Animated Solution for Chemistry - Periodicity in Properties: The correct option with respect to the Pauling electronegativity values of the elements is

Imagine a covalent bond as a tug-of-war between two atoms over a shared pair of electrons. Electronegativity is simply the strength of an atom in this microscopic game. The stronger the pull, the higher the electronegativity.

To master questions like this, we don't need to memorize the exact Pauling scale values. We just need to understand the two golden rules of the periodic table.

Across a Period

As we travel from left to right across a period, the atomic number increases. This means more protons are added to the nucleus, increasing the effective nuclear charge ($Z_{\text{eff}}$).

Because the new electrons are added to the same principal shell, the shielding effect doesn't increase enough to counteract the stronger nuclear pull. The nucleus grips the outer electrons tighter, shrinking the atomic radius and significantly boosting the atom's ability to attract shared electrons.

Therefore, electronegativity increases across a period.

Down a Group

Now, let's take an elevator down a group. With each step down, a completely new electron shell is added. The atom gets bulkier.

Even though the number of protons increases, the inner shells act like a thick blanket, shielding the outermost electrons from the nucleus's pull. Because the valence electrons are further away and well-shielded, the atom's grip on shared electrons weakens.

Therefore, electronegativity decreases down a group.

Armed with our two rules, let's dissect the options given in the problem.

Option (a): $P>S$ Phosphorus ($P$) and Sulphur ($S$) are neighbors in Period 3. Sulphur sits to the right of Phosphorus. According to our horizontal rule, Sulphur must have a higher electronegativity. So, $S>P$, making this option incorrect.

Option (b): $Si<Al$ Silicon ($Si$) and Aluminum ($Al$) also reside in Period 3. Silicon is to the right of Aluminum. Thus, Silicon pulls harder on electrons, meaning $Si>Al$. This option is also incorrect.

Option (c): $Te>Se$ Tellurium ($Te$) and Selenium ($Se$) belong to Group 16 (the Oxygen family). Tellurium is located below Selenium. Following our vertical rule, the larger Tellurium atom has a weaker pull. Therefore, $Se>Te$, making this option incorrect.

Option (d): $Ga<Ge$ Finally, we look at Gallium ($Ga$) and Germanium ($Ge$) in Period 4. Germanium is positioned to the right of Gallium. The increased effective nuclear charge means Germanium is more electronegative.

This gives us the relationship $Ga<Ge$, which perfectly matches the option!

By simply visualizing the periodic table and applying the fundamental trends of effective nuclear charge and atomic size, we can confidently conclude that Gallium is less electronegative than Germanium.