Coaching Centre
Covalent bonding and molecular shape
Covalent bonds as shared electron pairs, single, double, triple and coordinate bonds, drawing Lewis (electron-dot) structures, naming covalent compounds, VSEPR theory for predicting molecular shapes and bond angles, and bond and molecular polarity.
Carbon dioxide and water are both made of three atoms, a central atom with two others attached. Carbon dioxide is a gas that will not dissolve much in water and is unaffected by a charged rod. Water is a liquid that dissolves salt and bends towards a charged rod held near a thin stream from a tap. The difference is shape: CO₂ is a straight line, water is bent. This lesson shows how to predict a molecule's shape from its electrons, and why shape decides polarity.
The covalent bond
When two non-metal atoms bond, neither can take an electron from the other because both have high electronegativity. Instead they share a pair of electrons. Both nuclei are attracted to the shared pair between them, and that electrostatic attraction is the covalent bond. Each atom usually ends up with a full outer shell (an octet; two electrons for hydrogen).
- A single bond is one shared pair (H–H), a double bond two pairs (O=O), a triple bond three pairs (N≡N).
- More shared pairs pull the atoms closer and hold them more strongly: C–C 154 pm and 346 kJ mol⁻¹; C=C 134 pm and 614 kJ mol⁻¹; C≡C 120 pm and 839 kJ mol⁻¹.
- In a coordinate (dative) covalent bond, both shared electrons come from the same atom. Ammonia's lone pair bonds to H⁺ to make the ammonium ion, NH₄⁺; once formed, all four N–H bonds are identical.
Lewis structures
A Lewis structure (electron-dot structure) shows all the valence electrons in a molecule, as bonds (lines or pairs of dots between atoms) and lone pairs (non-bonding pairs on one atom). A reliable method:
- Add up the valence electrons of all the atoms (add one per negative charge, subtract one per positive charge).
- Write the skeleton: the least electronegative atom (never H) usually goes in the centre.
- Join each outer atom to the centre with a single bond (2 electrons each).
- Give the outer atoms lone pairs until each has an octet (H needs only 2).
- Put any remaining electrons on the central atom.
- If the central atom still lacks an octet, turn lone pairs on outer atoms into double or triple bonds.
Draw the Lewis structure of methanal (formaldehyde), CH₂O.
Valence electrons: C 4 + 2 × H 1 + O 6 = 12.
Carbon is central, bonded to H, H and O with single bonds: 6 electrons used, 6 left. Give O three lone pairs (6 electrons): all 12 used.
Carbon now has only 6 electrons around it. Convert one lone pair on O into a second C–O bond: H₂C=O. Now C has 8, O has 8 (two bonds + two lone pairs), each H has 2.
Exceptions to the octet rule. A few molecules have fewer than eight electrons around the central atom: BF₃ (boron has 6) and BeCl₂ (beryllium has 4). Elements from period 3 onwards are large enough to fit more than four electron pairs around them (often explained using their empty 3d orbitals): phosphorus in PCl₅ has 10 and sulfur in SF₆ has 12. These are called expanded octets.
Naming binary covalent compounds: name the less electronegative element first, end the second with -ide, and use prefixes for the number of atoms (mono-, di-, tri-, tetra-, penta-, hexa-). Drop "mono" on the first element. Examples: CO carbon monoxide, CO₂ carbon dioxide, N₂O₄ dinitrogen tetroxide, PCl₅ phosphorus pentachloride, SF₆ sulfur hexafluoride. Some keep traditional names: water, ammonia, methane.
VSEPR: predicting shape
Valence shell electron pair repulsion (VSEPR) theory says that the electron pairs around a central atom repel each other and move as far apart as possible. To use it, count the electron domains on the central atom: each lone pair is one domain, and each bond is one domain whether it is single, double or triple.
- 2 domains: linear (180°). 3 domains: trigonal planar (120°). 4 domains: tetrahedral (109.5°).
- The shape of the molecule is named from the positions of the atoms only. Lone pairs affect the shape but are not "seen".
- Lone pairs are held by only one nucleus, so they spread out more and repel more strongly than bonding pairs. Each lone pair squeezes the bond angles a little: CH₄ 109.5°, NH₃ 107°, H₂O 104.5°.
| Domains | Bonding | Lone pairs | Shape | Angle | Examples |
|---|---|---|---|---|---|
| 2 | 2 | 0 | linear | 180° | CO₂, HCN, BeCl₂ |
| 3 | 3 | 0 | trigonal planar | 120° | BF₃, CH₂O, SO₃ |
| 3 | 2 | 1 | bent | ≈ 117–119° | SO₂, O₃ |
| 4 | 4 | 0 | tetrahedral | 109.5° | CH₄, CCl₄, NH₄⁺ |
| 4 | 3 | 1 | trigonal pyramidal | 107° | NH₃, PH₃, H₃O⁺ |
| 4 | 2 | 2 | bent (V-shaped) | 104.5° | H₂O, H₂S |
Predict the shape and approximate bond angle of H₃O⁺.
Valence electrons: O 6 + 3 × H 1 − 1 (positive charge) = 8, which is 4 pairs.
Three pairs bond to H; one is a lone pair on O. Four domains, so a tetrahedral arrangement.
Three atoms and one lone pair: trigonal pyramidal, angle slightly less than 109.5° (about 107°, like NH₃, with which it is isoelectronic).
Polarity
In a bond between two different atoms, the more electronegative atom pulls the shared electrons closer and carries a small negative charge, δ−; the other end is δ+. This is a polar bond, and the separated charge is a bond dipole, drawn as an arrow pointing to the δ− end. Bonds with an electronegativity difference below about 0.4 (such as C–H) are treated as non-polar.
A molecule is polar if its bond dipoles do not cancel. That depends on shape:
- Symmetrical shapes with identical outer atoms cancel: CO₂ (linear), BF₃ (trigonal planar), CH₄ and CCl₄ (tetrahedral) are non-polar even though their bonds may be polar.
- Lone pairs on the central atom (bent, pyramidal) or different outer atoms break the symmetry: H₂O, NH₃, HCN and CHCl₃ are polar.
Explain why CCl₄ is non-polar but CHCl₃ (chloroform) is polar.
Both are tetrahedral with no lone pairs on carbon. C–Cl bonds are polar (ΔEN = 0.61); C–H is almost non-polar (ΔEN = 0.35).
In CCl₄ four identical dipoles point to the corners of a tetrahedron and cancel exactly. In CHCl₃ one corner is H instead of Cl, so the dipoles do not cancel and there is a net dipole towards the chlorine side.
- Start on NH₃. Drag to rotate it and find the view in which the molecule looks like a pyramid. Then turn off lone pairs: what shape do the atoms alone make?
- Compare CH₄, NH₃ and H₂O. All have 4 domains. Why do the bond angles fall from 109.5° to 107° to 104.5°?
- Turn on Bond dipoles and compare CCl₄ with CH₃Cl. Which is polar, and why?
- Look at CO₂ and SO₂. Both are XO₂. Explain why one is linear and the other bent.
- Extension: open SF₆ ★. How many electrons surround sulfur, and why can sulfur do this when oxygen cannot?
A quick classroom test for polarity: charge a plastic rod by rubbing it with a cloth and hold it near a thin stream of liquid from a burette. A stream of water (polar) bends noticeably towards the rod, whatever the sign of its charge, because the molecules rotate so their oppositely charged end faces it. A stream of hexane (non-polar) is barely affected.
Test yourself
What is the shape of a PH₃ molecule?
P has 3 bonds and 1 lone pair: 4 domains, 3 atoms.
How many lone pairs are on the central atom of H₂S?
S has 6 valence electrons: 2 used in bonds, 4 left as 2 lone pairs.
Which molecule is non-polar?
BF₃ is trigonal planar with three identical B–F dipoles at 120°, which cancel.
The H–N–H angle in NH₃ (107°) is less than the H–C–H angle in CH₄ (109.5°) because
Both have four domains; the lone pair occupies more space and pushes the bonds together.
The correct name for N₂O₅ is
Prefixes give the number of each atom: di + penta.
Name each molecular shape.
Classify each molecule.
A covalent bond is a shared pair of electrons attracted to two . A pair of valence electrons not involved in bonding is a . In VSEPR theory a double bond counts as electron domain. A molecule with four domains and two lone pairs on the central atom is . A bond in which both electrons come from one atom is a bond.
Draw Lewis structures for NH₃ and BF₃. Use VSEPR theory to predict the shape and bond angle of each molecule, and explain why NH₃ is polar but BF₃ is not.
NH₃: 8 valence electrons. N bonded to three H with single bonds and one lone pair on N. Four electron domains around N give a tetrahedral arrangement; with three atoms the shape is trigonal pyramidal. The lone pair repels the bonding pairs more strongly, so the angle is about 107°.
BF₃: 24 valence electrons. B bonded to three F with single bonds; each F has three lone pairs; B has no lone pair (only 6 electrons). Three domains give a trigonal planar shape with 120° angles.
Polarity: N–H (ΔEN 0.84) and B–F (ΔEN 1.94) bonds are both polar. In BF₃ the three identical dipoles are spaced evenly at 120° in a plane and cancel, so the molecule is non-polar. In NH₃ the dipoles all point towards N from below and the lone pair sits on top, so they add to a net dipole: NH₃ is polar.