Chemical Bonding
Bonding models, molecular shapes, intermolecular forces and solid structures.
Before you begin
Bonding models, molecular shapes, intermolecular forces and solid structures.
Learning goals
- Ionic and Covalent Bonding Models
- Dative Bonding and Common Examples
- Sigma and Pi Bonds (Orbital Overlap)
- Bond Energy, Bond Length and Bond Polarity
- Molecular Shapes and Bond Angles (VSEPR)
- Intermolecular Forces and Properties
- Solids: Structure and Physical Properties
Syllabus statements covered
- show understanding that all chemical bonds are electrostatic in nature and describe: — ionic bond as the electrostatic attraction between oppositely charged ions
- show understanding that all chemical bonds are electrostatic in nature and describe: — covalent bond as the electrostatic attraction between a shared pair of electrons and positively charged nuclei
- show understanding that all chemical bonds are electrostatic in nature and describe: — metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons
- describe, including the use of ‘dot-and-cross’ diagrams, — ionic bonding as in sodium chloride and magnesium oxide
- describe, including the use of ‘dot-and-cross’ diagrams, — covalent bonding as in hydrogen; oxygen; nitrogen; chlorine; hydrogen chloride; carbon dioxide; methane; ethene
- describe, including the use of ‘dot-and-cross’ diagrams, — co-ordinate (dative covalent) bonding, as in formation of the ammonium ion and in the Al2Cl6 molecule
- describe covalent bonding in terms of orbital overlap (limited to s and p orbitals only), giving σ and π bonds (see also Section 11.1)
- explain and deduce bond polarity using the concept of electronegativity [quantitative treatment of electronegativity is not required]
- deduce the polarity of a molecule using bond polarity and its molecular shape (analogous to those specified in (d))
- explain the terms bond energy and bond length for covalent bonds
- compare the reactivities of covalent bonds in terms of bond energy, bond length and bond polarity
- explain the shapes of, and bond angles in, molecules such as BF3 (trigonal planar); CO2 (linear); CH4 (tetrahedral); NH3 (trigonal pyramidal); H2O (bent); SF6 (octahedral) by using the Valence Shell Electron Pair Repulsion theory
- predict the shapes of, and bond angles in, molecules analogous to those specified in (d)
- describe the following forces of attraction (electrostatic in nature): — intermolecular forces, based on permanent and induced dipoles, as in CHCl3(l); Br2(l) and the liquid noble gases
- describe the following forces of attraction (electrostatic in nature): — hydrogen bonding, using ammonia and water as examples of molecules containing –NH and –OH groups
- outline the importance of hydrogen bonding to the physical properties of substances, including ice and water
- describe, in simple terms, the lattice structure of a crystalline solid which is: — ionic, as in sodium chloride and magnesium oxide
- describe, in simple terms, the lattice structure of a crystalline solid which is: — simple molecular, as in iodine
- describe, in simple terms, the lattice structure of a crystalline solid which is: — giant molecular, as in graphite and diamond
- describe, in simple terms, the lattice structure of a crystalline solid which is: — hydrogen-bonded, as in ice
- describe, in simple terms, the lattice structure of a crystalline solid which is: — metallic, as in copper [the concept of the ‘unit cell’ is not required]
- describe, interpret and/or predict the effect of different types of structure and bonding on the physical properties of substances
- suggest the type of structure and bonding present in a substance from given information
Lessons
Work through them in order.
- Ionic and Covalent Bonding ModelsDefine ionic, covalent and metallic bonding, and draw dot-and-cross diagrams for the required examples.
- Dative Bonding and Common ExamplesExplain dative bonding in the ammonium ion and the aluminium chloride dimer.
- Sigma and Pi Bonds (Orbital Overlap)Relate head-on and sideways overlap of s and p orbitals to sigma and pi bonds.
- Bond Energy, Bond Length and Bond PolarityRelate bond energy, length and electronegativity to bond polarity, molecular polarity and reactivity.
- Molecular Shapes and Bond Angles (VSEPR)Use VSEPR to predict the shapes and bond angles of molecules and ions.
- Intermolecular Forces and PropertiesExplain physical properties from dipole forces and hydrogen bonding, including those of water and ice.
- Solids: Structure and Physical PropertiesDeduce a solid's lattice type from physical data and explain its properties.
Practise and check
Recommended nextChemical Bonding topic check
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Topic reference
Chemical bonding connects electron behaviour to molecular shape, intermolecular forces and bulk properties. Each explanation follows one chain: identify the particles and structure, name the attraction, then explain the property.
Be comfortable with:
- Atomic Structure (A Level): orbitals and electron configuration language.
- Ionic Bonding and Covalent Bonding (O Level).
- Metallic Bonding (O Level).
Quick Reference
| Item | What to write in exams |
|---|---|
| ionic bond | electrostatic attraction between oppositely charged ions |
| covalent bond | electrostatic attraction between a shared pair of electrons and nuclei |
| dative bond | shared pair where both electrons come from the same atom (donor → acceptor) |
| single/double/triple | 1σ; 1σ + 1π; 1σ + 2π |
| hydrogen bond | attraction between δ⁺ H in N-H / O-H / F-H and a lone pair on N/O/F |
| VSEPR workflow | count electron regions → state shape → state angle → explain LP effect |
| Structure → property template | state structure → state forces/charge carriers → “therefore” property |
| Conductivity wording | ionic: mobile ions (molten/aq only); metallic/graphite: delocalised electrons; simple molecular: no charge carriers |
| Solubility wording | compare energy to break IMFs vs hydration/solvation energy (and mention polarity / H-bonding where relevant) |
What You Must Memorise
- Ionic bonding: electrostatic attraction between oppositely charged ions.
- Covalent bonding: electrostatic attraction between a shared pair and nuclei.
- Metallic bonding: attraction between positive ions and delocalised electrons.
- Dative covalent bond: both electrons in shared pair from one atom.
- Hydrogen bonding: attraction between δ⁺ H in N-H / O-H / F-H and a lone pair on N/O/F.
- σ vs π: σ is end-on overlap; π is side-on p–p overlap.
Common Exam Traps
- Confusing intermolecular forces with covalent bonds (boiling breaks IMFs, not covalent bonds).
- Conductivity explanations that don’t name charge carriers (mobile ions vs delocalised electrons).
- Treating “polar bond” as the same as “polar molecule” (shape matters).
- Using “stronger covalent bonds” to explain boiling points of simple molecules (it’s about IMFs).
- For dative bonds, drawing the arrow from acceptor → donor (wrong direction).
- Definitions that miss “electrostatic attraction” (key phrase for bonds).
- London forces explanations that don’t mention instantaneous dipole–induced dipole interactions.
- Treating a dative bond as a different type after formation: its electron pair has a different origin, but the resulting bond is covalent.
- Naming a structure from melting point alone instead of combining melting point, conductivity and solubility evidence.