Edexcel International A Level Chemistry: Ionic and covalent bonding (YCH11)

This study guide covers Topic 3 parts 3A (Ionic bonding) and 3B (Covalent bonding) of the Pearson Edexcel International Advanced Subsidiary/Advanced Level in Chemistry (XCH11/YCH11) specification, Issue 1, September 2017. It teaches learning outcomes 3.1 to 3.15. This is International AS (IAS) content, examined in Unit 1: Structure, Bonding and Introduction to Organic Chemistry. Shapes of molecules (3C) and metallic bonding (3D) are not taught here.

Condensed recall is in the ionic and covalent bonding revision notes, and the practice set tests every outcome. Other topics are on the Edexcel IAL Chemistry hub; track outcomes with the printable checklist. A free 10-minute diagnostic shows which outcomes to start with.

What this unit covers

Outcomes	What you must be able to do
3.1	Interpret evidence for ions
3.2–3.3	Describe ion formation; draw dot-and-cross diagrams of ions
3.4–3.6	Describe ionic lattices and what makes ionic bonding strong
3.7	Explain ionic radius trends, including N³⁻ to Al³⁺
3.8–3.9	Explain polarisation, polarising power and polarisability
3.10–3.11	Describe covalent bonding and its evidence; draw dot-and-cross diagrams, including dative bonds in Al₂Cl₆ and NH₄⁺
3.12	Describe diamond, graphite and graphene and their uses
3.13–3.15	Explain electronegativity, the bonding continuum and polarity

3A Ionic bonding

Evidence for ions (3.1)

• Physical properties. Ionic compounds have high melting temperatures, are hard but brittle, and conduct when molten or dissolved but not as solids. That fits charged particles locked in place in the solid and free to move in the liquid.
• Electron density maps. X-ray diffraction of an ionic crystal gives a contour map of electron density. Contours around each ion are nearly circular, and between neighbouring ions the density falls almost to zero: separate, roughly spherical particles of two sizes.
• Migration of ions. Coloured ions move towards the oppositely charged electrode. In the further suggested practical, green copper(II) chromate solution sits in a U-tube under a colourless electrolyte; with a voltage applied, blue (Cu²⁺) rises on the negative side and yellow (CrO₄²⁻) on the positive side. A potassium manganate(VII) crystal on moist filter paper on a microscope slide gives a purple streak drifting towards the positive electrode, because MnO₄⁻ is an anion.

Forming ions and dot-and-cross diagrams (3.2, 3.3)

Metals lose outer electrons to form cations; non-metals gain electrons to form anions. In simple cases both ions end up with the electron arrangement of a noble gas.

Worked example: potassium sulfide, K₂S

1. Potassium is 2.8.8.1 and sulfur is 2.8.6.
2. Each potassium atom loses its single outer electron: K → K⁺ + e⁻. The ion is 2.8.8.
3. One sulfur atom gains two electrons, one from each potassium: S + 2e⁻ → S²⁻. The ion is 2.8.8.
4. Draw each K⁺ with an empty outer shell, in square brackets with + outside.
5. Draw S²⁻ with its own six outer electrons as crosses and the two gained electrons as dots (eight, in four pairs), in square brackets with 2− outside.
6. Show two K⁺ ions so the charges balance: 2(+1) + (−2) = 0.

Dots and crosses only show where electrons came from; in the ion they are identical.

Lattices and the strength of ionic bonding (3.4–3.6)

An ionic crystal is a giant lattice: a regular three-dimensional arrangement of very many cations and anions. Each ion is surrounded by ions of opposite charge. Ionic bonding is the strong net electrostatic attraction between the ions. “Net” matters: like charges repel, but each ion is closer to oppositely charged neighbours, so attraction wins overall, acting in all directions.

Ionic bonding is stronger when the ions have higher charges and smaller radii: larger charges attract more strongly, and smaller ions have their centres closer together. A rough indicator: multiply the two ionic charges, then divide by the sum of the ionic radii.

Worked example: potassium fluoride against calcium oxide

Use these values: K⁺ 138 pm, F⁻ 133 pm, Ca²⁺ 100 pm, O²⁻ 140 pm. Melting temperatures: KF 858 °C, CaO 2613 °C.

1. KF: charges 1 × 1 = 1; sum of radii 138 + 133 = 271 pm; indicator 1 ÷ 271 = 0.00369 pm⁻¹.
2. CaO: charges 2 × 2 = 4; sum of radii 100 + 140 = 240 pm; indicator 4 ÷ 240 = 0.0167 pm⁻¹.
3. The CaO value is about 4.5 times larger: double charges on both ions, and closer ion centres.
4. More energy is needed to separate the ions in CaO, so it melts at a far higher temperature.

The indicator only compares compounds with the same lattice type; it is not a bond energy.

Trends in ionic radius (3.7)

Effective ionic radii (pm) for Group 1 cations are Li⁺ 76, Na⁺ 102, K⁺ 138, Rb⁺ 152 and Cs⁺ 167. For the isoelectronic set:

Isoelectronic ion	N³⁻	O²⁻	F⁻	Na⁺	Mg²⁺	Al³⁺
Protons	7	8	9	11	12	13
Electrons	10	10	10	10	10	10
Radius	146	140	133	102	72	53.5

Different data sources give slightly different radii, so learn the trends, not the numbers.

• Down a group ionic radius increases: each step down adds an occupied shell, placing the outermost electrons further out.
• Across N³⁻ to Al³⁺ every ion has 10 electrons (1s² 2s² 2p⁶) but protons rise from 7 to 13. Same electrons and shielding, larger nuclear charge: the electrons are pulled in more tightly and the radius falls.

Polarisation (3.8, 3.9)

Polarisation is the distortion of an anion’s electron cloud by a nearby cation. Electron density is pulled into the space between the ions, so the bond gains some covalent character.

• Polarising power of a cation is high when it is small and highly charged (charge concentrated in a small volume).
• Polarisability of an anion is high when it is large and highly charged (outer electrons far from the nucleus and easily distorted).

Worked example: ranking polarising power

Use these values: Na⁺ 102 pm, Mg²⁺ 72 pm, Al³⁺ 53.5 pm.

1. Charge ÷ radius: Na⁺ 1 ÷ 102 = 0.0098; Mg²⁺ 2 ÷ 72 = 0.0278; Al³⁺ 3 ÷ 53.5 = 0.0561 (units: charge per pm).
2. Order of polarising power: Al³⁺ > Mg²⁺ > Na⁺.
3. Al³⁺ with a large anion such as I⁻ gives much covalent character; Na⁺ with small F⁻ is close to purely ionic.

3B Covalent bonding

What a covalent bond is (3.10)

A covalent bond is the strong electrostatic attraction between two nuclei and the shared pair of electrons between them. The specification names two pieces of evidence.

1. Giant atomic structures, such as diamond, are very hard with very high melting temperatures. Melting breaks a great many covalent bonds, so the bonds must be strong.
2. Electron density maps of simple molecules show contours enclosing both nuclei, with significant electron density between them, unlike the near-zero density between ions.

Dot-and-cross diagrams for covalent species (3.11)

Count outer electrons, place shared pairs, then lone pairs. Check each atom other than hydrogen has eight outer electrons, unless the species is a known exception.

Worked example: hydrogen cyanide, HCN, and dinitrogen difluoride, F–N=N–F

Species	Outer electrons	Bonding pairs	Lone pairs
HCN	1 + 4 + 5 = 10	H–C single (1), C≡N triple (3)	1 on N
N₂F₂	2 × 5 + 2 × 7 = 24	two N–F (2), N=N double (2)	1 on each N, 3 on each F

In HCN, carbon shares one pair with hydrogen and three with nitrogen (eight outer electrons); nitrogen has three shared pairs and one lone pair (also eight).

Dative covalent (coordinate) bonds. Both electrons of the shared pair come from one atom. Once formed, it is the same as any other covalent bond.

• Ammonium, NH₄⁺. The lone pair on nitrogen in NH₃ bonds to H⁺, which has no electrons. All four N–H bonds are then identical. The ion has 10 electrons (7 + 4 − 1); draw it in square brackets with + outside.
• Al₂Cl₆. Aluminium in AlCl₃ has only six outer electrons. In the vapour at lower temperatures, a chlorine of each AlCl₃ unit donates a lone pair to the aluminium of the other. Each Al then has four bonds and eight outer electrons. Check: 48 outer electrons; 8 Al–Cl bonds use 16, leaving 16 lone pairs (three on each of four terminal Cl, two on each of two bridging Cl). Draw each dative bond as an arrow from Cl to Al.

Giant lattices of carbon (3.12)

Structure	Bonding and arrangement	Key properties	Applications
Diamond	Each C bonded to four others, tetrahedral, 3D lattice	Very hard, very high melting temperature, does not conduct (no free electrons)	Cutting, drilling and grinding tools
Graphite	Each C bonded to three others in flat hexagonal layers; one electron per atom delocalised; weak forces between layers	Conducts along layers; layers slide; very high melting temperature	Electrodes, pencil “leads”, dry lubricant
Graphene	A single layer of graphite, one atom thick	Very strong for its mass, conducts electricity, almost transparent	Being developed for flexible electronics, touchscreens, sensors and strong lightweight composites

All three have high melting temperatures (strong covalent bonds); conduction depends on delocalised electrons.

Electronegativity and the bonding continuum (3.13, 3.14)

Electronegativity measures how strongly an atom attracts the bonding pair of electrons in a covalent bond. Values rise across a period and fall down a group; fluorine has the highest value.

If two bonded atoms differ in electronegativity, the shared pair is pulled towards the more electronegative one, which becomes δ− while the other becomes δ+: the bond is polar. Ionic and covalent bonding are the ends of one continuum.

Equal electronegativity (Cl–Cl) gives pure covalent bonding; a small difference gives polar covalent; a very large difference (K and F) means the pair is transferred almost completely, so the bonding is ionic. There is no sharp cut-off, and polarisation (3.8) describes the same continuum from the ionic end.

Worked example: placing bonds on the continuum

Use these Pauling values: H 2.20, C 2.55, O 3.44, Cl 3.16, Al 1.61, K 0.82, F 3.98.

1. C–H: 2.55 − 2.20 = 0.35. Nearly non-polar.
2. O–H: 3.44 − 2.20 = 1.24. Strongly polar covalent, O is δ−.
3. Al–Cl: 3.16 − 1.61 = 1.55. Highly polar, intermediate character.
4. K–F: 3.98 − 0.82 = 3.16. Ionic.

Order of increasing ionic character: C–H < O–H < Al–Cl < K–F.

Polar bonds and polar molecules (3.15)

A molecule with polar bonds is only polar if the bond dipoles do not cancel. That depends on shape (from 3C).

• CCl₄: four polar C–Cl bonds, tetrahedral and symmetrical, so the dipoles cancel. Non-polar.
• CHCl₃: one C–H replaces a C–Cl, so the dipoles do not cancel. Polar.
• BCl₃ (trigonal planar) is non-polar; NH₃ (pyramidal) is polar.

The further suggested practical tests this directly. A charged rod held next to a thin jet of liquid deflects water and ethanol (polar molecules) but leaves cyclohexane (non-polar) undeflected.

Common errors

• Writing that ionic bonding is “a transfer of electrons”. Transfer forms the ions; the bond is the attraction between them.
• Explaining the N³⁻ to Al³⁺ trend by “more electrons”. All six have 10 electrons; proton number changes.
• Saying polarisability depends only on size. The specification includes charge for both ions.
• Drawing ions without brackets or charges, or pointing an Al₂Cl₆ dative arrow from Al to Cl.
• Calling a molecule polar just because it contains polar bonds.
• Saying graphite conducts because of “free ions”; it has delocalised electrons.

Where next

Try the practice questions. Wider Unit 1 reading: the overview guide to structure, bonding and organic chemistry and the overview bonding notes.

Official syllabus

Pearson Edexcel International Advanced Subsidiary/Advanced Level in Chemistry (XCH11/YCH11), Specification, Issue 1, September 2017, published by Pearson Education Limited. Topic 3: Bonding and Structure, parts 3A Ionic bonding and 3B Covalent bonding, outcomes 3.1 to 3.15.