Electron affinity is the energy change when an electron is added to a neutral gas-phase atom to form a negative ion. In AP Chem (Topic 1.7), it generally becomes more negative (more energy released) across a period because increasing effective nuclear charge pulls the incoming electron in more strongly.
Electron affinity (EA) is the energy change when a neutral atom in the gas phase picks up an electron and becomes an anion. If the atom "wants" the electron, energy is released and the EA is negative (exothermic). Chlorine is the classic example. It's one electron short of a full octet, so adding an electron releases a lot of energy. That makes the halogens the electron-affinity champions of the periodic table.
The trend is explained the same way as every other periodic trend in Topic 1.7, through Coulomb's law and effective nuclear charge. Across a period, protons are added but the incoming electron joins the same shell, so shielding barely changes. The effective nuclear charge felt by that new electron goes up, the attraction gets stronger, and more energy is released. The trend down a group is weaker and messier, but the big idea is that a larger atomic radius means the added electron sits farther from the nucleus and feels less pull. Watch for the exceptions, though. Noble gases and elements with filled or half-filled subshells (like Group 15) resist gaining an electron, because the new electron would have to start pairing up or open a new subshell.
Electron affinity lives in Unit 1: Atomic Structure and Properties, Topic 1.7 (Periodic Trends), under learning objective 1.7.A: explain the relationship between trends in atomic properties and electronic structure and periodicity. The essential knowledge (1.7.A.2) names it explicitly as one of the trends you explain using Coulomb's law, the shell model, and shielding/effective nuclear charge. It's also one of the few trends with a famous exception (Groups 14 vs. 15), and the AP exam loves testing whether you can explain anomalies with electron configurations instead of just memorizing the arrow on the periodic table. Beyond Unit 1, electron affinity is half the story of why ionic bonds form at all. Metals with low ionization energy hand electrons to nonmetals with very negative electron affinities.
Keep studying AP Chemistry Unit 1
Ionization Energy (Unit 1)
These two are mirror images. Ionization energy is the cost of removing an electron; electron affinity is the payoff (or cost) of adding one. Both trends come from the same cause, increasing effective nuclear charge across a period, which is why fluorine and chlorine score high on both.
Coulomb's Law & Effective Nuclear Charge (Unit 1)
Electron affinity isn't a fact to memorize, it's Coulomb's law in action. A higher effective nuclear charge and a smaller distance between the nucleus and the incoming electron mean a stronger attraction and more energy released. Every "explain the trend" answer should come back to charge and distance.
Atomic Radius (Unit 1)
Radius and electron affinity run in opposite directions across a period. Smaller atoms put the incoming electron closer to the nucleus, so they release more energy when they gain it. If you know the radius trend, you can predict the EA trend.
Ionic Bonding (Unit 2)
Electron affinity explains who becomes the anion. When sodium meets chlorine, sodium's low ionization energy and chlorine's very negative electron affinity make the electron transfer favorable, setting up the Coulombic attraction that holds NaCl together in Unit 2.
Electron affinity shows up almost entirely in multiple-choice and short-answer questions asking you to explain or predict a trend. Typical stems ask you to explain the trend across Period 3 from Na to Cl (answer: increasing effective nuclear charge with roughly constant shielding) or to explain the anomalous dip between Groups 14 and 15 (answer: Group 15 has a half-filled p subshell, so an added electron must pair up, costing repulsion energy). You may also get an unknown-element setup, like estimating properties of an element between gallium and indium, where electron affinity is one of the properties you interpolate from its periodic table position. No released FRQ has used the term verbatim, but the skill it tests, justifying a periodic trend with Coulomb's law and electron configurations, is a staple of Unit 1 free-response prompts. Never answer with just "it's a trend." Always name the cause.
Ionization energy is the energy required to REMOVE an electron from a neutral atom (always endothermic, always positive). Electron affinity is the energy change when an atom GAINS an electron (usually exothermic for nonmetals, so usually negative). Both increase in magnitude across a period for the same Coulombic reason, but they describe opposite processes. A quick check: halogens have huge electron affinities because they want one more electron; noble gases have huge ionization energies because they refuse to give one up.
Electron affinity is the energy change when a neutral gas-phase atom gains an electron to form an anion, and it's usually exothermic (negative) for nonmetals.
Across a period, electron affinity generally becomes more negative because effective nuclear charge increases while shielding stays about the same, so the incoming electron feels a stronger pull.
Halogens like chlorine have the most negative electron affinities because gaining one electron completes their octet.
The Group 14 to Group 15 anomaly happens because Group 15 elements have a half-filled p subshell, so an added electron must pair up and the extra electron-electron repulsion makes gaining it less favorable.
Noble gases have electron affinities near zero or positive because an added electron would have to start a whole new shell far from the nucleus.
On the exam, always explain electron affinity trends using Coulomb's law, effective nuclear charge, and electron configurations, never just 'because of the trend.'
It's the energy change when a neutral gas-phase atom gains an electron to form a negative ion. It's part of Topic 1.7 (Periodic Trends) and is explained using Coulomb's law and effective nuclear charge under learning objective 1.7.A.
It generally becomes more negative (more energy released) from left to right, peaking at the halogens. Across Period 3, chlorine has a much more negative electron affinity than sodium because its effective nuclear charge is higher with similar shielding.
More negative means more energy is released, so the atom attracts electrons more strongly. Chemists often say chlorine has a 'high' electron affinity even though the value is very negative, which trips a lot of people up. Focus on the magnitude of energy released.
Nitrogen has a half-filled 2p subshell (2p³), which is a relatively stable configuration. An added electron would have to pair up in an occupied orbital, and that electron-electron repulsion offsets the nuclear attraction. This Group 14 vs. 15 anomaly is a favorite AP multiple-choice question.
Ionization energy is the energy needed to remove an electron (always positive); electron affinity is the energy change when an atom gains one (usually negative for nonmetals). They're opposite processes driven by the same Coulombic logic, so both grow in magnitude across a period.