3.1
Stability of Atoms & Why They Bond
The Drive Towards Stability
Atoms form bonds because doing so lowers their overall energy, producing a more stable arrangement. The driving force is the tendency to achieve a full outer electron shell — the same electron configuration as the nearest noble gas (Group 18). Noble gases are chemically inert precisely because they already have full outer shells (2 electrons for He, 8 for the rest).
This idea is formalised as the octet rule: atoms tend to gain, lose, or share electrons to achieve 8 electrons in their outer shell (or 2 for hydrogen and lithium, which aim for He configuration).
The octet rule has exceptionsSome atoms (e.g. Be, B, P, S, and transition metals) can have fewer or more than 8 electrons in their outer shell when bonded. The octet rule is a useful guide, not an absolute law.
Types of Chemical Bond
| Bond Type | Formed Between | How | Example |
|---|---|---|---|
| Ionic | Metal + non-metal | Transfer of electrons | NaCl, MgO, CaCl₂ |
| Covalent | Non-metal + non-metal | Sharing of electrons | H₂O, CO₂, CH₄ |
| Metallic | Metal + metal (or pure metal) | Delocalised electron sea | Fe, Cu, Al, brass |
Electronegativity and Bond Type
Electronegativity is the ability of an atom to attract shared electrons towards itself. The difference in electronegativity (ΔEN) between two bonded atoms determines the bond type:
ΔEN > 1.7 → Ionic bond (electron transfer)
ΔEN 0.4–1.7 → Polar covalent bond
ΔEN < 0.4 → Non-polar covalent bond
Electronegativity increases across a period (left → right) and decreases down a group (top → bottom). Fluorine is the most electronegative element (EN = 4.0 on Pauling scale).
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Section 3.1 — Ionic Bonding
10 QuestionsQ1 of 10
Ionic bonding involves:
Ionic bond: metal loses electrons → cation; non-metal gains → anion. Electrostatic attraction between opposite charges holds ions in lattice.
Q2 of 10
Which pair forms an ionic compound?
Na (metal, low IE) + Cl (non-metal, high EA) → Na⁺ and Cl⁻ → ionic bond. Metal-nonmetal combinations generally form ionic compounds.
Q3 of 10
The coordination number of Na⁺ in NaCl structure is:
NaCl: rock salt structure. Each Na⁺ is surrounded by 6 Cl⁻ and each Cl⁻ by 6 Na⁺. Coordination number = 6:6.
Q4 of 10
Why do ionic compounds have high melting points?
Strong electrostatic attraction between oppositely charged ions in the crystal lattice → large energy needed to melt → high melting points. E.g. NaCl m.p. = 801°C.
Q5 of 10
Ionic compounds conduct electricity:
Solid: ions fixed in lattice → cannot carry charge. Molten or dissolved: ions free to move → conducts. This is electrolytic conduction.
Q6 of 10
Lattice energy is defined as:
Lattice energy: M⁺(g) + X⁻(g) → MX(s), ΔH negative (exothermic). Greater lattice energy = stronger ionic lattice. Depends on charge and size of ions.
Q7 of 10
Which ionic compound has the greatest lattice energy?
Lattice energy ∝ q⁺×q⁻/r. MgO: Mg²⁺ and O²⁻ (charges 2+,2−), small ions → much greater electrostatic attraction than NaCl (1+,1−, larger ions). MgO lattice energy ≈ −3791 kJ/mol vs NaCl ≈ −787 kJ/mol.
Q8 of 10
The Born-Haber cycle is used to calculate:
Born-Haber: applies Hess's law to ionic compound formation cycle. ΔHf° = IE + EA + ΔHat + lattice energy (rearranged to find LE since LE can't be measured directly).
Q9 of 10
As ionic radius increases (larger ions), lattice energy:
Lattice energy ∝ 1/r. Larger ions → greater r → weaker electrostatic attraction → less negative (smaller magnitude) lattice energy. Example: LiF has greater lattice energy than CsI.
Q10 of 10
CaO has much higher lattice energy than NaF because:
Lattice energy ∝ q⁺×q⁻. CaO: 2+×2− = 4 times the charge product of NaF (1+×1−). Also similar ionic sizes. Much stronger attraction → larger lattice energy.
3.2
Ionic Bonding
Ionic BondThe electrostatic force of attraction between oppositely charged ions (cation and anion) formed by the complete transfer of one or more electrons from a metal atom to a non-metal atom.
Formation of Ions
Metal atoms (low IE, low EN) lose their outer electrons to form cations (positive ions). They achieve the electron configuration of the preceding noble gas.
Non-metal atoms (high IE, high EN) gain electrons to form anions (negative ions). They achieve the electron configuration of the next noble gas.
WORKED EXAMPLE
Formation of Sodium Chloride (NaCl)
1
Na (Z=11): 1s²2s²2p⁶3s¹ → Na loses its 3s¹ electron → Na⁺: 1s²2s²2p⁶ (isoelectronic with Ne)
2
Cl (Z=17): 1s²2s²2p⁶3s²3p⁵ → Cl gains one electron → Cl⁻: 1s²2s²2p⁶3s²3p⁶ (isoelectronic with Ar)
3
Na⁺ and Cl⁻ are attracted by electrostatic forces → ionic bond formed. The compound NaCl is electrically neutral (charges balance).
More Examples of Ionic Bond Formation
WORKED EXAMPLE
Formation of Magnesium Oxide (MgO)
1
Mg (Z=12): [Ne]3s² → loses 2 electrons → Mg²⁺: [Ne] (2,8)
2
O (Z=8): [He]2s²2p⁴ → gains 2 electrons → O²⁻: [He]2s²2p⁶ = [Ne] (2,8)
3
Mg²⁺ + O²⁻ → MgO. The 2+ and 2− charges balance → formula MgO.
WORKED EXAMPLE
Formation of Calcium Chloride (CaCl₂)
1
Ca (Z=20): [Ar]4s² → loses 2 electrons → Ca²⁺: [Ar] (2,8,8)
2
Cl (Z=17): [Ne]3s²3p⁵ → gains 1 electron → Cl⁻: [Ne]3s²3p⁶ = [Ar]
3
Ca²⁺ needs 2 Cl⁻ to balance charge → formula CaCl₂. Ca transfers one electron to each Cl atom.
Ionic Crystal Structure
Giant Ionic Lattice
In the solid state, ionic compounds form a giant ionic lattice — a regular, repeating 3-dimensional arrangement of cations and anions held together by strong electrostatic forces in all directions. There are no discrete molecules in an ionic solid.
In NaCl, each Na⁺ is surrounded by 6 Cl⁻ ions, and each Cl⁻ is surrounded by 6 Na⁺ ions. This is called a 6:6 coordination.
Lattice Energy
Lattice Energy (ΔH𝑳𝑮𝑻)The energy released when one mole of an ionic compound is formed from its gaseous ions under standard conditions.
Na⁺(g) + Cl⁻(g) → NaCl(s) ΔH = −787 kJ mol⁻¹ (lattice energy of NaCl)
Always negative (exothermic) — energy is released as ions come together.
Na⁺(g) + Cl⁻(g) → NaCl(s) ΔH = −787 kJ mol⁻¹ (lattice energy of NaCl)
Always negative (exothermic) — energy is released as ions come together.
Factors Affecting Lattice Energy
Lattice energy depends on the strength of the electrostatic attraction between ions:
E ∝ (q⁺ × q⁻) / r
q⁺ = charge on cation, q⁻ = charge on anion, r = ionic radius
- Higher ionic charge → stronger attraction → more exothermic lattice energy (larger magnitude).
e.g. MgO (−3791 kJ/mol) >> NaCl (−787 kJ/mol) because Mg²⁺ and O²⁻ have greater charge than Na⁺ and Cl⁻. - Smaller ionic radius → ions closer together → stronger attraction → larger (more negative) lattice energy.
Physical Properties of Ionic Compounds
| Property | Observation | Explanation |
|---|---|---|
| Melting/boiling point | High (e.g. NaCl: 801°C) | Strong electrostatic forces between oppositely charged ions in the giant lattice require large amounts of energy to overcome. |
| Electrical conductivity (solid) | Non-conductor | Ions are fixed in the lattice — they cannot move to carry charge. |
| Electrical conductivity (molten/aq) | Conductor | When melted or dissolved in water, ions are free to move and carry charge. |
| Solubility in water | Many are soluble | Polar water molecules can attract and separate ions from the lattice (hydration). Lattice energy must be overcome by hydration energy. |
| Solubility in organic solvents | Generally insoluble | Non-polar solvents cannot interact strongly enough with ions to break the lattice. |
| Hardness and brittleness | Hard but brittle | Lattice is rigid (hard). When stressed, layers of ions shift — like charges align and repel, causing the crystal to shatter (brittle). |
Writing ionic formulae from chargesUse the "cross-over" rule: the magnitude of the cation's charge becomes the subscript of the anion and vice versa, then simplify. E.g. Al³⁺ + O²⁻ → Al²O³.
3.3
Metallic Bonds & Physical Properties of Metals
Metallic BondThe electrostatic attraction between a lattice of positive metal ions (cations) and a surrounding "sea" of delocalised electrons. The outer electrons of metal atoms are released into a shared pool that flows freely throughout the structure.
The Metallic Bonding Model
When metal atoms come together, their outer (valence) electrons leave the individual atoms and become delocalised — free to move throughout the entire metallic structure. This leaves behind positively charged metal ions arranged in a regular lattice. The delocalised electrons are attracted to all the positive ions simultaneously, holding the structure together.
This can be described as: "positive ions immersed in a sea of delocalised electrons."
Physical Properties of Metals Explained
| Property | Observation | Explanation (metallic bonding) |
|---|---|---|
| High melting/boiling point | Most metals have high m.p. (Fe: 1538°C) | Strong electrostatic attraction between the lattice of cations and the delocalised electron sea requires large energy to disrupt. More valence electrons per atom = stronger bonding (e.g. Al > Na). |
| Electrical conductivity | Good conductors (solid and liquid) | Delocalised electrons are free to move throughout the metal. When a voltage is applied, electrons flow, creating a current. (Unlike ionic compounds, metals conduct in the solid state.) |
| Thermal conductivity | Good heat conductors | Mobile delocalised electrons transfer kinetic energy rapidly throughout the metal. Vibrations also propagate through the lattice. |
| Malleability & ductility | Can be hammered into sheets (malleable) or drawn into wires (ductile) | Metal ion layers can slide past each other without breaking the bond — the delocalised electron sea adjusts around the new positions. No localised bonds to break. |
| Metallic lustre | Shiny appearance | Delocalised electrons absorb and re-emit light across a broad range of frequencies, giving the characteristic shiny surface. |
| High density | Generally dense (Fe: 7.87 g cm⁻³) | Metal atoms are closely packed in a regular lattice with little empty space. |
| Strength | High tensile strength | Strong, non-directional metallic bonds hold the lattice together in all directions; difficult to pull apart. |
Factors Affecting Strength of Metallic Bonding
Metallic bond strength increases with:
- More delocalised electrons per atom — Group 2 metals bond more strongly than Group 1 metals (e.g. Mg has higher m.p. than Na). Transition metals are often very strong due to d-electron involvement.
- Smaller ionic radius — smaller ions allow electrons to be held more tightly, increasing bond strength and m.p.
- Higher charge on metal ion — more positive cations attract the electron sea more strongly.
Comparing metallic bond strengthGoing across Period 3 (Na → Mg → Al), the number of delocalised electrons increases (1, 2, 3), the ionic charge increases, and the ionic radius decreases. All three factors strengthen the metallic bond, so melting point increases: Na (98°C) < Mg (650°C) < Al (660°C). (Note: Al and Mg are close because the higher charge of Al³⁺ is partly offset by its slightly larger effective radius in context.)
Comparison of Bond Types
| Property | Ionic | Metallic | Covalent (simple) |
|---|---|---|---|
| Melting point | High | Variable (often high) | Low (weak IMF) |
| Electrical conductivity | Only when molten/dissolved | Always (solid & liquid) | None (usually) |
| Solubility in water | Often soluble | Insoluble | Variable |
| Mechanical | Hard, brittle | Malleable, ductile | Soft (molecular) |
| Bonding particle | Ions (cation + anion) | Cation + delocalised e⁻ | Molecules |
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Section 3.2 — Metallic Bonding
10 QuestionsQ1 of 10
Metallic bonding is described as:
Metallic bond: cations (positive ions) in regular lattice; valence electrons delocalised — free to move throughout the metal. Attraction between cations and electron sea holds metal together.
Q2 of 10
Why are metals good conductors of electricity?
Delocalised electrons move freely → carry electric current. This is electronic conduction — different from electrolytic conduction in ionic solutions.
Q3 of 10
Why are metals malleable and ductile?
Metal layers slide → cations rearranged but electron sea still holds structure → doesn't break. Non-directional bonding allows deformation. Contrast: ionic crystals crack (ion repulsion when layers shift).
Q4 of 10
As you go across Period 3 (Na to Al), metallic bond strength:
Na: 1 valence e; Mg: 2; Al: 3. More electrons in sea + higher nuclear charge → stronger metallic bond → higher melting points: Na (98°C) < Mg (650°C) < Al (660°C).
Q5 of 10
Which best explains why metals have high melting points?
Metallic bonds are strong and non-directional. Large energy needed to free cations from electron sea. Giant metallic structure → high melting point. Tungsten (W) has m.p. 3422°C — 6 valence electrons, very strong metallic bond.
Q6 of 10
Why does mercury (Hg) have a very low melting point for a metal?
Hg: [Xe]4f¹⁴5d¹⁰6s². Fully filled d (5d¹⁰) is compact and doesn't participate in metallic bonding. Only 2 valence 6s electrons contribute. Weak metallic bond → low m.p. (−39°C).
Q7 of 10
Why are metals lustrous (shiny)?
Free electrons in metals absorb photons and immediately re-emit them → metals reflect light across all visible wavelengths → appear shiny. Same electrons explain electrical and thermal conductivity.
Q8 of 10
Alloys are harder than pure metals because:
Alloy: mixture of metals. Different-sized atoms disrupt regular lattice → harder for layers to slide → stronger, harder material. Example: steel (Fe + C) is harder than pure iron.
Q9 of 10
Why do metals conduct heat well?
Thermal conduction in metals: energetic electrons at the hot end move quickly through the metal, colliding with ions and transferring energy to cooler regions. Much faster than phonon conduction in non-metals.
Q10 of 10
Which property would ionic NaCl and metallic Na share?
Both have regular/ordered structures (lattice). NaCl = ionic lattice; Na = metallic lattice with electron sea. Differences: NaCl doesn't conduct when solid, is brittle; Na conducts always, is malleable.
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Exercises
- Using dot-and-cross diagrams, show the formation of: (a) KBr, (b) MgCl₂, (c) Al₂O₃. In each case state the charge on each ion.
(a) K ([Ar]4s¹) loses 1e → K⁺ [Ar]. Br ([Ar]3d¹⁰4s²4p⁵) gains 1e → Br⁻ [Kr]. Formula KBr (1:1).
(b) Mg ([Ne]3s²) loses 2e → Mg²⁺ [Ne]. Cl ([Ne]3s²3p⁵) gains 1e → Cl⁻ [Ar]. 1 Mg²⁺ needs 2 Cl⁻ → MgCl₂.
(c) Al ([Ne]3s²3p¹) loses 3e → Al³⁺ [Ne]. O ([He]2s²2p⁴) gains 2e → O²⁻ [Ne]. Cross-over: 2Al³⁺ + 3O²⁻ → Al₂O³. - State and explain four physical properties of ionic compounds in terms of their giant ionic lattice structure.
(1) High m.p.: strong electrostatic forces between oppositely charged ions throughout the lattice require large energy to overcome. (2) Non-conductor (solid): ions are fixed in place — cannot move to carry charge. (3) Conductor (molten/dissolved): ions become free to move. (4) Brittle: when stress shifts layers, like charges align and repel strongly → crystal shatters. Also: many soluble in water because polar water molecules hydrate ions and overcome lattice energy.
- Compare the lattice energies of NaCl (−787 kJ/mol) and MgO (−3791 kJ/mol). Explain the large difference in terms of ionic charge and radius.
MgO has a much larger (more negative) lattice energy. (1) Charge: MgO contains Mg²⁺ (+2) and O²⁻ (−2), giving a charge product of 4. NaCl has Na⁺ (+1) and Cl⁻ (−1), charge product = 1. Higher charge → much stronger electrostatic attraction. (2) Radius: Mg²⁺ (72 pm) is smaller than Na⁺ (102 pm), and O²⁻ (140 pm) is smaller than Cl⁻ (181 pm). Smaller ions are closer together → stronger attraction → more exothermic lattice energy.
- Describe the structure and bonding in a metal and use this to explain why metals are (a) good electrical conductors and (b) malleable.
Structure: regular lattice of positive metal cations surrounded by a sea of delocalised (mobile) electrons. The metallic bond is the electrostatic attraction between cations and this electron sea. (a) Conductivity: delocalised electrons are free to move throughout the metal. When a potential difference is applied, electrons move in one direction, creating an electric current. (b) Malleability: layers of metal cations can slide past each other when a force is applied. The non-directional, delocalised electron sea adjusts to maintain bonding around the new positions — no specific bonds are broken, so the metal bends/deforms rather than shattering.
- Explain why the melting point increases across Na → Mg → Al. Use metallic bonding theory in your answer.
Na has 1, Mg has 2, and Al has 3 delocalised electrons per atom. More delocalised electrons → stronger, denser electron sea → stronger attraction to cations. Additionally, the cationic charge increases (Na⁺, Mg²⁺, Al³⁺) and ionic radii decrease across the period. All these factors increase the strength of metallic bonding → more energy required to disrupt → higher melting points (Na: 98°C, Mg: 650°C, Al: 660°C).
- Give the formula of the ionic compound formed between: (a) Ca and N; (b) Al and S; (c) Fe³⁺ and O²⁻.
(a) Ca²⁺ + N³⁻: cross-over gives Ca³N₂.
(b) Al³⁺ + S²⁻: cross-over gives Al₂S³.
(c) Fe³⁺ + O²⁻: cross-over gives Fe₂O³.
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Multiple Choice Quiz — 25 Questions
Unit 3 Quiz
Select one answer per questionQ1
Ionic bonds form by:
Ionic bonds form by electron transfer from metal to non-metal, producing oppositely charged ions that attract each other.
Q2
Which compound contains the strongest ionic bond?
MgO: Mg²⁺ and O²⁻ both have charge 2 → much stronger electrostatic attraction. NaF/NaCl/KBr are all ions.
Q3
What is the correct formula of the ionic compound formed between Al³⁺ and O²⁻?
Cross-over rule: charge on O²⁻ (2) gives subscript 2 for Al; charge on Al³⁺ (3) gives subscript 3 for O → Al₂O³.
Q4
Why does solid NaCl not conduct electricity?
In solid NaCl, ions are held rigidly in the giant ionic lattice — they cannot move to carry charge. (D is also partially correct but B is the most precise reason.)
Q5
Metallic bonding is best described as:
Metallic bond = electrostatic attraction between lattice of positive metal cations and surrounding delocalised (mobile) electron sea.
Q6
Why are metals malleable?
In metals, cation layers can slide relative to each other. The non-directional delocalised electron sea adjusts and maintains attraction, so the metal deforms rather than shattering.
Q7
Why are ionic compounds brittle?
Ionic crystals: when an external force shifts a layer, originally alternating +/− ions now have +/+ and −/− neighbours → strong repulsion → crystal cleaves/shatters.
Q8
Which species has the electron configuration of argon (Ar)?
Ar = 18 electrons. Cl⁻ (Z=17, +1e) = 18e ✓. Mg²⁺ (Z=12, −2e) = 10e (neon config). Na⁺ (Z=11, −1e) = 10e. Only Cl⁻ = 18e = Ar config.
Q9
Lattice energy is always:
Lattice energy is always exothermic (ΔH negative): energy is released when gaseous ions come together to form the solid lattice due to electrostatic attraction.
Q10
Why does Al have a higher melting point than Na?
Al contributes 3 delocalised electrons per atom (vs 1 for Na) and has Al³⁺ (vs Na⁺). Greater charge density and electron density → stronger metallic bond → higher m.p.
Q11
Which correctly explains why metals conduct electricity?
The mobile, delocalised electrons in the metallic structure can flow freely when a voltage is applied, carrying charge and creating a current.
Q12
The octet rule states that atoms tend to:
Octet rule: atoms gain, lose, or share electrons to achieve 8 electrons in the outer shell (noble gas configuration). Exceptions exist (H aims for 2).
Q13
Which property of ionic compounds is best explained by strong electrostatic forces in the lattice?
High m.p.: large amounts of energy needed to overcome strong electrostatic forces between oppositely charged ions in the giant lattice.
Q14
When Na forms Na⁺, it becomes isoelectronic with:
Na (Z=11) loses 1e → Na⁺ has 10 electrons = same as Ne (Z=10). Na⁺ is isoelectronic with Ne.
Q15
The coordination number of Na⁺ in the NaCl lattice is:
In NaCl, each Na⁺ is surrounded by 6 Cl⁻ ions (top, bottom, left, right, front, back). Coordination number = 6. This is 6:6 coordination.
Q16
Which statement about metallic bonding is correct?
Metallic bonds are non-directional: the delocalised electron sea acts uniformly in all directions around each cation. This non-directionality is why metals are malleable.
Q17
MgO has a higher lattice energy than MgS because:
Both O²⁻ and S²⁻ have the same −2 charge, so charge is not the difference. O²⁻ (ionic radius 140 pm) is much smaller than S²⁻ (184 pm) → Mg²⁺ and O²⁻ are closer → stronger attraction → larger (more −ve) lattice energy.
Q18
Which correctly describes how ionic compounds dissolve in water?
Water is polar; its δ⁻ oxygen is attracted to cations and δ⁺ hydrogens to anions. Hydration energy released compensates for energy needed to break the lattice (lattice energy).
Q19
The formula of the ionic compound formed between Fe³⁺ and SO₄²⁻ is:
Fe³⁺ (charge 3+) and SO₄²⁻ (charge 2−). Cross-over: 2 Fe³⁺ and 3 SO₄²⁻ → Fe₂(SO₄)³. Total charge: 2(+3) + 3(−2) = 0 ✓.
Q20
Which property do metals and ionic compounds share?
Both metals and ionic compounds generally have high melting points due to strong bonding (metallic or electrostatic). Metals conduct in solid state (ionic do not); metals are malleable (ionic are brittle).
Q21
Which correctly describes the electron configuration of O²⁻?
O (Z=8) has 8 electrons: 1s²2s²2p⁴. O²⁻ gains 2 electrons: 8+2=10 electrons → 1s²2s²2p⁶ (isoelectronic with Ne).
Q22
The shiny lustre of metals is due to:
The sea of delocalised electrons can absorb photons of many frequencies and re-emit them, giving metals their characteristic lustre (shininess) and reflectivity.
Q23
Which element has the highest electronegativity?
Fluorine has the highest electronegativity of all elements (4.0 on the Pauling scale). Electronegativity increases across periods and up groups → F is top-right of the periodic table.
Q24
An element X forms an ion X³⁺. It is most likely:
X³⁺ loses 3 electrons → the element has 3 valence electrons → Group 13 (e.g. Al³⁺). Group 1 loses 1, Group 2 loses 2, Group 17 gains 1.
Q25
Molten NaCl conducts electricity because:
When NaCl melts, the giant lattice breaks down. Na⁺ and Cl⁻ ions become free to migrate when a voltage is applied → ions carry charge → electrical conduction.
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Unit Test — 50 Marks
Section A — Short Answer
20 marksQ1 [4 marks]
Describe the formation of the ionic bond in magnesium chloride (MgCl₂). Include: the electron transfer that occurs [2], the electron configurations of the ions formed [1], and why the formula is MgCl₂ and not MgCl [1].
Mg ([Ne]3s²) loses 2 electrons → Mg²⁺ [Ne]; each of two Cl ([Ne]3s²3p⁵) atoms gains 1 electron → Cl⁻ [Ne]3s²3p⁶ = [Ar]. [2] Mg²⁺ has configuration 1s²2s²2p⁶ (isoelectronic with Ne); Cl⁻ has 1s²2s²2p⁶3s²3p⁶ (isoelectronic with Ar). [1] The formula is MgCl₂ because Mg loses 2 electrons and each Cl gains only 1, so 2 Cl atoms are needed to accept both electrons and balance the charges (2+ from Mg and 2✕1− from two Cl). [1]
Q2 [4 marks]
State and explain four physical properties of a giant ionic compound (e.g. NaCl). [4]
(1) High m.p.: strong electrostatic forces between oppositely charged ions throughout the giant lattice require large energy to overcome. [1] (2) Non-conductor (solid): ions are fixed rigidly in the lattice; they cannot move to carry electrical charge. [1] (3) Conductor when molten or dissolved: lattice breaks down; ions become mobile and can migrate under voltage. [1] (4) Brittle: when a shearing force displaces a layer, like charges align → strong repulsion → crystal shatters. [1] (Also acceptable: solubility in water — polar water molecules hydrate ions, overcoming lattice energy.)
Q3 [4 marks]
Compare the lattice energies of NaF and MgO. Explain in detail why they differ so much, using the relationship E ∝ (q⁺ × q⁻)/r. [4]
NaF lattice energy ≈ −918 kJ/mol; MgO ≈ −3791 kJ/mol — MgO has a much larger (more −ve) lattice energy. [1] Charge effect [2]: In NaF, ions are Na⁺ (+1) and F⁻ (−1) → charge product = 1. In MgO, Mg²⁺ (+2) and O²⁻ (−2) → charge product = 4. The electrostatic force (E ∝ q⁺×q⁻/r) is four times greater for MgO just from charge alone. [2] Radius effect [1]: Mg²⁺ (72 pm) is smaller than Na⁺ (102 pm) and O²⁻ (140 pm) is smaller than F⁻ (133 pm in context of the compound). Smaller interionic distance r → stronger attraction → more exothermic lattice energy. [1]
Q4 [4 marks]
Describe the metallic bonding model. Explain how it accounts for (a) electrical conductivity [1] and (b) malleability [1] of metals, and explain why aluminium has a higher melting point than sodium [2].
Metallic bonding: metal atoms lose their outer (valence) electrons into a delocalised “sea”, leaving behind a regular lattice of positive cations. The metallic bond is the electrostatic attraction between the cation lattice and this mobile electron sea. [1 implicit] (a) Electrical conductivity: delocalised electrons are free to move throughout the metal under an applied voltage, carrying charge and creating a current. [1] (b) Malleability: cation layers can slide past each other under stress; the non-directional electron sea adjusts and maintains bonding in the new configuration, so the metal deforms plastically rather than fracturing. [1] Al vs Na m.p. [2]: Al contributes 3 delocalised electrons per atom; Na contributes only 1. More electrons in the sea + higher ionic charge (Al³⁺ vs Na⁺) + smaller ionic radius of Al³⁺ vs Na⁺ → stronger electrostatic attraction between cations and electron sea in Al → more energy required to disrupt → higher melting point (Al 660°C vs Na 98°C).
Q5 [4 marks]
Define electronegativity and explain how it relates to bond type. Using electronegativity, predict the bond type in: (a) HF, (b) Cl₂, (c) CaO. [4]
Electronegativity: the ability of an atom to attract the shared pair of electrons in a covalent bond towards itself. [1] Bond type depends on ΔEN: ΔEN > 1.7 → ionic; ΔEN 0.4–1.7 → polar covalent; ΔEN < 0.4 → non-polar covalent. [1] (a) HF: EN(H)=2.2, EN(F)=4.0, ΔEN=1.8 → ionic/strongly polar covalent (borderline, often called polar covalent). [0.67] (b) Cl₂: both Cl atoms identical, ΔEN=0 → non-polar covalent. [0.67] (c) CaO: EN(Ca)=1.0, EN(O)=3.5, ΔEN=2.5 → ionic. [0.67]
Section B — Extended Answer
30 marksQ6 [8 marks]
Describe the giant ionic lattice structure of NaCl. Draw a diagram of the lattice. Explain how each of the following properties arises from the structure: (a) high melting point; (b) non-conductor in the solid state but conductor when molten; (c) brittleness; (d) solubility in water. [8]
Structure [2]: NaCl has a giant ionic lattice: Na⁺ and Cl⁻ ions are arranged in a regular, repeating 3D pattern. Each Na⁺ is surrounded by 6 Cl⁻ (6:6 coordination) and vice versa, held together by strong electrostatic forces (ionic bonds) acting in all directions. There are no discrete NaCl molecules. Diagram: shows alternating Na⁺ and Cl⁻ in a grid/3D arrangement with lines indicating bonds. [1 for diagram]
(a) High m.p. [1]: strong electrostatic forces throughout the entire lattice must be overcome. Large amounts of energy are required → high melting point (801°C for NaCl).
(b) Non-conductor (solid)/conductor (molten) [2]: In solid NaCl, all ions are fixed in lattice positions — they cannot move to carry electrical charge. When melted, the lattice breaks down and ions are free to migrate towards oppositely charged electrodes under an applied voltage → electrical current flows.
(c) Brittleness [1]: Under a mechanical stress, one layer of ions shifts. Ions that were alternating +/− now align +/+ and −/− in the shifted layers → strong repulsion between like charges → the crystal cleaves/shatters.
(d) Solubility in water [2]: Water is a polar molecule (δ⁻ oxygen, δ⁺ hydrogen). The δ⁻ oxygen is attracted to Na⁺ cations; the δ⁺ hydrogens are attracted to Cl⁻ anions. These interactions (hydration energy) provide enough energy to pull ions away from the lattice surface, overcoming the lattice energy. Ions become surrounded by water molecules (hydrated) and disperse throughout the solution.
(a) High m.p. [1]: strong electrostatic forces throughout the entire lattice must be overcome. Large amounts of energy are required → high melting point (801°C for NaCl).
(b) Non-conductor (solid)/conductor (molten) [2]: In solid NaCl, all ions are fixed in lattice positions — they cannot move to carry electrical charge. When melted, the lattice breaks down and ions are free to migrate towards oppositely charged electrodes under an applied voltage → electrical current flows.
(c) Brittleness [1]: Under a mechanical stress, one layer of ions shifts. Ions that were alternating +/− now align +/+ and −/− in the shifted layers → strong repulsion between like charges → the crystal cleaves/shatters.
(d) Solubility in water [2]: Water is a polar molecule (δ⁻ oxygen, δ⁺ hydrogen). The δ⁻ oxygen is attracted to Na⁺ cations; the δ⁺ hydrogens are attracted to Cl⁻ anions. These interactions (hydration energy) provide enough energy to pull ions away from the lattice surface, overcoming the lattice energy. Ions become surrounded by water molecules (hydrated) and disperse throughout the solution.
Q7 [8 marks]
Explain the metallic bonding model in detail. Describe the physical properties of metals (high m.p., conductivity, malleability, lustre, density) in terms of this model. Compare metals with ionic compounds in terms of: melting point, electrical conductivity, and mechanical properties. [8]
Metallic bonding model [2]: Metal atoms release their valence electrons into a “sea” of delocalised electrons that extends throughout the entire metallic structure. Left behind are positive metal cations arranged in a close-packed regular lattice. The metallic bond is the strong, non-directional electrostatic attraction between these cations and the mobile electron sea. The bond acts in all directions equally.
Physical properties [4]: (1) High m.p.: strong electrostatic forces between cations and dense electron sea require large energy to disrupt. More valence electrons and higher ionic charge → stronger bond → higher m.p. (2) Electrical conductivity: delocalised electrons flow freely when a potential difference is applied, carrying charge. Metals conduct in both solid and liquid states (unlike ionic compounds which only conduct when molten/dissolved). (3) Malleability/ductility: cation layers slide past each other under stress; the non-directional electron sea continuously maintains bonding around displaced ions → metal deforms without fracturing. (4) Metallic lustre: delocalised electrons absorb and re-emit photons across a broad range of frequencies → shiny surface. (5) High density: cations are closely packed in the lattice with little wasted space.
Comparison with ionic [2]: Melting point: both can be high (ionic due to lattice energy; metallic due to electron sea attraction), but ionic compounds vary widely; transition metals often have very high m.p. Conductivity: metals always conduct (solid state); ionic only when molten/dissolved (fixed ions in solid). Mechanical: metals are malleable/ductile (non-directional bonds); ionic compounds are hard but brittle (layer shift causes like-charge repulsion and fracture).
Physical properties [4]: (1) High m.p.: strong electrostatic forces between cations and dense electron sea require large energy to disrupt. More valence electrons and higher ionic charge → stronger bond → higher m.p. (2) Electrical conductivity: delocalised electrons flow freely when a potential difference is applied, carrying charge. Metals conduct in both solid and liquid states (unlike ionic compounds which only conduct when molten/dissolved). (3) Malleability/ductility: cation layers slide past each other under stress; the non-directional electron sea continuously maintains bonding around displaced ions → metal deforms without fracturing. (4) Metallic lustre: delocalised electrons absorb and re-emit photons across a broad range of frequencies → shiny surface. (5) High density: cations are closely packed in the lattice with little wasted space.
Comparison with ionic [2]: Melting point: both can be high (ionic due to lattice energy; metallic due to electron sea attraction), but ionic compounds vary widely; transition metals often have very high m.p. Conductivity: metals always conduct (solid state); ionic only when molten/dissolved (fixed ions in solid). Mechanical: metals are malleable/ductile (non-directional bonds); ionic compounds are hard but brittle (layer shift causes like-charge repulsion and fracture).
Q8 [6 marks]
Using the concept of lattice energy and ionic charge/radius, explain why: (a) MgO has a higher m.p. than NaCl [3]; (b) LiF has a higher m.p. than LiI [2]; (c) CaO has a higher m.p. than BaO [1]. [6]
(a) MgO vs NaCl [3]: MgO contains Mg²⁺ (+2) and O²⁻ (−2) while NaCl has Na⁺ (+1) and Cl⁻ (−1). The charge product for MgO (4) is four times that of NaCl (1). Additionally, Mg²⁺ is smaller than Na⁺ and O²⁻ is smaller than Cl⁻ → ions are closer together. Both effects give MgO a much more exothermic lattice energy (−3791 vs −787 kJ/mol) → much more energy needed to melt → higher m.p. (MgO 2852°C vs NaCl 801°C). [3]
(b) LiF vs LiI [2]: Both contain Li⁺ (+1). F⁻ (ionic radius 133 pm) is much smaller than I⁻ (220 pm). Smaller anion → shorter Li⁺–anion distance → stronger electrostatic attraction → more exothermic lattice energy → higher m.p. for LiF. [2]
(c) CaO vs BaO [1]: Both contain O²⁻ (−2) and +2 cations. Ca²⁺ (ionic radius 100 pm) is smaller than Ba²⁺ (135 pm). Smaller interionic distance in CaO → stronger attraction → larger lattice energy → higher m.p. [1]
(b) LiF vs LiI [2]: Both contain Li⁺ (+1). F⁻ (ionic radius 133 pm) is much smaller than I⁻ (220 pm). Smaller anion → shorter Li⁺–anion distance → stronger electrostatic attraction → more exothermic lattice energy → higher m.p. for LiF. [2]
(c) CaO vs BaO [1]: Both contain O²⁻ (−2) and +2 cations. Ca²⁺ (ionic radius 100 pm) is smaller than Ba²⁺ (135 pm). Smaller interionic distance in CaO → stronger attraction → larger lattice energy → higher m.p. [1]
Q9 [8 marks]
(a) Define ionic bonding and explain why atoms tend to form ions. [2] (b) Describe what is meant by a “giant ionic lattice” with reference to NaCl. [2] (c) Using the Born-Haber cycle concept (without drawing it in full), explain how lattice energy can be determined indirectly and why it cannot be measured directly. [2] (d) Compare the melting points of NaF, NaCl, NaBr and NaI. Explain the trend. [2]
(a) [2]: Ionic bond = electrostatic force of attraction between oppositely charged ions formed by electron transfer from metal to non-metal. Atoms form ions to achieve a full outer electron shell (noble gas configuration), which is a lower-energy, more stable arrangement than an incomplete shell. The energy released when the ionic lattice forms (lattice energy) provides the thermodynamic driving force. [2]
(b) [2]: In NaCl, Na⁺ and Cl⁻ ions are arranged in an infinite, regular, three-dimensional repeating pattern. Each Na⁺ is surrounded by 6 Cl⁻ neighbours and each Cl⁻ by 6 Na⁺ (6:6 coordination). The lattice extends in all three dimensions and is held together by strong, omnidirectional electrostatic forces. There are no discrete NaCl molecules — the formula unit represents the simplest ratio. [2]
(c) [2]: Lattice energy cannot be measured directly because it is impossible to bring together one mole of gaseous ions and measure the energy experimentally — producing pure gaseous ions at standard conditions is not feasible. Instead, the Born-Haber cycle uses Hess’s law to calculate lattice energy indirectly: the formation enthalpy of the ionic compound is related to enthalpy of atomisation, ionisation energies, electron affinities, and lattice energy via a thermodynamic cycle. Since all other quantities can be measured, lattice energy is calculated as the “missing” step. [2]
(d) [2]: NaF > NaCl > NaBr > NaI (m.p. decreases down the series). All contain Na⁺. The halide ion increases in size: F⁻ < Cl⁻ < Br⁻ < I⁻. As ionic radius increases, the interionic distance r increases → electrostatic attraction E ∝ 1/r decreases → lattice energy becomes less negative → less energy needed to melt → lower m.p. [2]
(b) [2]: In NaCl, Na⁺ and Cl⁻ ions are arranged in an infinite, regular, three-dimensional repeating pattern. Each Na⁺ is surrounded by 6 Cl⁻ neighbours and each Cl⁻ by 6 Na⁺ (6:6 coordination). The lattice extends in all three dimensions and is held together by strong, omnidirectional electrostatic forces. There are no discrete NaCl molecules — the formula unit represents the simplest ratio. [2]
(c) [2]: Lattice energy cannot be measured directly because it is impossible to bring together one mole of gaseous ions and measure the energy experimentally — producing pure gaseous ions at standard conditions is not feasible. Instead, the Born-Haber cycle uses Hess’s law to calculate lattice energy indirectly: the formation enthalpy of the ionic compound is related to enthalpy of atomisation, ionisation energies, electron affinities, and lattice energy via a thermodynamic cycle. Since all other quantities can be measured, lattice energy is calculated as the “missing” step. [2]
(d) [2]: NaF > NaCl > NaBr > NaI (m.p. decreases down the series). All contain Na⁺. The halide ion increases in size: F⁻ < Cl⁻ < Br⁻ < I⁻. As ionic radius increases, the interionic distance r increases → electrostatic attraction E ∝ 1/r decreases → lattice energy becomes less negative → less energy needed to melt → lower m.p. [2]