Q.(a) Give reasons:

The Oxygen Family: An Intuitive Start

Imagine you are building a molecule. You have a central atom, and around it you want to attach other atoms. Some atoms are greedy for electrons — they want to pull negative charge toward themselves. Others are more willing to share.

Now think about the element oxygen. It is the most electron-hungry element in its entire family. When oxygen bonds, it almost always grabs two electrons and settles into a −2 oxidation state. It does not easily give up its own electrons because it holds them so tightly.

But move down the family — to sulfur, selenium, tellurium, polonium — and something changes. These larger atoms are less greedy. They can afford to share some of their own electrons with highly electronegative partners like oxygen or fluorine. That is why sulfur can form $+4$ (as in $SO_2$) and $+6$ (as in $SO_3$) states, while oxygen itself almost never does.

The Precise Statement

Group 16 of the periodic table — oxygen (O), sulfur (S), selenium (Se), tellurium (Te), and polonium (Po) — is called the oxygen family or chalcogens. Their general electronic configuration is $ns^2 np^4$, meaning they have six valence electrons and need two more to complete an octet.

The common oxidation states are:

Why Oxygen is Anomalous

Oxygen breaks the family pattern for two interconnected reasons.

Small size. Oxygen is the second-smallest atom in the periodic table (after fluorine). Its valence shell is the $n=2$ level, which has no available $d$-orbitals. To reach a $+4$ or $+6$ state, an atom must expand its octet — promote electrons to empty $d$-orbitals. Oxygen simply cannot do this. Sulfur and below have accessible $3d$, $4d$, or $5d$ orbitals, so they can.

High electronegativity. Oxygen (3.44 on the Pauling scale) is the second most electronegative element. It pulls bonding electrons so strongly toward itself that it almost always ends up negative. When bonded to fluorine — the only element more electronegative — oxygen can show a positive oxidation state ($+2$ in $OF_2$), but this is the exception that proves the rule. …

What Makes a Halogen a Halogen?

Imagine you are a fluorine atom. You have seven electrons in your outermost shell. You are one electron short of a full, stable octet — that perfect, noble-gas configuration that every atom craves. That single missing electron makes you intensely hungry. You will grab an electron from almost anything that moves.

That hunger is the entire story of the halogens.

The Intuition: The One-Electron Gap

All elements in Group 17 — fluorine, chlorine, bromine, iodine, and astatine — share the same electronic signature: their outermost shell has seven electrons. The nearest noble gas has eight. So each halogen is exactly one electron short of stability.

This is not a small difference. It is the most powerful chemical drive in the periodic table. It means:

• Halogens are the most reactive non-metals in their respective periods.
• They exist naturally as diatomic molecules ($\text{F}_2$, $\text{Cl}_2$, $\text{Br}_2$, $\text{I}_2$) because two halogen atoms can share one electron each, giving both a pseudo-octet.
• When they react with metals, they gain one electron to become halide ions ($\text{F}^-$, $\text{Cl}^-$, $\text{Br}^-$, $\text{I}^-$), achieving the noble-gas configuration. The resulting compounds are called salts — sodium chloride, potassium iodide, calcium fluoride.

The Precise Statement

Group 17 elements (fluorine, chlorine, bromine, iodine, astatine) are highly reactive non-metals with the general electronic configuration $ns^2 np^5$ in their valence shell. Their characteristic oxidation state is −1, achieved by gaining one electron to form a halide ion. In their elemental form, they exist as diatomic molecules ($\text{X}_2$). Their reactivity decreases down the group: fluorine is the most reactive, iodine the least.

The Trend Down the Group

Why −1 and Not +1 or +7?

You might ask: if halogens have seven valence electrons, could they not lose seven electrons and show a +7 oxidation state? In principle, yes — chlorine, bromine, and iodine do show positive oxidation states (+1, +3, +5, +7) when bonded to more electronegative elements like oxygen. But the characteristic oxidation state, the one that defines their salt-forming behaviour, is −1.

The reason is simple: gaining one electron is energetically far cheaper than losing seven. The energy required to remove seven electrons is enormous; the energy released when one electron is added is substantial. So whenever a halogen meets a metal, the metal loses electrons and the halogen gains one. That is the fundamental exchange. …