Classification of Elements and Periodicity in Properties

Question 1

s-block

Answer 1

Groups 1 and 2

Question 2

p-block

Answer 2

Groups 13-18

Question 3

d-block

Answer 3

Groups 3-12

Question 4

f-block

Answer 4

Lanthanoids and Actinoids at the bottom of the periodic table

Question 5

Normal/respresentative elements

Answer 5

s-block and p-block elements

Question 6

Atomic radius- definition

Answer 6

The distance between the centre of the nucleus and the valence shell of electrons is known as the atomic radius

Question 7

Crystal radius

Answer 7

One half of the distance between the centre of the nuclei of two adjacent atoms in a metallic nucleus

Question 8

Covalent radius

Answer 8

One half of the distance between the centres of two nuclei of two similar atoms bonded together by a single covalent bond.

Question 9

van der Waals’ radius

Answer 9

One half of the distance between the nuclei of two atoms of the same substance at their closest approach

Question 10

How does atomic radius vary
1. across a period
2. down a group
+ why?

Answer 10

1. Atomic radius decreases because no. of shells remain the same but the nuclear charge increases, pulling the electron cloud closer to the nucleus.
2. Atomic radius increases because no. of shells increase and this takes the valence shell further away from the nucleus, decreasing the effective nuclear charge.

Question 11

Ionic radius- definition

Answer 11

The distance between the centre of the nucleus and the point up to which the nucleus has an influence on the electron cloud of the ion.

Question 12

Why is the radius of a cation always less than the radius of the atom from which it is formed?

Answer 12

• Cations are positively charged ions and are formed when a neutral atom loses electrons.
• The number of valence electrons decrease but the positive charge in the nucleus remains the same.
• Due to less number of electrons and same nuclear charge, the effective nuclear charge on the remaining electrons is more, causing them to come closer to the nucleus.

Question 13

Why is the radius of an anion always more than the radius of the atom from which it is formed?

Answer 13

• An anion possesses same nuclear charge but more electrons than the neutral/parent atom.
• This leads to a decrease in the effective nuclear charge i.e. the nucleus exerts less influence on the valence electrons.
• The valence electrons move away and the size increases due to the expansion of the electron cloud.

Question 14

Isoelectronic ions

Answer 14

Ions of different elements which contain the same number of electrons are called isoelectronic elements

Question 15

Relationship between the ionic radii of isoelectronic ions and nuclear charge

Answer 15

ionic radii of isoelectronic compounds decrease with increase in nuclear charge.

Question 16

Ionisation enthalpy or ionisation energy

Answer 16

The amount of energy required to remove a loosely bound electron from the outermost shell of an isolated gaseous atom to form a gaseous ion

Question 17

Unit of ionisation enthalpy

Answer 17

kJ/mol

Question 18

Ionisation potential

Answer 18

The voltage at which the ionisation of a gas occurs

Question 19

Why is the second ionisation energy higher than the first ionisation energy?

Answer 19

This is because the removal of a loosely bound electron from the valence shell of the ion forms a monopositive ion, whose ionic radius is much less than the atomic radius of the parent atom. Due to smaller size of the cation, the remaining electrons will experience a greater pull of the nucleus. Therefore, higher energy will be required to remove another electron.

Question 20

How does ionisation energy depend upon:

1) atomic radius
2) nuclear charge

Answer 20

1) It decreases with an increase in atomic radius

2) It decreases with a decrease in nuclear charge

Question 21

What is screening/shielding effect?

Answer 21

In a multielectron atom, the electrons present between the valence electron and the nucleus shield the valence electron from the nucleus.
This decreases the effect of the nucleus on the valence electron i.e. the nucleus exerts less force of attraction on the valence electron.

Question 22

Dependence of ionisation energy on shielding effect.

Answer 22

If other factors remain the same, ionisation energy increases with a decrease in shielding effect.

Question 23

What is penetration effect?

Answer 23

In a multielectron atom, the electrons do not maintain distinct boundaries. Instead the electron cloud of one electron penetrates into the electron cloud of another inner electron. This is known as penetration effect.
Due to this effect, electrons move towards the nucleus.

Question 24

Dependence of ionisation energy on penetration power.

Answer 24

Increases with increase in penetration power

Question 25

Why are alkali metals highly reactive?

Answer 25

They have extremely low ionisation energies and get easily ionised, making them highly reactive.

Question 26

Variation of ionisation energy across a period + why

Answer 26

Increases across a period because:

i) nuclear charge increases
ii) size of the atom decreases

Question 27

Variation of of ionisation energy down a group + why

Answer 27

Ionisation energy decreases down a group because:

i) atomic radius increases
ii) no. of inner electrons increase so shielding effect increases

Question 28

Electron affinity

Answer 28

The amount of energy released on the addition of an electron to the outermost shell of an isolated gaseous atom to form a gaseous anion.

Question 29

Dependence of electron gain enthalpy on:

1. nuclear charge
2. size of the atom

Answer 29

1. electron gain enthalpy increases with increase in nuclear charge
2. electron gain enthalpy increases with decrease in atomic size

Question 30

Variation of electron gain enthalpy:

1. across a period
2. down a group

Answer 30

1. becomes more negative as we move across the period

2. becomes less negative as we move down a group

Question 31

Why do halogens have the most negative electron gain enthalpies?

Answer 31

• Halogens are elements of group 17.
• Their valence electronic configuration is of ns²np⁵ type which is similar to that of noble gases (ns²np⁶)
• Halogens have a very strong tendency to accept electrons.
• Therefore they have the most negative electron gain enthalpies.

Question 32

Why is the electron gain enthalpy of chlorine more negative than that of fluorine?

Answer 32

• In F, the added electron goes to the 2p subshell while in Cl it goes to the 3p subshell.
• The 2p subshell is more compact compared to the 3p subshell as the 2p subshell is closer to the nucleus.
• Due to the small size of the F atom the inter-electronic repulsion in 2p subshell of F is much larger as compared to that of the 3p subshell of Cl.
• Therefore, the added electrons find it easier to enter into the Cl atom as compared to the F atom.

Question 33

Why is the electron gain enthalpy of nitrogen positive?

Answer 33

N possesses a half filled shell. Due to this relatively stable electronic configuration, nitrogen atom does not accept the added electron easily.
the addition of an electron to a nitrogen atom is only possible when external energy is supplied to it

Question 34

Electronegativity

Answer 34

The tendency of an atom to attract the bonding or shared pair of electrons towards itself in a covalent bond is called the electronegativity of that atom.

Question 35

Dependence of electronegativity on:

1. Effective nuclear charge
2. Size of the atom
3. Oxidation state
4. State of hybridization

Answer 35

1. More the effective nuclear charge, more the electronegativity
2. Less the size of the atom, more the electronegativity
3. Greater the oxidation state. more the electronegativity
4. Greater the s-character in the hybridization state of an atom, more the electronegativity