Chemical bonding

Use H1 Chemistry bonding models to explain molecular shapes, polarity, intermolecular forces and the physical properties of solids.

GCE A-Level H1 Chemistry 2027
8 lessons

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Before you begin

Every chemical bond is an electrostatic attraction, and the way electrons are shared or transferred decides how a substance behaves. This topic starts with ionic, covalent, dative and metallic bonds and the σ and π overlap behind covalent bonds. It then predicts molecular shapes, covers bond energy and polarity, explains intermolecular forces and the liquefaction of gases, and finishes by linking the structure of a solid to its physical properties.

Be comfortable with: O-Level bonding and structure, atomic orbitals and electron configuration.

Learning goals

Describe the ionic bond as the electrostatic attraction between oppositely charged ions
Describe the covalent bond as the electrostatic attraction between a shared pair of electrons and the positive nuclei
Describe the metallic bond as the electrostatic attraction between a lattice of positive ions and delocalised electrons
Describe ionic bonding in sodium chloride and magnesium oxide, using dot-and-cross diagrams
Describe covalent bonding in H₂, O₂, N₂, Cl₂, HCl, CO₂, CH₄ and C₂H₄, using dot-and-cross diagrams
Describe dative covalent bonding in the ammonium ion and Al₂Cl₆, using dot-and-cross diagrams
Describe covalent bonding as overlap of s and p orbitals, forming σ and π bonds
Use VSEPR theory to explain the shapes and bond angles of BF₃, CO₂, CH₄, NH₃, H₂O and SF₆
Predict the shapes and bond angles of molecules similar to BF₃, CO₂, CH₄, NH₃, H₂O and SF₆
Explain and deduce bond polarity from electronegativity
Deduce whether a molecule is polar from its bond polarities and shape
Describe intermolecular forces from permanent and induced dipoles, as in CHCl₃, Br₂ and the noble gases
Describe hydrogen bonding, using ammonia and water as examples of molecules with –NH and –OH groups
Outline how intermolecular forces allow gases to liquefy at high pressure or low temperature
Outline how hydrogen bonding affects physical properties, including those of ice and water
Explain bond energy and bond length for covalent bonds
Compare how reactive covalent bonds are using bond energy, bond length and bond polarity
Describe the ionic lattice of sodium chloride and magnesium oxide
Describe the simple molecular lattice of iodine
Describe the giant molecular lattices of graphite and diamond
Describe the hydrogen-bonded lattice of ice
Describe the metallic lattice of copper
Describe, interpret and predict how structure and bonding affect a substance's physical properties
Suggest the structure and bonding in a substance from given information

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 9)
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)
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))
describe the following forces of attraction (electrostatic in nature): — intermolecular forces, based on permanent and induced dipoles, as in liquid and gaseous CHCl3, Br2 and the 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 intermolecular forces to the liquefaction of gases when subjected to high pressure and/or low temperature
outline the importance of hydrogen bonding to the physical properties of substances, including ice and water
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
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

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Topic reference

Quick reference

Question	First move	What to show
Define a bond	name the particles that attract	ionic: oppositely charged ions; covalent: a shared pair and both nuclei; metallic: positive ions and delocalised electrons
Dot-and-cross diagram	count the outer electrons	ions in brackets with charges; shared pairs and every lone pair
Dative bond	find a lone pair and an empty orbital	an arrow from the donor atom to the acceptor atom
σ and π bonds	count the bonded pairs of atoms	one σ per pair; one π per extra bond
Shape and bond angle	count the regions around the central atom	a multiple bond is one region; each lone pair closes the angle
Polar molecule?	bond polarity, then shape	whether the bond dipoles cancel
Boiling-point comparison	list the intermolecular forces in each	hydrogen bonding, electron-cloud size and molecular shape
Liquefaction	effect on kinetic energy and on separation	cooling slows molecules; compression brings them closer
Unknown solid	combine all the evidence	the structure, then each property from its particles and attractions

Core knowledge to remember

Ionic bond: the electrostatic attraction between oppositely charged ions.
Covalent bond: the electrostatic attraction between a shared pair of electrons and the nuclei of the two bonded atoms.
Metallic bond: the electrostatic attraction between a lattice of positive ions and delocalised electrons.
In a dative bond, both electrons of the shared pair come from one atom. Once formed, it is an ordinary covalent bond.
A single bond is 1σ; a double bond is 1σ + 1π; a triple bond is 1σ + 2π.
Electron pairs repel: lone pair–lone pair > lone pair–bond pair > bond pair–bond pair. Shapes: linear 180°, trigonal planar 120°, tetrahedral 109.5°, trigonal pyramidal about 107°, bent about 104.5°, octahedral 90°.
Bond energy is the energy needed to break one mole of a covalent bond in gaseous molecules. Shorter bonds are usually stronger.
Electronegativity is the ability of an atom in a covalent bond to attract the shared pair.
Hydrogen bonding needs an H bonded to N, O or F and a lone pair on N, O or F in another molecule.
A substance conducts only if it has charged particles that can move: mobile ions or delocalised electrons.

Common mistakes

Defining a bond by how it formed (“transfer”, “sharing”) instead of as an electrostatic attraction.
Saying boiling or melting a simple molecular substance breaks covalent bonds.
Counting a multiple bond as more than one region in VSEPR, or naming the electron-pair arrangement instead of the shape.
Saying a molecule is polar because its bonds are polar, without checking whether the dipoles cancel.
Naming hydrogen bonding whenever a molecule contains hydrogen.
Saying pressure makes intermolecular forces stronger. It brings the molecules closer together.
Explaining conductivity without naming the mobile charged particles in the stated physical state.