
Answer: c
Explanation: The alkali metal atom show
only +1 Oxidation State, because of their
unipositive Ion at the time the stable noble
gas configuration. The atomic volume of
alkali metals is the highest in its period and
goes on increasing down the group from
top to bottom.
7. Does the degree of hydration depend
upon the size of the cation?
a) Yes
b) No
c) Maybe
d) Cannot say
Answer: a
Explanation: The degree of hydration
depends upon the size of the cation, smaller
the size of the cation, greater is its

hydration enthalpy. The relative degree of
hydration in an increasing order is Li+
>
Na+ > K+ > Rb+ > Cs+.
8. The flame of caesium is in the colour
_____________
a) white
b) red violet
c) yellow
d) blue
Answer: d
Explanation: Alkali metals and their salt
impart characteristic colours to the flame
because the outer electrons get excited to
higher energy levels. When the electrons
return to the original state, it releases
visible light of a characteristic wavelength
which provides colour to the flame. The
colour of the Flame of the caesium is blue.
9. Caesium has the highest electrical
conductivity in its group.
a) true
b) false
Answer: a
Explanation: Due to the presence of loosely
held Valence Electrons which are free to
move throughout the metal structure, the
alkali metals are good conductors of heat
and electricity. Electrical conductivity
increases from top to bottom in the order,
so caesium has the highest electrical
conductivity in its group.
10. All alkali metals are good dash agents?
a) oxidizing
b) reducing
c) both oxidising and reducing
d) neither oxidizing not reducing
Answer: b
Explanation: All the alkali metals are good
reducing agents due to their low ionization
energies. The reducing character of group 1
elements follows the increasing order of
Sodium, Potassium, rubidium, Caesium and
lithium.
11. What happens when alkali metals are
exposed to moist air?
a) formation of nitrates
b) formation of oxides
c) formation of chlorides
d) formation of sulphates
Answer: b
Explanation: On exposure to moist air, the
surface gets tarnished due to the formation
of oxides, hydroxide and carbonates. Few
examples are sodium hydroxide, sodium
carbonate, potassium hydroxide etc.
12. Sodium Peroxide is _____________ in
colour and potassium superoxide is used as
a source of _____________
a) blue, yellow
b) yellow, hydrogen
c) blue, oxygen
d) yellow, oxygen
Answer: d
Explanation: Sodium Peroxide acquires
yellow colour due to the presence of
superoxide as an impurity. Potassium
superoxide is used as a source of oxygen in
submarines, space shuttles and an
emergency breathing apparatus such as
oxygen masks.
13. Which of the following is true regarding
the reactivity order of alkali metals towards
hydrogen?
a) Li < Na < K > Rb
b) Lithium < Na < K < Rb < Cs
c) Li > Na < Cs
d) Li < Rb > Cs
Answer: b
Explanation: Two moles of alkali metal
reacts with one mole of hydrogen molecule
in order to form 2 moles of alkali metal
hydride. The correct order of reactivity of
alkali metals towards hydrogen is Li < Na < K
< Rb < Cs.
14. Lithium fluoride is _____________ in
water.

4) Slaked lime + sand + H2O
Answer: Lime mortar is a mixture of one
part of Slaked lime (Ca(OH)2), three parts of
Sand (SiO2) and water. It is converted into
calcium silicate with time and becomes
hardened. Hence it is used as building
material.
Ca(OH)2 + SiO2 --------> CaSiO3 + H2O
138. The correct order of thermal stability
of following carbonates is:
1. BaCO3 > CaCO3 > SrCO3 > MgCO3
2. BaCO3 > SrCO3 > CaCO3 > MgCO3
3. MgCO3 > CaCO3 > SrCO3 > BaCO3
4. MgCO3 > CaCO3 > BaCO3 > SrCO3
Logic and Solution: If the carbonate ion is
polarized by the cation, it undergoes
decomposition easily upon heating. The
polarizing power of alkali metal cations
decreases with increase in the size.
i.e. the order of polarizing power of cations
is:
Mg2+ > Ca2+ > Sr2+ > Ba2+
Hence the correct order of thermal stability
of carbonates will be:
BaCO3 > SrCO3 > CaCO3 > MgCO3
139. The substance not likely to contain
CaCO3 is:
1) a marble statue
2) calcined gypsum
3) sea shells
4) dolomite
Solution:
 Marble, sea shells and dolomite
contain calcium carbonate.
 Dolomite is CaCO3 + MgCO3
 Calcined gypsum is dehydrated
gypsum otherwise known as plaster
of paris. It does not contain any
calcium carbonate.
140. ‘Magnalium’ is an alloy of :
1) Mg + Zn
2) Mg + Al
3) Cu + Zn
4) Mg + Cu
Answer: Magnalium is an alloy of
magnesium (1.5 - 2%) and aluminium
metals. Due to high strength, low density
and ability to resist corrosion; it is used in
amking aircraft and automobile parts.
141. A burning strip of Magnesium in
introduced into a jar containing a gas. After
sometime the walls of the container are
coated with carbon. The gas in the
container is:
1) O2
2) N2
3) CO2
4) H2O
Explanation: The gas in the container is
carbon sicne magnesium can reduce carbon
dioxide to carbon.
2Mg + CO2 ---------> 2MgO + C
CHEPTER – 5 – p – BLOCK ELEMENTS I MCQs
1. Are group 13 elements a part of p block
elements?
a) Yes
b) No
c) Only a few
d) Only one
Answer: a
Explanation: The last electron enters in the
outermost p-orbital in the p block elements,
from group 13 to group 18 the entire
elements belong to p-block, whereas group
13 is called a boron family. It includes the
elements boron, aluminum, gallium, indium
and thallium.
2. Which of the following group’s elements
have smaller atomic radii?
a) Group 1 elements
b) Group 2 elements

c) Group 13 elements
d) All have the same atomic radii
Answer: c
Explanation: Group 13 elements have
smaller atomic radii and ionic radii than
those of alkaline earth metals and alkali
metals due to the greater effective nuclear
charge, atomic radii increases on going
down the group with an abnormality at
gallium.
3. The atomic radius of gallium is greater
than that of aluminum.
a) True
b) False
Answer: b
Explanation: Though the atomic radii
increase on going down the group, the
radius of gallium decreases unexpectedly
because of the presence of electrons in the
orbitals which do not screen the attraction
of the nucleus effectively. So the atomic
radius of gallium is less than that of
aluminium.
4. Gallium remains liquid up to __________
Kelvin.
a) 2176
b) 2376
c) 2476
d) 2276
Answer: d
Explanation: Low melting point of gallium is
due to the fact that it consists of
Ga2 molecules and gallium remains liquid up
to 2276 k. Hence it is used in high-
temperature thermometer. Gallium as a
chemical symbol that is Ga and its atomic
number is given as 31.
5. The ionization enthalpy _________ down
the group in the family.
a) Increases
b) Decreases
c) Constant
d) Is a regular
Answer: d
Explanation: On moving down the group,
ionization enthalpy decreases from Boron
to aluminium, but the next element gallium
has slightly higher ionization enthalpy than
aluminium due to the poor shielding of
intervening d-electrons, it again increases in
indium and then decreases in the last
element thallium.
6. Inert pair affect __________ down the
group.
a) Increases
b) Decreases
c) Constant
d) Is a regular
Answer: a
Explanation: Inert pair effect is the
reluctance of the selections of the valence
shell to take part in bonding, it occurs due
to pore shielding of ns2
electrons by
intervening d-electrons and f-electrons,
down the group, it increases. The below
elements of the group exhibit lower
oxidation States.
7. Which of the following element exhibits +
3 Oxidation State only?
a) Gallium
b) Thallium
c) Indium
d) Aluminium
Answer: d
Explanation: Boron and aluminium exhibit
oxidation state of + 3 only, while gallium,
indium and thallium exhibit oxidation states
of both +1 and +3. As we move down the
group, the tendency to exhibit + 3 Oxidation
State decreases this occurs due to the inert
pair effect.
reducing character?
a) Gallium < aluminium > indium > thallium
b) Aluminium > gallium > indium > thallium

c) Aluminium > gallium < indium > thallium
d) Gallium > aluminium > indium > thallium
Answer: b
Explanation: Reducing character of the
boron family decreases down the group
from aluminium to thallium because of the
increase in electrode potential value for
M3+
/M, therefore, the correct order is given
as aluminium > gallium > indium > thallium.
9. Complex formation is more likely to be
possible in __________
a) alkali metals
b) alkaline earth metals
c) boron family
d) equally likely
Answer: c
Explanation: The complex formation in the
boron family is greater than the S block
elements due to their smaller size and
greater charge. So they can form complexes
more likely than alkali metals and alkaline
earth metals.
10. The compounds formed by the Boron
family are __________
a) ionic
b) covalent
c) both ionic and covalent
d) neither ionic nor covalent
Answer: c
Explanation: Ionic compound formation’s
tendency increases from Boron to thallium.
Boron can only form covalent compounds,
whereas aluminium can form both covalent
as well as ionic compounds. Gallium forms
mainly ionic compounds.
11. What is the chemical formula of
aluminium carbide?
a) AlC
b) AlC3
c) AlC2
d) AC3
Answer: b
Explanation: 4 moles of aluminium atom
combines with 3 moles of carbon atom on
heating, in order to form aluminium
carbide. Aluminium carbide is ionic in
nature and it also forms methane with
water. Its chemical formula is given by AlC3.
12. When boron reacts with nitrogen which
of the following compound is formed?
a) Boron oxide
b) Boron nitrate
c) Boron hydrides
d) Boron nitride
Answer: d
Explanation: On heating, two moles of
boron atom combine with one mole of a
nitrogen molecule in order to form 2 moles
of boron nitride. Aluminium also when
reacted with nitrogen forms aluminium
nitride in the same way.
13. What forms when boron combines with
caustic soda?
a) Formation of oxygen
b) Formation of washing soda
c) Formation of Boron nitride
d) Formation of sodium borate
Answer: d
Explanation: Two moles of boron atoms
fuse with 6 moles of sodium hydroxide in
order to form 2 moles of sodium borate and
three moles of hydrogen molecules. Sodium
hydroxide is also known as caustic soda.
14. The metallic character of __________ is
less than that of alkaline earth metals.
a) Boron family
b) Alkali metals
c) Magnesium
d) Hydrogen
Answer: a
Explanation: The elements of the Boron
family are less electropositive than the
alkaline earth metals due to their smaller
size and higher ionization enthalpies. On
moving down the group, the electropositive
character first increases from Boron to

[Mn(CO)5, Cl] are isolobal (Deficiency of 1 e–
).
[Mn(CO)5]: VEC = 17 (Deficiency of 1)
O: VEC = 6 (Deficiency of 2)
Thus, [Mn(CO)5], O are not isolobal.
103. According to wade’s Rule, [C2B10H12]
adopts which type of structure?
a) closo structure
b) nido structure
c) archano structure
d) hypo structure
Answer: a
Explanation: C2B10H12 → (BH)10H2
(2(10)+2)/2 = 10+x
x = 1, Closo structure.
104. Which property is the same for isolobal
molecules?
a) e–
capture
b) Boiling point
c) Melting point
d) Solubility
Answer: a
Explanation: Isolobal molecules have same
deficiency of e–
from their stable no. Thus,
they have same e–
.
105. Which of the following is not
considered as an organometallic
compound?
a) Ferrocene
b) Cis-platin
c) Ziese’s salt
d) Grignard reagent
Answer: b
Explanation: Cis-platin does not have metal-
carbon bond. Therefore, cis platin not
considered as organometallic chemistry.
CHEPTER – 6, 7 – p – BLOCK ELEMENTS II, NOBLE GASES MCQs
Group 16 Elements MCQs
1. The ionization enthalpy and density
increase in the group from top to bottom.
a) true
b) false
Answer: b
Explanation: Density increases with increase
in atomic number due to the increase in
mass per unit volume down the group and
the ionization enthalpy decreases from
carbon to stannum, for plumbum it is
slightly higher than stannum. So the above
statement is considered to be false.
2. Why hydrides of Germanium are known
as _____________
a) silanes
b) germanes
c) stannum
d) plumbane
Answer: b
Explanation: The hydrides of carbon are
called hydrocarbons alkanes, alkenes or
alkynes, whereas the hydrides of silicon are
called silanes and the hydrides of
germanium are called Germanes the only
hydrides of stannane and plumbum are
stand and plumbane.
3. The group 14 elements form
_____________ hydrides.
a) metallic
b) ionic
c) covalent
d) both covalent and ionic
Answer: c
Explanation: All the members of group 14
form covalent hydrides, their number and
ease of formation decreases down the
group along with their thermal stability
while their reducing character increases
down the group.
4. Which of the following elements does not
belong to the Carbon family?
a) aluminium
b) silicon
c) plumbum
d) stannum

Answer: a
Explanation: The elements of Carbon family
are the elements of group 14. They are
carbon, Silicon, Germanium, stannum and
plumbum. Their valence shell configuration
is ns2np2 and their valency is four. But
aluminium belongs to group 13.
5. Which of the following group 14
elements is a metal?
a) Stannum
b) Carbon
c) Germanium
d) Silicon
Answer: a
Explanation: There are mainly five elements
in carbon family; carbon, silicon,
germanium, stannum and plumbum. The
carbon and silicon are non-metals,
germanium is a metalloid whereas
stannum, plumbum are metals.
6. What is the colour of silicon?
a) blue
b) silver
c) black
d) light brown
Answer: d
Explanation: One of the main general
physical properties of group 14 elements is
their colour. The colour of carbon is black,
silicon is light brown, germanium is greyish,
stannum is silvery white and plumbum is
also silvery white in colour.
7. What is the fajan’s rule about?
a) electronegativity
b) ionic compounds
c) Oxidation State
d) covalent compounds
Answer: c
Explanation: The Fajan’s rule is that the
compounds in +2 oxidation state are ionic in
nature and + 4 oxidation state is covalent in
nature, therefore the Fajan’s rule is about
Oxidation state and their nature of the
compounds.
8. Do Carbon family elements show multiple
bonding?
a) Yes
b) Maybe
c) No
d) Cannot say
Answer: a
Explanation: Yes, carbon forms pπ-pπ bonds
with itself and with sulphur, nitrogen and
oxygen other elements show the negligible
tendency of this type due to their larger
size. Other elements form dπ-pπ multiple
bonds.
9. All the elements in group 14 exhibit
tetravalency.
a) true
b) false
Answer: a
Explanation: In the case of carbon, 406 KJ
per mole of energy is required for
promotion of 2s electron to 2p. The
formation of two extra bonds provides this
energy, therefore we can say that all the
elements exhibit tetravalency in group 14.
So the above statement is true.
10. Is catenation possible in carbon?
a) Yes
b) Maybe
c) No
d) Cannot say
Answer: a
Explanation: Catenation is a tendency of
elements to form long chains with repeated
units of the same element. The greater the
strength of element, the greater the
strength of catenation. In the carbon family,
the catenation strength is in the decreasing
order of carbon, silicon, germanium =
stannum and plumbum.
the thermal stability of halides of Carbon

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