AP Chemistry Electronegativity

What it measures, why the trend runs the way it does, a values table, and worked practice.

Electronegativity shows up whenever a bond is classified, a molecule's polarity is argued, or an oxidation number is assigned, yet it is easy to mix up with its neighbors, ionization energy and electron affinity. On the exam you will not be handed a table of values, so the useful skill is explaining the ranking from position and Coulomb's law. The periodic table guide shows what the exam sheet does and does not print; this page is about the property itself.

What Electronegativity Is and How It Differs From Ionization Energy and Electron Affinity

Electronegativity is how strongly an atom in a bond attracts the shared pair of electrons. The Pauling scale expresses it as a number with no units, running from about 0.8 (cesium) to 3.98 (fluorine). Two other properties sound alike but measure something different:

All three come from the same causes, so they usually trend together, but they are not interchangeable. Chlorine, for example, releases slightly more energy when it gains an electron (about 349 kJ/mol) than fluorine does (about 328 kJ/mol), yet fluorine is the more electronegative element. For the trends in ionization energy and radius, see periodic trends.

Pauling Electronegativity Values for Main-Group Elements

Read the table the way the trend describes it: values climb as you move right along a row and fall as you move down a column. In the lower p-block there are small reversals, such as gallium (1.81) coming out higher than the aluminum (1.61) above it, which come from the filled d subshell in period 4. Do not rely on these irregular pairs in an explanation; use pairs where the trend is clean, or use the numbers a question gives you.

Why Electronegativity Rises Across a Period and Falls Down a Group

The reasoning is the same Coulomb's law argument used for other trends: an atom pulls harder on a shared pair when the effective nuclear charge is high and the distance to its valence electrons is short.

• Across a period, protons are added but the valence shell stays the same, so shielding barely changes and the effective nuclear charge rises. The atom pulls shared electrons harder.
• Down a group, the valence electrons are in a higher shell, farther from the nucleus and shielded by more inner electrons. The pull on shared electrons weakens.

Worked Example 1: Comparing Two Elements in the Same Period

Which is more electronegative, aluminum or chlorine? Explain.

Both are in period 3, so their valence electrons are in the same shell with the same 10 core electrons shielding them. Chlorine has 17 protons and aluminum has 13, so the approximate effective nuclear charge is +7 for Cl against +3 for Al. At about the same distance, chlorine pulls shared electrons much harder: chlorine is more electronegative (3.16 against 1.61).

Worked Example 2: Comparing Two Elements in the Same Group

Which is more electronegative, fluorine or iodine? Explain.

Both have an approximate effective nuclear charge of +7, so the difference is distance and shielding. Fluorine's valence electrons are in the second shell, close to the nucleus. Iodine's are in the fifth shell, far away and screened by many more inner electrons. The attraction for shared electrons is weaker at the greater distance: fluorine is more electronegative (3.98 against 2.66).

Worked Example 3: Comparing Two Elements When the Trends Conflict

Which is more electronegative, oxygen or chlorine?

Here the two trends pull in opposite directions. Chlorine is farther right, which favors chlorine, but oxygen is higher in the table, which favors oxygen. Position alone cannot settle it. The data say oxygen is more electronegative (3.44 against 3.16), because at the top of the table the short distance outweighs the extra proton count.

On an exam, a question that needs this comparison will give you the values or choose a clean pair. If you must commit to an answer without data, say which trend you used and why.

Worked Example 4: Ranking Bonds by Polarity and Finding the Partial Negative Atom

Rank the bonds C–H, C–N, C–O, and C–F from least to most polar, and say which atom carries the partial negative charge in each.

A bigger electronegativity difference means a more polar bond. Using the table:
C–H: 2.55 − 2.20 = 0.35 (so little difference that the bond counts as effectively nonpolar).
C–N: 3.04 − 2.55 = 0.49.
C–O: 3.44 − 2.55 = 0.89.
C–F: 3.98 − 2.55 = 1.43.
Order, least to most polar: C–H < C–N < C–O < C–F. In C–N, C–O, and C–F the partial negative charge is on the more electronegative atom (N, O, and F). Carbon is the partial positive end.

Whether those polar bonds make a polar molecule depends on shape, which is the subject of molecular polarity. The course also treats bond type as a continuum rather than a sharp cutoff on the difference, a point made in ionic versus covalent bonds.

Five Electronegativity Practice Questions With Answers

1. Rank Na, Mg, and Cl from least to most electronegative.
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   Na < Mg < Cl. All are in period 3; the effective nuclear charge rises from Na to Cl.
2. Which is more electronegative, oxygen or sulfur? Explain briefly.
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   Oxygen. It is in the same group, with its valence electrons in a closer shell with less shielding, so it pulls shared electrons harder.
3. In an H–Cl bond, which atom carries the partial negative charge?
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   Chlorine, which is more electronegative than hydrogen, so the shared electrons are drawn toward Cl.
4. Rank H–F, H–Cl, and H–Br by bond polarity, least to most.
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   H–Br < H–Cl < H–F. The halogen's electronegativity falls down the group, so the difference from hydrogen is largest for fluorine.
5. Explain why bromine is less electronegative than chlorine.
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   Both have the same approximate effective nuclear charge, but bromine's valence electrons are in a higher shell, farther from the nucleus and more shielded. By Coulomb's law the attraction for shared electrons is weaker.

Common Electronegativity Mistakes

• Treating it as the same thing as ionization energy or electron affinity. It describes bonded atoms and has no units.
• Saying electronegativity increases down a group. It decreases, because distance and shielding increase.
• Explaining the trend with “more protons” alone. A full explanation names shielding and distance as well.
• Forgetting the noble gases. They are left out of the trend because they form almost no bonds.
• Assuming a polar bond means a polar molecule. Shape decides whether the bond dipoles cancel.

Electronegativity differences also drive how oxidation numbers are assigned in oxidation numbers and which atoms carry formal charge in formal charge problems. To check how your preparation is going overall, the AP Chem Score Calculator is a quick way to see where a practice score lands.

Related Resources

• AP Chem Score Calculator
• Unit 1 Review: Atomic Structure and Properties
• AP Chemistry Periodic Table
• AP Chemistry Periodic Trends
• AP Chemistry Molecular Polarity
• AP Chemistry Ionic vs Covalent Bonds
• AP Chemistry Oxidation Numbers
• AP Chemistry Study Guide

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