Chemistry · Periodic Classification Of Properties · NEET
It increases across a period (left to right). As you move right, the nuclear charge (number of protons) goes up but electrons are added to the same shell. So the atom gets smaller and the nucleus pulls the outer electrons harder. Harder pull means more energy is needed to remove an electron, so ionization enthalpy rises. Example: across period 2, Li has the lowest and Ne (a noble gas) has the highest.
Going down a group, each element has one more electron shell. The outer electron is now farther from the nucleus, and the inner filled shells shield (screen) it from the nuclear pull. A far-away, well-shielded electron is loosely held, so less energy is needed to remove it. That is why Cs (bottom of group 1) loses its electron much more easily than Li (top).
First ionization enthalpy is the energy to remove the FIRST (most loosely held) electron: X(g) → X⁺(g) + e⁻. Second ionization enthalpy removes the next electron from the +1 ion: X⁺(g) → X²⁺(g) + e⁻. The second is always larger than the first, because pulling an electron away from a positive ion is harder than from a neutral atom. Third is larger than second, and so on.
Because you must always SUPPLY energy to tear an electron away from an atom — the nucleus attracts the electron, so you fight that attraction. Energy going IN means the value is positive (endothermic). It is never negative. For NEET, remember: ionization enthalpy is always positive; electron gain enthalpy can be negative, but ionization enthalpy cannot.
They are opposite. Bigger atom = outer electron is far from nucleus = weak pull = low ionization enthalpy. Smaller atom = outer electron is close to nucleus = strong pull = high ionization enthalpy. So across a period size decreases and ionization enthalpy increases; down a group size increases and ionization enthalpy decreases. Size and ionization enthalpy move in opposite directions.
The general rule (increase across a period) has two exceptions caused by stable electron arrangements. Be > B because Be has a full 2s² which is stable. N > O because N has a half-filled 2p³ which is stable. These give the order Li < B < Be < C < O < N < F < Ne. NEET tests this exact order often — the detailed reason is covered on the anomalies page.
For the second period elements the correct increasing order of first ionisation enthalpy is:
Arrange the following elements in increasing order of first ionization enthalpy: Li, Be, B, C, N.
Which arrangement does NOT agree with the property indicated: B < C < N < O (increasing first ionization enthalpy)?
Try the real previous-year questions from this chapter — each with the answer and a full solution.
It is the energy you must give to remove one electron from a single gaseous atom in its ground state. The reaction is X(g) → X⁺(g) + e⁻. It is measured in kJ/mol and is always positive.
Helium has the highest ionization enthalpy of all elements. Among period 2 elements, neon is highest. In general, noble gases at the top-right have the highest values because they are small and have stable full shells.
Caesium (Cs) has one of the lowest, and francium the very lowest. Elements at the bottom-left of the periodic table lose electrons most easily because their outer electron is far from the nucleus and well shielded.
They are almost the same idea. Ionization enthalpy is the enthalpy change (ΔiH) for removing an electron; NCERT and NEET use the term ionization enthalpy. In practice, for NEET, treat ionization energy and ionization enthalpy as the same quantity.
Noble gases have completely filled outer shells, which is a very stable arrangement. Removing an electron would break that stability, so it needs a lot of energy. That is why they are chemically inert and rarely react.
