Atomic size, or atomic radius, generally decreases as you move from left to right across a period on the periodic table. This is a consistent trend that is primarily due to the increasing nuclear charge experienced by the electrons.
Understanding the Trend
Nuclear Charge
- As you move across a period, the number of protons in the nucleus increases. This increased positive charge draws the electrons closer to the nucleus.
- Even though the number of electrons also increases, the added electrons go into the same energy level, or shell.
- The increase in nuclear charge is more significant than the increase in electron-electron repulsion, resulting in an overall shrinking of the electron cloud.
Electron Shells
- Within a period, the number of electron shells remains constant. The electrons are all being added to the same principal energy level.
- This lack of additional shells means that the electrons are not being added farther from the nucleus, meaning they are pulled tighter as the nuclear charge increases.
Table of Atomic Radius Trend Within a Period
| Direction | Atomic Radius Change | Reason |
|---|---|---|
| Left to Right (Across a Period) | Decreases | Increasing Nuclear Charge, constant electron shells |
Example: Period 3
Looking at Period 3, we observe this trend:
- Sodium (Na) has a larger atomic radius than magnesium (Mg).
- Magnesium (Mg) has a larger atomic radius than aluminum (Al).
- This trend continues until we get to chlorine (Cl), which has the smallest atomic radius in the period (not including noble gases).
In summary
The primary reason for the decrease in atomic radius across a period is the increasing positive charge of the nucleus pulling the electrons more tightly. The addition of electrons to the same shell does not offset this effect. As stated in the reference, "atomic radius generally decreases as you move from left to right across a period (due to increasing nuclear charge)".
