Electron affinity describes the energy change when an isolated gaseous atom gains an electron, and this property varies in a predictable way across a period. Understanding why does electron affinity increase across a period helps explain reactivity trends in the periodic table.
Across a period, rising nuclear charge, constant shielding, and decreasing atomic size combine to make it easier for atoms to attract added electrons. The structured overview below summarizes the main factors driving this trend.
| Factor | What It Means | Effect on Electron Affinity | Example Elements |
|---|---|---|---|
| Nuclear Charge | Number of protons increases by one each step | Stronger pull on incoming electron | Li → Ne |
| Shielding | Inner electrons remain the same across a period | Limited extra shielding for added electron | Same core, valence e− feel more Zeff |
| Atomic Radius | Distance from nucleus decreases | Electron added closer to nucleus, less repulsion | Smaller atoms more negative EA |
| Electron Repulsion | Extra electron enters same shell | Small increase, often outweighed by nuclear gain | Exceptions at Group 15, Group 18 |
Rising Nuclear Charge Across a Period
As you move from left to right across a period, the number of protons in the nucleus increases by exactly one for each next element. This rising nuclear charge pulls on all electrons more strongly, including the added electron that defines electron affinity.
The increased positive charge enhances the attraction between the nucleus and the incoming electron, making the addition process more energetically favorable. When energy is released in larger amounts, the electron affinity value becomes more negative, indicating a stronger tendency to gain an electron.
Minimal Increase in Shielding Effect
Why shielding does not keep pace with nuclear charge
Across a period, new electrons are added to the same principal energy level, so inner electron shells remain unchanged. Because inner electrons do not increase, the shielding effect on valence electrons stays nearly constant.
With little extra shielding and a steadily increasing nuclear charge, the effective nuclear charge felt by the added electron rises. This stronger effective pull reinforces the trend of increasing electron affinity across the period.
Decreasing Atomic Radius Enhances Attraction
Closer electrons experience stronger pull
The higher effective nuclear charge also causes the atomic radius to shrink as protons accumulate. The valence electron cloud contracts, bringing the incoming electron closer to the nucleus on average.
A shorter distance between the nucleus and the added electron reduces potential energy release when the electron is captured. This closer approach amplifies the negative energy change and contributes directly to why does electron affinity increase across a period for most elements.
Exceptions and Small Repulsion Effects
Although the overall trend is an increase in electron affinity, some elements show small dips due to electronic configuration details. Group 15 atoms have half filled p subshells, which adds stability and makes electron addition slightly less favorable.
Group 18 elements already have fully filled valence shells, so capturing an extra electron requires pairing in a new shell, which requires energy and results in a positive or near zero electron affinity. These exceptions refine the pattern without overturning the main trend driven by nuclear charge and radius changes.
Key Takeaways on Periodic Trends in Electron Affinity
- Rising nuclear charge across a period is the primary driver of increasing electron affinity.
- Shielding remains nearly constant since new electrons enter the same shell.
- Decreasing atomic radius allows added electrons to experience stronger attraction.
- Exceptions occur at Group 15 and Group 18 due to stable or paired electron configurations.
- Understanding these factors clarifies reactivity patterns and chemical behavior.
FAQ
Reader questions
Why does electron affinity generally become more negative across a period despite added electron repulsion?
The increase in nuclear charge and decrease in atomic radius outweigh the small rise in electron repulsion, so energy released still increases overall.
Are there elements that do not follow the increase in electron affinity across a period?
Yes, nitrogen has lower electron affinity than carbon, and noble gases often have positive or very low electron affinities due to stable electron configurations.
How does effective nuclear charge explain the upward trend in electron affinity across a period?
Effective nuclear charge rises across a period because protons increase while shielding stays similar, pulling added electrons more strongly and releasing more energy.
Why is atomic radius shrinking important for the increase in electron affinity across a period?
A smaller radius means the added electron is closer to the nucleus, which increases attraction and leads to a larger release of energy when the electron is gained.