Electronegativity and bonding

1. Overview

When two atoms share a pair of electrons, one atom usually pulls that pair harder than the other. Electronegativity measures that pull. This section explains which features of an atom control its electronegativity, why values rise across a period and fall down a group, and how the difference in electronegativity between two atoms predicts whether they form ionic or covalent bonds. The same ideas come back in bond polarity, hydrogen bonding, oxidation numbers and organic reaction mechanisms.

Key Definitions

• Electronegativity: the power of an atom to attract electrons to itself; in a covalent bond, these are the bonding (shared) pair of electrons.
• Pauling electronegativity value: a number on a relative scale, with no unit, that compares how strongly atoms attract bonding electrons; fluorine, the most electronegative element, has the highest value, 4.0.
• Nuclear charge: the total positive charge on the nucleus of an atom, set by its number of protons (its atomic number).
• Shielding: the repulsion from electrons in inner shells and inner sub-shells, which reduces the attraction between the nucleus and the outer electrons.
• Atomic radius: half the distance between the nuclei of two neighbouring atoms of the same element that are bonded together; it is a measure of the size of an atom.
• Covalent bond: the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons.
• Polar covalent bond: a covalent bond between atoms with different electronegativities, so the bonding pair is shared unequally; the more electronegative atom carries a partial negative charge (δ−) and the other atom a partial positive charge (δ+).
• Ionic bond: the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).

Content

What electronegativity is

Electronegativity is the power of an atom to attract electrons to itself. In a covalent bond, those electrons are the bonding pair: the two electrons the atoms share. The more electronegative atom pulls the bonding pair closer to its own nucleus.

Three points make the idea precise:

• It describes an atom in a bond, not an ion or a whole molecule.
• It is not an energy. Values come from the Pauling scale, a relative scale with no unit. Fluorine is the most electronegative element, at 4.0.
• Helium, neon and argon have no Pauling value, because they form no stable compounds. The data booklet leaves them out.

Values are given in questions or in the data booklet, so learn the trends and the reasons, not the numbers. If a question prints its own values, use those, even where they differ slightly from the booklet's.

The same factors control electronegativity and first ionisation energy, so the two are easy to confuse. Keep them apart:

	Electronegativity	First ionisation energy
What it describes	how strongly an atom in a bond attracts the bonding pair	the energy needed to remove one mole of electrons from one mole of gaseous atoms, forming one mole of gaseous 1+ ions
Unit	none (relative scale)	kJ mol⁻¹
Across Period 3	rises at every step, from Na (0.9) to Cl (3.0)	rises overall, but dips at Al and at S

The three factors that control electronegativity

The bonding pair is in the outer shell of each bonded atom. An atom's electronegativity depends on how strongly its nucleus attracts electrons in that shell, and three factors decide this:

Factor	If it is larger	Attraction for the bonding pair	Electronegativity
Nuclear charge (number of protons)	more positive charge pulling	stronger	higher
Atomic radius (distance from the nucleus to the outer shell)	the bonding pair is further away	weaker	lower
Shielding by inner shells and sub-shells	more inner electrons repel the bonding pair	weaker	lower

The factors are linked, so an explanation must weigh them against each other. A quick check is the net charge felt by the outer shell:

$$\text{net charge} \approx \text{number of protons} - \text{number of inner electrons}$$

Chlorine (17 protons, 10 inner electrons) has a net charge of about +7; sodium (11 protons, 10 inner electrons) about +1. This count is a simple model, not an exact value: inner electrons do not shield perfectly, and electrons in the same shell also repel each other a little. It still gives the right order across Periods 2 and 3. In exam answers, still name the three factors: nuclear charge, shielding and distance.

Inner sub-shells shield unequally. Electrons in inner s and p sub-shells shield the outer electrons well. Electrons in a filled inner d sub-shell shield them less well. This matters for the Period 4 elements that follow the d-block (see "A Period 4 twist" below).

The trend across a period

Electronegativity increases across a period (from left to right). The data booklet values for Periods 2 and 3 show it:

Group	1	2	13	14	15	16	17
Period 2	Li 1.0	Be 1.6	B 2.0	C 2.5	N 3.0	O 3.5	F 4.0
Period 3	Na 0.9	Mg 1.3	Al 1.5	Si 1.9	P 2.2	S 2.6	Cl 3.0

The explanation is a chain of four links:

1. Nuclear charge increases. Each element has one more proton than the one before it (sodium 11 up to chlorine 17).
2. Shielding stays about the same. Each new electron goes into the same outer shell, so the inner electrons do not change. Every Period 3 atom has the same 10 inner electrons.
3. Atomic radius decreases. The larger net charge pulls the outer shell closer to the nucleus.
4. So the bonding pair is closer to a nucleus that attracts it more strongly, and electronegativity increases.

Unlike first ionisation energy, electronegativity has no dips at aluminium or sulfur. Those ionisation-energy dips depend on exactly which electron is removed: the single 3p electron in aluminium, or one of a pair of 3p electrons in sulfur. Electronegativity does not involve removing an electron, so do not use spin-pair repulsion or the 3p sub-shell to explain the trend across a period.

The trend down a group

Electronegativity decreases down a group. For example, Group 1 falls from Li 1.0 to Na 0.9 to K 0.8, Group 2 from Be 1.6 to Mg 1.3 to Ca 1.0, and Group 17 from F 4.0 to Cl 3.0, with bromine and iodine lower still.

1. Atomic radius increases. Each element down the group has one more occupied shell, so the bonding pair is further from the nucleus.
2. Shielding increases. There are more inner shells of electrons between the nucleus and the bonding pair.
3. Nuclear charge increases too, but the extra protons are largely cancelled by the extra inner electrons. Lithium, sodium and potassium all have a net charge of about +1.
4. The greater distance and extra shielding outweigh the larger nuclear charge, so the attraction for the bonding pair is weaker and electronegativity decreases.

A Period 4 twist. In Groups 13 to 16 of the data booklet, the fall stops, or even reverses, after Period 3. Gallium (1.8) is more electronegative than aluminium (1.5), even though it is lower in Group 13, and germanium (2.0) is slightly above silicon (1.9). Arsenic equals phosphorus (2.2), and selenium equals sulfur (2.6).

Here the simple count fails, and gallium shows why. Gallium has 18 more protons than aluminium and 18 more inner electrons, but ten of those inner electrons are in the 3d sub-shell, filled across the d-block. 3d electrons shield less well than inner s and p electrons, so the extra protons are not fully cancelled, and gallium's outer electrons feel a larger net attraction than aluminium's. Gallium's atoms are about the same size as aluminium's, so the larger net attraction gives gallium the higher value. The same poor shielding by the 3d sub-shell holds up the values of germanium, arsenic and selenium. The main trends are enough for most questions; this is the case that shows why shielding by sub-shells matters.

Worked example 1: explaining three values

Question: The Pauling electronegativities of nitrogen, phosphorus and silicon are 3.0, 2.2 and 1.9. Explain why phosphorus is (a) less electronegative than nitrogen and (b) more electronegative than silicon.

Step 1: collect the facts for each atom.

Atom	Protons	Electronic configuration	Inner electrons	Outer shell	Net charge (protons − inner electrons)
N	7	1s² 2s² 2p³	2	second	+5
P	15	1s² 2s² 2p⁶ 3s² 3p³	10	third	+5
Si	14	1s² 2s² 2p⁶ 3s² 3p²	10	third	+4

Step 2: compare phosphorus with nitrogen (same group).

• Phosphorus has a larger nuclear charge (15 protons against 7).
• It also has 8 more inner electrons, so its bonding pair is more shielded. The net charge is about +5 for both atoms.
• Phosphorus's bonding electrons are in the third shell, not the second, so they are further from the nucleus.
• The greater distance and extra shielding outweigh the larger nuclear charge.

Step 3: compare phosphorus with silicon (same period).

• Phosphorus has one more proton (15 against 14).
• Both atoms have the same 10 inner electrons, so the shielding is about the same.
• The net charge on phosphorus's outer shell is larger (+5 against +4), so its atomic radius is smaller.

Answer: (a) Phosphorus is less electronegative than nitrogen because its bonding pair is in a shell further from the nucleus and is more shielded by inner electrons; these effects outweigh its larger nuclear charge, so it attracts the bonding pair less strongly. (b) Phosphorus is more electronegative than silicon because it has a larger nuclear charge with about the same shielding and a smaller atomic radius, so it attracts the bonding pair more strongly.

Using electronegativity differences to predict bonding

Subtract the smaller Pauling value from the larger one. The size of the difference tells you how unequally the two atoms share electrons.

• Difference of 0: both atoms attract the bonding pair equally. The bond is non-polar covalent, as in Cl₂ (3.0 − 3.0 = 0).
• Small difference: the more electronegative atom pulls the bonding pair closer, giving a polar covalent bond. The more electronegative atom carries a partial negative charge, δ−, and the other atom a partial positive charge, δ+. In H-Cl the difference is 3.0 − 2.1 = 0.9, and chlorine is δ−.
• Large difference: the more electronegative atom attracts electrons so strongly that they transfer completely from the other atom. Ions form, and the bonding is ionic. In NaCl the difference is 3.0 − 0.9 = 2.1.

The larger the difference, the more polar the bond, until the bonding becomes ionic. The change is gradual, and there is no fixed difference at which covalent bonding turns into ionic bonding. Instead, compare the values you are given: of several pairs of elements, the pair with the largest difference is the one most likely to bond ionically.

A number on its own cannot decide, because the same difference can belong to an ionic compound and to a covalent one:

Difference	Ionic compound	Covalent compound
2.0	CaCl₂ (3.0 − 1.0)	BF₃ (4.0 − 2.0)
2.1	NaCl (3.0 − 0.9)	SiF₄ (4.0 − 1.9)

So use the difference together with the kind of elements:

• A bond between two non-metal atoms is covalent, even when the difference is large, as in B-F (2.0). Count boron and silicon with the non-metals here: they do not form simple positive ions.
• A Group 1 metal, or a Group 2 metal other than beryllium, forms ionic compounds with oxygen and the halogens, even when the difference is moderate, as in MgCl₂ (1.7).
• For other metals, such as aluminium, compare the differences. Aluminium chloride (1.5) is covalent, while aluminium oxide (2.0) is ionic.

The chlorides of Period 3 show the gradual change. The difference shrinks from sodium to phosphorus, and the bonding switches from ionic to covalent between MgCl₂ and AlCl₃:

Chloride	NaCl	MgCl₂	AlCl₃	SiCl₄	PCl₅
Difference	2.1	1.7	1.5	1.1	0.8
Bonding	ionic	ionic	covalent	covalent	covalent

Some ionic compounds have partly covalent bonding, but at this level treat each compound as either ionic or covalent.

Worked example 2: predicting bond types

Question: Pauling values: K 0.8, B 2.0, Si 1.9, N 3.0, O 3.5, Cl 3.0. Predict the type of bonding between (a) potassium and oxygen, (b) boron and chlorine, (c) silicon and oxygen, (d) nitrogen and chlorine. For each covalent bond, state which atom, if any, carries δ−, and identify the most polar covalent bond.

Step 1: find each difference (larger value minus smaller value).

Pair	Working	Difference
K and O	3.5 − 0.8	2.7
B and Cl	3.0 − 2.0	1.0
Si and O	3.5 − 1.9	1.6
N and Cl	3.0 − 3.0	0.0

The data have one decimal place, so give each difference to one decimal place.

Step 2: use each difference together with the kind of elements.

• (a) 2.7 is a large difference, and potassium is a Group 1 metal: ionic. Potassium forms K⁺ ions and oxygen forms O²⁻ ions (K₂O).
• (b) 1.0 is a small difference between two non-metal atoms (boron forms no simple positive ions): polar covalent, with chlorine δ− and boron δ+. (The BCl₃ molecule as a whole is non-polar, because its symmetrical shape cancels the bond polarities.)
• (c) 1.6 is a moderate difference, but silicon forms no simple positive ions, so this is a bond between two non-metal atoms: polar covalent, with oxygen δ−.
• (d) With these values the difference is zero, so the N-Cl bond is non-polar covalent, although nitrogen and chlorine are different elements.

Step 3: rank the covalent bonds. The largest difference among them is Si-O (1.6), then B-Cl (1.0), then N-Cl (0.0).

Answer: (a) ionic; (b) polar covalent, Cl δ−; (c) polar covalent, O δ−; (d) non-polar covalent on these values, no partial charges. The most polar covalent bond is Si-O, with a difference of 1.6.

Key Equations

Electronegativity difference (EN = Pauling electronegativity, which has no unit; give the difference to the same number of decimal places as the data, usually one): $$\Delta\text{EN} = \text{EN}_{\text{larger}} - \text{EN}_{\text{smaller}}$$

Reading the difference (a comparison, not a cut-off: there is no fixed boundary value):

• $\Delta\text{EN} = 0$: non-polar covalent
• the larger $\Delta\text{EN}$, the more polar the covalent bond
• a large $\Delta\text{EN}$ between a Group 1 or Group 2 metal (not beryllium) and oxygen or a halogen: ionic
• two non-metal atoms (boron and silicon included): covalent, whatever the value of $\Delta\text{EN}$

Net charge felt by the outer shell (a simple model for comparing atoms across a period, not an exact value): $$\text{net charge} \approx \text{number of protons} - \text{number of inner electrons}$$

Common Mistakes to Avoid

• Mistake: defining electronegativity as an energy, or as the power of an ion, a molecule or an element instead of an atom. Fix: it is the power of an atom to attract electrons (the bonding pair) to itself, and it has no unit. In H-Cl, the chlorine atom attracts the shared pair more strongly, so chlorine is the more electronegative atom.
• Mistake: explaining electronegativity with ionisation energy or spin-pair repulsion, or bringing the 3p sub-shell into the trend across a period. Fix: talk about the attraction between the nucleus and the bonding pair. First ionisation energy falls from P (1060) to S (1000 kJ mol⁻¹), yet electronegativity rises from 2.2 to 2.6, so a spin-pair argument cannot explain electronegativity.
• Mistake: mentioning the distance or the number of shells without saying what it does. Fix: finish the link. "Chlorine's bonding pair is in the third shell, further from the nucleus than fluorine's in the second, so it is attracted less strongly" explains why Cl (3.0) is below F (4.0).
• Mistake: saying that nuclear charge decreases down a group, or leaving it out. Fix: nuclear charge increases down a group (fluorine has 9 protons, chlorine 17). State that the greater distance and extra shielding outweigh it.
• Mistake: saying that shielding increases across a period. Fix: the new electrons go into the same outer shell, so the inner electrons stay the same. From sodium to chlorine, every atom has 10 inner electrons while the nuclear charge rises from 11 to 17.
• Mistake: deciding the bond type from one cut-off number, or from "metal with non-metal" alone. Fix: compare the differences and check the kind of elements. NaCl and SiF₄ both have a difference of 2.1, yet NaCl is ionic and SiF₄ is covalent, because silicon forms no simple positive ions. AlCl₃ (1.5) is covalent even though aluminium is a metal.
• Mistake: using electronegativity to explain properties that depend on forces between molecules. Fix: bromine is less volatile than chlorine because Br₂ molecules have more electrons than Cl₂ molecules (70 against 34), so the instantaneous dipole-induced dipole forces are stronger. Electronegativity plays no part.

Exam Tips

• Define electronegativity in the set form: "the power of an atom to attract electrons to itself". Say "atom", and in a bond add "the bonding pair of electrons".
• Explain a trend as a chain of links that ends in the attraction for the bonding pair. Across a period: nuclear charge increases, shielding stays about the same, atomic radius decreases. Down a group: more shells, more shielding and a larger radius outweigh the larger nuclear charge.
• Compare two named elements with comparative words and real numbers: "phosphorus has 15 protons and silicon 14; both have 10 inner electrons". Avoid vague phrases such as "the charge is big".
• Use or predict with electronegativity values: write each subtraction, state the difference, then name the bond type, giving the kind of elements as well as the number. For a polar covalent bond, say which atom is δ−. When comparing, say which difference is larger.
• Suggest means applying ideas to an unfamiliar case. If a value breaks the trend just after the d-block (gallium above aluminium), point to the poor shielding by the inner 3d sub-shell.