Q.Which of the following statements are correct? (Note: more than one of the given options may be correct.)

The key idea here involves understanding Periodic Properties and Trends of elements.

Let's evaluate each statement:

(A) Ionisation enthalpy generally increases across a period and decreases down a group. Helium, being the smallest atom with the highest effective nuclear charge and a stable duplet configuration, requires the most energy to remove an electron. This statement is correct.

(B) Electron gain enthalpy generally becomes more negative across a period. However, due to the very small size of fluorine, there are significant interelectronic repulsions in its compact $2p$ subshell when an electron is added. This makes the addition of an electron less favourable compared to chlorine, where the incoming electron goes into the larger $3p$ subshell. Therefore, chlorine has a more negative electron gain enthalpy than fluorine (e.g., $\text{Cl} = -349 \text{ kJ/mol}$, $\text{F} = -328 \text{ kJ/mol}$). The statement claims chlorine has less negative electron gain enthalpy, which is incorrect. …

Understanding periodic trends is fundamental in chemistry. These trends allow us to predict and explain the properties of elements based on their position in the periodic table. The statements in this question directly relate to these trends and some important exceptions or specific facts.

Here's a step-by-step evaluation of each statement:

1. Statement (A): Helium has the highest first ionisation enthalpy in the periodic table.

   • Concept: First ionization enthalpy ($\Delta_i H_1$) is the minimum energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state.
   • Periodic Trend: Ionization enthalpy generally increases across a period from left to right due to increasing effective nuclear charge and decreasing atomic size, which means the valence electrons are held more tightly. It generally decreases down a group due to increasing atomic size and shielding effect, which means the valence electrons are further from the nucleus and experience less attraction.
   • Noble Gases: Noble gases (Group 18) have completely filled valence electron shells, giving them exceptional stability. This stability means a very large amount of energy is required to remove an electron from them.
   • Helium's Position: Helium is the first element in Group 18 and the smallest noble gas. Its two electrons are in the 1s orbital, very close to the nucleus, and experience the highest effective nuclear charge for their principal energy level. This combination of small size, high effective nuclear charge, and stable electron configuration makes it extremely difficult to remove an electron from helium.
   • Conclusion: Therefore, helium indeed has the highest first ionization enthalpy among all elements in the periodic table.
   • Statement (A) is correct.

2. Statement (B): Chlorine has less negative electron gain enthalpy than fluorine.

   • Concept: Electron gain enthalpy ($\Delta_{eg} H$) is the energy change when an electron is added to a neutral gaseous atom to form a negative ion. A more negative value indicates a greater tendency for the atom to accept an electron.
   • Periodic Trend: Electron gain enthalpy generally becomes more negative across a period (due to increasing effective nuclear charge) and less negative down a group (due to increasing atomic size). Halogens (Group 17) typically have very negative electron gain enthalpies because they are one electron short of a stable noble gas configuration.
   • Fluorine vs. Chlorine Anomaly: While fluorine is above chlorine in Group 17 and is expected to have a more negative electron gain enthalpy due to its smaller size, an anomaly exists. Fluorine has a very small atomic size. When an incoming electron approaches a fluorine atom, it experiences significant interelectronic repulsion from the already existing electrons in the compact 2p subshell. This repulsion makes the addition of an electron less favorable (less exothermic, or less negative) compared to chlorine. In chlorine, the incoming electron enters the larger 3p subshell, where electron-electron repulsion is less significant.
   • Values: The electron gain enthalpy of fluorine is approximately $-328 \text{ kJ/mol}$, while that of chlorine is approximately $-349 \text{ kJ/mol}$.
   • Comparison: Since $-328 \text{ kJ/mol}$ is numerically smaller (closer to zero) than $-349 \text{ kJ/mol}$, fluorine has a less negative electron gain enthalpy than chlorine.
   • The statement says "Chlorine has less negative electron gain enthalpy than fluorine". This means $\Delta_{eg} H(\text{Cl}) > \Delta_{eg} H(\text{F})$. However, $-349 \text{ kJ/mol}$ is not greater than $-328 \text{ kJ/mol}$. In fact, chlorine has a more negative electron gain enthalpy than fluorine.
   • Statement (B) is incorrect.

3. Statement (C): Mercury and bromine are liquids at room temperature. …