Covalent bonding
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Physics 9702
Covalent Bonding Covalent Bonding The electrostatic attraction between the nuclei of two atoms and a shared pair of electrons. Covalent bonds form between non-metal atoms. Each atom contributes one electron to a shared pair (bonding pair).
The shared electrons are attracted to both nuclei simultaneously, holding the atoms together. A single bond = 1 shared pair of electrons A double bond = 2 shared pairs of electrons A triple bond = 3 shared pairs of electrons Lone pairs = non-bonding pairs on an atom Dot-and-Cross Diagrams for Covalent Molecules Only outer (valence) shell electrons are shown.
Shared pairs sit between the two atoms; lone pairs sit on one atom only. Molecule Bond type Bonding pairs Lone pairs on central atom H₂ Single H–H 1 0 Cl₂ Single Cl–Cl 1 3 on each Cl HCl Single H–Cl 1 3 on Cl O₂ Double O=O 2 1 on each O (+ 1 unpaired - free radical character) N₂ Triple N≡N 3 1 on each N CO₂ Two double C=O 2 per bond 2 on each O NH₃ Three single N–H 3 1 on N CH₄ Four single C–H 4 0 C₂H₆ (ethane) C–C and four C–H per C 7 total 0 C₂H₄ (ethene) C=C and two C–H per C 6 total 0 Figure 3.6: Dot-and-Cross Diagrams for H₂, Cl₂, HCl, O₂, N₂ (Row of five dot-and-cross diagrams.
H₂: two H atoms sharing one pair. Cl₂: two Cl atoms with 3 lone pairs each, sharing one bonding pair. HCl: H and Cl sharing one pair, 3 lone pairs on Cl. O₂: two O atoms sharing 2 pairs, 2 lone pairs each.
N₂: two N atoms sharing 3 pairs, 1 lone pair each.) Figure 3.7: Dot-and-Cross Diagrams for NH₃, CH₄, CO₂, H₂O (NH₃: N central with 3 N–H bonds and 1 lone pair on N. CH₄: C central with 4 C–H bonds, no lone pairs.
CO₂: O=C=O with 2 lone pairs on each O. H₂O: O central with 2 O–H bonds and 2 lone pairs on O.) Figure 3.8: Dot-and-Cross Diagrams for C₂H₆ and C₂H₄ (C₂H₆: two C atoms each bonded to 3 H and to each other by single C–C bond, no lone pairs.
C₂H₄: two C atoms joined by a double bond (C=C), each also bonded to 2 H atoms, no lone pairs.) Expanded Octet - Period 3 Elements Elements in Period 3 (and beyond) have available 3d orbitals that can be used for bonding, allowing them to accommodate more than 8 electrons in their valence shell - an expanded octet .
Compound Central atom Valence electrons around central atom Octet expanded? SO₂ S 10 (one double bond to O, one dative bond, one lone pair) Yes (10 electrons) PCl₅ P 10 (five P–Cl bonds) Yes (10 electrons) SF₆ S 12 (six S–F bonds) Yes (12 electrons) Figure 3.9: Dot-and-Cross Diagrams - PCl₅ and SF₆ (PCl₅: P central with 5 Cl atoms bonded by single bonds.
Each Cl has 3 lone pairs. P has 10 electrons total - expanded octet. SF₆: S central with 6 F atoms, each F has 3 lone pairs. S has 12 electrons total.) Key Point Elements in Period 2 (C, N, O, F) cannot expand their octet - they have no available d orbitals in the second shell.
Only Period 3+ elements (P, S, Cl, etc.) can do so. Coordinate (Dative Covalent) Bonding Coordinate (Dative Covalent) Bond A covalent bond in which both electrons of the shared pair are donated by the same atom.
The donor atom must have a lone pair; the acceptor atom must have a vacant orbital. Once formed, a coordinate bond is identical to an ordinary covalent bond - the distinction is only in how it was formed.
It is represented by an arrow (→) pointing from the donor to the acceptor. Example 1: Formation of the Ammonium Ion, NH₄⁺ NH₃ has a lone pair on the nitrogen atom. H⁺ (a proton) has a vacant orbital and no electrons.
The lone pair on N is donated to H⁺, forming a new N→H coordinate bond. The resulting NH₄⁺ ion has four identical N–H bonds - all equivalent despite one being dative in origin. Figure 3.10: Formation of NH₄⁺ by Coordinate Bonding (Left: NH₃ with lone pair on N and H⁺ with empty orbital.
Arrow from lone pair on N to H⁺. Right: NH₄⁺ in square brackets with + charge, all four N–H bonds shown equivalently. The coordinate bond may be shown with an arrow N→H.) Example 2: Al₂Cl₆ (Aluminium Chloride Dimer) AlCl₃ has only 6 electrons around Al - an electron-deficient molecule with a vacant orbital.
A lone pair from Cl of one AlCl₃ molecule donates into the vacant orbital of Al in another AlCl₃. Two coordinate bonds form, creating the dimer Al₂Cl₆ (two bridging Cl atoms). Figure 3.11: Structure of Al₂Cl₆ (Two AlCl₃ units joined by two bridging Cl atoms.
Each Al bonded to 2 terminal Cl and 2 bridging Cl. Coordinate bonds shown as arrows from bridging Cl lone pairs to Al. Al now has 8 electrons - octet complete.) Sigma (σ) and Pi (π) Bonds σ Bond (Sigma Bond) Formed by the direct (head-on) overlap of orbitals along the internuclear axis.
Electron density is concentrated between the two nuclei. All single bonds are σ bonds. π Bond (Pi Bond) Formed by the sideways overlap of adjacent p orbitals, above and below the σ bond axis. Electron density is concentrated above and below the internuclear axis.
π bonds only exist alongside a σ bond. Molecule Bond(s) σ bonds π bonds H₂ H–H single bond 1 (s-s overlap) 0 C₂H₆ (ethane) C–C + C–H 7 0 C₂H₄ (ethene)…
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