3.1 Electronegativity & Bonding
🎯 Learning outcomes
After completing this lesson, you should be able to:
- define electronegativity
- explain the factors influencing the electronegativities of the elements
- state and explain the trends in electronegativity across a period and down a group of the Periodic Table
- use differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds
1🧲What is electronegativity?
Electronegativity is the power of an atom to attract the electrons in a covalent bond towards itself.
The greater an atom’s electronegativity value, the greater its power to pull the shared pair of bonding electrons towards itself.
For Groups 1 to 17, the pattern is straightforward:
- electronegativity increases across a period (left to right)
- electronegativity increases up each group (bottom to top)
These two trends combine to put fluorine — top-right of the periodic table (excluding the noble gases) — at the very top of the electronegativity scale.
2📈Trends across a period and down a group
Plotting electronegativity values for the elements of Period 3 shows a clear rising trend from left to right:
The same pattern appears going down a group — but here, electronegativity decreases, since “increasing up a group” is the same as “decreasing down a group”. The halogens (Group 17) show this clearly:
3⚛️Factors influencing electronegativity
Three factors together determine how strongly an atom attracts a shared pair of bonding electrons:
⚡ Nuclear charge
Atoms in the same period with a greater (more positive) nuclear charge attract the bonding pair more strongly.
📏 Atomic radius
Atoms in the same group where the outer electrons sit farther from the nucleus attract the bonding pair less strongly — the pull of the nucleus weakens with distance.
🛡️ Shielding effect
The more inner shells and sub-shells of electrons an atom has, the lower the effective nuclear charge felt by the bonding electrons.
Moving across a period, nuclear charge increases while shielding stays roughly constant (same number of inner shells) — so electronegativity rises. Moving down a group, atomic radius and shielding both increase, outweighing the rise in nuclear charge — so electronegativity falls.
4⚖️Predicting ionic vs. covalent bonds
The most widely used scale of electronegativity values is the Pauling electronegativity scale, NP.
| Li | Be | B | C | N | O | F |
|---|---|---|---|---|---|---|
| 1.0 | 1.5 | 2.0 | 2.5 | 3.0 | 3.5 | 4.0 |
| Na | Mg | Al | Si | P | S | Cl |
| 0.9 | 1.2 | 1.5 | 1.8 | 2.1 | 2.5 | 3.0 |
Pauling electronegativity has no units. Carbon (NP = 2.5) and hydrogen (NP = 2.1) both have electronegativities lower than many other non-metallic elements — worth remembering, since it’s easy to assume carbon is highly electronegative.
The difference between the electronegativity values of two bonded atoms tells you what kind of bond is likely to form:
| Electronegativity difference | Bond type | Example |
|---|---|---|
| high (2.0 or more) | ionic | NaCl: Na = 0.9, Cl = 3.0 → difference = 2.1 |
| intermediate (around 1.0) | polar covalent, with some ionic character | — |
| low (below 1.0) | covalent | CH₄: C = 2.5, H = 2.1 → difference = 0.4 |
| zero | pure covalent, no ionic character | Cl–Cl → difference = 0 |
Sodium chloride: Na (0.9) and Cl (3.0) differ by 2.1 — a large difference, so NaCl is predicted to be ionic.
Methane: C (2.5) and H (2.1) differ by only 0.4 — a small difference, so the C–H bond is predicted to be covalent.
These boundaries (≈2.0 for ionic, <1.0 for covalent) are guidelines, not sharp cut-offs — bonding exists on a continuous spectrum from purely covalent to purely ionic. A compound with an intermediate difference (around 1.0) is best described as having a polar covalent bond with some ionic character, rather than being firmly one or the other.
📌 Exam tips & quick summary
✅ Do remember
- electronegativity = power of an atom to attract the bonding electron pair in a covalent bond
- increases across a period, increases up a group — fluorine is the most electronegative element
- three factors: nuclear charge, atomic radius, shielding
- Pauling scale has no units; difference in values predicts bond type
- large difference (≥2.0) → ionic; small difference (<1.0) → covalent; zero → pure covalent
🚫 Do not confuse
- “increases up a group” is the same as “decreases down a group” — don’t state both as if they’re separate trends
- carbon and hydrogen have relatively low electronegativities compared to other non-metals — a common misconception
- the ionic/covalent boundary values are guidelines — bonding is a spectrum, not a strict on/off switch
🔁 Quick recap
- Electronegativity: power to attract bonding electrons in a covalent bond
- Trend: ↑ across a period, ↑ up a group; F is highest
- Factors: nuclear charge, atomic radius, shielding
- Bond prediction: difference ≥2.0 → ionic; <1.0 → covalent; 0 → pure covalent