Chemical bonding

• A bond is an attraction, never “a way to get a full shell”; and a shape names the atoms, not the pairs.

• 3.1.1 I can define — define electronegativity as the power of an atom to attract electrons to itself
• 3.1.2 I can explain — explain the factors influencing the electronegativities of the elements in terms of nuclear charge, atomic radius and shielding by inner shells and sub-shells
• 3.1.3 I can state — state and explain the trends in electronegativity across a period and down a group of the Periodic Table
• 3.1.4 I can use — use the differences in Pauling electronegativity values to predict the formation of ionic and covalent bonds (the presence of covalent character in some ionic compounds will not be assessed) (Pauling electronegativity values will be given where necessary)
• 3.2.1 I can define — define ionic bonding as the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions)
• 3.2.2 I can describe — describe ionic bonding, including the examples of sodium chloride, magnesium oxide and calcium fluoride
• 3.3.1 I can define — define metallic bonding as the electrostatic attraction between positive metal ions and delocalised electrons
• 3.4.1 I can define — define covalent bonding as the electrostatic attraction between the nuclei of two atoms and a shared pair of electrons (a) describe covalent bonding in molecules including: hydrogen, H₂; oxygen, O₂; nitrogen, N₂; chlorine, Cl₂; hydrogen chloride, HCl; carbon dioxide, CO₂; ammonia, NH₃; methane, CH₄; ethane, C₂H₆; ethene, C₂H₄ (b) understand that elements in period 3 can expand their octet including in the compounds sulfur dioxide, SO₂, phosphorus pentachloride, PCl₅, and sulfur hexafluoride, SF₆ (c) describe coordinate (dative covalent) bonding, including in the reaction between ammonia and hydrogen chloride gases to form the ammonium ion, NH₄⁺, and in the Al₂Cl₆ molecule
• 3.4.2 I can — (a) describe covalent bonds in terms of orbital overlap giving σ and π bonds: σ bonds are formed by direct overlap of orbitals between the bonding atoms; π bonds are formed by the sideways overlap of adjacent p orbitals above and below the σ bond (b) describe how the σ and π bonds form in molecules including H₂, C₂H₆, C₂H₄, HCN and N₂ (c) use the concept of hybridisation to describe sp, sp² and sp³ orbitals
• 3.4.3 I can — (a) define the terms: bond energy as the energy required to break one mole of a particular covalent bond in the gaseous state; bond length as the internuclear distance of two covalently bonded atoms (b) use bond energy values and the concept of bond length to compare the reactivity of covalent molecules
• 3.5.1 I can state — state and explain the shapes of, and bond angles in, molecules by using VSEPR theory, including as simple examples: BF₃ (trigonal planar, 120°); CO₂ (linear, 180°); CH₄ (tetrahedral, 109.5°); NH₃ (pyramidal, 107°); H₂O (non-linear, 104.5°); SF₆ (octahedral, 90°); PF₅ (trigonal bipyramidal, 120° and 90°)
• 3.5.2 I can predict — predict the shapes of, and bond angles in, molecules and ions analogous to those specified in 3.5.1
• 3.6.1 I can — (a) describe hydrogen bonding, limited to molecules containing N–H and O–H groups, including ammonia and water as simple examples (b) use the concept of hydrogen bonding to explain the anomalous properties of H₂O (ice and water): its relatively high melting and boiling points; its relatively high surface tension; the density of solid ice compared with liquid water
• 3.6.2 I can use — use the concept of electronegativity to explain bond polarity and dipole moments of molecules
• 3.6.3 I can — (a) describe van der Waals' forces as the intermolecular forces between molecular entities other than those due to bond formation, and use the term van der Waals' forces as a generic term to describe all intermolecular forces (b) describe the types of van der Waals' forces: instantaneous dipole–induced dipole (id–id) forces, also called London dispersion forces; permanent dipole–permanent dipole (pd–pd) forces, including hydrogen bonding (c) describe hydrogen bonding and understand that hydrogen bonding is a special case of permanent dipole–permanent dipole forces between molecules where hydrogen is bonded to a highly electronegative atom
• 3.6.4 I can state — state that, in general, ionic, covalent and metallic bonding are stronger than intermolecular forces
• 3.7.1 I can use — use dot-and-cross diagrams to illustrate ionic, covalent and coordinate bonding including the representation of any compounds stated in 3.4 and 3.5 (dot-and-cross diagrams may include species with atoms which have an expanded octet or species with an odd number of electrons)

Why a lattice, not molecules. Nothing in an ionic solid pairs one Na⁺ with one Cl⁻. The formula NaCl gives the ratio of ions, and each ion is attracted equally to its six neighbours. How the structure explains conductivity, brittleness and solubility is topic 4.

• “Atoms bond to get a full shell.” Correct A bond is an electrostatic attraction. Ionic, metallic and covalent bonding are each defined as one; none of the three definitions mentions shells. Write the definition, not the motive.
• “NH₃ is tetrahedral, because it has four pairs.” Correct The four pairs are arranged tetrahedrally, but the shape names the atoms: NH₃ is pyramidal, 107°; H₂O is non-linear, 104.5°.
• “Ionic bonding is the transfer of electrons.” Correct The transfer is how the ions form. The bond is the electrostatic attraction between the oppositely charged ions that result.
• “CO₂ has polar bonds, so it is a polar molecule.” Correct The molecule is linear and symmetrical, so its two bond dipoles cancel: no dipole moment. Polar bonds are necessary for a dipole moment but not sufficient; the shape decides.
• “Boiling water breaks the O–H bonds.” Correct Boiling a molecular substance overcomes the intermolecular forces. The covalent O–H bonds inside each molecule are untouched; steam is still H₂O.
• “The double bond in ethene is two identical bonds.” Correct It is one σ bond plus one π bond. The π bond is weaker (610 − 350 = 260 kJ mol⁻¹ from Data-section values) and has two lobes, above and below the σ bond, which together are one bond.
• “Van der Waals’ forces are the weak forces between non-polar molecules.” Correct In this syllabus van der Waals’ forces is the generic term for all intermolecular forces: id–id and pd–pd, with hydrogen bonding a special case of pd–pd.
• “Electronegativity is how much an atom wants electrons.” Repair The power of an atom to attract electrons to itself. Quote Data-section values when comparing two atoms.
• “Electronegativity decreases across a period because the atoms get smaller.” Repair It increases across a period. The nuclear charge rises while the shielding stays almost the same, so the bonding pair is attracted more strongly; the smaller radius is a consequence of the same change.
• “Ionic bonding is the transfer of electrons from a metal to a non-metal.” Repair The transfer forms the ions. The bond is the electrostatic attraction between oppositely charged ions (positively charged cations and negatively charged anions).
• “Metallic bonding is the attraction between positive ions and negative ions.” Repair There are no anions in a metal. Metallic bonding is the electrostatic attraction between positive metal ions and delocalised electrons.
• “A dative bond is weaker than a normal covalent bond.” Repair Once formed it is identical to the others: all four N–H bonds in NH₄⁺ are the same. “Dative” describes where the electrons came from, not how strong the bond is.
• “Sulfur cannot have twelve electrons around it.” Repair Period 3 atoms can expand the octet, because their outer shell has a 3d sub-shell; SF₆ has six bonding pairs. Period 2 atoms cannot.
• “A π bond is formed by end-on overlap of p orbitals.” Repair End-on overlap gives a σ bond. A π bond is the sideways overlap of adjacent p orbitals, above and below the σ bond.
• “Ethene has two π bonds, one above and one below the molecule.” Repair Ethene has one π bond, with two lobes, one above and one below the plane of the molecule.
• “A double bond is twice as strong as a single bond.” Repair C=C is 610 kJ mol⁻¹ and C–C 350 (Data section): the π component adds only 260.
• “NH₃ is tetrahedral, with a bond angle of 109.5°.” Repair The pairs are arranged tetrahedrally; the shape names the atoms: pyramidal, 107°, because the lone pair repels more strongly than the bonding pairs.
• “CO₂ is polar because C=O bonds are polar.” Repair The bonds are polar, but the linear molecule is symmetrical and the two bond dipoles cancel: no dipole moment.
• “Hydrogen bonds are the bonds inside a water molecule.” Repair Hydrogen bonds are intermolecular: between the δ+ H of one molecule and a lone pair on the O of another. The O–H bonds inside the molecule are covalent.
• “Any molecule containing hydrogen can form hydrogen bonds.” Repair The hydrogen must be bonded to N or O (in this syllabus), and the other molecule must have a lone pair on N or O. CH₄ and HCl do not form hydrogen bonds.
• “When water boils, the O–H bonds break.” Repair Boiling overcomes the hydrogen bonds and other intermolecular forces between molecules. The covalent bonds stay; steam is H₂O.
• “Van der Waals’ forces are only the weak forces between non-polar molecules.” Repair In this syllabus van der Waals’ forces is the generic term for all intermolecular forces: id–id is one type, pd–pd (including hydrogen bonding) the other.

• Read the command word before you decide how much to write. This syllabus has twenty-two of them: analyse, calculate, compare, consider, contrast, deduce, define, demonstrate, describe, determine, discuss, evaluate, examine, explain, give, identify, justify, predict, show (that), sketch, state and suggest. Comment, estimate, name and outline are not among them: where a question wants something named it says identify, which the syllabus glosses as “name/select/recognise”. State and give want a fact and nothing more. Describe wants the points or the features. Explain wants the reasons and the relationships — a describe-level answer to an explain question is incomplete however well written it is. Deduce and determine want a conclusion reached from the information given, with the reasoning visible.
• Interleave with the chapters that use this one. Topic 4 builds giant and simple lattices from these bonds: re-answer “why does MgO melt higher than NaCl?”. Topic 5 turns bond energies into enthalpy changes: re-answer “why is C=C not twice C–C?”. Topic 9 uses metallic bonding across Period 3; topic 11 the H–X bond energies; topics 13 and 14 hybridisation, shape and the π bond in every organic molecule; topic 28 the coordinate bond in complexes. Recalling a topic inside a new context is worth more than another pass over this chapter on its own; at A Level, Paper 4 assumes the whole of the AS content, so nothing here is ever finished with.

• Cambridge International AS & A Level Chemistry 9701 has five components. Topic 3 is AS Level content, so it is examined in Papers 1, 2 and 3. AS Level content: examined in Paper 1 (multiple choice), Paper 2 (AS structured) and, as practical context, Paper 3. Assumed knowledge for Papers 4 and 5. AS Level candidates take Papers 1, 2 and 3; A Level candidates take all five, either staged over two years (Papers 1–3 in year one, Papers 4 and 5 in year two) or together in one series. Examinations are available in the June and November series, and in March in India.
• Across both the AS Level and the A Level the assessment objectives are weighted AO1 40% (knowledge and understanding), AO2 40% (handling, applying and evaluating information) and AO3 20% (experimental skills and investigations). AS candidates are graded a–e; A Level candidates A*–E.
• A Paper 1 item on this topic can turn on a single discrimination: which molecule has no dipole moment, which shape and angle go together, how many σ and π bonds a molecule has, which substance forms hydrogen bonds. A Paper 2 structured question asks you to define a bond or electronegativity in the syllabus’s words, draw a dot-and-cross diagram, state and explain a shape and angle, or explain a boiling point or a trend. Explanations earn their marks point by point: the pairs, the repulsion, the lone-pair effect; the force, where it acts, why it is stronger.
• Topic 3 has almost no arithmetic: electronegativity differences and a bond-energy comparison such as 610 − 350 = 260 kJ mol⁻¹. Pauling electronegativities and bond energies are printed in the Data section and supplied with the paper; the syllabus says electronegativity values will be given where necessary. Boiling points and bond lengths are not in the Data section: a question that needs them supplies them. What you must know by heart are the seven shapes and their angles.
• No quantitative procedure belongs to this topic, so it reaches Papers 3 and 5 as data interpretation and observation: plotting and explaining boiling points of a series of compounds, testing a liquid for polarity with a charged rod, and the precision of a thermometer reading. The data and practical focus sets out a Paper 5-style item.
• Read the command word before you decide how much to write. This syllabus has twenty-two of them: analyse, calculate, compare, consider, contrast, deduce, define, demonstrate, describe, determine, discuss, evaluate, examine, explain, give, identify, justify, predict, show (that), sketch, state and suggest. Comment, estimate, name and outline are not among them: where a question wants something named it says identify, which the syllabus glosses as “name/select/recognise”. State and give want a fact and nothing more. Describe wants the points or the features. Explain wants the reasons and the relationships — a describe-level answer to an explain question is incomplete however well written it is. Deduce and determine want a conclusion reached from the information given, with the reasoning visible.