Quick answer: Electronegativity is the ability of a bonded atom to attract the shared pair of electrons towards itself. It generally increases across a period — from 0.93 (sodium) to 3.16 (chlorine) on the Pauling scale — because nuclear charge rises while shielding barely changes. It generally decreases down a group, since added shells increase both shielding and distance faster than the rising nuclear charge can compensate. Fluorine, at 3.98, is the most electronegative element on the periodic table. The electronegativity difference between two bonded atoms determines bond polarity, and the resulting shift toward more covalent bonding across Period 3 is a specific, examinable learning outcome in H2 Chemistry.
Most students can recite that electronegativity increases across a period — far fewer can explain why using nuclear charge and shielding together, or apply the idea to bond and molecular polarity under exam conditions. This guide covers the trend with real Pauling values, five worked Cambridge-style questions, and the Period 3 bonding shift the syllabus specifically expects. This kind of reasoned, exam-focused explanation is what H2 Chemistry tuition at Pamela’s Place is built around, in line with the H2 Chemistry syllabus (9476).
Table of Contents
What Is Electronegativity?
Electronegativity is the ability of an atom in a molecule to attract the bonding pair of electrons towards itself — how strongly an atom pulls shared electrons when bonded to another atom. This makes it distinct from ionisation energy (removing an electron from an isolated atom) and electron affinity (adding an electron to an isolated atom): electronegativity specifically describes attraction for a shared pair within a bond.
Picture two students sharing a whiteboard marker. If one student keeps pulling it closer, that student has the stronger pull. In a polar bond, the more electronegative atom attracts the shared electrons more strongly and becomes slightly more negative.
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Electronegativity and Nuclear Attraction
Electronegativity depends on how strongly the nucleus attracts the bonding electrons — the same nuclear charge and shielding model that drives every other periodic trend in H2 Chemistry:
- Nuclear charge — a higher nuclear charge generally produces a stronger attraction for electrons.
- Shielding — inner electrons reduce the attraction an outer electron (or bonding pair) feels from the nucleus.
- Distance — a bonding electron farther from the nucleus feels a weaker attractive force, which matters most when comparing elements down a group.
Electronegativity Across Period 3: Real Data
| Element | Na | Mg | Al | Si | P | S | Cl |
| Electronegativity (Pauling scale) | 0.93 | 1.31 | 1.61 | 1.90 | 2.19 | 2.58 | 3.16 |
Across the period, protons increase while electrons are added to the same principal shell, so shielding barely changes. The nucleus attracts the bonding pair progressively more strongly, and atomic size also decreases, holding the bonding electrons closer — both push electronegativity up from sodium to chlorine.
A common mistake is attributing this entirely to “more protons.” A complete answer states that nuclear charge increases while shielding stays roughly constant, so the attraction for the bonding pair strengthens.
Electronegativity Down Groups 2 and 17: Real Data
| Group 2 | Mg | Ca | Sr | Ba |
| Electronegativity | 1.31 | 1.00 | 0.95 | 0.89 |
| Group 17 | F | Cl | Br | I |
| Electronegativity | 3.98 | 3.16 | 2.96 | 2.66 |
Down a group, each step adds an occupied electron shell. Nuclear charge does increase, but the added shielding and greater distance reduce the attraction for the bonding pair more than the higher nuclear charge can compensate — so electronegativity falls, as shown by chlorine (3.16) being more electronegative than bromine (2.96) and iodine (2.66).
Why Fluorine Is the Most Electronegative Element
Fluorine’s value of 3.98 is the highest on the periodic table: a strong nuclear attraction combined with a very small atomic size and minimal shielding between the nucleus and the bonding pair. This strong pull explains why fluorine forms highly polar bonds — in H–F, electron density shifts strongly towards fluorine.
Electronegativity and Bond Polarity
Two identical atoms have equal electronegativity and share the bonding pair equally — a non-polar bond. Two different atoms usually have different electronegativities, so the more electronegative atom pulls the shared pair more strongly, producing a polar bond. Think of it as a tug-of-war: equal pull means non-polar, unequal pull shifts electron density towards the stronger puller.
- H–H is non-polar — identical electronegativity.
- H–Cl is polar — chlorine (3.16) attracts the pair more strongly than hydrogen (2.20).
- O–H is strongly polar — oxygen (3.44) is well above hydrogen.
- C=O is polar — oxygen pulls the bonding electrons more strongly than carbon (2.55).
The more electronegative atom carries a partial negative charge (δ−); the other carries a partial positive charge (δ+).
Bond Polarity vs Molecular Polarity
A molecule can contain polar bonds and still have no overall dipole, if the bond dipoles cancel due to molecular shape. Carbon dioxide has two polar C=O bonds, but its linear shape means the dipoles point in opposite directions and cancel — CO₂ is non-polar overall. Water also has polar O–H bonds, but its bent shape means the dipoles don’t cancel — water is polar overall.
In exam questions, check both the electronegativity difference and the molecular shape before concluding whether a molecule has an overall dipole.
Electronegativity and Bonding Across Period 3
As electronegativity rises across Period 3, the electronegativity difference between bonded elements changes too, contributing to a gradual shift from more ionic towards more covalent bonding in many Period 3 compounds. Sodium and chlorine have a large electronegativity difference, so sodium chloride is a clearly ionic lattice of Na⁺ and Cl⁻. Silicon and chlorine have a much smaller difference, so silicon tetrachloride is a simple molecular covalent structure instead.
The H2 Chemistry syllabus specifically expects students to relate this variation in bonding across Period 3 oxides and chlorides to electronegativity — with aluminium chloride treated as a named exception in that learning outcome, so it’s worth learning the syllabus-specific treatment rather than assuming it fits the general pattern cleanly.
Where This Page Fits
The complete Periodic Table guide for H2 Chemistry covers electronegativity alongside atomic radius, ionic radius, melting point, and conductivity at an overview level across the full property set. This page is the fuller breakdown of electronegativity specifically — real Pauling data, worked exam questions, and the Period 3 bonding application — the same relationship Ionisation Energy Explained has to that same pillar for its own trend.
Common Electronegativity Mistakes
- Saying electronegativity increases down a group because nuclear charge increases — shielding and distance increase faster and dominate the trend.
- Ignoring shielding when explaining a group trend — nuclear charge alone is an incomplete answer.
- Confusing electronegativity with ionisation energy or electron affinity — these describe different processes (bonding attraction vs isolated-atom electron removal vs isolated-atom electron addition).
- Assuming any molecule with polar bonds must be polar overall — always check molecular shape.
- Labelling a bond ionic or covalent without considering the electronegativity difference behind that classification.
How to Answer Electronegativity Questions
The same charge-shielding-distance reasoning behind the effective nuclear charge framework applies here directly:
Across a period: state the trend → note the rise in nuclear charge → note that shielding barely changes (same principal shell) → conclude the bonding pair feels stronger attraction.
Down a group: state the trend → note the extra occupied shell → note the resulting rise in shielding and distance → conclude the attraction weakens despite the higher nuclear charge.
Cambridge-Style Exam Questions
Original practice questions written to reflect H2 Chemistry reasoning — not reproduced from Cambridge examination papers.
Question 1. Electronegativity generally increases from sodium to chlorine across Period 3. Explain this trend.
Answer: Nuclear charge increases as protons are added. The added electrons enter the same principal shell, so shielding does not increase significantly. The bonding pair therefore experiences a stronger attraction towards the nucleus, so electronegativity increases.
Question 2. Explain why electronegativity generally decreases from chlorine to iodine down Group 17.
Answer: An additional occupied electron shell is added at each step down the group. This increases shielding and increases the distance between the nucleus and the bonding pair, reducing the attraction despite the higher nuclear charge.
Question 3. Hydrogen chloride contains a polar bond. State which atom carries the partial negative charge and explain why.
Answer: Chlorine, because chlorine is more electronegative than hydrogen and attracts the bonding pair more strongly.
Question 4. Carbon dioxide contains polar C=O bonds but has no overall dipole. Explain why.
Answer: The molecule is linear, so the two equal bond dipoles point in opposite directions and cancel, leaving no overall molecular dipole.
Question 5. A student writes: “Electronegativity increases down Group 17 because each atom has more protons.” Identify the error.
Answer: The explanation considers nuclear charge but ignores the increase in shielding and electron-shell distance down the group — these effects dominate, so electronegativity actually decreases.
How Pamela’s Place Supports H2 Chemistry Revision
At Pamela’s Place, electronegativity is taught as one application of the same nuclear-charge-and-shielding model used across every periodic trend, with real data and worked Cambridge-style questions rather than a memorised arrow diagram. This sits inside the same small Omakase groups, capped at 7 students, used across our Chemistry programmes.
Electronegativity FAQs
What is electronegativity?
Electronegativity is the ability of an atom in a molecule to attract the shared bonding pair of electrons towards itself.
- It’s distinct from ionisation energy and electron affinity, which describe isolated-atom processes rather than attraction within a bond.
What happens to electronegativity across Period 3?
It generally increases — from 0.93 (sodium) to 3.16 (chlorine) — because nuclear charge rises while shielding barely changes.
- Atomic size also decreases across the period, holding the bonding pair closer to the nucleus.
Why does electronegativity decrease down a group?
Added electron shells increase both shielding and distance from the nucleus, and these outweigh the rise in nuclear charge.
- Chlorine (3.16) is more electronegative than bromine (2.96) and iodine (2.66) for exactly this reason.
Which element is the most electronegative, and why?
Fluorine, at 3.98 — a strong nuclear attraction, very small atomic size, and minimal shielding combine to make it the most electronegative element on the periodic table.
Are all molecules with polar bonds polar overall?
No — molecular shape determines whether the individual bond dipoles cancel.
- Carbon dioxide has polar bonds but is non-polar overall because its linear shape cancels the dipoles; water’s bent shape means its dipoles don’t cancel.
Why is aluminium chloride treated as an exception in Period 3 bonding questions?
The H2 Chemistry syllabus specifically names aluminium chloride as an exception in the learning outcome on bonding variation across Period 3 oxides and chlorides, rather than expecting it to fit the general ionic-to-covalent pattern cleanly.
- Worth learning the syllabus-specific treatment directly rather than inferring it from the general trend.
Does this apply to IP Chemistry too?
Yes — electronegativity and periodic trends are introduced on the IP track using the same reasoning, without an O-Level checkpoint forcing early mastery of the model.
- See IP Chemistry tuition for how this foundation is built earlier on that track.
Do O-Level Chemistry students need this level of detail?
Not to the same depth — O-Level Chemistry touches electronegativity and bond polarity more descriptively, without Period 3 bonding-variation or Pauling-scale precision.
- See O-Level Chemistry tuition for the equivalent foundation at O-Level depth.
