Electron affinity is defined as the energy change that occurs when an electron is added to a neutral atom in the gas phase to form a negative ion. It is commonly expressed in units of electronvolts (eV) or kilojoules per mole (kJ/mol). A more negative electron affinity indicates a greater tendency for an atom to gain an electron, meaning it is more favorable for the atom to form an anion.
Trend Across a Period
As you move from left to right across a period in the periodic table, the electron affinity generally becomes more negative (i.e., it increases in magnitude). This trend can be explained as follows:
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Nuclear Charge: As you move across a period, the number of protons in the nucleus increases, leading to a higher positive charge. This increased nuclear charge attracts electrons more strongly.
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Atomic Radius: The atomic size decreases across a period due to increased nuclear charge which draws the electron cloud closer to the nucleus. A smaller radius means that when the atom gains an electron, it is in closer proximity to the nucleus, and hence the attraction is stronger.
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Effective Nuclear Charge: The effective nuclear charge (the net positive charge experienced by valence electrons) increases across a period, making it energetically more favorable for the atom to attract and hold onto an additional electron.
Trend Down a Group
As you move down a group in the periodic table, the electron affinity generally becomes less negative (i.e., it decreases in magnitude). This trend can be explained as follows:
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Increased Atomic Size: The atomic radius increases as you move down a group due to the addition of electron shells. This means that the outer electrons are further from the nucleus and experience a weaker attraction from the protons.
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Shielding Effect: Additional inner electron shells create shielding, which reduces the effective nuclear charge felt by the outer electrons. This makes it less favorable for the atom to gain an electron, as the increased distance and shielding reduce the attraction.
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Energy Consideration: The energy associated with adding an electron decreases because the added electron is further from the nucleus and has less interaction with it, resulting in a less negative (or more positive) electron affinity.
Summary
- Across a Period: Electron affinity generally becomes more negative as nuclear charge increases and atomic size decreases.
- Down a Group: Electron affinity generally becomes less negative due to increased atomic size and shielding effects, leading to a weaker attraction of added electrons.
These trends are not always consistent (e.g., noble gases generally have positive electron affinities), but they provide a good general framework for understanding how electron affinity changes in the periodic table.