Q.Define electronegativity. How does it differ from electron gain enthalpy?

Electronegativity is the tendency of a bonded atom to attract the shared pair of electrons towards itself. It is a relative, dimensionless property that depends on the atom’s state in a molecule (e.g., Pauling scale).

Electron gain enthalpy ($\Delta_{eg}H$) is the energy change when an isolated gaseous atom gains an electron to form a negative ion. It is an absolute, measurable quantity with units (kJ/mol).

The key differences are:

1. Context: Electronegativity applies to atoms in a chemical bond; electron gain enthalpy applies to isolated gaseous atoms.
2. Nature: Electronegativity is a relative tendency (no units); electron gain enthalpy is an actual energy change (kJ/mol). …

Why This Distinction Matters

Many students confuse electronegativity with electron gain enthalpy because both involve an atom’s attraction for electrons. But they describe completely different situations. One is about an atom already bonded to another atom, pulling on a shared pair. The other is about an isolated atom capturing a free electron. Mixing them up leads to wrong predictions about bond polarity and reactivity.

Let’s define each clearly, then compare them point by point.

1. Defining Electronegativity

Electronegativity is a relative, dimensionless measure. It tells you how strongly an atom, when it is part of a covalent bond, pulls the bonding electrons toward itself.

• It is not a physical property you can measure directly. It is a derived scale (most commonly Pauling’s scale).
• It depends on the atom’s bonding environment — the same element can have slightly different electronegativity in different compounds (though we usually use a fixed value for simplicity).
• Example: In HCl, chlorine has higher electronegativity than hydrogen, so the shared pair is closer to Cl, giving the bond a partial ionic character.

2. Defining Electron Gain Enthalpy

Electron gain enthalpy ($\Delta_{eg}H$) is a thermodynamic quantity — the enthalpy change when one mole of isolated gaseous atoms gains one electron each to form gaseous anions.

• It is an absolute, measurable value (in kJ/mol or eV/atom).
• It applies to an isolated atom in the gas phase, not to an atom in a molecule.
• Most elements release energy when gaining an electron (exothermic, negative $\Delta_{eg}H$), but some (like noble gases) require energy input (endothermic, positive).

3. Key Differences at a Glance