Electron Affinity

If ionization energy measures how hard it is to take an electron away, electron affinity measures how eagerly an atom accepts one. The two concepts are complementary, and understanding both gives you a complete picture of how elements interact with electrons - the currency of all chemical bonding.

Defining Electron Affinity

Electron affinity (EA) is the energy change that occurs when a neutral, gaseous atom gains one electron:

The Sign Convention Trap

The sign convention for electron affinity confuses many students, and the MCAT knows it. Here is the key distinction:

• The thermodynamic enthalpy change (Delta H) for gaining an electron is negative for most nonmetals (energy is released, exothermic).
• The electron affinity value reported in tables is usually given as a positive number representing the magnitude of that energy release.

So when someone says “chlorine has a high electron affinity of 349 kJ/mol,” they mean that 349 kJ of energy is released per mole when gaseous chlorine atoms each gain one electron. The process is exothermic (Delta H = -349 kJ/mol), but the EA is reported as positive (349 kJ/mol).

If an MCAT question asks whether the process is exothermic or endothermic, the answer is exothermic for elements with positive EA values. Do not let the positive number fool you into thinking it is endothermic.

Trend Across a Period: Generally Increases Left to Right

As you move left to right across a period, electron affinity generally increases (becomes more exothermic). Elements on the right side of the table are closer to achieving a full octet, so gaining an additional electron is energetically favorable. The increasing Z_eff across the period also means the incoming electron is attracted to a stronger nuclear pull.

Halogens (Group 17) have the highest electron affinities of any group. They need just one more electron to complete their octet and achieve noble gas configuration. Chlorine holds the record with an EA of 349 kJ/mol.

Trend Down a Group: Generally Decreases Top to Bottom

As you move down a group, electron affinity generally decreases. The incoming electron enters a shell that is farther from the nucleus, so it feels less nuclear attraction. Even though the nuclear charge increases going down, the added shielding from inner shells more than offsets it.

There is one notable exception: fluorine’s EA (328 kJ/mol) is actually slightly lower than chlorine’s (349 kJ/mol), despite fluorine being higher in the group. This occurs because fluorine’s 2p orbitals are so small and compact that the incoming electron experiences significant repulsion from the electrons already crammed into that tiny space. Chlorine’s larger 3p orbitals accommodate the extra electron more comfortably.

Elements with Very Low or Negative EA

Some elements have electron affinities near zero or are actually endothermic (the atom must absorb energy to accept an electron). These include:

Noble gases (Group 18): Their valence shells are completely full. An added electron would have to enter the next higher energy level, which is energetically unfavorable. Noble gas EA values are effectively zero or slightly positive (endothermic).

Alkaline earth metals (Group 2) and nitrogen: These elements have filled or half-filled subshells (Be: 2s², N: 2p³) that are relatively stable. Adding an electron disrupts that stability, making the process much less favorable.

Electron Affinity of Selected Elements

| Element | Group | EA (kJ/mol) | Notes |
|---------|-------|-------------|-------|
| Na | 1 | 53 | Low - prefers to lose electrons |
| Mg | 2 | ~0 | Filled 2s subshell, very low |
| Al | 13 | 42 | Low, starting to fill 3p |
| Si | 14 | 134 | Moderate |
| P | 15 | 72 | Lower than Si (half-filled 3p stability) |
| S | 16 | 200 | High, one electron short of half-filled plus one |
| Cl | 17 | 349 | Highest of all elements |
| Ar | 18 | ~0 | Full octet, does not want electrons |

Notice the irregularities: P has a lower EA than Si because phosphorus has a stable half-filled 3p³ configuration. These exceptions are less predictable than IE exceptions and are tested less frequently, but the overall left-to-right increase is reliable.

Comparing EA, IE, and Atomic Radius

All three trends are driven by the same underlying factors - effective nuclear charge and atomic size - but they point in different directions:

| Property | Across a Period (L → R) | Down a Group (Top → Bottom) |
|----------|------------------------|-----------------------------|
| Atomic Radius | Decreases | Increases |
| Ionization Energy | Increases | Decreases |
| Electron Affinity | Generally Increases | Generally Decreases |

Atomic radius is the odd one out. IE and EA both increase toward the upper right of the periodic table (where atoms are small and Z_eff is high). Atomic radius increases toward the lower left (where atoms are large and Z_eff is low).

EA vs. Electronegativity

Students sometimes confuse electron affinity with electronegativity. They are related but distinct:

• Electron affinity is the energy released when an isolated gaseous atom gains an electron. It is a measurable thermodynamic quantity with units of kJ/mol.
• Electronegativity is the tendency of an atom to attract shared electrons toward itself within a chemical bond. It is a relative scale (Pauling scale) with no units.

Both trend in the same direction (increase toward the upper right), and both reflect an atom’s desire for electrons. But EA applies to isolated atoms gaining electrons, while electronegativity applies to bonded atoms pulling on shared electrons.