Bonding and Structure
Bonding and structure covers how atoms are held together by ionic, covalent and metallic bonding, how electron-pair repulsion theory predicts molecular shapes, and how intermolecular forces (London forces, dipole-dipole forces and hydrogen bonding) explain physical properties. It links directly to melting point, boiling point and conductivity questions common in A Level Chemistry exams.
Method
- Identify the type of bonding present (ionic, covalent, dative covalent or metallic) by considering whether the substance is a metal with a non-metal (ionic), two non-metals (covalent) or a metal alone (metallic).
- To predict molecular shape, count the bonding pairs and lone pairs around the central atom, then apply electron-pair repulsion theory: electron pairs repel and arrange as far apart as possible, with lone pair-lone pair repulsion greater than lone pair-bonding pair repulsion, which is greater than bonding pair-bonding pair repulsion.
- Compare electronegativity values across a bond to judge its polarity; a large difference gives a polar bond, but check the molecule's overall shape and symmetry before concluding the whole molecule is polar, as dipoles can cancel.
- Identify the strongest intermolecular force present: hydrogen bonding (N-H, O-H or H-F bonds) is strongest, then permanent dipole-dipole forces (in polar molecules), then London (dispersion) forces (present in all molecules, increasing with more electrons/greater Mr).
- Link structure to property: giant structures (ionic, giant covalent, metallic) have high melting points because many strong bonds/attractions must be broken throughout the lattice, while simple molecular substances have low melting points because only weak intermolecular forces need to be overcome.
- When explaining conductivity, check for mobile charge carriers (delocalised electrons in metals and graphite, or freely moving ions in molten/aqueous ionic compounds); no mobile charge carriers means no conductivity.
Worked example
Predict the shape and bond angle of a molecule of phosphine, PH3, and explain why its bond angle differs from the bond angle in methane, CH4 (109.5 degrees).
- Count the electron pairs around the central phosphorus atom: phosphorus has 5 outer electrons, 3 are used in P-H bonds, leaving 1 lone pair, so there are 3 bonding pairs and 1 lone pair (4 electron pairs in total).
- Four electron pairs arrange themselves in a tetrahedral electron-pair arrangement to minimise repulsion.
- Removing the lone pair to describe the shape of the atoms only gives a trigonal pyramidal molecular shape.
- The lone pair repels the three bonding pairs more strongly than the bonding pairs repel each other, compressing the H-P-H bond angle below the tetrahedral angle.
- In methane there are 4 bonding pairs and 0 lone pairs, so all repulsions are equal (bonding pair-bonding pair), giving the full 109.5 degree angle; the extra lone pair repulsion in PH3 reduces its bond angle to about 93.5 degrees.
- Final answer: trigonal pyramidal, bond angle approximately 93.5 degrees (less than 109.5 degrees because of the extra lone pair repulsion)
Practice questions
Try each question, then tap to reveal the answer.
Q1Define the term ionic bonding.Show answer
Answer: The electrostatic force of attraction between oppositely charged ions
Q2State the shape and bond angle of a molecule with 4 bonding pairs and 0 lone pairs around the central atom.Show answer
Answer: Tetrahedral, 109.5 degrees
Q3State the three conditions required for hydrogen bonding to occur between molecules.Show answer
Answer: A hydrogen atom bonded directly to a highly electronegative N, O or F atom; a lone pair of electrons on an N, O or F atom of a neighbouring molecule; the hydrogen atom and the lone pair align to form the bond
Q4Explain why diamond has a very high melting point.Show answer
Answer: Diamond is a giant covalent (macromolecular) structure in which every carbon atom is joined to four others by strong covalent bonds in all directions; melting requires breaking many of these strong bonds, needing a large amount of energy
Q5Using electronegativity values H = 2.1, Cl = 3.0, Br = 2.8 and F = 4.0, rank the hydrogen halides HBr, HCl and HF in order of increasing bond polarity, showing your working.Show answer
Answer: HBr < HCl < HF (electronegativity differences: HBr = 0.7, HCl = 0.9, HF = 1.9; a larger difference means a more polar bond)
Q6A student claims that 'silicon dioxide, SiO2, conducts electricity because it contains covalent bonds between silicon and oxygen.' Evaluate this claim, referring to structure and bonding.Show answer
Answer: The claim is incorrect; SiO2 is a giant covalent (macromolecular) structure with no free ions or delocalised electrons, so it has no mobile charge carriers and does not conduct electricity in any state
Exam-style questions
Written in the style of a A Level Science exam paper, with a full mark scheme.
Explain, in terms of structure and bonding, why sodium chloride conducts electricity when molten but not when solid.
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A student states: 'Sulfur dioxide, SO2, has polar S=O bonds, so it must be a polar molecule.' Evaluate this statement.
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Explain why the boiling point of ammonia, NH3 (boiling point -33 degrees C), is much higher than the boiling point of phosphine, PH3 (boiling point -88 degrees C), even though phosphine has a greater relative molecular mass (Mr = 34 compared with Mr = 17 for ammonia).
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