Atomic number = protons = identifies the element. Mass number = protons + neutrons. Neutral atom: protons = electrons. Electron config fills 2, 8, 8. Group number = valence electrons. Metals lose electrons (more reactive DOWN groups). Non-metals gain electrons (more reactive UP groups). Isotopes = same protons, different neutrons, same CHEMICAL properties.
1.1 Inside the Atom
Particle
Charge
Mass (amu)
Location
Proton (p⁺)
+1
1
Nucleus
Neutron (n⁰)
0
1
Nucleus
Electron (e⁻)
−1
≈ 0 (1/1836)
Electron shells
Key Numbers
Atomic number (Z) = number of protons = identifies the element
Mass number (A) = protons + neutrons
Neutrons = mass number − atomic number
In a neutral atom: protons = electrons
TRAP: Atomic number = protons ONLY, not protons + neutrons. Students constantly confuse atomic number with mass number. Atomic number is the SMALL number on the periodic table.
Isotopes
Isotopes are atoms of the same element (same number of protons) with different numbers of neutrons. This gives them different mass numbers but identical chemical properties.
Example: Carbon-12 (6p, 6n) and Carbon-14 (6p, 8n) are both carbon, but C-14 is radioactive.
TRAP: "Isotopes have different chemical properties" = WRONG. They have the same chemical properties because they have the same electron configuration. Only physical properties (mass, radioactivity) differ.
1.2 Electron Configuration (Shell Model)
Electrons fill shells from the inside out. The maximum electrons per shell:
Shell
Max Electrons
Rule
1st (K)
2
2n² where n=1
2nd (L)
8
2n² where n=2
3rd (M)
8 (for Year 10)
Up to 18, but fills 8 first
Worked Example: Electron Configuration of Sodium (Na, Z=11)
11 protons → 11 electrons
Shell 1: 2 electrons
Shell 2: 8 electrons
Shell 3: 1 electron
Configuration: 2, 8, 1
Sodium has 1 valence electron → it's in Group 1 → it loses 1 electron to form Na⁺
Electron shell diagram for Sodium — the single outer electron makes it highly reactive
What is the electron configuration of chlorine (Z=17)?
2, 8, 7 — 17 electrons: fill shell 1 (2), shell 2 (8), shell 3 (7). Chlorine has 7 valence electrons → Group 17 → gains 1 electron to form Cl⁻.
An atom has 12 protons, 12 neutrons, and 12 electrons. What element is it? What's its mass number?
Magnesium (Mg). Atomic number = 12 protons. Mass number = 12 + 12 = 24. Config: 2, 8, 2.
1.3 Electron Configuration of IONS
When an atom gains or loses electrons to become an ion, its electron configuration CHANGES. This is the mechanism that links reactivity, bonding, and periodic trends — so it's worth getting solid.
Atom (config)
Gains/Loses
Ion (config)
Why
Na (2, 8, 1)
Loses 1 e⁻
Na⁺ (2, 8)
Empties outer shell → now has stable Ne-like config
Mg (2, 8, 2)
Loses 2 e⁻
Mg²⁺ (2, 8)
Empties outer shell → Ne-like config
Al (2, 8, 3)
Loses 3 e⁻
Al³⁺ (2, 8)
Empties outer shell → Ne-like config
Cl (2, 8, 7)
Gains 1 e⁻
Cl⁻ (2, 8, 8)
Fills outer shell → Ar-like config
O (2, 6)
Gains 2 e⁻
O²⁻ (2, 8)
Fills outer shell → Ne-like config
Why this matters: Every ion in this table ends up with a full outer shell — the same electron config as the nearest noble gas. This is WHY atoms form ions: they're chasing the stability of a full outer shell. It's also why Group 1 always forms 1+ ions, Group 2 always forms 2+, Group 17 always forms 1−, etc.
TRAP: Na⁺ and Ne have the SAME electron configuration (both 2, 8) but they are NOT the same thing. Na⁺ still has 11 protons (it's sodium). Ne has 10 protons. Same electrons ≠ same element. The nucleus defines the element.
Write the electron configuration of Ca²⁺ (Ca has Z=20).
Ca atom = 2, 8, 8, 2. Loses 2 electrons → Ca²⁺ = 2, 8, 8 (same as argon). Group 2 metals lose their 2 outer electrons to form 2+ ions.
1.4 The Periodic Table — Structure
Feature
What It Tells You
Groups (columns 1–18)
Number of valence (outer) electrons. Same group = similar chemical properties.
Periods (rows 1–7)
Number of electron shells. Period 3 = 3 shells.
Metals (left side)
Lose electrons → form positive ions (cations)
Non-metals (right side)
Gain or share electrons → form negative ions (anions) or covalent bonds
Metalloids (staircase line)
Properties of both metals and non-metals (e.g., silicon)
1.5 Periodic Trends
Trend
Across a Period →
Down a Group ↓
Why
Atomic radius
Decreases
Increases
→ More protons pull electrons closer. ↓ More shells = bigger atom.
Electronegativity
Increases
Decreases
→ Stronger pull on shared electrons. ↓ Outer electrons further from nucleus.
Metal reactivity
Decreases
Increases
↓ Outer electron is further away → easier to lose.
Non-metal reactivity
Increases
Decreases
→ Closer to full shell → stronger pull on electrons.
Ionisation energy
Increases
Decreases
→ Harder to remove electron. ↓ Outer electron easier to remove.
WHY metals get more reactive DOWN: The outer electron is further from the nucleus, shielded by inner shells → weaker attraction → easier to lose → more reactive. Think: potassium is more reactive than sodium because K has one more shell.
WHY non-metals get more reactive UP: Fewer shells = outer electrons closer to nucleus = stronger pull on incoming electrons. Fluorine is the most reactive non-metal because its outer shell is closest to the nucleus.
TRAP: "Why is potassium more reactive than sodium?" — Don't just say "it's lower in the group." You MUST explain: more electron shells → greater shielding → weaker nuclear attraction on outer electron → easier to lose → more reactive. The exam wants the MECHANISM, not just the fact.
The 2,8,8 shell model is the Year 10 baseline. But shells are divided into SUBSHELLS: s (holds 2 e⁻) and p (holds 6 e⁻). Notation: 1s² 2s² 2p³ means "2 in 1s, 2 in 2s, 3 in 2p". Shorthand like [He] 2s² means "same core as helium, plus these outer electrons". This is Year 11 content — but understanding it now will save you months later.
1B.1 Why Subshells Exist
The shell model (2, 8, 8) is a simplification. In reality, each shell is split into smaller subshells. For Year 10 you only need to know two types:
Subshell
Holds Max e⁻
Which Shells Have It?
s
2 electrons
Every shell (1s, 2s, 3s...)
p
6 electrons
Shell 2 onwards (2p, 3p...)
This is why shell 1 holds only 2 electrons (it only has 1s), but shell 2 holds 8 (2s + 2p = 2 + 6).
1B.2 How to Write Subshell Notation
Fill subshells in this order: 1s → 2s → 2p → 3s → 3p
Worked Example: Nitrogen (N, Z=7)
7 electrons to place.
Fill 1s first → holds 2 → 1s² (5 electrons left)
Fill 2s next → holds 2 → 2s² (3 electrons left)
Fill 2p next → can hold 6, but we only have 3 → 2p³
Full notation: 1s² 2s² 2p³
Check: 2 + 2 + 3 = 7 electrons ✓
Shell model equivalent: (2, 5) — 2 in shell 1, and 2+3=5 in shell 2
How to read "2p³": The number before the letter (2) is the shell. The letter (p) is the subshell type. The superscript (³) is how many electrons are in that subshell. So "2p³" = "3 electrons in the p subshell of shell 2".
1B.3 The Noble Gas Shorthand — Why "[He]" Shows Up
Writing out full configurations gets tedious for larger atoms. Chemists use noble gas shorthand: replace the filled inner shells with the symbol of the previous noble gas in square brackets.
Element
Full Notation
Shorthand
What [X] Means
Lithium (Z=3)
1s² 2s¹
[He] 2s¹
[He] = the 1s² core (2 electrons)
Carbon (Z=6)
1s² 2s² 2p²
[He] 2s² 2p²
[He] = 1s² (saves writing it)
Nitrogen (Z=7)
1s² 2s² 2p³
[He] 2s² 2p³
Same [He] core, different outer
Sodium (Z=11)
1s² 2s² 2p⁶ 3s¹
[Ne] 3s¹
[Ne] = 1s² 2s² 2p⁶ (10 electrons)
Why [He] shows up for nitrogen: It's not saying nitrogen IS helium. It's saying "nitrogen has the same inner core as helium (1s²), PLUS the outer electrons after that". [He] is just a shortcut for "the first 2 electrons in 1s². Don't bother rewriting them — focus on what's different." Every element has a noble gas core because noble gases have exactly filled shells, which are the natural stopping points.
TRAP: [He] doesn't mean the atom contains helium. It's a notation shortcut meaning "1s² — same as helium has". Same for [Ne], [Ar], etc. The brackets flag the CORE electrons, and what follows is the VALENCE (outer) electrons — the ones that do all the chemistry.
Write the full and shorthand electron configuration for fluorine (Z=9).
Full: 1s² 2s² 2p⁵. Shorthand: [He] 2s² 2p⁵. Check: 2 + 2 + 5 = 9 ✓. Fluorine has 7 valence electrons (2s² 2p⁵) — which matches Group 17.
What does [Ne] 3s² 3p¹ represent?
Aluminium (Al, Z=13). [Ne] = 10 electrons (1s² 2s² 2p⁶), plus 3s² 3p¹ = 3 more. Total 13. Valence electrons = 3 (in shell 3), which matches Group 13.
Orbital filling order — the full picture
After 3p, things get weird because 4s fills BEFORE 3d (it's lower in energy). The full order is: 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p...
You won't be tested on this in Year 10, but recognising the pattern makes Year 11 transition-metal chemistry much easier. The d subshell holds 10 electrons, which is why the transition metal block is 10 columns wide.
2. Chemical Bonding
Metal + non-metal = ionic (electron transfer). Non-metal + non-metal = covalent (electron sharing). Metal + metal = metallic (sea of electrons). Ionic: high MP, conduct when molten/dissolved, NOT solid. Covalent bonds are STRONG — it's the intermolecular forces that are weak. Metals conduct because delocalised electrons are free to move.
2.1 Three Types of Bonding
Bond Type
Between
What Happens
Example
Ionic
Metal + Non-metal
Electrons TRANSFERRED from metal to non-metal
NaCl, MgO, CaF₂
Covalent
Non-metal + Non-metal
Electrons SHARED between atoms
H₂O, CO₂, CH₄
Metallic
Metal + Metal
"Sea of delocalised electrons" around positive metal ions
Fe, Cu, Al, alloys
QUICK RULE: Look at what's bonding. Metal + non-metal = ionic. Non-metal + non-metal = covalent. Metal + metal = metallic. This rule works for 95% of exam questions.
2.2 Ionic Bonding — Electron Transfer
Metals LOSE electrons → become positive ions (cations). Non-metals GAIN electrons → become negative ions (anions). The opposite charges attract → ionic bond.
Worked Example: Formation of Sodium Chloride (NaCl)
Na has config 2, 8, 1 → loses 1 electron → Na⁺ (2, 8)
Cl has config 2, 8, 7 → gains 1 electron → Cl⁻ (2, 8, 8)
Na⁺ and Cl⁻ attract each other → ionic bond
They form a giant ionic lattice (regular 3D arrangement of alternating ions)
Connects to Section 1.3: Notice Na⁺ ends up with config (2, 8) — a full outer shell. Cl⁻ ends up with (2, 8, 8) — also a full outer shell. Ionic bonding IS atoms swapping electrons to reach noble-gas-like stability. Charge of the ion = number of electrons lost/gained.
Properties of Ionic Compounds
Property
Explanation
High melting/boiling points
Strong electrostatic attraction between ions requires lots of energy to overcome
TRAP: "Ionic compounds conduct electricity" — this is only HALF right. They conduct when molten or dissolved (ions can move), but NOT as solids (ions locked in lattice). The exam will specifically test this distinction. Always state the condition.
2.3 Covalent Bonding — Electron Sharing
Non-metal atoms SHARE pairs of electrons to achieve a stable outer shell (usually 8 electrons — the octet rule, or 2 for hydrogen).
Covalent bonding (sharing) vs Ionic bonding (transfer) — know the difference visually
Properties of Covalent Compounds
Property
Simple Molecular (e.g., H₂O, CO₂)
Giant Covalent (e.g., diamond, SiO₂)
Melting point
LOW (weak intermolecular forces)
VERY HIGH (strong covalent bonds throughout)
Conductivity
Do NOT conduct (no free charges)
Do NOT conduct (except graphite)
Solubility
Often insoluble in water
Insoluble
TRAP: "Covalent bonds are weak" = WRONG. Covalent BONDS are strong. It's the INTERMOLECULAR FORCES (forces between molecules) that are weak in simple molecular substances. This is the #1 bonding mistake in exams. Say: "The intermolecular forces are weak, so less energy is needed to separate the molecules."
TRAP: Diamond and graphite are BOTH carbon but have completely different properties because of their structure. Diamond = each C bonded to 4 others (hard, no conductivity). Graphite = layers with delocalised electrons (soft, conducts). The exam loves asking why two forms of the same element differ.
Why does solid NaCl NOT conduct electricity?
In solid NaCl, ions are locked in fixed positions in the crystal lattice and cannot move to carry charge. When molten or dissolved, ions are free to move → conducts.
MgO has a higher melting point than NaCl. Why?
Mg²⁺ and O²⁻ have higher charges than Na⁺ and Cl⁻, so the electrostatic attraction between ions is stronger, requiring more energy to overcome.
2.4 Metallic Bonding
Metal atoms release their outer electrons into a shared "sea of delocalised electrons." The positive metal ions are held together by their attraction to this electron sea.
Properties of Metals
Property
Explanation
Good conductors of electricity
Delocalised electrons are FREE TO MOVE and carry charge
Good conductors of heat
Delocalised electrons transfer kinetic energy quickly
Malleable & ductile
Layers of ions can slide over each other without breaking the bond (electron sea adjusts)
High melting points (usually)
Strong attraction between positive ions and electron sea
Shiny (lustrous)
Delocalised electrons reflect light
EXAM PATTERN: "Explain why copper is a good electrical conductor" — Model answer: "Copper has metallic bonding with delocalised electrons that are free to move through the structure, carrying electrical charge." Always mention: (1) delocalised electrons, (2) free to move, (3) carry charge.
3. Chemical Reactions & Equations
Balance equations by changing coefficients ONLY (never subscripts). Reaction types: synthesis (A+B→AB), decomposition (AB→A+B), displacement (check reactivity series), combustion (fuel+O₂→CO₂+H₂O). A more reactive metal displaces a less reactive one. State symbols: (s) solid, (l) liquid, (g) gas, (aq) dissolved.
3.1 Writing Chemical Equations
A chemical equation shows the reactants (left) and products (right) with an arrow between them.
Word equation: magnesium + oxygen → magnesium oxide
Formula equation:2Mg + O₂ → 2MgO
3.2 Balancing Equations
The law of conservation of mass: atoms are never created or destroyed — they just rearrange. So the number of each type of atom must be equal on both sides.
Worked Example: Balance Fe + O₂ → Fe₂O₃
Count: Left has 1 Fe, 2 O. Right has 2 Fe, 3 O. Not balanced.
Fix iron: put 4Fe on left → need 2Fe₂O₃ on right (4 Fe each side)
Now right has 6 O → need 3O₂ on left (6 O each side)
Balanced: 4Fe + 3O₂ → 2Fe₂O₃ ✓
Check: 4 Fe = 4 Fe ✓ | 6 O = 6 O ✓
TRAP: NEVER change subscripts to balance an equation. Only change the big numbers (coefficients) in front. Changing H₂O to H₂O₂ changes the substance entirely (water → hydrogen peroxide). This costs marks every single exam.
STRATEGY: Balance in this order: (1) metals first, (2) non-metals next, (3) oxygen and hydrogen LAST. This makes balancing much easier.
3.3 Types of Chemical Reactions
Type
Pattern
Example
How to Spot It
Synthesis (combination)
A + B → AB
2Na + Cl₂ → 2NaCl
Two or more substances combine into ONE product
Decomposition
AB → A + B
CaCO₃ → CaO + CO₂
ONE substance breaks into two or more products
Single displacement
A + BC → AC + B
Zn + CuSO₄ → ZnSO₄ + Cu
A more reactive element replaces a less reactive one
Double displacement
AB + CD → AD + CB
NaOH + HCl → NaCl + H₂O
Two compounds swap partners (neutralisation is a type)
Combustion
Fuel + O₂ → CO₂ + H₂O
CH₄ + 2O₂ → CO₂ + 2H₂O
Substance burns in oxygen → always makes CO₂ + H₂O
3.4 The Activity Series (Reactivity Series)
Metals are ranked by how easily they lose electrons. A more reactive metal displaces a less reactive metal from a compound.
Zn can displace Cu (Zn is above Cu), but Cu CANNOT displace Zn. Gold & platinum are so unreactive they're found as pure metals in nature.
TRAP: "Will copper react with zinc sulfate solution?" = NO. Copper is BELOW zinc in the reactivity series, so it cannot displace zinc. Students often forget to check the series before predicting reactions.
EXAM PATTERN: "Predict whether a reaction will occur between [metal] and [metal salt solution]." — Check the reactivity series. If the metal is ABOVE the metal in the compound → reaction occurs. If below → no reaction.
Will silver (Ag) react with copper sulfate (CuSO₄) solution?
No. Silver is BELOW copper in the reactivity series, so it cannot displace copper. A less reactive metal cannot displace a more reactive one.
Balance this equation: Al + O₂ → Al₂O₃
4Al + 3O₂ → 2Al₂O₃. Check: 4 Al each side ✓, 6 O each side ✓.
3.5 State Symbols
Symbol
Meaning
Example
(s)
Solid
NaCl(s)
(l)
Liquid
H₂O(l)
(g)
Gas
CO₂(g)
(aq)
Aqueous (dissolved in water)
NaCl(aq)
3B. Redox — Oxidation & Reduction
OIL RIG: Oxidation Is Loss (of electrons), Reduction Is Gain (of electrons). Oxidation and reduction ALWAYS happen together — if one atom loses electrons, another must gain them. Metals are oxidised when they react (lose electrons). Non-metals are reduced when they react with metals (gain electrons).
3B.1 The Electron View of Reactions
When a metal reacts with a non-metal, electrons move. This movement of electrons is called redox (reduction + oxidation).
Term
Meaning
Who Does This?
Oxidation
LOSS of electrons
Metals (they give up electrons)
Reduction
GAIN of electrons
Non-metals (they take electrons)
Redox reaction
A reaction where electrons transfer
Both happen together — always
MEMORY TOOL — OIL RIG: Oxidation Is Loss (of electrons) Reduction Is Gain (of electrons)
Write "OIL RIG" at the top of your exam paper the moment you sit down.
3B.2 Worked Example — NaCl Formation as Redox
Identify the oxidation and reduction in: 2Na + Cl₂ → 2NaCl
Na starts as atom (0 charge), ends as Na⁺ (+1 charge) → LOST an electron → Na is oxidised
Cl starts as atom (0 charge, in Cl₂), ends as Cl⁻ (−1 charge) → GAINED an electron → Cl is reduced
Half-equations show the electron movement clearly:
Na → Na⁺ + e⁻ (oxidation — electron leaves)
Cl + e⁻ → Cl⁻ (reduction — electron arrives)
Why this matters: Redox is the mechanism behind ionic bond formation, metal reactivity, combustion, rusting, and batteries. The reactivity series is essentially a ranking of "how easily is this metal oxidised". Year 11 electrochemistry is built entirely on this concept — get it locked now.
3B.3 Spotting Redox in Common Reactions
Reaction
What's Oxidised
What's Reduced
2Mg + O₂ → 2MgO
Mg (0 → +2, loses 2e⁻)
O (0 → −2, gains 2e⁻)
Zn + CuSO₄ → ZnSO₄ + Cu
Zn (0 → +2)
Cu²⁺ (+2 → 0)
CH₄ + 2O₂ → CO₂ + 2H₂O (combustion)
C in CH₄
O in O₂
4Fe + 3O₂ → 2Fe₂O₃ (rusting)
Fe (0 → +3)
O (0 → −2)
TRAP: Oxidation does NOT require oxygen. The name is historical — we kept it even after we realised the real definition is "loss of electrons". A metal reacting with chlorine is oxidation, even though no oxygen is involved.
In the reaction Fe + Cu²⁺ → Fe²⁺ + Cu, which species is oxidised?
Fe is oxidised (goes from 0 charge to +2, losing 2 electrons). Cu²⁺ is reduced (gains 2 electrons to become Cu). This is also why iron can displace copper — iron is more easily oxidised, so it "wants" to give up electrons more than copper does.
3C. Gas Tests — Easy Marks
Know the four standard gas tests cold. They come up every single exam and are almost always worth straightforward marks. H₂ = squeaky pop (lit splint). CO₂ = turns limewater milky. O₂ = relights a glowing splint. Cl₂ = bleaches damp litmus paper.
Gas
Test
Positive Result
Where It's Produced
Hydrogen (H₂)
Hold a LIT splint at the mouth of the tube
Squeaky "pop" sound
Acid + metal → salt + H₂
Carbon dioxide (CO₂)
Bubble gas through LIMEWATER (Ca(OH)₂ solution)
Limewater turns milky / cloudy
Acid + carbonate, combustion, respiration
Oxygen (O₂)
Insert a GLOWING (not lit) splint
Splint relights into a flame
Decomposition of H₂O₂, electrolysis of water
Chlorine (Cl₂)
Hold DAMP blue litmus paper near the tube
Paper turns red, then bleaches white
Electrolysis of brine, reactive halogen reactions
The four classic gas tests — memorise the test AND the positive result.
TRAP: Students mix up which splint for which gas. Remember: H₂ needs a LIT splint (you're igniting the gas → pop). O₂ needs a GLOWING splint (there's no flame; oxygen relights it). Getting these backwards loses an easy mark.
TRAP: For CO₂, the limewater test: the milkiness is caused by calcium carbonate (CaCO₃) forming as a precipitate. If you keep bubbling more CO₂ through, the milkiness can disappear (CaCO₃ + CO₂ + H₂O → Ca(HCO₃)₂, which is soluble). So the milkiness is the KEY positive result.
Why this matters: Every "describe the observations" question in an exam expects you to identify gases by their tests. If acid + metal produces a gas and you don't say "test with a lit splint — squeaky pop indicates hydrogen", you're leaving marks on the table.
A student drops a piece of marble chip (CaCO₃) into hydrochloric acid. How would they test the gas produced?
Bubble the gas through limewater. If it turns milky/cloudy, the gas is CO₂. This makes sense: acid + carbonate → salt + water + CO₂.
4. Acids, Bases & pH
Acids produce H⁺ (pH < 7). Bases produce OH⁻ (pH > 7). Strong ≠ concentrated (strong = fully ionised). Acid + base → salt + water. Acid + metal → salt + H₂. Acid + carbonate → salt + water + CO₂. Salt name: metal from base + ending from acid (HCl→chloride, H₂SO₄→sulfate, HNO₃→nitrate).
The difference between "strong" and "weak" acids (or bases) is the degree of ionisation — i.e. what percentage of the molecules actually split into ions when dissolved in water.
TRAP: Strong ≠ concentrated. These are DIFFERENT concepts. Strong = fully dissociates (all molecules split into ions). Concentrated = lots of particles per volume. You can have a dilute strong acid (few molecules, but all ionised) or a concentrated weak acid (many molecules, only some ionised). A 0.1 mol/L HCl solution is DILUTE but STRONG.
Why this matters: Two acids with the same pH can behave very differently. A strong acid at pH 2 will react fast and completely. A weak acid at pH 2 has LOTS more acid molecules in reserve — as H⁺ gets used up, more ionises to replace it. This matters for reaction rate, neutralisation capacity, and biological buffers.
4.4 Key Acid Reactions
Reaction Type
General Equation
Products
Acid + Base (neutralisation)
Acid + Base → Salt + Water
HCl + NaOH → NaCl + H₂O
Acid + Metal
Acid + Metal → Salt + Hydrogen gas
2HCl + Zn → ZnCl₂ + H₂↑
Acid + Carbonate
Acid + Carbonate → Salt + Water + CO₂
2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂↑
NAMING SALTS: The metal comes from the base/metal/carbonate. The ending comes from the acid: HCl → chloride, H₂SO₄ → sulfate, HNO₃ → nitrate. So NaOH + HCl → sodium chloride. Mg + H₂SO₄ → magnesium sulfate.
TRAP: "What gas is produced when acid reacts with a carbonate?" = Carbon dioxide (CO₂). Students often say hydrogen. Hydrogen is produced when acid reacts with a METAL. Test for CO₂ = turns limewater milky. Test for H₂ = squeaky pop with a lit splint. (See Section 3C.)
Name the salt produced when magnesium reacts with sulfuric acid.
A solution has pH 2. Is it a strong acid or concentrated acid?
You can't tell from pH alone. pH 2 means high H⁺ concentration, but it could be a small amount of strong acid (fully ionised) OR a large amount of weak acid. Strong = fully ionised. Concentrated = many particles. They're independent concepts.
5. Rates of Reaction
Rate depends on: temperature, concentration, surface area, catalyst, pressure (gases). ALWAYS explain using collision theory: more/faster collisions = faster rate. Catalyst lowers activation energy, is NOT used up. Powder = MORE surface area (not less). Template: "Increasing [X] causes particles to [collide more/with more energy] because [reason]."
5.1 Collision Theory
For a reaction to happen, particles must:
Collide with each other
With enough energy (≥ activation energy)
In the correct orientation
Anything that increases the frequency or energy of collisions will increase the rate of reaction.
5.2 Factors Affecting Rate
Factor
Change
Effect on Rate
Why (Collision Theory)
Temperature
Increase ↑
Faster
Particles move faster → more frequent collisions + more energy per collision
Concentration
Increase ↑
Faster
More particles per volume → more frequent collisions
Surface area
Increase ↑ (smaller pieces)
Faster
More exposed surface → more particles available to collide
Catalyst
Add catalyst
Faster
Provides alternative pathway with LOWER activation energy
Pressure (gases)
Increase ↑
Faster
Particles pushed closer together → more frequent collisions
TRAP: A catalyst is NOT used up in the reaction. It speeds up the reaction by providing an alternative pathway with lower activation energy, but it is chemically unchanged at the end. If an answer says "the catalyst is consumed" → wrong.
TRAP: "More surface area" doesn't mean "bigger pieces." It means SMALLER pieces (powder vs chunks). A powder has MORE surface area exposed than a single lump because it's broken into thousands of tiny pieces. Students often get this backwards.
EXAM ANSWER TEMPLATE: Every rate question should be answered using collision theory. Use this structure:
"Increasing [factor] causes particles to [collide more frequently / with more energy] because [reason]. This increases the rate of reaction."
Example: "Increasing temperature causes particles to move faster, resulting in more frequent collisions with greater energy. More collisions exceed the activation energy, so the rate of reaction increases."
5.3 Reading Rate Graphs (Concentration/Volume vs Time)
Rate experiments usually plot product formed (volume of gas, mass lost, etc.) on the y-axis against time on the x-axis. Exam questions LOVE asking you to interpret these graphs.
Three classic rate curves. Steepest gradient = fastest rate. Plateau = reaction complete. Higher plateau = more reactant / more product.
How to Read a Rate Graph
Graph Feature
What It Means
Steep gradient (near origin)
Fast rate of reaction at that time
Gradient decreasing over time
Reaction slowing down (reactant being used up)
Plateau / flat line
Reaction has finished — no more change
Higher plateau
More product formed overall → more reactant used
Steeper AND same plateau
Faster reaction, same amount of reactant (e.g. catalyst, higher temp, higher conc of same amount)
Steeper AND higher plateau
More reactant used (e.g. more of the reactant added)
EXAM KEY QUESTION: "Curve A reaches the same plateau as Curve B but faster. What could cause this?" → Something that speeds up the rate without changing the amount of reactant: catalyst, higher temperature, higher concentration, or larger surface area. NOT "more reactant" (that would change the plateau).
TRAP: Students confuse "faster" with "more". A faster reaction reaches completion sooner — but if the amount of reactant is the same, the FINAL amount of product is the same. A steeper curve that plateaus at the SAME height = faster. A curve that plateaus HIGHER = more reactant.
A student repeats an experiment with the same mass of magnesium but uses powder instead of ribbon. The graph of hydrogen volume vs time changes — how?
The curve is steeper at the start (faster rate due to more surface area), but reaches the same plateau (same mass of Mg → same total H₂ produced). Same reactant, faster reaction.
6. Energy Changes in Reactions
Exothermic = releases energy (surroundings get hotter, products LOWER on diagram). Endothermic = absorbs energy (surroundings get colder, products HIGHER). Breaking bonds = absorbs energy. Making bonds = releases energy. EXO = EXIT = energy leaves = products lower. Combustion + neutralisation = exothermic. Photosynthesis + thermal decomposition = endothermic.
6.1 Exothermic vs Endothermic
Feature
Exothermic
Endothermic
Energy
RELEASES energy to surroundings
ABSORBS energy from surroundings
Temperature of surroundings
Goes UP ↑
Goes DOWN ↓
Energy diagram
Products LOWER than reactants
Products HIGHER than reactants
Bond energy
Energy released forming bonds > energy absorbed breaking bonds
Energy absorbed breaking bonds > energy released forming bonds
Exothermic: products lower (energy released). Endothermic: products higher (energy absorbed). The "hump" is the activation energy.
TRAP: "Exothermic = hot, endothermic = cold" — this is too simple and will lose marks. Say: "Exothermic releases energy to the surroundings, increasing the temperature." The temperature of the SURROUNDINGS changes, not the reaction itself.
TRAP: Students confuse which way the energy diagram goes. Remember: EXO = EXIT = energy leaves = products LOWER. If energy is released, the products end up at a lower energy level than the reactants.
BOND ENERGY RULE: Breaking bonds = ABSORBS energy (endothermic process). Making bonds = RELEASES energy (exothermic process). In an exothermic reaction overall, MORE energy is released making new bonds than was absorbed breaking old ones.
6B. Reading Energy Profile Diagrams
Energy diagrams show energy (y-axis) vs reaction progress (x-axis). The "hump" = activation energy (Ea). The gap between reactants and products = energy change (ΔH). Products LOWER than reactants = exothermic (ΔH negative). Products HIGHER = endothermic (ΔH positive). A catalyst LOWERS the hump but doesn't change reactant or product energies.
6B.1 Key Features of Energy Profiles
Exothermic profile. Red solid = no catalyst. Green dashed = with catalyst (lower hump, same ΔH).
What Each Part Tells You
Feature
What It Is
Exam Question It Answers
The hump
Activation energy (Ea) — minimum energy needed to start the reaction
"Label the activation energy"
Height of reactant line
Total bond energy in the reactants
"Identify the reactants"
Height of product line
Total bond energy in the products
"Identify the products"
Gap between reactants & products (ΔH)
Overall energy change of the reaction
"Calculate/describe the energy change"
Products LOWER than reactants
Exothermic — energy released to surroundings
"Is this exo or endothermic? Justify."
Products HIGHER than reactants
Endothermic — energy absorbed from surroundings
Same as above
Smaller hump (dashed)
Catalyst present — alternative pathway with lower Ea
"What has been added to the reaction?"
CLASSIC EXAM QUESTION: "A catalyst is added to the reaction. Draw the new energy profile." → Draw the SAME reactant line, the SAME product line, but a SMALLER hump between them. The catalyst doesn't change ΔH — it only lowers the activation energy.
TRAP: Students draw the catalysed curve with the product line shifted. WRONG. A catalyst does NOT change how much energy is released or absorbed overall. It only changes how much energy is needed to START the reaction. Same reactants, same products, same ΔH — only the hump is smaller.
On an energy profile, the reactants are at 150 kJ and the products are at 80 kJ. The hump peaks at 210 kJ. What are ΔH and Ea?
ΔH = 80 − 150 = −70 kJ (negative because products are lower → exothermic, energy released). Ea = 210 − 150 = 60 kJ (activation energy is measured from reactants up to peak).
An energy profile shows products higher than reactants. Is the reaction exothermic or endothermic? How would the surroundings feel?
Endothermic. Energy is absorbed from the surroundings → the surroundings would feel colder. Example: dissolving ammonium nitrate in water (used in instant cold packs).
7. Moles & Calculations (Year 11 Prep)
n = m/M (moles = mass ÷ molar mass). c = n/V (concentration = moles ÷ volume in LITRES). C₁V₁ = C₂V₂ (dilution). The pipeline: Mass → Moles → Ratio → Moles → Answer. NEVER skip the mole step. Volume must be in LITRES for c = n/V. Coefficients in balanced equation = mole ratio. Limiting reagent = whichever runs out first.
7.1 The Mole Concept
A mole is a counting unit for atoms/molecules — like "dozen" means 12, a mole means 6.022 × 10²³ particles (Avogadro's number).
The Three Mole Formulas
Worked Example 1: How many moles in 46g of sodium (Na)?
Formula: n = m / M
m = 46 g, M = 23 g/mol (from periodic table)
n = 46 / 23 = 2 moles
Worked Example 2: What is the concentration of a solution with 0.5 mol NaOH in 250 mL?
Formula: c = n / V
n = 0.5 mol, V = 250 mL = 0.25 L (must convert!)
c = 0.5 / 0.25 = 2 mol/L
Worked Example 3: 100 mL of 2 mol/L HCl is diluted to 500 mL. New concentration?
Formula: C₁V₁ = C₂V₂
C₁ = 2, V₁ = 100, C₂ = ?, V₂ = 500
2 × 100 = C₂ × 500
C₂ = 200 / 500 = 0.4 mol/L
TRAP: Volume must be in LITRES for c = n/V. Students lose marks by forgetting to convert mL to L. 250 mL = 0.25 L. 500 mL = 0.5 L. Always check units before substituting.
TRAP: In dilution problems (C₁V₁ = C₂V₂), the units must be CONSISTENT on both sides. Either both in mL or both in L — but don't mix them.
7.2 Percentage Composition
Percentage composition tells you the mass percentage of each element in a compound.
% element = (mass of element in formula ÷ molar mass of compound) × 100
Worked Example: % of oxygen in water (H₂O)
Molar mass of H₂O = (2 × 1) + (1 × 16) = 18 g/mol
Mass of O = 16 g/mol
% O = (16/18) × 100 = 88.9%
7.3 The Stoichiometry Pipeline — The ONE Flow That Solves 90% of Calculations
This is the single most important concept for Year 11 prep. Every calculation follows this pipeline:
The Master Pipeline: Mass → Moles → Ratio → Moles → Answer. If you understand this flow, you can solve 90% of calculation questions.
THE BIG INSIGHT: You can NEVER skip the mole step. Mass ratios don't work — mole ratios do. Every stoichiometry problem goes: convert GIVEN → moles → apply ratio → convert to what's ASKED.
Full Stoichiometry Worked Example
Q: What mass of carbon dioxide is produced when 10g of calcium carbonate reacts with excess HCl?
Step 1: Write balanced equation
CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂
Step 2: Convert given → moles
M(CaCO₃) = 40 + 12 + (3×16) = 100 g/mol
n(CaCO₃) = 10 / 100 = 0.1 mol
Step 3: Apply mole ratio
Ratio CaCO₃ : CO₂ = 1 : 1
So n(CO₂) = 0.1 mol
Step 4: Convert moles → answer
M(CO₂) = 12 + (2×16) = 44 g/mol
m(CO₂) = 0.1 × 44 = 4.4 g
MEGA TRAP: Students skip the mole step and try to use mass ratios directly: "10g in, so 10g out" = WRONG. Different substances have different molar masses. You MUST go through moles. If you didn't convert to moles, your answer is almost certainly wrong.
7.4 Empirical Formula
The empirical formula is the simplest whole-number ratio of atoms in a compound.
Worked Example: A compound is 40% carbon, 6.7% hydrogen, 53.3% oxygen. Find the empirical formula.
C = 3.33/3.33 = 1 | H = 6.7/3.33 = 2 | O = 3.33/3.33 = 1
Step 4: Write formula → CH₂O (ratio 1:2:1)
TRAP: If you get ratios like 1:1.5:1, you can't have half an atom. Multiply ALL values by 2 to get whole numbers (2:3:2). If you get X.33, multiply by 3. Always convert to the simplest WHOLE number ratio.
BONUS WEAPON: If the exam also gives you the molar mass of the compound, you can find the MOLECULAR formula: divide the molar mass by the empirical formula mass, then multiply all subscripts by that number. E.g., if empirical = CH₂O (mass 30) and actual molar mass = 180 → 180/30 = 6 → molecular formula = C₆H₁₂O₆ (glucose!).
7.5 Limiting Reagent — Core Skill
In real reactions, you're almost never given the exact amounts needed. One reactant runs out first — that's the limiting reagent. It determines how much product you can make. The other reactant is said to be "in excess".
3-STEP METHOD: 1. Calculate moles of EACH reactant. 2. Divide each by its coefficient in the balanced equation. 3. The SMALLER number = limiting reagent. Use that reactant's moles to calculate product.
Worked Example: 6g of Mg reacts with 8g of O₂. Which is limiting? What mass of MgO forms?
Balanced equation: 2Mg + O₂ → 2MgO
Step 1 — Moles of each:
n(Mg) = 6 / 24 = 0.25 mol
n(O₂) = 8 / 32 = 0.25 mol
Step 2 — Divide by coefficient:
Mg: 0.25 ÷ 2 = 0.125
O₂: 0.25 ÷ 1 = 0.25
Step 3 — Smaller = limiting: Mg is the limiting reagent (0.125 < 0.25). O₂ is in excess.
Calculate product: Use n(Mg) with the Mg:MgO ratio = 2:2 = 1:1
n(MgO) = 0.25 mol (same as Mg used)
M(MgO) = 24 + 16 = 40 g/mol
m(MgO) = 0.25 × 40 = 10 g
Why this matters: Limiting reagent problems come up in every exam from Year 11 onwards, and they're increasingly common in Year 10 too. They're also the foundation of real-world chemistry — in industry, you almost always deliberately use one reactant in excess to make sure the expensive reactant is fully used.
TRAP: Students pick the limiting reagent based on moles alone, without dividing by coefficients. WRONG. If the balanced equation is 2A + 3B → products and you have 2 mol A and 2 mol B, then A: 2÷2 = 1, B: 2÷3 = 0.67. B is limiting even though you have equal moles.
TRAP: "The limiting reagent runs out last" — WRONG. It runs out FIRST (hence "limiting" — it limits how much product you can make). The excess reagent is the one left over at the end.
In 2H₂ + O₂ → 2H₂O, you have 4 mol H₂ and 1 mol O₂. Which is the limiting reagent?
H₂: 4 ÷ 2 = 2. O₂: 1 ÷ 1 = 1. Smaller = O₂ is the limiting reagent. You'd produce n(H₂O) based on O₂: the ratio O₂ : H₂O is 1:2, so 1 mol O₂ → 2 mol H₂O. The leftover H₂ (2 mol) is in excess.
7B. States of Matter & Particle Theory
Solids: particles vibrate in fixed positions, closely packed, strong forces. Liquids: particles slide past each other, close but not fixed, moderate forces. Gases: particles move freely and fast, widely spaced, very weak forces. Gases compress because there's lots of space between particles.
Property
Solid
Liquid
Gas
Particle arrangement
Regular, closely packed
Random, close together
Random, widely spaced
Particle movement
Vibrate in fixed positions
Slide past each other
Move fast in all directions
Forces between particles
Strong
Moderate
Very weak
Shape
Fixed
Takes shape of container
Fills entire container
Volume
Fixed
Fixed
Expands to fill space
Compressible?
No
No (almost)
Yes (lots of space between particles)
Density
High
High
Low
Solid: tightly packed, fixed positions. Liquid: close but sliding. Gas: widely spaced, fast movement.
TRAP: "Why can gases be compressed but liquids can't?" — Because gas particles have LARGE SPACES between them that can be reduced. Liquid particles are already close together with minimal space. Don't say "gas particles are smaller" — they're not; there's just more SPACE between them.
A student heats a solid and it becomes a liquid. What happens to the particles?
The particles gain kinetic energy and vibrate more. Eventually they overcome the forces holding them in fixed positions and begin to slide past each other. The arrangement changes from ordered to random, but particles remain close together.
7C. Separation Techniques
Filtration = insoluble solid from liquid. Evaporation = dissolved solid from solution. Simple distillation = liquid from solution (collect the liquid). Chromatography = separate mixtures of dissolved substances. Choose technique based on WHAT you want to keep and the PROPERTIES of the mixture.
Technique
Separates
When to Use
What You Collect
Filtration
Insoluble solid from liquid
Sand from saltwater, precipitate from solution
Residue (solid) in filter paper, filtrate (liquid) passes through
Evaporation
Dissolved solid from solution
Getting salt from saltwater
Solid crystals remain after water evaporates
Simple Distillation
Liquid from a solution
Getting PURE WATER from saltwater
The liquid (condensed vapour). Solid remains behind.
Chromatography
Dissolved substances from each other
Separating colours in ink, identifying unknown substances
Spots on chromatography paper showing different components
Filtration keeps the solid. Evaporation keeps the solid. Distillation keeps the liquid.
KEY DECISION: "Which technique?" depends on WHAT you want to keep:
→ Want the solid? Evaporation (dissolved solid) or Filtration (insoluble solid).
→ Want the liquid? Distillation (collects the condensed liquid).
→ Want to identify components in a mixture? Chromatography.
Chromatography — How It Works
A spot of mixture is placed on paper. A solvent travels up the paper carrying dissolved components. Different substances travel different distances based on their solubility in the solvent. Pure substances produce ONE spot. Mixtures produce MULTIPLE spots.
Rf VALUE: Rf = distance travelled by substance ÷ distance travelled by solvent. Each substance has a unique Rf value, so you can identify unknowns by comparing Rf values to known standards.
Fractional Distillation — separates a mixture of liquids with different boiling points (e.g., crude oil, alcohol from water).
How it differs from simple distillation: uses a fractionating column that provides multiple condensation/evaporation cycles. Liquids with lower boiling points rise higher and are collected first. Liquids with higher boiling points condense lower and are collected later.
Classic exam question: "Explain the difference between simple and fractional distillation."
→ Simple: separates a liquid from a dissolved solid (one boiling point).
→ Fractional: separates two or more liquids with different boiling points using a fractionating column.
Paper vs Thin-Layer Chromatography (TLC): Paper chromatography uses paper as the stationary phase. TLC uses a thin layer of silica on a glass plate — it's faster, more precise, and used in forensics/pharma. Same principle, different surface.
EXCEPT silver chloride (AgCl) and lead chloride (PbCl₂)
Most sulfates are soluble
EXCEPT barium sulfate (BaSO₄) and lead sulfate (PbSO₄)
Most carbonates are insoluble
EXCEPT Group 1 and ammonium carbonates
Most hydroxides are insoluble
EXCEPT Group 1 and Ca(OH)₂ (slightly)
TRAP: "Predict whether a precipitate forms when solutions are mixed" — use solubility rules. If the potential product is insoluble → it forms a precipitate (solid). If soluble → no precipitate. Example: NaCl(aq) + AgNO₃(aq) → AgCl(s)↓ + NaNO₃(aq). AgCl is insoluble → white precipitate forms.
You have saltwater and need to collect pure water. Which technique?
Simple distillation. Heat the saltwater → water evaporates → travels through condenser → condenses back to pure liquid water. The salt remains behind in the flask. (NOT evaporation — that keeps the salt, not the water.)
How would you test if a black ink contains one or more dyes?
Chromatography. Place a spot of ink on chromatography paper, dip edge in solvent. If the ink is a mixture, it will separate into multiple spots at different heights. A pure substance produces only one spot.
8. Working Scientifically — Experimental Skills
8.1 Variables
Variable Type
Definition
Example (Rate of Reaction Experiment)
Independent
What YOU change
Temperature of the acid
Dependent
What you MEASURE (the result)
Time for reaction to complete
Controlled
What you keep the SAME
Concentration, volume, type of metal, surface area
TRAP: "What is the controlled variable?" — Students often confuse this with "controlling the experiment." Controlled variables are the things you keep CONSTANT so the test is fair. If you're testing temperature, you must keep concentration, volume, and surface area the same.
8.2 Reliability, Accuracy & Validity
Concept
Meaning
How to Achieve It
Reliability
Getting consistent, repeatable results
Repeat trials (at least 3) and calculate average
Accuracy
How close results are to the TRUE value
Use calibrated equipment, correct technique
Validity
Whether the experiment actually tests what it claims to
Control all variables except the one you're testing
EXAM LANGUAGE: "How could you improve reliability?" → "Repeat the experiment multiple times and calculate a mean average." "How could you improve validity?" → "Ensure all controlled variables are kept constant." Know these phrases — they're worth easy marks.
8.3 Common Experimental Errors
Systematic error: Affects all results the same way (e.g., zero error on a balance). Results are precise but NOT accurate.
Random error: Varies unpredictably between measurements. Reduced by repeating trials.
Parallax error: Reading a meniscus from the wrong angle. Always read at EYE LEVEL from the bottom of the meniscus.
9. Exam Technique & How to Answer
9.1 How to Answer — Mark-Scoring Templates
These templates show you EXACTLY how to structure answers for full marks. Use them every time.
Bonding / Properties Questions
TEMPLATE: [Substance] has [type] bonding/structure → [describe the forces/structure] → therefore [property] because [link to particles/energy].
Example: "Explain why NaCl has a high melting point."
Structure: NaCl has ionic bonding in a giant ionic lattice
Forces: Strong electrostatic attraction between oppositely charged Na⁺ and Cl⁻ ions
Link to property: A large amount of energy is needed to overcome these strong attractions, so the melting point is high
Rate of Reaction Questions
TEMPLATE: Increasing [factor] causes particles to [what changes] → [effect on collisions] → [effect on rate]. ALWAYS use collision theory words.
Effect on collisions: Particles collide more frequently
Effect on rate: More collisions per second that exceed the activation energy → rate increases
Calculation Questions
TEMPLATE: (1) Write formula → (2) Substitute values → (3) Calculate → (4) Write answer with UNITS. Show ALL steps. Even if you get the number right, missing units or missing working = lost marks.
Example: "Calculate moles in 20g of NaOH (M = 40 g/mol)."
Formula: n = m / M
Substitute: n = 20 / 40
Calculate: n = 0.5
Answer with units: n = 0.5 mol
Compare / Contrast Questions
TEMPLATE: [Thing A] has [feature], whereas [Thing B] has [different feature]. This is because [reason]. Always state BOTH sides — don't just describe one.
Command Words — What They Actually Want
Command Word
What It Means
How Many Marks (Usually)
State / Name / Give
Just say it. No explanation needed.
1 mark
Describe
Say WHAT happens (observations, trends)
1–2 marks
Explain
Say what happens AND WHY (link to theory)
2–4 marks
Compare
State similarities AND differences
2–3 marks
Evaluate
Pros and cons, then give a judgment
3–4 marks
Justify
Give reasons to support a conclusion
2–3 marks
Calculate
Show all working, include units
2–3 marks
BIGGEST MARK KILLER: Writing "Describe" answers for "Explain" questions. If it says "Explain", you MUST say WHY. "The rate increases" = 1 mark. "The rate increases BECAUSE particles have more kinetic energy, leading to more frequent collisions with enough energy to overcome the activation energy" = 3 marks. The WHY is where the marks are.
9.2 The 5 Golden Rules
RULE 1: Use scientific language. Don't write "the stuff gets hot." Write "the temperature of the surroundings increases, indicating an exothermic reaction."
RULE 2: Always explain WHY, not just WHAT. "The rate increases" gets 1 mark. "The rate increases because particles have more kinetic energy, leading to more frequent and energetic collisions" gets full marks.
RULE 3: Link to particles. Chemistry answers almost always need to refer to atoms, ions, electrons, molecules, or particles. If your answer doesn't mention particles, it's probably incomplete.
RULE 4: Read the command word. "Describe" = say what happens. "Explain" = say what happens AND why. "Compare" = state similarities AND differences. "Evaluate" = give positives and negatives then a judgment.
RULE 5: Check your units. Marks are lost for wrong units more than wrong numbers. Temperature in °C. Mass in grams. Volume in mL or L (check which one the formula needs). Concentration in mol/L.
9.3 Common Exam Traps — Master Table
Trap
What Students Write ❌
What You Should Write ✅
Atomic vs mass number
"Atomic number = protons + neutrons"
"Atomic number = protons only. Mass number = protons + neutrons."
Ionic conductivity
"Ionic compounds conduct electricity"
"Ionic compounds conduct when molten or dissolved, not as solids"
Covalent bond strength
"Covalent bonds are weak"
"Covalent bonds are strong. The intermolecular forces are weak."
Strong vs concentrated
"Strong acid = lots of acid"
"Strong = fully ionised. Concentrated = high number of particles per volume."
Catalyst use
"The catalyst is used up"
"The catalyst provides an alternative pathway with lower activation energy and is not consumed."
Surface area
"Bigger pieces = more surface area"
"Smaller pieces (powder) = more surface area exposed."
Exo/endo confusion
"Exothermic = cold"
"Exothermic releases energy → surroundings get hotter."
Balancing equations
Changes subscripts (H₂O → H₃O)
Only change coefficients (2H₂O, not H₃O).
Carbonate + acid
"Produces hydrogen gas"
"Produces CO₂ + water + salt. H₂ is from acid + metal."
Isotopes
"Different chemical properties"
"Same chemical properties (same electrons). Different physical properties (different mass)."
Reactivity series
"Cu displaces Zn"
"Cu is BELOW Zn → cannot displace it. Zn displaces Cu."
Mole calculations
Uses mL instead of L for c = n/V
"Convert mL to L first. 250 mL = 0.25 L."
Energy diagrams
"Products higher = exothermic"
"Products LOWER = exothermic (energy was released)."
Bond breaking/making
"Breaking bonds releases energy"
"Breaking bonds ABSORBS energy. Making bonds RELEASES energy."
Limiting reagent
"Smaller moles = limiting"
"Divide moles by coefficient first — smallest result = limiting."
Oxidation = oxygen
"Oxidation needs oxygen"
"Oxidation = loss of electrons. Oxygen not required (OIL RIG)."
MOST MARKS ARE LOST FROM: Incomplete explanations, not wrong knowledge. You usually know the right answer — you just don't write ENOUGH. Always ask yourself: "Have I explained WHY? Have I mentioned particles? Have I used scientific terms?"
10. Practice Questions & Worked Solutions
Try each question before reading the answer. Cover the solution column first. Questions are grouped by type so you can target weak areas.
10.1 Short Answer — Explain/Compare (2–3 marks each)
#
Question
Model Answer
1
Explain why sodium is more reactive than lithium.
Sodium has more electron shells than lithium (3 vs 2), so the outer electron is further from the nucleus with more shielding. This means weaker attraction → easier to lose the electron → more reactive.
2
Why do ionic compounds have high melting points?
Ionic compounds have strong electrostatic attraction between oppositely charged ions arranged in a giant lattice. A large amount of energy is needed to overcome these strong attractions.
3
Explain why solid NaCl does not conduct electricity but molten NaCl does.
In solid NaCl, ions are fixed in the lattice and cannot move. When molten, ions are free to move and carry charge, allowing electrical conductivity.
4
Explain why increasing temperature increases the rate of reaction.
Higher temperature gives particles more kinetic energy. They move faster, collide more frequently, and more collisions have enough energy to overcome the activation energy barrier.
5
What is the difference between a strong acid and a concentrated acid?
A strong acid fully dissociates (ionises) in water — all molecules break into ions. A concentrated acid has a large number of acid particles per unit volume. A dilute strong acid is still fully ionised; a concentrated weak acid is only partially ionised.
6
Predict whether iron will react with copper sulfate solution. Explain why.
Yes, iron will react. Iron is above copper in the reactivity series, so it is more reactive and can displace copper. Fe + CuSO₄ → FeSO₄ + Cu. The iron dissolves and copper metal is deposited.
7
Explain why diamond has a very high melting point.
Diamond is a giant covalent structure where each carbon atom is bonded to four others by strong covalent bonds in a rigid tetrahedral arrangement. Breaking these many strong bonds requires a very large amount of energy.
8
A student adds hydrochloric acid to calcium carbonate. Describe what they would observe and write the equation.
Observation: Fizzing/bubbles (CO₂ gas released), solid dissolves. Equation: 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂. Test for CO₂: pass through limewater → turns milky.
10.2 Electron Configuration & Ions
#
Question
Model Answer
9
Write the full and shorthand electron configuration of oxygen (Z=8).
Write the electron configuration (shell model) of Mg²⁺ and explain why it is stable.
Mg atom: 2, 8, 2. Loses 2 electrons → Mg²⁺: 2, 8. It is stable because it now has a full outer shell — the same electron configuration as neon (a noble gas).
11
Na⁺ and Ne both have the electron configuration 2, 8. Are they the same element? Explain.
No. Same electron configuration does NOT mean same element. Na⁺ has 11 protons (still sodium, just missing one electron). Ne has 10 protons. The number of protons (atomic number) defines the element, not the number of electrons.
10.3 Redox & Gas Tests
#
Question
Model Answer
12
In the reaction 2Mg + O₂ → 2MgO, identify what is oxidised and what is reduced.
Mg is oxidised (loses 2 electrons, 0 → +2 charge). O is reduced (gains 2 electrons, 0 → −2 charge). Remember OIL RIG: Oxidation Is Loss, Reduction Is Gain.
13
A gas is produced when zinc reacts with dilute HCl. Describe how you would test this gas and state the expected result.
Hold a lit splint at the mouth of the tube. A squeaky "pop" indicates hydrogen gas (H₂). This is consistent with the reaction: Zn + 2HCl → ZnCl₂ + H₂.
14
How would you distinguish between a test tube of O₂ and a test tube of CO₂?
Use a glowing splint: it will relight in O₂, but be extinguished in CO₂. Alternatively, bubble each through limewater: CO₂ turns it milky, O₂ does not.
10.4 Calculations (Including Limiting Reagent)
#
Question
Worked Solution
15
Calculate the number of moles in 80g of NaOH (M = 40 g/mol).
n = m/M = 80/40 = 2.0 mol
16
What mass of magnesium contains 0.25 mol? (M = 24 g/mol)
m = n × M = 0.25 × 24 = 6.0 g
17
Calculate the concentration of a solution containing 0.1 mol HCl in 500 mL.
Convert: 500 mL = 0.5 L. c = n/V = 0.1/0.5 = 0.2 mol/L
18
50 mL of 4 mol/L H₂SO₄ is diluted to 200 mL. What is the new concentration?
What mass of CO₂ is produced when 25g of CaCO₃ reacts with excess HCl? (Ca=40, C=12, O=16)
Equation: CaCO₃ + 2HCl → CaCl₂ + H₂O + CO₂. M(CaCO₃) = 100. n = 25/100 = 0.25 mol. Ratio 1:1. n(CO₂) = 0.25 mol. M(CO₂) = 44. m = 0.25 × 44 = 11 g
20
In the reaction N₂ + 3H₂ → 2NH₃, you have 2 mol N₂ and 3 mol H₂. Identify the limiting reagent and calculate the moles of NH₃ produced.
Divide by coefficients: N₂: 2/1 = 2. H₂: 3/3 = 1. Smaller = H₂ is limiting. Ratio H₂:NH₃ = 3:2. n(NH₃) = 3 × (2/3) = 2 mol NH₃. N₂ is in excess.
21
A compound is 75% carbon and 25% hydrogen. Find the empirical formula. (C=12, H=1)
Assume 100g: 75g C, 25g H. n(C) = 75/12 = 6.25. n(H) = 25/1 = 25. Divide by smallest: C = 6.25/6.25 = 1, H = 25/6.25 = 4. Empirical formula = CH₄ (methane)
10.5 Graph Interpretation
#
Question
Model Answer
22
Two reaction curves (gas volume vs time) are drawn. Curve A is steeper than Curve B but they reach the same final volume. Suggest what could cause this difference.
A factor that increases rate WITHOUT changing the amount of reactant: higher temperature, higher concentration, larger surface area, or presence of a catalyst. Same plateau = same amount of reactant used.
23
An energy profile shows reactants at 120 kJ, peak at 200 kJ, products at 60 kJ. State (a) the activation energy, (b) the energy change ΔH, and (c) whether the reaction is exo- or endothermic.
(a) Ea = 200 − 120 = 80 kJ. (b) ΔH = 60 − 120 = −60 kJ. (c) Exothermic (products lower than reactants; ΔH is negative).
24
On the energy profile above, how would the curve change if a catalyst were added?
The hump would be LOWER (smaller activation energy) — a catalyst provides an alternative pathway with lower Ea. The reactant and product energies (and therefore ΔH) would stay the SAME.
A reaction is exothermic. The student says "energy is created." Is this correct? Explain.
No. Energy cannot be created or destroyed (Law of Conservation of Energy). The energy released was stored as chemical energy in the bonds of the reactants. More energy is released forming new bonds than was absorbed breaking old bonds.
26
Why does a catalyst NOT change the products of a reaction?
A catalyst provides an alternative pathway with lower activation energy. It doesn't change the reactants or products — only the route between them. The same bonds are broken and formed, just more easily.
27
Graphite conducts electricity but diamond doesn't. Both are carbon. Explain.
In graphite, each carbon bonds to 3 others (not 4), leaving one electron per carbon delocalised and free to move — these carry charge. In diamond, all 4 electrons are used in bonding — none are free to move.
28
A student mixes acid with metal. No reaction occurs. Suggest two possible reasons.
(1) The metal is below hydrogen in the reactivity series (e.g., copper, silver, gold) so it cannot displace hydrogen. (2) The acid is too dilute / weak to initiate the reaction at a measurable rate.
29
Explain using redox why a more reactive metal displaces a less reactive metal from solution.
The more reactive metal is more easily oxidised (loses electrons more readily). Those electrons are taken by the less reactive metal's ion in solution, which is reduced to its metallic form. So the more reactive metal "steals" the place of the less reactive one.
📋 Exam Night Summary — 1-Page Brain Primer
Read this the night before or morning of the exam. Everything you need in one view.
⚛️ Atomic Structure
Z = protons (identifies element). A = protons + neutrons. Neutral atom: protons = electrons. Config fills 2, 8, 8. Group = valence electrons. Period = shells. Isotopes = same Z, different neutrons, same chemical properties. Subshells: s (2 e⁻), p (6 e⁻). Shorthand [He] = 1s² core. Ions have different electron counts (Na⁺ = 2,8 = Ne config, but still sodium).
🔗 Bonding
Metal + non-metal = ionic (transfer, lattice, high MP, conduct when molten). Non-metal + non-metal = covalent (sharing, molecules, low MP, no conduction). Metal = metallic (sea of electrons, conduct, malleable). Covalent bonds are STRONG — intermolecular forces are weak. Ionic solids DON'T conduct.
⚗️ Reactions & Redox
Synthesis: A+B→AB. Decomposition: AB→A+B. Displacement: check reactivity series. Combustion: fuel+O₂→CO₂+H₂O. Balance by changing coefficients ONLY. Acid+base→salt+water. Acid+metal→salt+H₂. Acid+carbonate→salt+water+CO₂. OIL RIG: Oxidation Is Loss, Reduction Is Gain (of electrons). Oxidation doesn't need oxygen.
Exo = releases energy, surroundings hotter, products LOWER on diagram. Endo = absorbs, surroundings colder, products HIGHER. Breaking bonds ABSORBS energy. Making bonds RELEASES energy. EXO = EXIT = energy leaves. Catalyst lowers the HUMP (Ea), not ΔH.
🔢 Key Formulas
n = m / M
moles = mass ÷ molar mass
c = n / V
concentration = moles ÷ volume (LITRES!)
C₁V₁ = C₂V₂
dilution formula
% = (part/total) × 100
percentage composition
Mass → Moles → Ratio → Moles → Answer
THE stoichiometry pipeline
Divide moles by coefficient
smallest = limiting reagent
🧪 Separation Quick-Pick
Want the solid (dissolved)?
Evaporation
Want the solid (insoluble)?
Filtration
Want the liquid (pure)?
Distillation
Identify components?
Chromatography
Separate liquids (diff. boiling pts)?
Fractional distillation
🚫 Top Traps
1. Atomic number ≠ mass number ·
2. Ionic solids DON'T conduct ·
3. Covalent bonds are strong (intermolecular forces weak) ·
4. Strong ≠ concentrated ·
5. Never change subscripts to balance ·
6. Catalyst is NOT used up ·
7. Powder = MORE surface area ·
8. Exo = products LOWER ·
9. Breaking bonds ABSORBS energy ·
10. Volume must be in LITRES for c = n/V ·
11. Limiting reagent = divide moles by coefficient, smallest wins ·
12. Oxidation doesn't need oxygen — it's loss of ELECTRONS ·
13. H₂ = LIT splint, O₂ = GLOWING splint ·
14. Catalyst lowers the hump, not ΔH
✍️ How to Answer
Bonding Q: Name structure → describe forces → link to property Rate Q: State change → collision theory → effect on rate Calculation: Write formula → substitute → calculate → units Redox Q: Identify electron movement → apply OIL RIG Graph Q: Identify Ea (hump), ΔH (gap), exo/endo (products higher or lower) "Explain" Q: Say WHAT happens AND WHY (link to particles). The WHY is where the marks are. "Compare" Q: State BOTH sides. "X has... whereas Y has..."
You've got this! 🧪
Remember: explain WHY, link to particles, use scientific language, check your units.