So far we have talked about a shared pair as if it sits neatly in the middle. Usually it doesn’t. One atom pulls harder than the other, drags the pair towards itself, and the bond ends up lopsided — that is a polar bond.
📚 What you need to know
Electronegativity is the ability of an atom to attract the shared pair of electrons in a covalent bond.
Values come from the Pauling scale, in Section 9 of the data booklet. Fluorine is the highest at 4.0.
Electronegativity increases across a period and decreases down a group.
Same electronegativity → equal sharing → nonpolar bond.
Different electronegativity → unequal sharing → polar bond, with δ+ on the weaker puller and δ− on the stronger one.
The bigger the difference (ΔEN), the more polar the bond.
Electronegativity
Every atom pulls on the electrons in a bond, but some pull harder than others. How hard an atom pulls is its electronegativity, and it depends on two things: how many protons are in the nucleus, and how far away the outer shell is.
Across a period the nuclear charge goes up but the electrons stay in the same shell, so the pull gets stronger.
Down a group the outer shell is further from the nucleus and better shielded, so the pull gets weaker.
Put those together and the strongest pullers sit in the top right of the periodic table. Fluorine is the champion.
Pauling electronegativity values for the first three periods. Notice how they climb from left to right and fall as you go down.
You don’t have to memorise these numbers — they are in the data booklet. What you do need is the trend, and the habit of always subtracting the two values when a question asks about polarity.
Nonpolar bonds: a fair fight
When the two atoms are the same element, they have identical electronegativities. Neither wins the tug of war, the shared pair sits exactly in the middle, and there are no partial charges anywhere. This is a nonpolar covalent bond.
This is always true for the diatomic elements: H2, O2, N2, Cl2.
Polar bonds: an unfair fight
When the two atoms are different, one of them pulls harder. The shared pair drifts towards the stronger puller, and that end of the bond becomes slightly negative while the other end becomes slightly positive. We write these partial charges as δ− and δ+.
The more electronegative atom gets δ− — extra electron density.
The less electronegative atom gets δ+ — it has been partly stripped.
In Cl₂ the shared pair is shared evenly. In HCl the chlorine wins the tug of war, so the electron density piles up on its side.
💡 What δ actually means
δ (delta) means partial or slightly. It is not a full charge.
δ+ is not the same as +. In HCl the hydrogen has not lost its electron — it has just lost its fair share of it.
A molecule with δ+ and δ− ends is still neutral overall. The charges add up to zero.
Dipoles and how to draw them
When charge is separated across a bond like this, we say the bond has a dipole. The dipole moment is just a measure of how big that separation is — the bigger the electronegativity difference, the bigger the dipole moment.
The dipole is drawn with a special arrow. Get the direction right or you lose the mark:
The crossed tail marks the δ+ end and the arrowhead points to the δ− end — the direction the electrons have shifted.
How polar is polar?
Bond polarity is not an on/off switch. Work out ΔEN — the difference between the two electronegativity values — and it tells you roughly where on the scale the bond sits. There is no sharp cut-off, but as a rough guide:
Ionic and covalent bonding are the two ends of one scale, not two separate categories.
This is a really useful way to think about Topic 2 as a whole. Ionic and covalent are not two separate boxes — they are the two ends of one continuous scale. A bond like H–Cl is genuinely somewhere in between.
WORKED EXAMPLE
Using the values C = 2.6, N = 3.0, O = 3.4 and F = 4.0, place the bonds in NO, OF2, CO and CF4 in order of increasing polarity.
Work out ΔEN for each bondN–O: 3.4 − 3.0 = 0.4O–F: 4.0 − 3.4 = 0.6C–O: 3.4 − 2.6 = 0.8C–F: 4.0 − 2.6 = 1.4Bigger difference = more polarNO < OF₂ < CO < CF₄
WORKED EXAMPLE
Explain why the O–H bond in water is polar, and state which atom carries the δ−.
Oxygen and hydrogen have different electronegativitiesO = 3.4, H = 2.2, so ΔEN = 1.2.Oxygen pulls the shared pair towards itselfδ− on the oxygen, δ+ on each hydrogenThis one bond dipole is behind almost everything water does — hydrogen bonding, solubility, its high boiling point.
⚠️ Common mix-up
Electronegativity is about bonds, not atoms on their own. It describes how an atom behaves inside a covalent bond.
Don’t confuse it with electron affinity or with ionisation energy. Electronegativity is a comparison, not a measured energy.
δ+ is not +. Partial charges are fractions of a charge, not whole ions.
A polar bond does not always mean a polar molecule. CO2 is full of polar bonds and is completely nonpolar — that is the whole point of the next set of notes.
Up next: Molecular Polarity — how bond dipoles combine with molecular shape to decide whether the whole molecule is polar.
Want this explained one-to-one?
Book a free session with an experienced IB Chemistry tutor and get your trickiest topics made simple.
