QUESTION NO. 6

(a) Explain:
(i) Electron affinity of chlorine is higher than fluorine.
(ii) First ionisation energy of nitrogen is higher than phosphorus.

(a) Explanation:

(i) Electron Affinity of Chlorine is Higher than Fluorine:

Electron Affinity is the amount of energy released when an atom gains an electron to form a negatively charged ion. For most elements, electron affinity increases across a period (from left to right) due to the increasing nuclear charge, which attracts the added electron more strongly. However, there are exceptions due to factors such as electron-electron repulsions and the electron configuration of the atoms.

For Chlorine (Cl) vs. Fluorine (F):

• Fluorine has a higher nuclear charge than chlorine because it is located higher up the periodic table and has more protons. This would suggest that fluorine should have a higher electron affinity.
• Chlorine, however, has a higher electron affinity than fluorine. This is due to several factors:
• Electron-Electron Repulsions: Fluorine has a smaller atomic radius than chlorine, meaning that when an electron is added to fluorine, it is added to a very small and densely packed electron cloud. This causes significant repulsion between the added electron and the electrons already present in the small 2p orbital. This repulsion makes it less favorable to add an extra electron.
• Electron Shells: Chlorine, being in the third period, has more electron shells compared to fluorine. As a result, the added electron in chlorine is less repelled by the existing electrons and is more effectively attracted to the nucleus, despite the lower nuclear charge compared to fluorine.
• Effective Nuclear Charge: Although chlorine has a lower nuclear charge than fluorine, the added electron in chlorine experiences a lower repulsion due to less crowding compared to fluorine’s smaller electron cloud.

(ii) First Ionisation Energy of Nitrogen is Higher than Phosphorus:

Ionization Energy is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom. Generally, ionization energy decreases as you move down a group in the periodic table because the outer electrons are farther from the nucleus and are shielded by inner electron shells.

For Nitrogen (N) vs. Phosphorus (P):

• Nitrogen has a higher ionization energy than phosphorus due to:
• Atomic Radius: Nitrogen is higher up in the periodic table and has a smaller atomic radius compared to phosphorus. The electrons in nitrogen are closer to the nucleus and experience a stronger attractive force, making them harder to remove.
• Shielding Effect: In phosphorus, the outer electrons experience more shielding from the inner electron shells compared to nitrogen. This reduces the effective nuclear charge felt by the outer electrons, making them easier to remove.
• Electron Configuration: Nitrogen has a half-filled p orbital configuration, which is relatively stable due to electron pairing and exchange energy. This stability contributes to its higher ionization energy as it requires more energy to disrupt this stable configuration. In contrast, phosphorus has a more complex electron configuration and less stability in its outer electrons.

In summary, while the general trend is that ionization energy decreases down a group and electron affinity increases across a period, these specific cases involve additional factors such as electron repulsion, atomic radius, and electron configuration that lead to the observed differences in electron affinity and ionization energy.

(b) Define ionisation energy and electron affinity. How these properties of an element varies in a group and period of the periodic table?

Definitions:

Ionization Energy (or Ionization Potential):

• Definition: Ionization energy is the amount of energy required to remove the most loosely bound electron from an isolated gaseous atom or ion in its ground state to form a cation.
• Symbol: Usually represented as ( \text{IE} ) or ( \text{I}_1 ) for the first ionization energy.
• Example: For an atom of sodium (Na), the first ionization energy is the energy required to remove one electron to form ( \text{Na}^+ ) (cation).

Electron Affinity:

• Definition: Electron affinity is the amount of energy released when an isolated gaseous atom gains an electron to form a negatively charged ion.
• Symbol: Usually represented as ( \text{EA} ) or ( \text{E}_a ).
• Example: For an atom of chlorine (Cl), the electron affinity is the energy released when it gains an electron to form ( \text{Cl}^- ) (anion).

Variation in a Group and Period:

**1. *Ionization Energy*:

• Across a Period (Left to Right):
• Trend: Ionization energy generally increases.
• Reason: As you move across a period from left to right, the number of protons in the nucleus increases, which increases the effective nuclear charge. This stronger attraction between the nucleus and the electrons makes it more difficult to remove an electron, thus increasing ionization energy. Additionally, the electrons are added to the same electron shell, so there is less shielding effect.
• Down a Group (Top to Bottom):
• Trend: Ionization energy generally decreases.
• Reason: As you move down a group, additional electron shells are added, increasing the atomic radius. The outer electrons are farther from the nucleus and experience increased shielding from the inner electrons, reducing the effective nuclear charge felt by the outermost electron. Consequently, it becomes easier to remove an electron, and ionization energy decreases.

**2. *Electron Affinity*:

• Across a Period (Left to Right):
• Trend: Electron affinity generally increases.
• Reason: As you move across a period, the effective nuclear charge increases, making the atom more attractive to additional electrons. This results in more energy being released when an electron is added, leading to higher electron affinity. The addition of electrons to the same shell results in less shielding effect.
• Down a Group (Top to Bottom):
• Trend: Electron affinity generally decreases.
• Reason: As you move down a group, the atomic radius increases, and the added electron is further from the nucleus and experiences increased shielding from the inner electrons. This reduces the nucleus’s ability to attract additional electrons, leading to a decrease in the amount of energy released (lower electron affinity).

In summary, ionization energy increases across a period and decreases down a group, while electron affinity typically increases across a period and decreases down a group. These trends are driven by changes in atomic structure, including nuclear charge, electron shielding, and atomic radius.