KAPPA CLASSES

All four species are isoelectronic and contain 10 electrons. For an isoelectronic series, ionic radius decreases as nuclear charge increases. Mg²⁺ has the highest nuclear charge and O²⁻ has the lowest.

Actinoids are elements from actinium (Z = 89) to lawrencium (Z = 103). Terbium (Z = 65) belongs to the lanthanoid series.

The shielding or screening ability depends on the penetration of the orbital. The penetration power follows s > p > d > f. Therefore, s-electrons shield most effectively.

Ionisation enthalpy generally increases across a period. However, Mg has a filled 3s subshell, so its ionisation enthalpy is higher than that of Al. Thus the order follows Na < Mg > Al < Si.

Gadolinium has a half-filled 4f subshell and one electron in the 5d subshell. Its configuration is [Xe] 4f⁷ 5d¹ 6s².

The 4s orbital is filled before the 3d orbitals in the Aufbau order. Therefore, statement (iii) is incorrect.

Chlorine has a more negative electron gain enthalpy than fluorine because the very small size of fluorine causes greater electron-electron repulsion in its compact 2p subshell. Hence chlorine releases more energy on gaining an electron.

The period number corresponds to the highest principal quantum number (n) present in the ground-state electronic configuration of the element.

The lanthanoid series involves progressive filling of the 4f subshell.

An anion is larger than its neutral atom because of increased electron-electron repulsion, whereas a cation is smaller because it has lost an electron.

The first electron addition to oxygen is favourable, but adding a second electron to O⁻ introduces a negative electron into an already negatively charged species. Strong electron-electron repulsion makes the second electron gain process endothermic.

Consider the modern periodic table and the position of elements according to their atomic numbers.

Answer: (iii) d-block

Answer: (iii)

Answer: (iii) 126

Answer: (i)

Answer: (ii) Halogens

Neon has a stable noble-gas configuration and has the least tendency to gain an electron. Sodium also has a low tendency, while oxygen and fluorine have greater tendencies. Thus the order is A < C < B < D.

Phosphorus and sulphur can show expanded valence shells and therefore can exhibit covalencies greater than 4.

Alkali and alkaline earth metal compounds commonly produce characteristic colours in flame tests.

Representative elements belong to the s- and p-blocks. Sets (i) and (iv) contain only s- and p-block elements.

Sulphur gains an electron more readily than oxygen in its group because the larger size of sulphur reduces electron-electron repulsion. Chlorine has a highly favourable electron gain enthalpy among halogens.

Helium has the highest ionisation enthalpy. Mercury and bromine are liquids at room temperature. Alkali metals have the largest atomic radii in their respective periods.

K⁺, Ca²⁺, Sc³⁺ and Cl⁻ each contain 18 electrons. Similarly, P³⁻, S²⁻, Cl⁻ and K⁺ each contain 18 electrons.

The sequence in (ii) ignores the known exception between nitrogen and oxygen because nitrogen has a half-filled 2p³ configuration. The sequence in (iii) also does not correctly represent the electron gain enthalpy trend because chlorine has a more negative value than fluorine.

Electronegativity and metallic character are relative properties and are expressed without units.

An increase in effective nuclear charge pulls electrons closer to the nucleus, reducing atomic or ionic radius. Greater screening reduces the effective attraction and increases the radius.

The period 3, group 13 element is aluminium. Aluminium is a solid metal and is a good conductor of electricity.

Fluorine has a very small atomic size. The incoming electron has to enter the compact 2p subshell where electron-electron repulsion is relatively high. In chlorine, the electron enters the larger 3p orbital, where repulsion is lower. Therefore, more energy is released when chlorine gains an electron, making its electron gain enthalpy more negative than that of fluorine.

Transition elements are those elements whose atoms or at least one of their ions have partially filled d-orbitals. They belong to the d-block. However, elements such as Zn, Cd and Hg have completely filled d¹⁰ configurations in their atoms and common ions, so they are d-block elements but are not transition elements.

The element with Z = 119 would begin the eighth period and have the outer configuration 8s¹.

Group: 1
Valency: 1
Outer electronic configuration: 8s¹
Oxide: M₂O

It would be an alkali-metal type element.

The first ionisation enthalpy generally increases from Li to Ne because effective nuclear charge increases and atomic size decreases. Two important deviations occur:

Be > B: B loses an electron from the higher-energy 2p orbital, whereas Be loses an electron from the stable 2s orbital.
N > O: N has a stable half-filled 2p³ configuration, while O has paired electrons in one 2p orbital.

Highest first ionisation enthalpy: Carbon (C).

Carbon is smaller than silicon and therefore its valence electron is held more strongly by the nucleus. Between B and C, carbon has greater effective nuclear charge.

Most metallic: Aluminium (Al).

Metallic character increases down a group and decreases from left to right across a period. Aluminium therefore has the greatest tendency to lose electrons among the four.

The important characteristics of p-block elements are:

1. Their valence-shell electronic configuration is generally ns²np^1–6.
2. The p-block contains metals, metalloids and non-metals.
3. They show variable oxidation states.
4. They form a large number of covalent compounds and molecular compounds.

The van der Waals radius of neon is much larger than the covalent radius of fluorine. Therefore, the values are approximately 72 pm for F and 160 pm for Ne.

Transition elements show variable oxidation states because the energies of ns and (n−1)d electrons are comparable. For example, iron shows +2 and +3 oxidation states, while manganese shows +2, +3, +4, +6 and +7.

Non-transition elements can also show variable oxidation states because both s and p electrons may participate in bonding. For example, phosphorus shows +3 and +5 oxidation states, while chlorine shows +1, +3, +5 and +7 oxidation states.

Nitrogen has the stable half-filled configuration 2p³. Addition of an electron disturbs this stability, so the process is less favourable.

Oxygen has configuration 2p⁴, so gaining an electron produces a more stable arrangement relative to nitrogen and its electron gain enthalpy is negative.

For ionisation enthalpy, nitrogen has a stable half-filled 2p³ configuration. Oxygen has 2p⁴, with one pair of electrons. Repulsion between paired electrons makes removal of one electron easier. Therefore, IE₁(O) < IE₁(N).

The first member of a group often behaves differently because of its very small size, high ionisation enthalpy and high electronegativity.

Example 1 – Lithium: Lithium differs from other alkali metals and shows diagonal relationship with magnesium.

Example 2 – Boron: Boron is a non-metallic/metalloid element, whereas the other members of group 13 are more metallic. Boron also forms mainly covalent compounds.

Across a period, the acidic character of oxides generally increases, while the basic character decreases.

Basic oxides: Na₂O and MgO.
Na₂O + H₂O → 2NaOH
MgO + H₂O → Mg(OH)₂

Acidic oxides: SO₂ and SO₃.
SO₂ + H₂O → H₂SO₃
SO₃ + H₂O → H₂SO₄

Amphoteric oxides: Al₂O₃ and ZnO.
Al₂O₃ and ZnO react with both acids and bases.

Na has configuration [Ne]3s¹, while Mg has [Ne]3s².

The first electron is easier to remove from Na, so IE₁(Na) < IE₁(Mg).

After losing one electron, Na⁺ has the stable noble-gas configuration [Ne]. Removing another electron from Na⁺ requires breaking this stable configuration. Therefore, IE₂(Na) is very high.

Mg⁺ still has a 3s electron, so its second electron is easier to remove than that of Na⁺.

Exothermic process: A process in which energy is released to the surroundings. Example: electron gain by chlorine:
Cl(g) + e⁻ → Cl⁻(g)

Endothermic process: A process in which energy is absorbed from the surroundings. Example:
Cl⁻(g) → Cl(g) + e⁻

(i) Increasing first ionisation enthalpy:
S < P < N < O

P has higher ionisation enthalpy than S due to its smaller size, while N has higher ionisation enthalpy than O because of its stable half-filled 2p³ configuration.

(ii) Increasing non-metallic character:
P < S < N < O

Non-metallic character increases from left to right across a period and from bottom to top in a group.

The general trend is an increase in first ionisation enthalpy from left to right across a period. However, there are exceptions.

Be > B: B loses a 2p electron, which is higher in energy and easier to remove than the 2s electron of Be.

N > O: Nitrogen has a stable half-filled 2p³ configuration. Oxygen has one paired set of electrons in 2p orbitals, causing additional repulsion and making electron removal easier.

(a) Across a period, effective nuclear charge increases and atomic size decreases. Therefore, the attraction between the nucleus and the shared pair of electrons increases, so electronegativity increases.

(b) Down a group, the number of electron shells increases and atomic size increases. Shielding also increases. Therefore, the attraction between the nucleus and the valence electron decreases, making electron removal easier. Hence ionisation enthalpy decreases.

Across a period from left to right, effective nuclear charge increases and atomic size decreases. The tendency to lose electrons decreases, so metallic character decreases.

At the same time, the tendency to gain electrons increases, so non-metallic character increases.

Na has electronic configuration [Ne]3s¹. When it loses its 3s electron to form Na⁺, the third shell is completely removed. Na⁺ therefore has fewer occupied shells and a smaller radius than the Na atom.

Cesium (Cs) is the least electronegative alkali metal among the commonly considered stable alkali metals. Down group 1, atomic size increases and shielding increases. The valence electron is therefore held less strongly by the nucleus, so the tendency to attract shared electrons decreases.

The correct matching is:

The values correspond to elements having different tendencies to lose or gain electrons.

(i) Most reactive non-metal: B
(ii) Most reactive metal: A
(iii) Least reactive element: D
(iv) Metal forming binary halide: C

A noble-gas configuration is highly stable and therefore has an unfavourable electron gain enthalpy. A halogen with 2p⁵ configuration has a strongly negative electron gain enthalpy. Oxygen has 2p⁴, while sodium has 3s¹.

Nitrogen has the stable half-filled 2p³ configuration, whereas oxygen has paired electrons in one of its 2p orbitals. Electron-electron repulsion makes electron removal from oxygen easier. Hence IE₁(N) > IE₁(O).

Across a period, nuclear charge increases while electrons are added to the same principal shell. Therefore, effective nuclear charge increases and attracts the electrons more strongly towards the nucleus. Consequently, atomic radius decreases.

Down a group, atomic size increases and the incoming electron experiences weaker attraction from the nucleus. Consequently, less energy is released on electron addition and electron gain enthalpy generally becomes less negative.

Electron gain enthalpy is the enthalpy change accompanying the addition of an electron to an isolated gaseous atom.

Factors affecting electron gain enthalpy:

1. Atomic size: As atomic size increases, the incoming electron is farther from the nucleus and the attraction decreases. Therefore, electron gain enthalpy generally becomes less negative.

2. Effective nuclear charge: Higher effective nuclear charge attracts the incoming electron more strongly and makes electron gain enthalpy more negative.

3. Electronic configuration: Stable configurations such as completely filled and half-filled subshells oppose the addition of an electron.

Periodic trend: Across a period, electron gain enthalpy generally becomes more negative because effective nuclear charge increases and atomic size decreases.

Down a group, electron gain enthalpy generally becomes less negative because atomic size and shielding increase.

Exception: The electron gain enthalpy of chlorine is more negative than that of fluorine because fluorine's very small 2p orbital produces considerable electron-electron repulsion.

Ionisation enthalpy is the minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.

For example:
M(g) → M⁺(g) + e⁻

Factors affecting ionisation enthalpy:

1. Atomic size: Larger atomic size means the valence electron is farther from the nucleus and is easier to remove. Thus ionisation enthalpy decreases.

2. Nuclear charge: Greater nuclear charge increases attraction for the electron and increases ionisation enthalpy.

3. Shielding effect: Greater shielding reduces the effective nuclear attraction and lowers ionisation enthalpy.

4. Electronic configuration: Half-filled and completely filled subshells are relatively stable and require more energy for electron removal.

Across a period: Ionisation enthalpy generally increases because effective nuclear charge increases and atomic radius decreases.

Down a group: Ionisation enthalpy generally decreases because atomic size and shielding increase.

Important exceptions include Be > B and N > O.

The modern periodic law states that the physical and chemical properties of elements are periodic functions of their atomic numbers.

Atomic number represents the number of protons in the nucleus and determines the electronic configuration of an element. Since the electronic configuration determines chemical properties, elements with similar outer electronic configurations show similar properties.

The arrangement according to atomic number also removes several anomalies present in Mendeleev's atomic-mass-based table.

For example, isotopes have different masses but the same atomic number and therefore occupy the same position in the modern periodic table.

Alkali metals have one electron in their outermost shell. Therefore, their general valence-shell configuration is ns¹.

This single valence electron can be easily lost to form a stable noble-gas configuration:

M → M⁺ + e⁻

As they have one valence electron and commonly show +1 oxidation state, they are placed in group 1 of the modern periodic table.

Examples include Li, Na, K, Rb and Cs.

The important limitations of Mendeleev's periodic table were:

1. Position of hydrogen: Hydrogen resembles both alkali metals and halogens, so its position was uncertain.

2. Anomalous pairs: Some elements had to be placed in positions that did not strictly follow increasing atomic mass.

3. Isotopes: Isotopes have different atomic masses but similar chemical properties. Mendeleev's table could not provide separate positions for isotopes.

4. Cause of periodicity: Mendeleev's periodic law did not explain the fundamental reason for periodic repetition of properties.

5. Rare earth elements: The proper position of lanthanoids was not clearly established.

The long-form periodic table is based on atomic number and electronic configuration, making it more logical and systematic.

Its advantages include:

1. Elements are arranged according to increasing atomic number.
2. Elements with similar valence-shell configurations are placed in the same group.
3. It explains periodicity in terms of electronic configuration.
4. Isotopes occupy the same position because they have the same atomic number.
5. The table clearly separates s-, p-, d- and f-block elements.
6. It accommodates transition and inner-transition elements systematically.
7. The anomalous pairs in Mendeleev's table are eliminated.

Group 1 elements have one valence electron and have relatively low first ionisation enthalpies. Group 17 elements have seven valence electrons and are much more strongly attracted to the nucleus, so they have high first ionisation enthalpies.

Across a period: Ionisation enthalpy generally increases from group 1 to group 17 because effective nuclear charge increases and atomic size decreases.

Down group 1: Ionisation enthalpy decreases because atomic size and shielding increase. The outer electron becomes easier to remove.

Down group 17: Ionisation enthalpy also decreases because the valence shell moves farther from the nucleus and shielding increases.

Thus, group 17 elements have much higher first ionisation enthalpies than group 1 elements in the same period, while both groups show a decreasing trend down their respective groups.