12.1 Halogens

The elements in Group 17 (or Group VII-A) of the periodic table are known as the halogens. The name is derived from the Greek words halous (meaning "salt") and gen (meaning "former"), as they readily react with metals to form salts.

The halogens include:

• Fluorine (F)
• Chlorine (Cl)
• Bromine (Br)
• Iodine (I)
• Astatine (At) (radioactive)
• Oganesson (Og) (artificial radioactive element; properties largely unknown)

All halogens are reactive non-metals and exist as diatomic molecules ($X_2$) at standard conditions. Their Electronic Configuration→ follows a specific pattern.

Physical States and Appearance

12.1.1 Volatility of Halogens

As you move down Group 17, the halogens become less volatile and their colours become darker.

Trend: Volatility decreases down the group: $F_2>C{l}_2>B{r}_2>I_2$

Reason: Halogen molecules ($X_2$) are non-polar. The only intermolecular forces present are weak instantaneous dipole–induced dipole forces (London dispersion forces).

• As you go down the group, the number of electrons and the size of the electron cloud increase.
• Larger electron clouds are more polarisable, creating stronger temporary dipoles.
• Stronger London forces require more energy to overcome, raising melting and boiling points.
• This is why $B{r}_2$ is a liquid and $I_2$ is a solid at room temperature, while $F_2$ and $C{l}_2$ are gases.

12.1.2 Bond Strength of Halogens

The bond energy is the energy required to break the covalent bond in one mole of diatomic halogen molecules ($X-X$).

General Trend: From $C{l}_2$ to $I_2$, bond energy decreases as atomic radius increases, making the bond longer and weaker.

Fluorine Anomaly: The $F-F$ bond energy (156 kJ/mol) is unexpectedly lower than the $Cl-Cl$ bond energy (243 kJ/mol). This is because fluorine atoms are very small, so the lone pairs on the two fluorine atoms are forced very close together, causing significant inter-electronic repulsion that weakens the covalent bond.