Electron affinity is a property of atoms that reflects their tendency to gain electrons. This characteristic is part of broader patterns observed across the periodic table. As one moves from left to right across a period, electron affinity increases. Atoms on the right side of a period have a greater inclination to accept an additional electron.
What is Electron Affinity?
Electron affinity describes the energy change that occurs when an electron is added to a neutral atom in its gaseous state to form a negative ion. It quantifies how much an atom “wants” an extra electron. When an atom gains an electron, energy is often released, making the process exothermic, and a more negative electron affinity value indicates a greater tendency for an atom to accept an electron.
Navigating the Periodic Table: Periods
The periodic table is organized into rows called periods and columns called groups. Moving across a period, from left to right, involves a systematic change in atomic structure. Each subsequent element in a period has one more proton in its nucleus than the one before it.
As protons are added to the nucleus, electrons are also added to the atom. These newly added electrons occupy the same outermost electron shell, or principal energy level, across a given period.
Key Atomic Properties at Play
Several atomic properties influence an atom’s electron affinity. The effective nuclear charge is the net positive charge experienced by an electron in an atom. It is not the full nuclear charge because inner electrons shield the outer electrons from the nucleus’s attractive force. Across a period, the nuclear charge increases, while the shielding effect from inner electrons remains relatively constant. This leads to an increase in the effective nuclear charge experienced by the valence electrons, pulling them more strongly towards the nucleus.
Atomic radius, which is the distance from the atom’s nucleus to its outermost electron cloud, also changes across a period. As the effective nuclear charge increases from left to right, the stronger pull from the nucleus draws the electron shells closer. This increased attraction results in a general decrease in atomic radius across a period. A smaller atomic radius means the valence electrons are held more tightly.
Electron shielding refers to the repulsion between electrons that reduces the attractive force of the nucleus on the outer electrons. Inner-shell electrons are responsible for this shielding. Across a period, the number of inner-shell electrons remains consistent because electrons are added to the same principal energy level. Therefore, the shielding effect does not significantly increase as one moves from left to right.
Explaining the Increase: How Properties Drive the Trend
The increase in electron affinity across a period is a direct consequence of the interplay between effective nuclear charge, atomic radius, and electron shielding. As one moves from left to right, the increasing number of protons in the nucleus leads to a rising effective nuclear charge. This stronger positive charge exerts a greater attractive force on all electrons, including any incoming electron.
The decreasing atomic radius further contributes to this trend. With a smaller radius, an additional electron approaching the atom will be closer to the positively charged nucleus. This closer proximity results in a stronger attraction between the nucleus and the incoming electron, making its acceptance more energetically favorable.
The relatively constant electron shielding across a period allows the increasing effective nuclear charge to dominate. Because the inner electrons do not significantly increase their shielding capacity, the added positive charge from the nucleus is more fully felt by the valence shell. This combination of a stronger nuclear pull and a more compact atomic size makes atoms on the right side of a period more receptive to gaining an electron.