- Reactivity trends
- Electronegativity and bond polarity
- How ionic bonds form
Sharing is not always equal
In a covalent bond, two atoms share a pair of electrons. In a hydrogen molecule, H–H, the two atoms are identical, so they share the pair equally. It sits exactly in the middle.
Now think about hydrogen chloride, H–Cl. The hydrogen nucleus and the chlorine nucleus both attract the shared pair. They are not equally strong. Chlorine pulls harder, so the pair spends more of its time nearer the chlorine atom. The bond is shared, but not shared fairly.
How much harder does one atom pull than another? Chemists answer with one number for each element: its electronegativity. This lecture explains what it measures, and why it follows the same pattern as atomic radius and ionisation energy. It also shows how the difference between two atoms tells you what kind of bond they form.
What electronegativity measures
Electronegativity is a measure of how strongly an atom attracts the shared pair of electrons in a covalent bond towards itself.
Notice the words "in a covalent bond". This is different from ionisation energy. Ionisation energy is about pulling an electron completely away from a single gaseous atom. Electronegativity is about a tug-of-war over electrons that two bonded atoms share.
The most widely used scale was published by the chemist Linus Pauling in 1932, in a paper about the energies of chemical bonds. On the Pauling scale, values have no units. They are relative numbers, chosen so that they can be compared. Fluorine has the highest value, 3.98. Caesium has one of the lowest, 0.79. (Francium is given about 0.7, but that value is only a rough estimate.)
Some books round fluorine to 4.0 and chlorine to 3.2. In this lecture, every value is a Pauling value from one table (LibreTexts, Table A2), given to two decimal places. As with atomic radii, compare numbers only when they come from the same table.
The same tug-of-war again
You already know what decides how strongly an atom holds electrons. The atomic radius and ionisation energy lectures used the same three things: the nuclear charge, the distance from the nucleus, and the shielding by inner electrons. A shared pair in a bond sits in the outer shell of each atom, so the same three things decide how strongly each atom pulls on it.
Across a period, the nuclear charge rises, the shielding stays about the same, and the atoms get smaller. All three changes make the pull on a shared pair stronger. So electronegativity increases across a period.
Down a group, each atom has an extra shell. The shared pair is further from the nucleus and better shielded from it. The extra protons do not make up for this. So electronegativity decreases down a group.
The trend down a group shows clearly in the halogens and the alkali metals:
| Group 17 | Electronegativity | Group 1 | Electronegativity |
|---|---|---|---|
| F | 3.98 | Li | 0.98 |
| Cl | 3.16 | Na | 0.93 |
| Br | 2.96 | K | 0.82 |
| I | 2.66 | Rb | 0.82 |
| Cs | 0.79 |
Notice that potassium and rubidium have the same value, 0.82. The trend is a strong general pattern, not a law that every step must obey. Small irregularities like this appear in real data.
That is it. That is the whole idea. Put the two trends together, and the most electronegative elements are at the top right of the periodic table, near fluorine, oxygen and nitrogen. The least electronegative are at the bottom left, near caesium. These are the same corners that held the smallest and largest atoms in the atomic radius lecture.
Neon has a higher first ionisation energy than fluorine. So why is fluorine, not neon, the most electronegative element?
Show the answer
Electronegativity describes an atom in a bond. Neon forms no stable compounds, so it never shares a pair of electrons, and the Pauling scale gives it no value. Most noble gases are left out of the scale for the same reason. Krypton and xenon do form a few compounds, and they have been given values (3.00 and 2.6), but these are not part of the usual trend. Among the atoms that do form bonds, fluorine pulls hardest.
When the two trends disagree
Compare nitrogen (Period 2, Group 15) with chlorine (Period 3, Group 17). Nitrogen is one period higher, which pushes its value up. Chlorine is two groups further right, which pushes its value up. The two trends point in opposite directions. The data decide: chlorine 3.16, nitrogen 3.04. Chlorine is slightly higher.
Now compare oxygen (Period 2, Group 16) with chlorine. Again the trends disagree, and this time the data favour oxygen: 3.44 against 3.16. For "diagonal" comparisons like these, the rules cannot decide. You need the numbers, just as for potassium and barium in the atomic radius lecture.
Electronegativity and the type of bond
The most useful thing about electronegativity is the difference between two bonded atoms. It tells you how unequal the sharing is.
- A difference near zero: both atoms pull about equally. The pair is shared almost equally, and the bond is non-polar covalent. Examples: H–H, Cl–Cl, and C–H (difference 0.35).
- A moderate difference: one atom pulls harder. The pair is shared unequally, and the bond is polar covalent. The more electronegative atom gets a small negative charge, written δ−, and the other a small positive charge, δ+. Example: H–Cl, where the difference is .
- A large difference: one atom pulls so much harder that it effectively takes the electron. The atoms become ions, and the bond is ionic. Example: Na–Cl, where the difference is .
As a rough guide, a difference below about 0.4 is treated as non-polar, and a difference above about 1.8 as ionic. But books place the boundaries differently, at 1.7, 1.8 or even 2.0, and real bonds do not change type suddenly at a line. Bond character changes gradually, from equal sharing to complete transfer. Treat the numbers as a guide, not a rule.
Explore it below. Choose any two atoms and watch where the shared pair sits.
Choose two atoms. The picture shows, roughly, where the shared pair sits between them. Read the difference in electronegativity and the likely type of bond.
Some bonds, their differences in electronegativity, and the likely type of bond (Pauling values):
| Bond | Difference | Likely type |
|---|---|---|
| H–H | 0 | non-polar covalent |
| C–H | 0.35 | non-polar covalent (almost) |
| H–I | 0.46 | slightly polar covalent, I is δ− |
| H–Cl | 0.96 | polar covalent, Cl is δ− |
| H–F | 1.78 | very polar covalent, F is δ− |
| Na–Cl | 2.23 | ionic: Na⁺ and Cl⁻ |
Did you try hydrogen with fluorine? The difference is 1.78, right at the boundary, yet hydrogen fluoride is made of molecules held together by a very polar covalent bond, not of ions. This is the clearest reason to treat the boundaries as a guide.
Worked examples
Arrange C, F, N, O and Si in order of increasing electronegativity.
C, N, O and F are all in Period 2, so across the period: C < N < O < F. Silicon is in the same group as carbon (Group 14), one period lower, so silicon is less electronegative than carbon. The order is Si < C < N < O < F.
Check: 1.90, 2.55, 3.04, 3.44 and 3.98.
Compare the bonds in hydrogen fluoride and hydrogen iodide.
H–F: . Fluorine pulls far harder than hydrogen, so the bond is very polar, with fluorine δ− and hydrogen δ+.
H–I: . Iodine pulls only a little harder, so the bond is only slightly polar.
The difference comes from the Group 17 trend. Fluorine is small, and the shared pair is close to its nucleus. Iodine is large, with three more shells, so the shared pair is far from its nucleus and well shielded. Differences like this affect how these substances behave, which the bonding lectures explain.
Use electronegativity to predict the bonding in sodium chloride and in tetrachloromethane, CCl₄.
Na–Cl: . This is a large difference, so chlorine effectively takes sodium's outer electron. The bonding is ionic: Na⁺ and Cl⁻. This matches what is known about sodium chloride, a typical ionic solid.
C–Cl: . This is a moderate difference, so each C–Cl bond is polar covalent, with chlorine δ−. (Whether the whole CCl₄ molecule is polar depends on its shape, which a later lecture explains.)
In each bond, which atom carries δ−? (a) O–H (b) C–O (c) Si–Cl
The more electronegative atom carries δ−. (a) O: 3.44 against 2.20. (b) O: 3.44 against 2.55. (c) Cl: 3.16 against 1.90.
Notice that oxygen is δ− in both (a) and (b). Oxygen is the second most electronegative element, so it is δ− in almost every bond it forms, except a bond to fluorine.
Practise classifying bonds below. One card is the classic borderline case.
Each card shows a bond and the difference in electronegativity. Place it in the most likely box. Read the reason on every card.
Each bond and its most likely type:
| Bond | Difference | Most likely type |
|---|---|---|
| H–H | 0 | non-polar covalent |
| Cl–Cl | 0 | non-polar covalent |
| C–H | 0.35 | non-polar covalent |
| C–Cl | 0.61 | polar covalent |
| H–Cl | 0.96 | polar covalent |
| O–H | 1.24 | polar covalent |
| H–F | 1.78 | polar covalent (borderline) |
| Na–Cl | 2.23 | ionic |
| K–F | 3.16 | ionic |
Three classic mistakes
Ionisation energy is the energy to remove an electron completely from an isolated gaseous atom, measured in kJ/mol. Electronegativity describes the pull on a shared pair in a bond, and it has no units. They follow similar trends, because the same tug-of-war drives both, but they are different quantities.
A difference of 1.79 does not make a bond covalent, nor does 1.81 make it ionic. Books disagree about where the lines go, and H–F sits right at one of them. Use the difference to describe how unequal the sharing is, not to sort bonds by a sharp rule.
Noble gases have the highest ionisation energies in each period, so it is tempting to give them the highest electronegativities too. But most of them form no bonds, so electronegativity does not apply. Fluorine is the most electronegative element.
Use the values in this lecture. Work each one on paper first. Open a hint only when you are stuck.
-
Arrange Na, Cl, Mg, S and P in order of increasing electronegativity.
Hint 1
All five are in Period 3.
Hint 2
Across a period, electronegativity increases. Put them in group order.
Full answer
Na < Mg < P < S < Cl. Check: 0.93, 1.31, 2.19, 2.58 and 3.16.
-
Use the differences in electronegativity to describe the bonds H–Br (H 2.20, Br 2.96) and Mg–O (Mg 1.31, O 3.44).
Hint 1
Subtract the smaller value from the larger one in each pair.
Hint 2
Compare each difference with the rough guide: about 0.4 and about 1.8.
Full answer
H–Br: . Polar covalent, with bromine δ−. Mg–O: . A large difference, so the bonding is likely ionic: Mg²⁺ and O²⁻.
-
Carbon (2.55) and sulfur (2.58) have almost the same electronegativity. What does this suggest about a C–S bond? Which atom, if either, is δ−?
Hint 1
Find the difference.
Hint 2
The difference is 0.03, far below 0.4.
Full answer
The difference is , so the two atoms pull almost exactly equally. The C–S bond is effectively non-polar. Sulfur pulls very slightly harder, but the difference is too small to give the bond any useful polarity.
-
Oxygen (Period 2, Group 16) is more electronegative than chlorine (Period 3, Group 17), but nitrogen (Period 2, Group 15) is less electronegative than chlorine. Explain why the trends alone cannot predict either comparison, and why oxygen ends up above chlorine.
Hint 1
For each pair, decide which way the period trend points and which way the group trend points.
Hint 2
Oxygen is one period up from chlorine and only one group to the left. Nitrogen is one period up and two groups to the left.
Full answer
In both pairs, the Period 2 atom is one period higher, which favours it, but further left, which favours chlorine. The trends point in opposite directions, so only data can decide. Oxygen is only one group to the left of chlorine, and its smaller size (a shared pair in shell 2, not shell 3) wins: 3.44 against 3.16. Nitrogen is two groups to the left, so its nuclear charge is lower still, and chlorine's two extra groups just win: 3.16 against 3.04. Diagonal comparisons always need the numbers.
Three lectures in a row have followed one idea: how strongly a nucleus holds outer electrons, against distance and shielding. Atomic radius showed it as size. Ionisation energy showed it as the energy to remove an electron. Electronegativity shows it as a pull on electrons that are shared. One cause, three measurements.
Electronegativity is also where the periodic table starts to explain bonding. The difference between two atoms tells you whether they share electrons equally, unequally, or not at all. The next lecture uses the same tug-of-war to explain why some metals and non-metals react so much more violently than others.
- Electronegativity measures how strongly an atom attracts the shared pair of electrons in a covalent bond, and on the Pauling scale it has no units.
- Electronegativity increases across a period, because the nuclear charge rises while shielding stays about the same and the atoms get smaller.
- Electronegativity decreases down a group, because the shared pair is further from the nucleus and better shielded.
- Fluorine is the most electronegative element, and most noble gases are left out of the scale because they form almost no bonds.
- The difference in electronegativity between two atoms shows how unequal the sharing is, from non-polar covalent through polar covalent to ionic, with boundaries that are only a rough guide.