Q.Arrange the elements N, P, O and S in the order of-

The key idea for arranging elements based on these properties is Periodic Trends.

(i) Increasing first ionisation enthalpy:

The first ionisation enthalpy generally increases across a period due to increasing effective nuclear charge and decreases down a group due to increasing atomic size and shielding effect.

1. Nitrogen (N) and Phosphorus (P) belong to Group 15, while Oxygen (O) and Sulfur (S) belong to Group 16.
2. Elements with stable half-filled p-orbitals have higher ionisation enthalpy than their immediate right neighbours. Thus, N ($2p^3$) has higher ionisation enthalpy than O ($2p^4$), and P ($3p^3$) has higher ionisation enthalpy than S ($3p^4$).
3. Down a group, ionisation enthalpy decreases. So, N $>$ P and O $>$ S.
4. Combining these trends, the overall order of increasing first ionisation enthalpy is $\text{S} < \text{P} < \text{O} < \text{N}$.

(ii) Increasing non-metallic character:

Non-metallic character generally increases across a period and decreases down a group.

1. Across a period, the effective nuclear charge increases, leading to a greater tendency to gain electrons. Therefore, O has higher non-metallic character than N, and S has higher non-metallic character than P.
2. Down a group, atomic size increases and the attraction for valence electrons decreases, reducing the tendency to gain electrons. Therefore, O has higher non-metallic character than S, and N has higher non-metallic character than P. …

Understanding the periodic trends for properties like ionisation enthalpy and non-metallic character is crucial for arranging elements. These trends are primarily governed by atomic size, effective nuclear charge, and electron configuration.

Let's first identify the positions of the given elements in the periodic table:

• Nitrogen (N): Group 15, Period 2
• Oxygen (O): Group 16, Period 2
• Phosphorus (P): Group 15, Period 3
• Sulfur (S): Group 16, Period 3

(i) Increasing First Ionisation Enthalpy

1. Concept of First Ionisation Enthalpy: First ionisation 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, forming a positive ion.

A higher ionisation enthalpy indicates that it is harder to remove an electron.

2. General Periodic Trends:

* Across a Period (left to right): First ionisation enthalpy generally increases. This is because the effective nuclear charge increases, and atomic size decreases, leading to a stronger attraction between the nucleus and the valence electrons.

* Down a Group (top to bottom): First ionisation enthalpy generally decreases. This is due to an increase in atomic size and the shielding effect of inner electrons, which reduces the attraction between the nucleus and the valence electrons.

3. Applying Trends to N, P, O, S:

   • Comparing elements in the same group:

     • For Group 15 (N and P): N is above P. So, $\Delta_i H_1(\text{N}) > \Delta_i H_1(\text{P})$.
     • For Group 16 (O and S): O is above S. So, $\Delta_i H_1(\text{O}) > \Delta_i H_1(\text{S})$. This means P and S will generally have lower ionisation enthalpies than N and O, respectively.

   • Comparing elements in the same period:

     • For Period 2 (N and O): According to the general trend, O (Group 16) should have a higher ionisation enthalpy than N (Group 15). However, there's an exception.
       Watch out

       Exception for Group 15 vs Group 16: Elements in Group 15 (like N and P) have half-filled p-orbitals ($np^3$ configuration), which are exceptionally stable. Removing an electron from this stable configuration requires more energy than removing an electron from the $np^4$ configuration of Group 16 elements (like O and S), where removing one electron leads to a stable half-filled $np^3$ configuration.

       • N has an electronic configuration of $[He]2s^22p^3$ (half-filled $2p$ orbitals).
       • O has an electronic configuration of $[He]2s^22p^4$. Therefore, $\Delta_i H_1(\text{N}) > \Delta_i H_1(\text{O})$.
     • For Period 3 (P and S): Similarly, P (Group 15) has a half-filled $3p^3$ configuration, while S (Group 16) has $3p^4$.
       • P has an electronic configuration of $[Ne]3s^23p^3$ (half-filled $3p$ orbitals).
       • S has an electronic configuration of $[Ne]3s^23p^4$. Therefore, $\Delta_i H_1(\text{P}) > \Delta_i H_1(\text{S})$.

4. Combining the Trends to Determine the Order:

   • From group trends, we know P and S are generally lower than N and O.
   • Within Period 3: $\Delta_i H_1(\text{S}) < \Delta_i H_1(\text{P})$ (due to half-filled stability of P).
   • Within Period 2: $\Delta_i H_1(\text{O}) < \Delta_i H_1(\text{N})$ (due to half-filled stability of N).
   • Now, we need to compare P with O. P is in Period 3, and O is in Period 2. Generally, elements in Period 2 have higher ionisation enthalpies than those in Period 3. So, $\Delta_i H_1(\text{O}) > \Delta_i H_1(\text{P})$.

   Putting it all together:

   • S is the lowest (Period 3, Group 16).
   • P is higher than S (Period 3, Group 15, due to half-filled stability). So, $\text{S} < \text{P}$.
   • O is higher than P (Period 2 vs Period 3). So, $\text{P} < \text{O}$.
   • N is the highest (Period 2, Group 15, due to half-filled stability). So, $\text{O} < \text{N}$.

   Thus, the increasing order of first ionisation enthalpy is $\text{S} < \text{P} < \text{O} < \text{N}$.

(ii) Increasing Non-metallic Character

1. Concept of Non-metallic Character:

   Non-metallic character refers to the tendency of an element to gain electrons to form negative ions or to share electrons. It is inversely related to metallic character. Elements with high ionisation enthalpies and high (more negative) electron gain enthalpies tend to be more non-metallic.

2. General Periodic Trends:

   • Across a Period (left to right): Non-metallic character generally increases. This is because the effective nuclear charge increases, and atomic size decreases, making it easier for the atom to attract and gain electrons. …