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Covalent bonding occurs when two non-metal atoms share one or more pairs of electrons. Unlike ionic bonding, where electrons are transferred, covalent bonding involves a mutual sharing that allows both atoms to achieve a stable electron configuration. This type of bonding is found in elements like H₂, O₂, and N₂, and in compounds such as water, carbon dioxide, and methane.
When two non-metal atoms approach each other, neither has a low enough ionisation energy to simply lose electrons. Instead, their atomic orbitals overlap, and each atom contributes one electron to a shared pair. The shared electrons are attracted to both nuclei simultaneously, and this attraction holds the atoms together.
The covalent bond is the electrostatic attraction between the shared pair of electrons and the nuclei of both atoms. This is the key definition — notice that it is still fundamentally an electrostatic interaction, just like ionic bonding, but the arrangement of charges is different.
Atoms can share one, two, or three pairs of electrons:
As the number of shared pairs increases, the bond becomes shorter and stronger. A C=C double bond is shorter and has a higher bond energy than a C–C single bond. A C≡C triple bond is shorter and stronger still.
| Bond | Bond Length / pm | Bond Energy / kJ mol⁻¹ | Bond Order |
|---|---|---|---|
| C–C | 154 | 347 | 1 |
| C=C | 134 | 614 | 2 |
| C≡C | 120 | 839 | 3 |
| N–N | 146 | 163 | 1 |
| N=N | 120 | 418 | 2 |
| N≡N | 110 | 945 | 3 |
Notice that for nitrogen the jump from single to triple bond energy is even more dramatic — the triple bond is nearly 6 times stronger than the single bond, which is why N₂ is so unreactive.
Dot-and-cross diagrams show the arrangement of electrons in a covalent bond. Electrons from one atom are shown as dots (•) and electrons from the other as crosses (×). Only outer shell (valence) electrons are shown.
For example, in a water molecule (H₂O):
When drawing these diagrams, ensure every atom achieves a full outer shell (2 for hydrogen, 8 for second-period elements — the octet rule).
In a normal covalent bond, each atom contributes one electron to the shared pair. In a dative bond (also called a coordinate bond), both electrons in the shared pair come from the same atom. The atom donating the pair is called the donor, and the atom accepting is the acceptor.
Once formed, a dative bond is identical in strength and length to an ordinary covalent bond. The only difference is the origin of the electrons.
The ammonium ion (NH₄⁺): Ammonia (NH₃) has a lone pair on nitrogen. When it reacts with H⁺ (which has an empty orbital), nitrogen donates its lone pair to form a dative bond. All four N–H bonds in NH₄⁺ are equivalent once formed.
The hydroxonium ion (H₃O⁺): Water has two lone pairs on oxygen. One lone pair can be donated to H⁺, forming H₃O⁺ with a dative bond from oxygen to hydrogen.
Aluminium chloride dimer (Al₂Cl₆): Aluminium chloride exists as a dimer because each Al atom is electron-deficient (only 6 electrons in its outer shell). A lone pair from a Cl atom on one AlCl₃ unit is donated to the Al on the other, forming two dative bonds that bridge the two units.
Carbon monoxide (CO): Carbon and oxygen share a triple bond, but one of these is a dative bond. Oxygen has a lone pair that it donates to carbon. This gives carbon a full octet: one lone pair, plus three bonding pairs (two normal covalent, one dative from oxygen).
Why does BF₃ react readily with NH₃?
BF₃ is electron-deficient — boron has only 6 electrons in its outer shell and an empty p orbital. NH₃ has a lone pair on nitrogen. The lone pair on nitrogen is donated into the empty orbital on boron, forming a dative bond: F₃B←NH₃. This is a classic Lewis acid–base reaction: BF₃ is the Lewis acid (electron pair acceptor) and NH₃ is the Lewis base (electron pair donor). The product has a tetrahedral arrangement around both B and N.
Bond energy (or bond enthalpy) is the energy required to break one mole of a particular bond in the gaseous state, with all species in the gas phase. Higher bond energy means a stronger bond.
Bond length is the distance between the nuclei of two bonded atoms. In general:
Bond energies are average values because the exact energy depends on the molecular environment. For example, the O–H bond energy in water is slightly different from the O–H bond energy in ethanol.
| Bond | Bond energy / kJ mol⁻¹ | Bond length / pm |
|---|---|---|
| H–H | 436 | 74 |
| C–H | 413 | 109 |
| C–C | 347 | 154 |
| C–O | 358 | 143 |
| C=O | 805 | 122 |
| O–H | 464 | 96 |
| N–H | 391 | 101 |
| C–Cl | 346 | 177 |
| C–F | 485 | 135 |
At A-Level, you need to understand the distinction between sigma and pi bonds:
Sigma (σ) bonds are formed by the head-on overlap of atomic orbitals. The electron density is concentrated along the axis between the two nuclei. Every single bond is a sigma bond. Sigma bonds allow free rotation around the bond axis.
Pi (π) bonds are formed by the sideways overlap of p orbitals. The electron density lies above and below the plane of the nuclei. Pi bonds are weaker than sigma bonds because the sideways overlap is less effective.
In a double bond, there is one σ bond and one π bond. In a triple bond, there is one σ bond and two π bonds.
The presence of a pi bond prevents free rotation around the bond because rotation would break the sideways overlap of the p orbitals. This has important consequences for the chemistry of alkenes (geometric isomerism).
How many σ and π bonds are in ethanoic acid, CH₃COOH?
Draw the structure: H–C(–H)(–H)–C(=O)(–O–H)
Count each bond:
Total: 7σ bonds and 1π bond.
Some molecules do not obey the octet rule:
Common exam mistake: Students sometimes claim that Period 2 elements like carbon or nitrogen can expand their octets. They cannot — Period 2 elements have no accessible d orbitals. Only Period 3 and below can exceed 8 electrons.
This lesson sits inside Edexcel 9CH0 Topic 2 — Bonding and Structure, drawn from sub-topic 2.2 (covalent bonding, dative covalent bonding, and the σ/π distinction). Although the primary assessment is on Paper 1 — Advanced Inorganic and Physical Chemistry, covalent bonding is one of the most heavily synoptic ideas in the specification.
The concepts re-surface across all three papers. On Paper 1 they reappear inside Topic 4 — Inorganic Chemistry and the Periodic Table, where the bonding character of the Period 3 oxides and chlorides is examined and candidates predict which compounds are predominantly ionic and which are predominantly covalent. On Paper 2 — Advanced Organic and Physical Chemistry, the entire organic syllabus rests on a confident understanding of covalent bonding: every functional group is defined by which covalent bonds are present and which can be broken in mechanism. On Paper 3 — General and Practical Principles in Chemistry, synoptic questions routinely ask candidates to predict bond polarity in unfamiliar molecules using electronegativity reasoning. The dative-bond extension also overlaps with Topic 15 — Transition Metals, while bond enthalpy data feed into Topic 8 — Energetics I (Hess cycles) and Topic 9 — Kinetics I (activation energies).
Refer to the official Pearson Edexcel specification document for exact wording.
Question (7 marks):
(a) Draw a dot-and-cross diagram for phosphorus trichloride, PCl3, showing only the outer-shell electrons. (2)
(b) Phosphorus pentachloride, PCl5, exists in the gas phase as a molecule with 5 P–Cl bond pairs. Explain why phosphorus is able to form 5 bonds in PCl5 whereas nitrogen cannot form an analogous NCl5. (2)
(c) State the number of σ and π bonds present in (i) hydrogen cyanide, HCN, and (ii) ethene, C2H4. (3)
Solution with mark scheme:
(a) Step 1 — count outer-shell electrons. Phosphorus has 5 outer electrons (group 5); each chlorine has 7. PCl3 therefore needs 3 P–Cl single bonds, leaving one lone pair on P and three lone pairs on each Cl.
M1 — three shared pairs drawn between P and each Cl using distinct symbols (e.g. dots for P, crosses for Cl) so the examiner can see which electron came from which atom. The single P–Cl bond is two electrons total.
A1 — one lone pair correctly shown on P and three lone pairs on each Cl. The diagram must show 8 electrons in total around the central P (1 lone pair + 3 bond pairs) and 8 around each Cl. Common slip: candidates draw only the bond pairs and omit the lone pair on P, or omit the lone pairs on Cl entirely.
(b) M1 — phosphorus is a Period 3 element and therefore has access to vacant 3d orbitals; this allows the 3s and 3p electrons to be promoted and hybridised so that 5 unpaired electrons become available for bonding.
A1 — nitrogen is a Period 2 element and has no accessible d sub-shell at the n = 2 energy level. The maximum number of unpaired electrons on N is therefore 3 and the octet around N cannot be expanded beyond 8 electrons.
Common slip: students draw PCl5 with only 4 bond pairs around P, forgetting that the 5 bond pairs constitute an expanded octet of 10 electrons.
(c)(i) M1 — HCN contains H–C and C≡N. The H–C is a single bond (1σ). The C≡N triple bond is 1σ + 2π. Total: 2σ and 2π.
(c)(ii) M1 — ethene H2C=CH2 contains four C–H bonds (4σ) and one C=C double bond (1σ + 1π). Total: 5σ and 1π.
A1 — both totals correct and clearly stated.
Total: 7 marks (M5 A2, split as shown).
Question (6 marks): The bond dissociation enthalpy of N≡N in N2(g) is +945 kJ mol⁻¹. The mean N=N bond enthalpy in diazene (N2H2) is approximately +418 kJ mol⁻¹ and the mean N–N bond enthalpy in hydrazine (N2H4) is approximately +158 kJ mol⁻¹. Explain, in terms of σ and π bonding and orbital overlap, why the N≡N bond enthalpy is more than twice the N=N value but the N=N value is less than twice the N–N value.
Mark scheme decomposition by AO:
| Mark | AO | Awarded for |
|---|---|---|
| M1 | AO1.1 | Identifying the N–N bond in hydrazine as a single σ bond (head-on overlap of sp3-type orbitals on N). |
| M1 | AO1.2 | Identifying N=N as 1σ + 1π and N≡N as 1σ + 2π. |
| M1 | AO2.1 | Recognising that π bonds are formed by sideways overlap of p orbitals and are individually weaker than σ bonds. |
| M1 | AO2.1 | Stating that adding a π bond to an existing σ shortens the bond, which in turn strengthens the σ component (through more effective overlap), so a triple bond is more than three times as strong as a single bond is one-third of it. |
| M1 | AO3.1a | Applying the reasoning quantitatively: 945 > 2 × 418 because the second π in N≡N benefits from the contracted internuclear distance imposed by the first π. |
| M1 | AO3.1b | Concluding that 418 < 2 × 158 because in N2H2 the σ component is itself weakened by lone-pair repulsion on adjacent N atoms (and by the sp2 vs sp3 character difference), so a "double bond is not simply two single bonds" reasoning applies in reverse. |
Total: 6 marks split AO1 = 2, AO2 = 2, AO3 = 2. This is a synoptic 6-marker — Edexcel uses this kind of question to test whether candidates can move beyond labelling σ and π bonds and reason about bond enthalpy as an emergent property of orbital overlap.
Covalent bonding is the most cross-referenced single idea in the 9CH0 specification. Five concrete connections worth memorising:
Topic 4 — Inorganic Chemistry (Period 3 chlorides): the chlorides of Period 3 show a continuous progression in bonding character. NaCl is essentially fully ionic; MgCl2 is largely ionic but with appreciable covalent character (Mg2+ polarises Cl⁻); AlCl3 exists as the covalent dimer Al2Cl6 (with two Cl bridging via dative bonds — exactly the example covered in Foundations); SiCl4, PCl3, PCl5, SCl2 and Cl2 are all simple covalent molecules. Knowing where the ionic-to-covalent transition lies, and being able to justify it via electronegativity differences, is a frequent 3–4 mark Paper 1 question.
Topics 6 and 16 — Organic Chemistry (alkene reactivity): every C–H, C–C and C–X bond in the organic syllabus is covalent, and the σ/π distinction is what underpins the entire reaction chemistry of alkenes. The π bond in C=C is the reactive site — it is more exposed, weaker and more polarisable than the σ framework, which is why alkenes undergo electrophilic addition while alkanes only undergo radical substitution. Without confident σ/π reasoning, alkene mechanism marks are inaccessible.
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