Inorganic Intro

Question 1

State Pauli’s exclusion principle

Answer 1

No two electrons in the same atom can have the same values for all the four quantum numbers

Question 2

State Hund’s rule of maximum multiplicity

Answer 2

Electrons occupy degenerate (equal energy) orbitals singly with parallel spins before electron pairing occurs in any of the orbitals

Question 3

State Aufbau’s principle

Answer 3

The orbitals are filled in order of increasing energy with the lowest filled first

Question 4

Define atomic radius

Answer 4

This is half the internuclear distance between two atoms of the same element joined by a single covalent bond

Question 5

Describe the factors that determine atomic radius of an element

Answer 5

Nuclear charge
- Increase in nuclear charge decreases atomic radius
- The higher the nuclear charge the more strongly electrons are attracted to the nucleus hence the smaller the atomic radius

Screening effect
- Increase in screening effect leads to increase in atomic radius
- The higher the screening effect, the lower the effective nuclear charge and the less strongly outer electrons attracted electrons are attracted to the nucleus; hence the larger the atomic radius

Question 6

What is the trend of atomic radius across the period?

Answer 6

Atomic radius decreases across the period from left to right

• Across the period from left to right, the nuclear charge progressively increases since a proton is added to the nucleus while screening effect remains constant since electrons are added to the same energy level
• Effective nuclear charge increases and outer electrons become more strongly attracted and closer to the nucleus hence decrease in atomic radius

Question 7

What is the trend of atomic radius down the group?

Answer 7

Atomic radius increases down the group

• Down the group, both the nuclear charge (due to addition of protons) and screening effect (due to addition of shells of electrons) increase progressively.
• However the increase in screening effect outweighs the increase in nuclear charge
• Consequently, the effective nuclear charge decreases down the group, the outer electrons become less strongly attracted to the nucleus and atomic radius increases

Question 8

Describe the trend of atomic radius among transition metals

Answer 8

• Generally, along a given series of transition elements there is only a gradual decrease in atomic radius from left to right
• Atomic radii of the first transition series decrease from Scandium to Chromium then remains almost constant till Nickel and increases from Copper to Zinc
• There in increase in nuclear charge in the beginning of the series but ad electrons continue to be filled in the d-orbitals they screen the outer 4s electrons from influence of the nuclear charge
• Increased nuclear charge and increased screening effect balance each other in the middle of the transition series and the atomic radius becomes constant
• Towards the end of the series the repulsive interaction between electrons in the orbitals becomes very dominant. There is expansion in the electron cloud and atomic radius increases

Question 9

Why is the radius of a cation smaller than that of the atom from which it is formed?

Answer 9

• In a neutral atom, the number of protons in the nucleus is equal to that of electrons outside the nucleus
• When a neutral atom loses one or more electrons in the outermost energy level, the electrons become fewer than protons
• The remaining few electrons experience greater nuclear attraction making the resulting cation have a smaller radius than that of the atom from which it is formed.
• For a given element, cationic radius decreases further with increase in number of electrons lost

Question 10

Why is the radius of an anion larger than that of the atom from which it is formed?

Answer 10

• When a neutral atom gains one or more electrons to form an anion, the number of electrons exceeds that of protons
• The electron cloud expands and electrons experience less nuclear attraction making the resulting anion have a larger radius than that of the atom from which it is formed

Question 11

Define ionization energy

Answer 11

The minimum energy required to remove one mole of electrons from one mole of free gaseous atoms.

Question 12

What is first ionization energy?

Answer 12

The minimum energy required to remove one mole of electrons from one mole of free gaseous atoms to form one mole of (unipositively charged ions) ions with a single positive charge

Question 13

Define second ionization energy

Answer 13

The minimum energy required to remove one mole of electrons from one mole of free gaseous unipositively charged ions to form one mole of ions with a double positive charge

Question 14

State the factors affecting the magnitude of ionization energy

Answer 14

1. Atomic radius
2. Screening effect
3. Nuclear Charge
4. Electronic structure/ configuration

Question 15

How does atomic radius affect first ionization energy of an element?

Answer 15

• First ionization energy decreases with increase in atomic radius
• Elements with big atomic radii> outermost electrons experience less nuclear attraction > little energy required to remove these electrons
• Elements with small atomic radii> outermost electrons experience strong nuclear attraction > a lot of energy required to remove these electrons

Question 16

How does nuclear charge affect first ionization energy of an element?

Answer 16

• First ionization energy increases with increase in nuclear charge
• High nuclear charge > outermost electrons are tightly held by the nucleus> high nuclear attraction = high ionization energy
• Low nuclear charge > outermost electrons are less strongly held by the nucleus> low nuclear attraction = low ionization energy

Question 17

How does screening effect affect first ionization energy of an element?

Answer 17

• First ionization energy decreases with increase in screening effect
• High screening effect> more effectively outer electrons are shielded from nuclear attraction> lower ionization energy
• Low screening effect> less effectively outer electrons are shielded from nuclear attraction> higher ionization energy

Question 18

How does electronic configuration affect first ionization energy of an element?

Answer 18

Completely filled or half filled orbitals tend to be very thermodynamically stable and require a lot of energy to remove an electron from them

Question 19

Describe the trend of first ionization energy in the periodic table down the group

Answer 19

• First ionization energy generally decreases down the group
• Down the group there is increase in both nuclear charge and screening effect. The progressive increase in screening effect outweighs the increase in nuclear charge causing a decrease in effective nuclear charge down the group
• Decrease in effective nuclear charge implies that the outermost shell electrons become less attracted to the nucleus requiring less energy to be removed thus low first ionization energy

Question 20

Describe the trend of first ionization energy in the periodic table across the period

Answer 20

• Generally, first ionization energy increases across the period from left to right
• From one element to another, across the period from left to right, an electron is added to the same outer shell and a proton added to the nucleus as well
• Screening effect of outer electrons by inner electrons remains almost constant while the nuclear charge progressively increases across the period
• Effective nuclear charge of outermost electrons increases across the period leading to increase in first ionization energy

Question 21

Why is the first ionization energy of Magnesium higher than that of aluminum?

Answer 21

• The electron in magnesium is obtained from a thermodynamically stable 3s sub-shell full of electrons which requires a lot of energy to be removed.
• In aluminum, the 3p subshell from which the electron is removed contains only one electron which is effectively screened from nuclear attraction by inner electrons. This makes it thermodynamically unstable and little energy is required to remove the electron.

Question 22

Why is the first ionization energy of phosphorus higher than that of sulphur?

Answer 22

• The electron in phosphorus is obtained from a thermodynamically stable half-full 3p sub-shell which requires a lot of energy to be removed.
• In sulphur the 3p sub-shell is neither half full nor full of electrons therefore it is thermodynamically unstable and little energy is required to remove the first electron.

Question 23

Define electron affinity

Answer 23

The energy liberated when each atom in one mole of free gaseous atoms gains an electron

Question 24

Explain why the second electron affinity is endothermic first is exothermic

Answer 24

When a negatively charged ion is formed its negative charge repels any further incoming electrons to be added to it such that a lot of energy must be supplied for the next electron to be added.

Question 25

Describe the factors that affect electron affinity

Answer 25

Atomic radius
- The smaller the atomic radius the stronger the incoming electron is attracted to the nucleus and the greater the electron affinity
- The larger the atomic radius the less nuclear attraction on the incoming electron and the less the electron affinity

Nuclear charge
- The greater the nuclear charge the more strongly the incoming electron is attracted to the nucleus and the greater the electron affinity
- The smaller the nuclear charge the less the incoming electron is attracted to the nucleus and the smaller the electron affinity

Screening effect
- the greater the screening effect the less the incoming electron is attracted to the nucleus and the smaller the electron affinity
- The less the screening effect the more strongly the incoming electron is attracted to the nucleus thus the larger the electron affinity

Question 26

Describe the trend of electron affinity in the periodic table down the group

Answer 26

Down the group electron affinity generally decreases

• Increase in both nuclear charge and screening effect > due to extra shells of electrons being added the screening effect outweighs the nuclear charge leading to decrease in the effective nuclear charge
• Reduced ability of atoms to attract electrons hence reduction in electron affinity

Question 27

Define electronegativity

Answer 27

The measure of the ability of an atom or element to attract the bonding pair of electrons in a covalent bond towards itself

Question 28

What factors affect electronegativity?

Answer 28

Nuclear charge
- The greater the nuclear charge the more strongly the bonding electrons are attracted and the greater the electronegativity

Screening effect
- The less the screening effect the stronger the nuclear attraction and effective nuclear charge on the bonding electrons, and the higher the electronegativity
- The higher the screening effect, the less the nuclear attraction and effective nuclear charge on the bonding electrons and the lower the electronegativity

Atomic radius
- The smaller the atomic radius, the more strongly the bonding electrons are attracted by the nucleus as the electrons are close to the nucleus and the greater the electronegativity

Question 29

What is the trend of electronegativity in the periodic table down the group?

Answer 29

Electronegativity decreases down the group

• Both nuclear charge and screening effect increase down the group, but the increase in screening effect due to extra shells outweighs the increase in nuclear charge, leading to decrease in effective nuclear charge
• This means the atoms ability to attract electrons decreases down the group

Question 30

What is the trend of electronegativity in the periodic table across the period?

Answer 30

Electronegativity increases across the period

• Across the period from left to right the nuclear charge increases from one elements to the next due to an extra proton added, but the screening effect remains constant as electrons are added to the same outer most shell.
• Therefore, the effective nuclear charge increases and the atoms ability to attract electrons increases as well hence the increase in electronegativity.

Question 31

True or false
Binary compounds are ionic if the electronegativity difference between the two bonded atoms is > 1.8

Answer 31

True

Question 32

True or false
Binary compounds are covalent if the electronegativity difference between the two bonded atoms is <=1.8

Answer 32

True

Question 33

True or false
Binary compounds are non-polar if the electronegativity difference between two bonded atoms is 0

Answer 33

True