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Chemistry · Classification of Elements and Periodicity

Electron Gain Enthalpy

Electron gain enthalpy, written $\Delta_{eg}H$, measures the energy change when a neutral gaseous atom captures an electron to become an anion. NCERT Class 11 Chemistry (Unit 3, §3.7.1) develops it as one of the four core periodic properties, and it is among the most error-prone of them: the sign convention reverses what students expect, and the textbook trend is broken by fluorine itself. For NEET this single subtopic supplies a recurring class of questions — the most-negative-$\Delta_{eg}H$ element and the F vs Cl ordering — so precision on signs and the anomaly is non-negotiable.

Definition and the Defining Equation

When an electron is added to a neutral gaseous atom $\ce{X}$ to convert it into a uninegative ion, the enthalpy change accompanying the process is defined as the electron gain enthalpy, $\Delta_{eg}H$. It is a quantitative measure of the ease with which an atom accepts an electron to form an anion. The defining process is written exactly as NCERT states it:

$$\ce{X(g) + e- -> X^-(g)} \qquad \Delta_{eg}H$$

The qualitative driving force behind the whole property is the universal tendency of an atom to acquire a noble-gas configuration. The NIOS text frames it directly: an atom with five, six or seven electrons in its outermost shell shows a tendency to accept electrons and reach the nearest noble-gas arrangement. Halogens, with seven valence electrons, need just one more electron, so they release the most energy on electron capture; metals, which would have to take several, instead show high positive values.

For the halogens, this is the step that takes a neutral atom to its stable closed-shell anion. Taking chlorine, the NCERT/NIOS-documented value is:

$$\ce{Cl(g) + e- -> Cl^-(g)} \qquad \Delta_{eg}H = -349~\text{kJ mol}^{-1}$$

Two features of this definition matter for every problem you will meet. First, $\Delta_{eg}H$ is defined for the gaseous state, so lattice or hydration effects are excluded — it is a property of the isolated atom. Second, the magnitude reflects how strongly the nucleus can hold one extra electron once it has been pulled in; the sign tells you whether the atom wanted that electron at all.

Figure 1 · Process schematic

A neutral gaseous atom captures a free electron and releases (or, for noble gases, absorbs) energy. The colour of the energy arrow encodes the sign of $\Delta_{eg}H$.

X neutral atom (g) e⁻ incoming X⁻ anion (g) energy released → ΔegH < 0

The Sign Convention — Read This Twice

This is the single most confused point in the entire subtopic, and it is where NEET sets its traps. Because adding an electron usually releases energy, and because $\Delta_{eg}H$ follows the standard thermodynamic convention in which released energy is negative, a more negative value means a greater tendency to accept an electron — the opposite of the intuitive "bigger number is better" reasoning.

NEET Trap · Sign of ΔegH

Negative means energy released, which means a stronger tendency to gain an electron.

An exothermic electron capture is assigned a negative $\Delta_{eg}H$. So when a question asks for the element with the most negative electron gain enthalpy, it is asking for the element that releases the most energy and accepts an electron most readily — that is chlorine, $-349~\text{kJ mol}^{-1}$. Do not pick fluorine just because it is the most electronegative element.

More negative $\Delta_{eg}H$ = more energy released = stronger electron-accepting tendency. A positive $\Delta_{eg}H$ (noble gases) means energy must be supplied — the atom resists the electron.

Group 17 elements have very high negative electron gain enthalpies precisely because by picking up one electron they reach a stable noble-gas configuration. By contrast, the closer an element already is to a filled or stable arrangement on its own, the weaker its drive to take an extra electron, pushing $\Delta_{eg}H$ toward less negative — and for the noble gases, positive — values.

Electron Gain Enthalpy Values (NCERT Table 3.7)

The numbers below are taken verbatim from NCERT Table 3.7. Commit the halogen and noble-gas rows to memory; almost every PYQ on this subtopic is decided by them. All values are in $\text{kJ mol}^{-1}$.

Group 1ΔegH Group 16ΔegH Group 17ΔegH Group 18ΔegH
H−73————He+48
Li−60O−141F−328Ne+116
Na−53S−200Cl−349Ar+96
K−48Se−195Br−325Kr+96
Rb−47Te−190I−295Xe+77
Cs−46Po−174At−270Rn+68

Read the table column by column. The halogens are sharply negative and the noble gases are uniformly positive — a contrast that captures the whole physical picture in one glance. Within group 1 and group 16 the values become less negative on descending (Li to Cs; O to Po), which is the normal group trend. Group 17 obeys the same descent except at its very top, where F is less negative than Cl — the anomaly examined below.

NCERT cautions that the variation in electron gain enthalpies is less systematic than that of ionisation enthalpies, so these are general rules with documented exceptions rather than rigid laws.

DirectionGeneral behaviour of ΔegHReason (per NCERT §3.7.1)
Left → right across a period Becomes more negative Effective nuclear charge increases and atomic size decreases, so the added electron sits closer to the nucleus and is bound more strongly.
Top → bottom down a group Becomes less negative Atomic size increases, so the incoming electron enters a level farther from the nucleus and is held more weakly.

NCERT also notes that electron gain enthalpies reach their largest negative values toward the upper-right region of the periodic table, just before the noble gases — the halogen corner. This is the same nuclear-pull logic that drives ionisation enthalpy and electronegativity trends, all three rooted in how tightly a shrinking atom grips its valence region.

Figure 2 · Trend directions

Arrows show the direction in which $\Delta_{eg}H$ becomes more negative. The strongest electron-accepting corner is upper-right, at the halogens; the broken arrow at the top of group 17 marks the F < Cl reversal.

across period → ΔegH more negative down group → less negative most −ve corner noble gas +ve dashed = F < Cl reversal

The Fluorine vs Chlorine Anomaly

On the simple group rule, fluorine — the topmost, smallest halogen — should release the most energy on gaining an electron. It does not. The NCERT values are $\Delta_{eg}H(\ce{F}) = -328$ and $\Delta_{eg}H(\ce{Cl}) = -349~\text{kJ mol}^{-1}$, so chlorine is more negative than fluorine, and chlorine in fact holds the most negative electron gain enthalpy of any element.

$$\ce{F(g) + e- -> F^-(g)}\,,\ \Delta_{eg}H = -328 \qquad \ce{Cl(g) + e- -> Cl^-(g)}\,,\ \Delta_{eg}H = -349$$

The reason is electron-electron repulsion in a compact orbital. Fluorine is so small that its incoming electron is forced into the tight $n = 2$ (2p) level, where it experiences significant repulsion from the electrons already crowded into that small region of space. In chlorine the added electron enters the larger $n = 3$ (3p) level, occupying a much bigger volume where electron-electron repulsion is far weaker. The reduced repulsion in chlorine more than compensates for its larger size, so chlorine releases more energy.

NEET Trap · F vs Cl

Chlorine, not fluorine, has the most negative electron gain enthalpy.

The standard answer key reasoning (NEET 2016) is: "Electron gain enthalpy of F is less than Cl because in F the small size of the 2p orbital results in high electron density, so inter-electronic repulsion is high." The correct increasing order of magnitude across the halogens is therefore $\ce{I < Br < F < Cl}$ — note F slots between Br and Cl, not at the top.

Same effect repeats one group left: O is less negative than S, and N is less negative than P, all because of crowding in the small $n = 2$ shell.

Figure 3 · F vs Cl magnitude bars

Bar heights are the magnitude $|\Delta_{eg}H|$ for the halogens (kJ mol⁻¹). Chlorine is the tallest bar — fluorine sits below it despite being smaller, the visual signature of the anomaly.

0 I 295 Br 325 F 328 Cl 349 (max) bar height = |ΔegH| / kJ mol⁻¹ · Cl > F > Br > I
⤳
Connect the dots

The same small-2p-orbital crowding that suppresses fluorine's electron gain enthalpy also makes the first ionisation enthalpies of period-2 elements unusually high. See Ionisation Enthalpy for the mirror-image trend.

Why Noble Gases Are Positive

The noble gases stand apart with positive electron gain enthalpies — He +48, Ne +116, Ar +96, Kr +96, Xe +77, Rn +68 (all $\text{kJ mol}^{-1}$). A positive value means energy must be supplied for the atom to accept an electron at all; the process is endothermic and the product anion is unstable.

The cause is the filled valence shell. A noble gas already has a complete $ns^2np^6$ octet (or $1s^2$ for He), so any extra electron cannot fit into that shell and is forced into the next higher principal quantum level. That places the electron far from the nucleus and well shielded, producing a very unstable configuration — exactly why the value is positive rather than merely small.

Notice too that the noble-gas values themselves drift toward less positive on descending the group (He +48 is actually lower than Ne +116, but from Ne downward the trend Ne > Ar ≈ Kr > Xe > Rn holds): the larger the atom, the less unfavourable it becomes to park an electron in the next shell. The qualitative message for an exam is simply that the noble-gas column is the one place in the table where $\Delta_{eg}H$ is positive, and that this positivity is a direct consequence of an already-stable octet.

Second Electron Gain Enthalpy

The first electron added to oxygen releases energy, but a second electron is a different matter. Forming $\ce{O^2-}$ requires pushing a negative electron onto an already-negative $\ce{O^-}$ ion:

$$\ce{O(g) + e- -> O^-(g)}\,,\ \Delta_{eg}H_1 < 0 \qquad \ce{O^-(g) + e- -> O^2-(g)}\,,\ \Delta_{eg}H_2 > 0$$

The electrostatic repulsion between the incoming electron and the $\ce{O^-}$ ion must be overcome, so the second electron gain enthalpy is positive (endothermic). Forming the free $\ce{O^2-}$ ion from a gaseous oxygen atom is therefore an overall energy-absorbing process; in real compounds the $\ce{O^2-}$ ion is stabilised only by the large lattice energy of the ionic solid, not by electron gain enthalpy itself.

Electron Gain Enthalpy vs Electron Affinity

Older texts use the term electron affinity ($A_e$) for this same idea, and the two carry opposite signs — a classic point of confusion. NCERT footnotes the relationship explicitly: electron affinity is defined as the negative of the enthalpy change, with a positive $A_e$ meaning energy is released, contrary to the thermodynamic convention. The link, at temperature $T$, is:

$$\Delta_{eg}H = -A_e - \tfrac{5}{2}RT$$

FeatureElectron gain enthalpy (ΔegH)Electron affinity (Ae)
ConventionStandard thermodynamic: energy released → negativeEnergy released → positive (contrary to thermodynamics)
Defined atTemperature $T$ (includes a $\tfrac{5}{2}RT$ term)Absolute zero
Sign for a halogenNegative (e.g. Cl $-349$)Positive (e.g. Cl $+349$)
Relationship$\Delta_{eg}H = -A_e - \tfrac{5}{2}RT$

For NEET, the safe habit is to work entirely in $\Delta_{eg}H$ with the thermodynamic sign convention, and to read any "electron affinity" phrasing as the magnitude with the sign flipped. Confusing the two is a deliberate trap in comparison questions.

Quick Recap

Electron Gain Enthalpy in One Screen

  • Definition: enthalpy change for $\ce{X(g) + e- -> X^-(g)}$, denoted $\Delta_{eg}H$, defined for the gaseous atom.
  • Sign convention: negative = energy released = exothermic = stronger tendency to gain an electron. Most negative ≠ "biggest tendency" being positive.
  • Period trend: more negative left→right (rising $Z_{eff}$, smaller atom). Group trend: less negative top→bottom (larger atom).
  • Anomaly: Cl ($-349$) is more negative than F ($-328$) — the small 2p orbital of F gives high electron-electron repulsion. Most negative element overall = Cl. Order: $\ce{I < Br < F < Cl}$.
  • Noble gases: positive $\Delta_{eg}H$ — filled shell forces the electron into the next quantum level.
  • Second ΔegH of O is positive; electron affinity $A_e$ has the opposite sign, with $\Delta_{eg}H = -A_e - \tfrac{5}{2}RT$.

NEET PYQ Snapshot — Electron Gain Enthalpy

Real NEET questions where the sign convention and the F vs Cl anomaly decide the answer.

NEET 2016

In which of the following options the order of arrangement does not agree with the variation of the property indicated against it?

  • (1) B < C < N < O (increasing first ionisation enthalpy)
  • (2) I < Br < Cl < F (increasing electron gain enthalpy)
  • (3) Li < Na < K < Rb (increasing metallic radius)
  • (4) Al³⁺ < Mg²⁺ < Na⁺ < F⁻ (increasing ionic size)
Answer: (1) and (2)

For option (2), the listed order wrongly places F as the most negative. Because the small 2p orbital of fluorine produces high electron density and strong inter-electronic repulsion, F is less negative than Cl. The correct order is $\ce{I < Br < F < Cl}$. (Option 1 also fails: N > O in first ionisation enthalpy due to the half-filled 2p³ shell, giving B < C < O < N.) This is the canonical electron-gain-enthalpy PYQ.

Concept · Most negative ΔegH

Among P, S, Cl and F, which element has the most negative electron gain enthalpy and which has the least negative? (NCERT Problem 3.7)

Answer: Most negative = Cl; least negative = P

$\Delta_{eg}H$ becomes more negative across a period, so among period-3 elements Cl > S > P in magnitude. Although F is above Cl in the group, adding an electron to F's small 2p orbital causes greater repulsion than adding it to Cl's larger 3p orbital, so Cl is more negative than F. Phosphorus, with a stable half-filled 3p³ configuration and lying furthest left, has the least negative value.

Concept · Sign of second ΔegH

Would you expect the second electron gain enthalpy of oxygen to be positive, more negative, or less negative than the first? (NCERT Exercise)

Answer: Positive (endothermic)

The first electron gain ($\ce{O -> O^-}$) is exothermic, but adding a second electron forces a negative electron onto the already-negative $\ce{O^-}$ ion. Overcoming this electrostatic repulsion requires energy, so $\Delta_{eg}H_2$ is positive. This is why isolated $\ce{O^2-}$ is unstable and exists only when stabilised by lattice energy in solids.

Concept · Noble gas sign

Why is the electron gain enthalpy of a noble gas such as neon positive?

Answer: Filled shell forces the electron to a higher level

Neon already has a complete $2s^2 2p^6$ octet. An added electron cannot enter the filled shell and must occupy the next principal quantum level ($n = 3$), producing a very unstable configuration. Energy must be supplied, so $\Delta_{eg}H$ is positive ($+116~\text{kJ mol}^{-1}$). All noble gases show this positive behaviour.

FAQs — Electron Gain Enthalpy

The sign-convention and anomaly questions that come up most often.

What does a negative electron gain enthalpy mean?
A negative electron gain enthalpy means that energy is released when an electron is added to a neutral gaseous atom, i.e. the process is exothermic and the resulting anion is more stable than the separated atom and electron. Group 17 halogens have large negative values because adding an electron lets them reach a stable noble-gas configuration. A positive electron gain enthalpy means energy must be supplied for the atom to accept an electron, as for the noble gases.
Why is the electron gain enthalpy of chlorine more negative than that of fluorine?
Fluorine is a very small atom, so its incoming electron must enter the compact n = 2 (2p) level, where it suffers significant electron-electron repulsion from the electrons already crowded into that small region. In chlorine the added electron enters the larger n = 3 (3p) level, occupying a bigger region of space where the repulsion is much smaller. The lower repulsion in chlorine outweighs its larger size, so Cl releases more energy on gaining an electron. The NCERT values are F = -328 kJ/mol and Cl = -349 kJ/mol, making chlorine the element with the most negative electron gain enthalpy in the whole periodic table.
Why do noble gases have positive electron gain enthalpies?
Noble gases already possess a completely filled, stable ns2np6 valence shell. An extra electron cannot enter that filled shell and is forced into the next higher principal quantum level, producing a very unstable electronic configuration. Energy therefore has to be supplied, so the electron gain enthalpy is positive (for example He = +48, Ne = +116, Ar = +96 kJ/mol).
What is the difference between electron gain enthalpy and electron affinity?
Electron gain enthalpy (Δ_egH) is the enthalpy change when one mole of electrons is added to one mole of gaseous atoms, and it follows the normal thermodynamic sign convention: negative when energy is released. Electron affinity (A_e) is defined as the negative of that enthalpy change at absolute zero, so a positive electron affinity corresponds to energy being released, opposite to the Δ_egH sign. The two are related by Δ_egH = -A_e - (5/2)RT, the (5/2)RT term accounting for heat capacities at temperature T.
How does electron gain enthalpy vary across a period and down a group?
Across a period, electron gain enthalpy generally becomes more negative from left to right because the effective nuclear charge increases and the atom gets smaller, so the added electron is held closer to the nucleus. Down a group, it generally becomes less negative because the atom gets larger and the incoming electron sits farther from the nucleus. Both trends are general rules with notable exceptions, such as fluorine being less negative than chlorine and oxygen being less negative than sulphur.
Why is the second electron gain enthalpy of oxygen positive?
Adding the first electron to a neutral oxygen atom releases energy, but adding a second electron means forcing a negative electron onto the already negatively charged O- ion. The electrostatic repulsion between the incoming electron and the negative ion must be overcome, so energy has to be supplied and the second electron gain enthalpy is positive (endothermic). This is why forming O2- from O(g) is overall endothermic, even though it is favoured in ionic solids by lattice energy.
Related Topics
Classification of Elements and Periodicity Back to the full chapter hub and all subtopics. Ionisation Enthalpy Energy to remove an electron — the mirror property of electron gain. Electronegativity An atom's pull on shared electrons in a bond. Atomic and Ionic Radius Size trends that drive every other periodic property.