TL;DR:
Across a period, nuclear charge rises while the electrons stay in the same shell, so atoms shrink and electronegativity rises.
Down a group, extra shells and shielding make atoms bigger and electronegativity lower.
Trends are generalisations with exceptions, so check a data table when it matters.
Sodium is a soft metal that reacts violently with water. Chlorine is a pale green gas that grabs electrons. Both are in period 3, only a few steps apart. The trends in the periodic table explain both behaviours with two ideas: how strongly the nucleus pulls on the outer electrons and how far away those electrons are.
Two forces to keep in mind
The nucleus has a positive charge equal to the atomic number Z. The inner electrons sit between the nucleus and the outer ones and partly cancel its pull. This is shielding. What an outer electron actually feels is the effective nuclear charge, Z_eff.
A simple estimate, good enough for a first look, is Z_eff ≈ Z − (number of inner core electrons). It ignores the fact that electrons in the same shell shield each other a little, so real values are somewhat lower. The trend, though, is right.
Worked example: across period 3
Every element from sodium to chlorine has the same 10 core electrons (1s² 2s² 2p⁶), and its valence electrons are in shell 3. Use the simple estimate, and Pauling electronegativities from a data table.
Element | Z | Core electrons | Simple Z_eff | Electronegativity |
|---|---|---|---|---|
Na | 11 | 10 |
| 0.93 |
Mg | 12 | 10 |
| 1.31 |
Al | 13 | 10 |
| 1.61 |
Si | 14 | 10 |
| 1.90 |
P | 15 | 10 |
| 2.19 |
S | 16 | 10 |
| 2.58 |
Cl | 17 | 10 |
| 3.16 |
Read the table. Valence electrons stay in the same shell, so the distance does not grow, but the pull grows from 1 to 7. Outer electrons are held closer and tighter, so the atomic radius decreases from sodium to chlorine. Electronegativity, the pull an atom exerts on shared electrons, increases, from 0.93 to 3.16.
Compare the ends. Chlorine feels about 7 times the effective pull of sodium in this estimate, since 7 ÷ 1 = 7, in the same shell. A Na-Cl pair differs in electronegativity by 3.16 − 0.93 = 2.23, big enough for chlorine to take the electron outright, as in ionic and covalent bonds.
Down a group
Moving down a column, each element adds a whole new shell. The outer electrons are farther from the nucleus, and the extra inner electrons shield the pull, so Z_eff for the outer electrons changes little. The larger distance wins: atomic radius increases down a group.
Electronegativity goes the other way. For the halogens: F 3.98, Cl 3.16, Br 2.96, I 2.66. The shared electrons sit farther from the nucleus in each, so the pull weakens. Fluorine is the most electronegative element.
Ionisation energy, the energy to remove an outer electron, follows the same pattern as electronegativity: larger across a period and smaller down a group. Metals on the left lose electrons easily, nonmetals on the right hold on tightly.
Exceptions to treat carefully
Noble gases have no electronegativity values in the usual table because they rarely bond.
Transition metals vary only slightly across a period, since the added electrons go into an inner d subshell and partly shield each other.
Gallium is more electronegative than aluminium above it, because of the d electrons added between them. Not every step down a group lowers the value.
Anions and cations differ from the neutral atom. A positive ion is smaller than its atom, and a negative ion is larger.
Trends, not laws: These are patterns that explain a lot, not exact rules. For a specific number, such as a radius in picometres, use a data table from your course, since values differ by how radius is defined.
A quick comparison uses the same tools. Which is larger, magnesium or calcium? Both are in column 2, and calcium is a row lower, so it has an extra shell and a bigger radius. Which is larger, magnesium or sulfur? Both are in period 3, and sulfur has more protons pulling on the same shell, so magnesium is larger.
How to check your answer
Same shell? If two atoms are in the same period, the one with more protons is smaller.
Different shell? The one in the lower period is usually larger.
Direction check. Fluorine, top right, is the most electronegative. Francium and caesium, bottom left, are the least.
Mention both factors. A full explanation names nuclear charge, shielding and distance.
Mistakes that cost marks
Saying atoms get bigger across a period because they have more electrons.
Forgetting shielding when explaining a trend down a group.
Comparing an atom with an ion as if they had the same size.
Treating electronegativity as the same as electron affinity.
Presenting a trend as having no exceptions.
Where you will use it next
Trends predict which atoms form ions, which bonds are polar and which elements are reactive. The arrangement of electrons behind them is in electron configurations and the periodic table. Gases behave more simply, as you will see in the gas laws in one equation.
Frequently asked questions
Why do noble gases not have electronegativity values?
They hold full outer shells and form almost no bonds, so there is little to measure.
Is a smaller atom always more electronegative?
Often, since a smaller atom holds shared electrons closer to its nucleus. It is a tendency, not a guarantee.
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