UNIT 2 CHEMISTRY · PEARSON AAQ · B1.3
Covalent Bonding and Structures
FREE CHEMISTRY REVISION · TOPIC 12
Shared electrons, orbital overlap and extended networks.
Connect electron-pair diagrams to sigma and pi bonding, coordinate bonds and simple versus giant covalent structures.
Pearson BTEC Level 3 AAQ Applied Science
Unit 2: Principles and Applications of Chemistry · B1.3
What you will learn
Describe covalent attraction; draw shared and lone electron pairs; distinguish single, double, triple and coordinate bonds; interpret sigma (σ) and pi (π) overlap; relate bond order to length and strength for comparable bonds; distinguish discrete molecules from giant covalent structures.
1. A shared pair attracted to two nuclei
A covalent bond is the electrostatic attraction between two positive nuclei and a shared pair of electrons between them. The negative shared electrons are attracted to both nuclei. A bond forms at a separation where attractions and repulsions balance, giving a lower energy than the separated atoms.
A bonding pair is shared between two atoms. A lone pair belongs to one atom in a Lewis model and is not shared in that bond. Dots and crosses track electron origin; all electrons have the same charge. Only outer-shell electrons are normally shown.
Hydrogen normally achieves two electrons in its first shell. Many simple main-group atoms form an octet, but this is a useful pattern rather than a universal rule: electron-deficient species and expanded-valence models also occur. Do not describe bond formation as atoms consciously “wanting” a full shell.
2. Single, double and triple bonds
A single bond contains one shared pair (2 electrons), a double bond two pairs (4 electrons), and a triple bond three pairs (6 electrons). Each shared pair contributes to the electron count around both bonded atoms. The total electrons in the diagram must still equal the sum of the atoms’ valence electrons, adjusted for any ionic charge.
In HCl, hydrogen contributes one electron and chlorine one to the shared pair; Cl retains three lone pairs. In O₂ each oxygen supplies two electrons to the double bond and retains two lone pairs. In N₂ each nitrogen supplies three to the triple bond and retains one lone pair. The Lewis description is useful for bonding counts; a fuller molecular-orbital treatment explains properties such as O₂’s paramagnetism.
3. Coordinate or dative covalent bonds
In a coordinate bond, both electrons in the shared pair are supplied by the same donor atom. The donor has a lone pair; the acceptor has an empty orbital available for the pair. Draw a formation arrow from the donor atom towards the acceptor. The arrow is not the direction of electron flow in an ordinary current.
In NH₃ + H⁺ → NH₄⁺, nitrogen donates its lone pair to H⁺. No extra pair appears from nowhere. NH₃ has 5 + 3 = 8 valence electrons, and NH₄⁺ has 5 + 4 − 1 = 8. After formation all four N–H bonds are equivalent. A coordinate origin does not make one bond permanently weaker or chemically different.
Another donor–acceptor example is NH₃ → BF₃: the nitrogen lone pair is accepted into an available boron orbital. Neutral BF₃ is electron-deficient in its simple Lewis structure. Ammonium is an ion, so show brackets and the overall positive charge.
4. Sigma and pi bonding
A sigma (σ) bond forms by end-on orbital overlap along the line joining the nuclei. Bonding density lies along the internuclear axis and has cylindrical symmetry about it. s–s, s–p and suitable hybrid-orbital overlaps can all form σ bonds.
A pi (π) bond forms by sideways overlap of parallel p orbitals. Bonding density lies on opposite sides of the nuclear axis; a nodal plane containing that axis separates the lobes. The two lobes together form one π bond. Orbital phase signs do not mean that one lobe is positively charged.
| Bond | σ bonds | π bonds | Shared electrons |
|---|---|---|---|
| Single | 1 | 0 | 2 |
| Double | 1 | 1 | 4 |
| Triple | 1 | 2 | 6 |
The two π components of a triple bond occupy perpendicular planes. A second bond between the same atoms cannot simply be another independent σ bond along the same axis. Rotation about an isolated σ bond can retain overlap; rotation about a double bond disrupts the parallel alignment needed for π overlap, so free rotation is restricted.
5. Bond length and bond strength
Bond length is the equilibrium distance between the bonded nuclei. Bond dissociation enthalpy measures the energy needed to break a specified gaseous bond per mole, usually forming gaseous radicals; larger values indicate stronger bonds. Bond breaking requires energy and bond formation releases energy.
For comparable bonds between the same elements, increased bond order usually means a shorter, stronger bond: C–C is generally longer and weaker than C=C, which is longer and weaker than C≡C. Compare like with like. Across unrelated atom pairs, atomic size, polarity and molecular environment also matter; “every triple bond is stronger than every double bond” is too broad.
A double bond is stronger overall than the corresponding single bond, but its energy is not exactly twice as large. The added π overlap is generally less effective than σ overlap. Do not confuse intramolecular bond strength with the melting point of a molecular substance.
6. Simple molecules and giant covalent networks
Simple molecular substances contain discrete molecules, such as H₂O, CH₄, CO₂ or I₂. Strong covalent bonds hold each molecule together; intermolecular attractions act between the molecules. A molecule can contain many atoms and still be “simple molecular” in the structural classification.
Giant covalent structures contain extended networks of covalent bonds. Diamond has four covalent bonds per carbon in a 3D tetrahedral network. Graphite has three σ bonds per carbon in planar hexagonal layers and delocalised electrons along the layers. Weaker attractions act between its layers. Graphene is a single such layer.
Silica, SiO₂, is another giant network: each Si is linked to four O atoms and each O bridges two Si atoms. Counting these shared connections gives the 1:2 bulk ratio. The formula is not a discrete bent SiO₂ molecule in crystalline silica. Graphite and diamond are allotropes of carbon with different structures.
Worked examples
Example 1 · Check an O₂ diagram
1. Each O contributes six outer electrons, total 12.
2. Put four electrons in two shared pairs between the O atoms.
3. Place the remaining eight as two lone pairs on each O.
4. Each O counts eight electrons around itself; the diagram contains only 12 electrons in total.
Example 2 · Count σ and π bonds in ethene
H₂C=CH₂ has four C–H single bonds and one C=C double bond. Each C–H contributes one σ; C=C contributes one σ and one π. Total = 5 σ and 1 π. Do not count the entire double bond as two π bonds.
Example 3 · Count σ and π bonds in ethyne
HC≡CH has two C–H single bonds and one C≡C triple bond. Total σ = 2 + 1 = 3; total π = 2. The two π orbitals are perpendicular. The triple bond contains six shared electrons.
Example 4 · Explain NH₄⁺ formation
NH₃ has a nitrogen lone pair. H⁺ accepts that pair, giving an N → H coordinate bond during formation. The valence count remains eight electrons. Draw the overall + charge and explain that the product has four equivalent N–H bonds.
Example 5 · Compare carbon bonds
Compare C–C, C=C and C≡C rather than unrelated atom pairs. Increasing bond order generally shortens the equilibrium separation and increases total bond strength. The additional bonds are π components, so the strength does not scale as exactly 1:2:3.
Example 6 · Recognise a giant structure
A solid has every Si joined to four O atoms and every O to two Si atoms. The connections continue through the crystal, so it is giant covalent. Four Si–O connections per Si are shared between two Si per O, giving two O per Si: SiO₂.
Common misconceptions and exam tips
“A double bond is two sigma bonds.” It is one σ and one π.
“Pi lobes are separate bonds.” Both lobes belong to one π orbital.
“A dative bond is permanently weaker.” The formation history does not distinguish the four N–H bonds in NH₄⁺.
“Dots and crosses have different charges.” They represent identical electrons.
“Low melting point means weak covalent bonds.” Molecular melting usually overcomes intermolecular attractions.
“All covalent substances are molecules.” Diamond, graphite and silica are extended networks.
In dot-and-cross work, count all outer electrons first, then shared pairs and lone pairs, and check charge. For orbital drawings label the nuclear axis, overlap direction and any nodal plane. For a structure explanation name the actual particles and distinguish bonds within a molecule from forces between molecules.
Your revision summary
Covalent bonds arise from shared electrons attracted to both nuclei. Single, double and triple bonds contain 1σ, 1σ + 1π and 1σ + 2π respectively. A coordinate bond uses a pair donated by one atom. Higher bond order generally gives shorter, stronger comparable bonds. Simple molecules are discrete units; giant covalent networks extend through a structure.
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