Would a neutral K atom or a K+ ion have a more negative value of electron affinity?

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Understand the concept of electron affinity: Electron affinity is the energy change that occurs when an electron is added to a neutral atom in the gaseous state to form a negative ion.

Consider the electron configuration of a neutral potassium (K) atom: Potassium has an electron configuration of [Ar] 4s^1.

Consider the electron configuration of a K+ ion: The K+ ion has lost one electron, so its electron configuration is [Ar], which is a stable noble gas configuration.

Analyze the electron affinity of a neutral K atom: Adding an electron to a neutral K atom would result in the configuration [Ar] 4s^2, which is less stable than the noble gas configuration.

Analyze the electron affinity of a K+ ion: Adding an electron to a K+ ion would result in the configuration [Ar] 4s^1, which is less stable than the noble gas configuration. Therefore, the K+ ion would have a less negative (or more positive) electron affinity compared to the neutral K atom.

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Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Electron Affinity

Electron affinity is the amount of energy released when an electron is added to a neutral atom in the gas phase, forming a negative ion. A more negative electron affinity indicates a greater tendency for an atom to gain an electron. This property varies across the periodic table, influenced by atomic size and effective nuclear charge.

Potassium Atom vs. Potassium Ion

A neutral potassium (K) atom has one more electron than the potassium ion (K+), which has lost one electron. The electron affinity of a neutral atom is typically more negative than that of its corresponding cation because the neutral atom is more stable and has a stronger attraction for an additional electron compared to the already positively charged ion.

Trends in Electron Affinity

Electron affinity generally increases (becomes more negative) across a period from left to right and decreases down a group in the periodic table. This trend is due to changes in atomic size and effective nuclear charge, which affect how strongly an atom can attract an additional electron. Understanding these trends helps predict the electron affinity of different elements.

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