B. Matter, Chemical Trends, and Chemical Bonding · Matter, Chemical Trends, and Chemical BondingPredicting Bond Type with Electronegativity
B2.5 B2.1 B3.4 A1.5About 65 minutes
Learning goals- I can calculate the electronegativity difference for a bond.
- I can classify a bond as non-polar covalent, polar covalent or ionic.
- I can show bond polarity with partial charges and dipole arrows.
- I can explain why bonding is a continuum and how molecular shape affects polarity.
Words to know
electronegativity difference (ΔEN)
The difference between the electronegativities of two bonded atoms.
non-polar covalent bond
A covalent bond in which electrons are shared equally or almost equally (ΔEN about 0 to 0.4).
polar covalent bond
A covalent bond in which electrons are shared unequally (ΔEN about 0.5 to 1.7).
partial charge (δ+ / δ−)
A small charge on an atom in a polar bond, caused by unequal sharing.
bond dipole
A bond with a partially positive end and a partially negative end.
polar molecule
A molecule with an overall uneven charge distribution, because its bond dipoles do not cancel.
Sharing is not always fair
In Lesson 8 you sorted bonds into two neat boxes: ionic (transferred electrons) and covalent (shared electrons). Real bonds are not so tidy. In a hydrogen chloride molecule, the electrons are shared, but chlorine pulls them much harder than hydrogen does. The result is a bond with a slightly negative end and a slightly positive end.
In this lesson you will use electronegativity values to predict where a bond falls on a continuum from perfectly equal sharing to complete transfer.
Electronegativity difference, ΔEN
To classify a bond, subtract the smaller electronegativity from the larger:
ΔEN = EN(higher) − EN(lower)
Then use these guidelines (the cut-off values vary a little from textbook to textbook):
- 0 to 0.4: non-polar covalent. Electrons are shared equally (or almost equally). Examples: H–H (0), Cl–Cl (0), C–H (2.6 − 2.2 = 0.4).
- 0.5 to 1.7: polar covalent. Electrons are shared unequally: the more electronegative atom gets a partial negative charge (δ−) and the other a partial positive charge (δ+). Examples: O–H (1.2), N–H (0.8), H–Cl (1.0), C–O (0.8).
- More than 1.7: ionic. The difference is so large that the electron is essentially transferred, forming ions. Examples: NaCl (3.2 − 0.9 = 2.3), MgO (3.4 − 1.3 = 2.1), KF (4.0 − 0.8 = 3.2).
The symbol δ (Greek letter delta) means "partial". A polar bond is a bond dipole: it has two poles, like a bar magnet has two ends.
| Element | H | C | N | O | F | Na | Mg | Al | Si | P | S | Cl | K | Ca | Br | I |
|---|
| EN | 2.2 | 2.6 | 3.0 | 3.4 | 4.0 | 0.9 | 1.3 | 1.6 | 1.9 | 2.2 | 2.6 | 3.2 | 0.8 | 1.0 | 3.0 | 2.7 |
Showing polarity
There are two common ways to show a polar bond:
- Partial charges: write δ+ over the less electronegative atom and δ− over the more electronegative one: δ+H–Clδ−.
- A dipole arrow: an arrow with a small cross at its tail (like a plus sign), pointing from the positive end to the negative end of the bond.
It's a continuum, not boxes
The boundaries 0.4 and 1.7 are guidelines, not laws of nature. Bonds change gradually from one type to another. Two good examples:
- H–F has ΔEN = 4.0 − 2.2 = 1.8, just over 1.7, yet hydrogen fluoride is a molecular compound (a gas at room temperature that forms molecules). It is best described as a very polar covalent bond. When the two atoms are both non-metals, treat the bond as covalent.
- AlCl₃ is made of a metal and a non-metal, but ΔEN = 3.2 − 1.6 = 1.6, so the bonds have a lot of covalent character. Indeed, aluminium chloride turns to vapour at a fairly low temperature (about 180 °C), unlike typical ionic compounds.
From polar bonds to polar molecules (a preview)
A molecule can contain polar bonds and still be non-polar overall if its shape is symmetrical, so the bond dipoles cancel out, like two equal teams pulling in opposite directions.
- CO₂ is linear (O=C=O). Each C=O bond is polar, but the two dipoles point in opposite directions and cancel: CO₂ is a non-polar molecule.
- H₂O is bent, not straight. The two O–H dipoles both point partly towards oxygen and do not cancel: water is a polar molecule, with a δ− oxygen side and a δ+ hydrogen side.
Water's polarity explains why it dissolves salt and sugar so well, a topic you will explore in the Solutions and Solubility unit. Molecular shapes are explained fully in Grade 12 Chemistry (SCH4U).
Worked exampleExample 1: Classifying bonds
Classify each bond as non-polar covalent, polar covalent or ionic: (a) N–H, (b) C–Cl, (c) Ca–O, (d) C–S, (e) Si–O.
- (a) N–H: 3.0 − 2.2 = 0.8 → polar covalent (N is δ−).
- (b) C–Cl: 3.2 − 2.6 = 0.6 → polar covalent (Cl is δ−).
- (c) Ca–O: 3.4 − 1.0 = 2.4 → ionic.
- (d) C–S: 2.6 − 2.6 = 0 → non-polar covalent.
- (e) Si–O: 3.4 − 1.9 = 1.5 → polar covalent (O is δ−).
N–H, C–Cl and Si–O are polar covalent; Ca–O is ionic; C–S is non-polar covalent.
Worked exampleExample 2: Ranking bond polarity
Rank these bonds from least to most polar: H–Br, H–F, H–Cl, H–I.
- ΔEN: H–I 2.7 − 2.2 = 0.5; H–Br 3.0 − 2.2 = 0.8; H–Cl 3.2 − 2.2 = 1.0; H–F 4.0 − 2.2 = 1.8.
- Bigger ΔEN means a more polar bond.
H–I < H–Br < H–Cl < H–F. Polarity increases up the halogen group as electronegativity increases.
Worked exampleExample 3: Partial charges
Label the partial charges on the atoms of (a) a C–O bond and (b) an N–Cl bond.
- (a) O (3.4) is more electronegative than C (2.6): δ+C–Oδ−.
- (b) N is 3.0 and Cl is 3.2 in our table: ΔEN = 0.2, so the bond is essentially non-polar. (If partial charges are shown, Cl would be very slightly δ−.)
(a) δ+C–Oδ−; (b) N–Cl is non-polar covalent by the guidelines.
Just for funA chemistry tug-of-war: when hydrogen (2.2) and chlorine (3.2) share a pair of electrons, chlorine wins, but not by enough to take the electrons home. That is why the bond is polar covalent and not ionic: chlorine gets the rope closer, but hydrogen never lets go.
Hands-on activityBend water with static electricity
You need: a plastic comb or a balloon, dry hair or a wool sweater, and a kitchen tap. (If you use a balloon, keep it away from small children; a popped balloon is a choking hazard.)
- Turn on the tap so a thin, steady stream of water flows (thinner than a pencil).
- Rub the comb through dry hair (or rub the balloon on wool) about 10 times to charge it.
- Slowly bring the comb close to the stream, without touching it.
- Observe. Then try again with the comb uncharged.
What to look for: the stream bends towards the charged comb. Chemistry textbooks use this demonstration as evidence that water molecules are polar: they turn so that their oppositely charged end faces the comb and are pulled towards it. Write a hypothesis: would a thin stream of vegetable oil (made of non-polar molecules) bend as much? (Expected answer: much less or not noticeably.) Optional with a grown-up: test it over a bowl and clean up the oil.
Your turn
1Calculate ΔEN and classify the bond: (a) O–H, (b) K–Cl, (c) C–C.
(a) 3.4 − 2.2 = 1.2, polar covalent; (b) 3.2 − 0.8 = 2.4, ionic; (c) 0, non-polar covalent.
2Which atom carries δ− in an N–O bond?
3Rank from least to most polar: C–O, C–F, C–N.
ΔEN: C–N 0.4, C–O 0.8, C–F 1.4: C–N < C–O < C–F.
4Why is the C–H bond considered non-polar covalent?
ΔEN = 2.6 − 2.2 = 0.4, small enough that the electrons are shared almost equally.
A partial positive charge: the atom has a slight positive charge because the shared electrons are pulled away from it.
6Classify the bonding in MgCl₂ using ΔEN.
7CO₂ has polar bonds but is a non-polar molecule. Explain.
CO₂ is linear and symmetrical, so the two C=O bond dipoles point in opposite directions and cancel.
More practice: warm-up
1What is the ΔEN of any bond between two atoms of the same element?
2Which bond type has ΔEN greater than 1.7?
3What symbol shows a partial charge?
4Is the O–H bond polar or non-polar?
5Which is more polar, H–F or H–Cl?
More practice: on your own
1Classify each bond and label partial charges where appropriate: (a) P–Cl, (b) S–O, (c) Li–F.
(a) 3.2 − 2.2 = 1.0, polar, δ+P–Clδ−. (b) 3.4 − 2.6 = 0.8, polar, δ+S–Oδ−. (c) 4.0 − 1.0 = 3.0, ionic (Li⁺ and F⁻).
2HF has ΔEN = 1.8 but is described as polar covalent. Explain why the guidelines should not be applied blindly.
The boundaries are guidelines on a continuum. HF is made of two non-metals, forms molecules and is a gas at room temperature, so the evidence shows it is molecular (very polar covalent), not ionic.
3Predict whether the bonding in AlCl₃ is more like that in NaCl or more like that in HCl, using ΔEN.
AlCl₃: 1.6 (polar covalent range); NaCl: 2.3; HCl: 1.0. AlCl₃ falls between, closer in ΔEN to HCl than NaCl, so it has a lot of covalent character.
4Draw (describe) a dipole arrow for the H–O bond.
Cross (plus) at the H end, arrowhead pointing to O.
5Explain, using electronegativity, why the bonds in Cl₂ are non-polar but those in HCl are polar.
In Cl₂, both atoms have the same EN (3.2), so they pull equally. In HCl, Cl (3.2) pulls more than H (2.2), so the electrons are shared unequally.
Challenge
1Arrange these compounds in order of increasing ionic character of their bonds: KF, CsF, NaCl, MgO, HCl.
ΔEN: HCl 1.0, MgO 2.1, NaCl 2.3, KF 3.2, CsF 3.2 (Cs 0.8). Order: HCl < MgO < NaCl < KF ≈ CsF. (With more precise values, Cs 0.79 vs K 0.82, CsF is slightly more ionic.)
2Methane (CH₄) and carbon tetrachloride (CCl₄) are both symmetrical (tetrahedral) molecules. Which has polar bonds, and which is a polar molecule? Explain.
C–H bonds are non-polar (0.4); C–Cl bonds are polar (0.6). Both molecules are non-polar overall: CH₄ has no bond dipoles to speak of, and CCl₄'s four equal dipoles cancel because of its symmetrical shape.
3Ammonia (NH₃) is pyramid-shaped, with the lone pair on top of the nitrogen. Predict whether NH₃ is a polar molecule, and justify.
N–H bonds are polar (0.8, N δ−). In a pyramid, the three N–H dipoles all point partly towards N and do not cancel, so NH₃ is a polar molecule (δ− at N, δ+ at the H side).
Quiz
1. What is the ΔEN for the O–H bond, using O = 3.4 and H = 2.2?
2. A bond with ΔEN = 0.9 is best classified as
3. Which bond is the most polar?
4. In a polar C–Cl bond, which atom has a partial negative charge?
5. Which compound would be classified as ionic by the ΔEN guidelines?
6. CO₂ is a non-polar molecule even though its C=O bonds are polar because
Big idea: Expectation B2.5 asks students to predict the nature of a bond (non-polar covalent, polar covalent or ionic) using electronegativity values. The key is the continuum: the bigger the difference, the more unequal the sharing.
How to help: Try the bending-water demonstration together; it is quick and memorable. Quiz your teen on a few ΔEN calculations, and ask them to explain why the more electronegative atom is δ−.
Common difficulties:
- Subtracting in the wrong order and getting a negative ΔEN. Always subtract smaller from larger.
- Treating the cut-offs (0.4, 1.7) as exact laws. Different textbooks use slightly different values; evidence (such as HF being molecular) matters.
- Confusing a polar bond with a polar molecule. A symmetrical molecule can have polar bonds and still be non-polar overall.