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

These notes condense 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: outcomes 3.1 to 3.15. This is International AS (IAS) content for Unit 1: Structure, Bonding and Introduction to Organic Chemistry.

Full explanations and worked examples are in the ionic and covalent bonding study guide; the practice set has fully worked answers. Other topics sit on the Edexcel IAL Chemistry hub and in the outcome checklist. Spend 10 minutes on a free diagnostic to see which outcomes need work.

Key definitions

Term	Definition to learn
Ionic bonding (3.5)	Strong net electrostatic attraction between oppositely charged ions
Giant ionic lattice (3.4)	Regular 3D arrangement of very many cations and anions, each ion surrounded by ions of opposite charge
Covalent bond (3.10)	Strong electrostatic attraction between two nuclei and the shared pair of electrons between them
Dative covalent bond (3.11)	Covalent bond where one atom supplies both electrons of the shared pair
Polarisation (3.8)	Distortion of an anion’s electron cloud by a neighbouring cation
Electronegativity (3.13)	Ability of an atom to attract the bonding pair of electrons in a covalent bond
Polar bond (3.14)	Covalent bond with unequal sharing, giving δ+ and δ− ends
Polar molecule (3.15)	Molecule whose bond dipoles do not cancel, so it has an overall dipole

3A Ionic bonding

Evidence for ions (3.1)

• Properties: high melting temperature; brittle; conducts only when molten or in solution, because the ions must be free to move.
• Electron density maps: near-circular contours around each ion; density falls almost to zero between ions.
• Migration: coloured ions drift to the oppositely charged electrode. Copper(II) chromate in a U-tube: blue Cu²⁺ to the negative electrode, yellow CrO₄²⁻ to the positive. Potassium manganate(VII) on a slide: purple MnO₄⁻ to the positive.

Ion formation and dot-and-cross diagrams (3.2, 3.3)

Method in steps

1. Write both electron arrangements.
2. Decide how many electrons each atom loses or gains.
3. Balance the charges to find the formula.
4. Put every ion in square brackets, write its charge at the top right, and use the other symbol for transferred electrons.

Worked reminder: calcium phosphide, Ca₃P₂. Ca is 2.8.8.2 and P is 2.8.5. Each Ca loses two electrons (Ca²⁺ is 2.8.8). Each P gains three (P³⁻ is 2.8.8). Six electrons move in total: three Ca²⁺ (+6) balance two P³⁻ (−6). Each phosphide ion shows five crosses and three dots in its outer shell.

Strength of ionic bonding (3.6)

Stronger with higher ionic charge and smaller ionic radius.

Worked reminder: NaF against NaI. Same charges. Sum of radii: Na⁺ 102 + F⁻ 133 = 235 pm; Na⁺ 102 + I⁻ 220 = 322 pm. The ions in NaF are closer, so attraction is stronger and NaF has the higher melting temperature.

Ionic radius trends (3.7)

Trend	Direction	Reason
Down a group (Li⁺ to Cs⁺; F⁻ to I⁻)	Increases	One more occupied shell each time
N³⁻ → O²⁻ → F⁻ → Na⁺ → Mg²⁺ → Al³⁺	Decreases	All have 10 electrons; protons rise 7 → 13, so the electrons are pulled in more

Halide radii (pm) as an example: F⁻ 133, Cl⁻ 181, Br⁻ 196, I⁻ 220.

Polarisation (3.8, 3.9)

	Cation: polarising power	Anion: polarisability
Increases with	Smaller radius, higher charge	Larger radius, higher charge
Strong example	Al³⁺, Mg²⁺	I⁻, S²⁻
Weak example	Cs⁺, K⁺	F⁻

More polarisation means more covalent character in an ionic bond.

Worked reminder: aluminium halides. The cation is the same (Al³⁺, high polarising power). The anion grows from F⁻ (133 pm) to Cl⁻ (181 pm) to I⁻ (220 pm), so polarisability rises. Covalent character therefore increases in the order AlF₃ < AlCl₃ < AlI₃.

3B Covalent bonding

Evidence (3.10)

• Giant atomic structures (diamond) are very hard with very high melting temperatures: many strong covalent bonds must break.
• Electron density maps of simple molecules show density concentrated between the two nuclei, with contours enclosing both.

Dot-and-cross diagrams (3.11)

Method in steps

1. Add up the outer electrons (adjust for charge: add one for each negative charge, remove one for each positive charge).
2. Join atoms with shared pairs; add double or triple bonds until each atom other than H has eight.
3. Place the remaining electrons as lone pairs.
4. Check the total matches step 1.

Species	Outer electrons	Bonding	Lone pairs
N₂	10	One triple bond (3 pairs)	1 on each N
CO₂	16	Two C=O double bonds (4 pairs)	2 on each O
H₃O⁺	8	Three O–H, one of them dative from O to H⁺	1 on O
NH₄⁺	8	Four N–H, one formed dative from N to H⁺	none
Al₂Cl₆	48	8 Al–Cl, two of them dative from bridging Cl to Al	16 in total

After a dative bond forms, all bonds of that type in the ion are identical (all four N–H in NH₄⁺; all three O–H in H₃O⁺).

Giant lattices of carbon (3.12)

	Diamond	Graphite	Graphene
Bonds per C	4	3	3
Arrangement	Tetrahedral 3D lattice	Hexagonal layers, weak forces between layers	One layer, one atom thick
Conducts?	No	Yes, along the layers	Yes
Uses	Cutting and drilling tools	Electrodes, lubricant, pencils	Being developed for electronics, sensors, composites

Electronegativity, continuum and polarity (3.13–3.15)

• Electronegativity rises across a period and falls down a group. Fluorine is highest.
• Zero difference: pure covalent. Moderate difference: polar covalent. Large difference: ionic. It is a continuum with no sharp boundary.
• The δ− end of a polar bond is the more electronegative atom.

Worked reminder: three points on the continuum. Use Pauling values C 2.55, S 2.58, Li 0.98, I 2.66, Cs 0.79, F 3.98. C–S: difference 0.03, essentially non-polar covalent. Li–I: 1.68, strongly polar, with marked covalent character (small Li⁺, large I⁻). Cs–F: 3.19, ionic.

Method in steps: is the molecule polar?

1. Are any bonds polar? If none, the molecule is non-polar.
2. Work out the shape (3C).
3. If the shape is symmetrical and all outer atoms are the same, the dipoles cancel: non-polar.
4. Otherwise there is a net dipole: polar.

Examples: SiCl₄ and BF₃ non-polar; PCl₃, CH₃Cl and H₂S polar.

Must-know distinctions

• Polarising power belongs to cations; polarisability belongs to anions.
• Polar bond is about two atoms; polar molecule is about the whole shape.
• Dative bond describes how a bond formed; it does not make the bond different afterwards.
• Graphite conducts along its layers; diamond does not conduct.
• Ionic electron density map: separate ions. Covalent map: density shared between nuclei.

Quick self-test

1. State what is meant by ionic bonding.
2. Which electrode does the purple colour move towards when potassium manganate(VII) is used in the migration experiment?
3. Give the electron arrangement of the ions in calcium chloride.
4. Which ion is larger, Mg²⁺ or Na⁺? Explain.
5. Use radii Li⁺ 76 pm and K⁺ 138 pm to say which has the greater polarising power.
6. Why does solid potassium bromide not conduct electricity?
7. How many lone pairs are there in one molecule of CO₂?
8. Describe the dative bonds in Al₂Cl₆.
9. Why does graphite conduct electricity?
10. Use Pauling values H 2.20, Br 2.96 and Cl 3.16 to say which is more polar, H–Br or H–Cl.
11. Explain why tetrachloromethane is non-polar.

Answers

1. Strong net electrostatic attraction between oppositely charged ions.
2. The positive electrode, because the coloured ion, MnO₄⁻, is negative.
3. Ca²⁺ 2.8.8 and Cl⁻ 2.8.8.
4. Na⁺. Both have 10 electrons, but Mg²⁺ has 12 protons against 11, so its electrons are pulled in more tightly.
5. Li⁺: same charge, smaller radius, so its charge is more concentrated.
6. Its ions are held in fixed positions in the solid lattice, so no charged particles can flow.
7. Four (two on each oxygen).
8. A chlorine atom on each AlCl₃ unit donates a lone pair to the aluminium of the other unit, giving two dative bonds, each from Cl to Al.
9. Each carbon uses three electrons in bonds; the fourth is delocalised and free to move along the layers.
10. H–Cl. Difference 0.96 against 0.76 for H–Br.
11. The C–Cl bonds are polar, but the molecule is tetrahedral and symmetrical, so the bond dipoles cancel.

Where marks are usually lost

• Describing ionic bonding as electron transfer instead of electrostatic attraction between ions.
• Missing “net” or “strong” from the ionic bond definition, or describing a bond between one pair of ions only.
• Explaining the isoelectronic trend with “more shells” or “more electrons”. The ions have identical electron arrangements; only proton number changes.
• Writing “polarisability of the cation” or “polarising power of the anion”.
• Leaving out charge when explaining polarisability; the specification links it to radius and charge.
• Drawing NH₄⁺ or H₃O⁺ without square brackets and a charge.
• Drawing the Al₂Cl₆ dative arrows from Al to Cl. The chlorine has the lone pair.
• Stating that graphite conducts because of “free ions” or that diamond conducts.
• Saying a molecule is polar because it has polar bonds, without considering shape.
• Placing δ− on the less electronegative atom.

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.