Chemical bonding
Explain H2 Chemistry bonding, orbital overlap, molecular shapes and intermolecular forces, then relate structures to physical properties.
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Before you begin
Identify the particles and attractions in a substance. Use orbital overlap, shape and polarity to explain its bonding, then connect intermolecular forces and structure to its properties.
Be comfortable with: O-Level bonding and structure, atomic orbitals and electron configuration.
Learning goals
- Ionic and Covalent Bonding Models
- Dative Bonding and Common Examples
- Sigma and Pi Bonds (Orbital Overlap)
- Molecular Shapes and Bond Angles (VSEPR)
- Bond Energy, Bond Length and Bond Polarity
- 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
How the ideas connect
- Identify the attraction
Name the particles that attract in ionic, covalent and metallic bonds. Use electron pairs and orbital overlap for covalent bonds.
Step 1 - Work out the shape
Count electron regions around the central atom. Use the shape to decide whether bond dipoles cancel.
Step 2 - Explain the property
Distinguish bonds from intermolecular forces. Use the actual structure and mobile charged particles to explain physical properties.
Step 3
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 overlapRelate head-on and sideways overlap of s and p orbitals to sigma and pi bonds.
- Molecular shapes and bond angles (VSEPR)Use VSEPR to predict the shapes and bond angles of molecules and ions.
- Bond energy, bond length and bond polarityRelate bond energy, length and electronegativity to bond polarity, molecular polarity and reactivity.
- 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
Or choose
Topic reference
Quick reference
Name the attraction
Ionic: opposite ions. Covalent: a shared electron pair and both nuclei. Metallic: positive ions and delocalised electrons.
Define a bond by the electrostatic attraction that holds it together.
Common pitfall
Transfer or sharing describes how a bond forms. It does not, by itself, define the bond.
Shape before polarity
A multiple bond counts as one electron region in VSEPR.
Determine the shape, then ask whether the bond dipoles cancel. Polar bonds can occur in a non-polar molecule.
Common pitfall
Molecular shape names the atom positions. It can differ from the electron-pair arrangement when lone pairs are present.
Explain boiling and conduction
Boiling a simple molecular substance overcomes intermolecular attractions, rather than breaking its covalent bonds.
Electrical conduction needs mobile charged particles: ions or delocalised electrons.
Common pitfall
State the physical state. An ionic solid has charged ions, but they are fixed in position.
Complete quick-reference table
| 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; lone pairs can reduce bond angles; name the molecular shape |
| 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 |
| Why ice floats | the hydrogen-bonded lattice | an open lattice, so a larger volume for the same mass |
| Unknown solid | combine all the evidence | the structure, then each property from its particles and attractions |
Core knowledge to remember
Definitions and core knowledge
- 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
More mistakes to check
- 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.
- Drawing a dative arrow from the acceptor to the donor.
- Explaining conductivity without naming the mobile charged particles in the stated physical state.