Periodic Trends in Properties of Elements: Affecting Factors, Examples and Applications

There are two periodic trends related to atomic radius which are discussed below. NCERT solutions of this chapter covers questions related to this topic:

1. Across a Period: Atomic radius decreases due to increasing effective nuclear charge ( $Z_{\text{eff}}$ ), which pulls electrons closer to the nucleus. For example, from Li to F in period 2, radius decreases from 152 pm to 64 pm (NCERT, p. 13.12, Fig. 13.3).

2. Down a Group: Atomic radius increases as the number of electron shells increases, outweighing the effect of increased nuclear charge. For instance, from Li to Cs in group 1, radius increases from 152 pm to 262 pm.

The following factors affect the atomic radius:

• Effective Nuclear Charge: Higher $Z_{\text{eff}}$ reduces radius by increasing nuclear attraction (Radius $\propto \frac{1}{Z_{\text{eff}}}$ ).
• Number of Shells: More shells increase the distance from the nucleus, enlarging the radius.
• Bond Multiplicity: Covalent radius decreases with bond order (e.g., $\mathrm{C}\mathrm{C}:77\mathrm{p}\mathrm{m},\mathrm{C}=\mathrm{C}$ : $67\mathrm{p}\mathrm{m},\mathrm{C}\mathrm{C}:60\mathrm{p}\mathrm{m}$ ).
• Ionic Character: Increased ionic character shortens bond length, reducing covalent radius.

Ionization energy (IE) is the energy required to remove the outermost electron from a neutral gaseous atom ( $M\left(g\right)->M+\left(g\right)+e-$ ). Successive ionization energies ( $I{E}_1,I{E}_2,I{E}_3$ ) increase as electrons are removed, with $I{E}_3>I{E}_2>I{E}_1$ (NCERT, p. 13.14).

The following trends is applicable to ionization energy:

1. Across a Period: IE increases due to higher $Z_{\text{eff}}$ and smaller atomic size. Exceptions occur due to stable configurations, e.g., Be (fully filled 2s) has higher IE than B, and N (half-filled 2p) has higher IE than O (NCERT, p. 13.15).

2. Down a Group: IE decreases as atomic size increases and shielding effect reduces $Z_{\text{eff}}$ . Exceptions include Al to Ga, where IE remains similar due to d-orbital shielding, and Ti to Hf, where lanthanide contraction increases IE for Hf.

Electron affinity (EA) is the energy released when a neutral gaseous atom gains an electron $\left(X\right(g)+e-->X-(g\left)\right)$ . The first EA ( $E{A}_1$ ) is positive, while subsequent EAs (e.g., $E{A}_2$ ) are negative due to electron repulsion. While we have provided the definition in this section, exams like NEET and IISER ask conceptual questions based on the topic.

Electronegativity measures an atoms ability to attract shared electrons in a covalent bond. Fluorine is the most electronegative element ( 4.0 on Pauling scale), while cesium is the least (0.7) (NCERT, p. 13.20).

In summary, the following are the remaining periodic trends followed across the table:

1. Metallic Character: This character decreases across a period, increases down a group.

2. Melting/Boiling Points: For metals, the boiling point and melting point increase across a period (e.g., Na: 98řC, Al: 660řC).

3. Oxide Nature: The basicity decreases, acidity increases across a period (e.g., Na 2 O basic, Cl2O7 acidic).