Tick off each objective as you master it. These are the exact learning objectives from the Cambridge 0620 syllabus (2026-2028).
All matter is made of particles (atoms and molecules). These particles are in constant motion, and the state of a substance—whether it's a solid, liquid, or gas—depends on three things: (1) how far apart the particles are, (2) how arranged they are, and (3) how fast they're moving. The kinetic particle theory explains everything about states of matter using this simple model.
Structure: Particles are tightly packed in a fixed, regular arrangement (a lattice). They vibrate in fixed positions but cannot move around.
Properties: Fixed shape, fixed volume, cannot be compressed, cannot flow. Hard and rigid.
Why these properties? Strong forces between particles keep them locked in place. Particles have very little kinetic energy—they only vibrate slightly about their fixed positions.
Examples: ice, iron, salt crystals, wood.
Structure: Particles are still close together but NOT in a fixed arrangement. They're touching or nearly touching, but can move around and swap positions with neighbours.
Properties: Fixed volume (liquids do NOT compress), but NO fixed shape—they take the shape of their container. Can flow. Denser than gases but less dense than solids.
Why these properties? Forces between particles are strong enough to keep them in contact (so volume is fixed) but weak enough to allow movement (so shape changes). Particles have more kinetic energy than in solids.
Examples: water, mercury, oil, alcohol.
Structure: Particles are far apart, with no fixed arrangement. They move rapidly and randomly in all directions, colliding with each other and container walls.
Properties: No fixed shape, no fixed volume—gases expand to fill any container. Very easily compressed. Can flow. Much less dense than solids and liquids.
Why these properties? Forces between particles are negligible. Particles have very high kinetic energy. They're far apart most of the time, constantly moving and colliding.
Examples: oxygen, nitrogen, steam, carbon dioxide.
When you heat or cool a substance, you're adding or removing energy. This changes how fast the particles move, which can change the state.
Melting: Heating a solid so particles vibrate faster and faster until they break free from fixed positions. Temperature stays constant during melting (energy goes into breaking bonds, not raising temperature). When melting finishes, all particles can move freely—the substance is now liquid.
Boiling: Heating a liquid so particles move fast enough to escape into the gas phase. Boiling happens at a specific temperature (the boiling point). During boiling, temperature stays constant—energy goes into separating particles, not raising temperature.
Evaporation: Particles at the surface of a liquid gain enough energy to escape as gas, even below boiling point. It's a slower process than boiling and happens at any temperature. Only the most energetic surface particles escape, so the liquid cools down.
Freezing: Cooling a liquid so particles slow down, move less, and eventually get locked into fixed positions. Temperature stays constant during freezing. When freezing finishes, all particles are vibrating in fixed positions—the substance is now solid.
Condensing: Cooling a gas so particles slow down and clump together into a liquid. Temperature stays constant during condensation. When condensation finishes, particles are touching but can move around.
Sublimation: Some solids (like dry ice and iodine) can transform directly to gases without melting first. Heating breaks the rigid structure, and particles escape as gas. Rare but important to know.
A heating curve shows temperature vs. time as you heat a substance. It has three key features:
1. Straight diagonal line (rising): Temperature increases as you heat. Particles move faster. This happens in the solid phase, then in the liquid phase, then in the gas phase.
2. Flat horizontal line (plateau): Temperature STOPS changing even though you keep heating. This is a change of state. Energy goes into separating particles, not into motion. Example: melting ice stays at 0°C until all ice is melted.
3. When the change finishes, temperature rises again in the new state.
Cooling curves are the reverse: falling diagonal lines with flat plateaus during condensing and freezing.
For a fixed amount of gas at constant temperature: if you decrease the volume (compress the gas), the pressure increases. If you increase the volume (expand the gas), the pressure decreases. This is Boyle's Law: pressure is inversely proportional to volume (pV = constant).
Why? In kinetic particle theory, pressure comes from particles colliding with the container walls. If you compress the gas into a smaller volume, the particles collide with the walls more frequently in the same time period → higher pressure. If you expand it, fewer collisions → lower pressure.
Temperature effect: Higher temperature → particles move faster → they collide harder and more frequently → pressure increases. Lower temperature → particles move slower → fewer, weaker collisions → pressure decreases.
Diffusion is the spreading out of particles from an area of higher concentration to an area of lower concentration, until they're evenly mixed.
In gases: Diffusion is fast because gas particles move quickly and are far apart. Ammonia gas (NH₃) and hydrogen chloride gas (HCl) diffuse through air. If you open a bottle of ammonia at one end of a room, you'll smell it throughout the room within minutes.
In liquids: Diffusion is slower because particles are closer together and moving more slowly than in gases. Dissolving a dye in water shows diffusion—the colour spreads through the water gradually.
Why does diffusion happen? Kinetic particle theory explains it: particles are in constant random motion. By chance, more particles move from the concentrated region to the dilute region than vice versa, so over time, the concentration evens out.
Rate of diffusion depends on molecular mass: Lighter molecules diffuse faster than heavier ones. For example, hydrogen (H₂, mass 2) diffuses much faster than oxygen (O₂, mass 32). At the same temperature, lighter particles move faster (higher average kinetic energy per particle), so they travel further and collide more frequently. This is why HCl gas diffuses faster than NH₃ when both are released—wait, actually NH₃ (mass 17) diffuses faster than HCl (mass 36.5) because NH₃ is lighter. Classical experiment: open bottles of concentrated ammonia and HCl at opposite ends of a tube; the white ring of ammonium chloride forms closer to the HCl end because ammonia diffuses faster.
These definitions use the exact wording expected in IGCSE mark schemes. Click each term to reveal.
The diagram below shows a heating curve for ice being heated at constant pressure until it becomes steam.
State the name of the process occurring between point B and point C, and explain why the temperature does not change during this process.
✓ The process is boiling / vaporization. [1 mark for process name]
Explanation: During boiling, energy is used to overcome the forces of attraction between particles, allowing them to escape from the liquid into the gas state. [1 mark for energy/forces] This energy does not increase the kinetic energy (motion) of the particles, so temperature remains constant. [1 mark for explaining temperature doesn't change]
Explain, using kinetic particle theory, why a gas exerts pressure on the walls of its container.
✓ Gas particles are in constant random motion [1 mark]
✓ They collide with the container walls; these collisions exert a force on the walls, creating pressure. [1 mark]
Two jars are opened—one containing ammonia gas (NH₃, relative molecular mass 17) and one containing hydrogen chloride gas (HCl, relative molecular mass 36.5)—at opposite ends of a long tube. A white ring of ammonium chloride forms closer to the HCl jar. Explain this observation using kinetic particle theory.
✓ Ammonia molecules diffuse faster than HCl molecules [1 mark for observation]
✓ At the same temperature, lighter molecules move faster than heavier ones [1 mark for kinetic theory principle]
✓ Ammonia (mass 17) is lighter than HCl (mass 36.5), so NH₃ molecules travel faster and diffuse further; they meet the slower HCl molecules closer to the HCl jar end, where they react to form the white ring. [1 mark for complete explanation]
Explain why a gas is easily compressed, but a liquid is almost incompressible.
✓ In a gas, particles are far apart with lots of empty space between them. When compressed, this empty space is reduced, allowing particles to move closer together. [1 mark]
✓ In a liquid, particles are already touching and close together. There is no empty space to remove, so the volume cannot be reduced further. [1 mark]
Explain why the pressure in a car tyre increases when the car is driven on a hot day.
✓ When temperature increases, gas particles move faster (increased kinetic energy). [1 mark]
✓ Faster-moving particles collide with the tyre walls more frequently and with greater force, increasing the pressure. [1 mark]
Draw and label a reaction pathway diagram for an exothermic reaction showing: reactants, products, activation energy (Ea), and enthalpy change (ΔH).
✓ Diagram shows energy (y-axis) vs reaction progress (x-axis). [1 mark for correct axes]
✓ Products positioned lower than reactants (showing ΔH is negative/exothermic). [1 mark]
✓ Peak between reactants and products labeled Ea; vertical arrow from reactants to peak showing activation energy. ΔH shown as vertical distance from reactants to products. [1 mark]
Photosynthesis is endothermic. Explain this in terms of bond breaking and bond making.
✓ In photosynthesis, bonds in CO₂ and H₂O are broken (requires energy). [1 mark]
✓ New bonds are formed in glucose (C₆H₁₂O₆). [1 mark]
✓ The energy required to break bonds in reactants is greater than the energy released in forming bonds in products, making the overall reaction endothermic (ΔH positive). [1 mark]
Mistake 1: Confusing boiling and evaporation. Remember: Boiling happens throughout the liquid at a fixed temperature; evaporation happens at the surface at any temperature below boiling point.
Mistake 2: Saying temperature increases during a change of state. Remember: Temperature is CONSTANT during melting, boiling, freezing, and condensation. The energy goes into separating or bringing together particles, not into speeding them up.
Mistake 3: Thinking particles stop moving in solids. Remember: Particles in solids vibrate constantly, just in fixed positions. They don't stop moving until absolute zero (−273°C).
Mistake 4: Forgetting that diffusion is random motion causing eventual even distribution. Remember: Particles diffuse because they move randomly in all directions; by chance, more move away from concentrated areas than towards them.
Mistake 5: Confusing "higher temperature" with "all particles moving faster." Remember: Higher temperature means particles move FASTER ON AVERAGE, but some are always moving slower and some faster.
Mistake 6: Not explaining why lighter gases diffuse faster. Remember: At the same temperature, lighter molecules have the same average kinetic energy but higher average speed, so they diffuse faster.
Always define using the key word. For example, "Melting is the process by which a solid changes into a liquid when heated." Use "process," "state," "change," or similar key terms. Examiners look for those specific words.
Always use kinetic particle theory: mention particles, their arrangement, their motion, and the energy involved. Three sentences is usually enough: (1) what the particles do, (2) why they do it (energy/temperature/forces), (3) the resulting property or phenomenon.
In questions with diagrams, particles in solids should be close and ordered; in liquids, close but jumbled; in gases, far apart and spread out. Labels are critical—show particle spacing differences clearly.
Always link molecular mass to speed: lighter molecules move faster at the same temperature (same average KE, but higher speed because KE = 1/2 m v²). This is a high-frequency exam point.
Tick off each objective as you master it. These are the exact learning objectives from the Cambridge 0620 syllabus (2026-2028).
Atoms are the basic building blocks of all matter. They're incredibly small—about 100 picometres (pm) across—yet they have internal structure: a central nucleus surrounded by electrons. The nucleus is made of protons and neutrons, which together make up almost all the atom's mass. Electrons are much lighter and orbit (or occupy) regions called shells around the nucleus.
This is the Rutherford-Bohr model of the atom, which is what Cambridge expects for IGCSE. It's not the full quantum mechanical picture, but it explains nearly everything in the 0620 syllabus.
All atoms are made of three types of particles:
| Particle | Location | Relative Charge | Relative Mass |
|---|---|---|---|
| Proton | Nucleus | +1 | 1 |
| Neutron | Nucleus | 0 | 1 |
| Electron | Shells (around nucleus) | −1 | Negligible (1/1836) |
Key facts:
• The nucleus is tiny (about 100,000 times smaller than the atom) but contains most of the atom's mass.
• Protons and neutrons have almost equal mass (1 atomic mass unit = 1 u), which is roughly 2000 times heavier than an electron.
• An atom is electrically neutral because the number of protons equals the number of electrons (positive charge = negative charge).
Atomic number (proton number): The number of protons in an atom. This is unique to each element—hydrogen has 1 proton, helium has 2, lithium has 3, and so on up to oganesson with 118 protons. Atomic number determines what element it is. Atomic number = number of protons = number of electrons (in a neutral atom).
Mass number (nucleon number): The total number of protons + neutrons in an atom. It's called the nucleon number because protons and neutrons are the nucleons (particles in the nucleus).
Calculating neutrons: Number of neutrons = Mass number − Atomic number
Notation: We write it as Mass NumberSymbol, for example:
• 16O = oxygen with mass number 16. Since oxygen's atomic number is 8, it has 8 protons and 8 neutrons (16 − 8 = 8).
• 12C = carbon with mass number 12. Carbon's atomic number is 6, so it has 6 protons and 6 neutrons.
• 23Na = sodium with mass number 23. Sodium's atomic number is 11, so it has 11 protons and 12 neutrons.
Electrons arrange themselves in shells (also called energy levels or orbitals) around the nucleus. The first shell holds a maximum of 2 electrons, the second shell holds maximum 8, and the third shell holds maximum 8 (for atoms up to atomic number 20, which is calcium).
We write electronic configuration as a sequence: 2,8,3 means 2 electrons in the first shell, 8 in the second, 3 in the third.
Building electronic configurations (filling order):
1. Fill the first shell: 1–2 electrons (H, He)
2. Fill the second shell: 3–10 electrons (Li, Be, B, C, N, O, F, Ne)
3. Fill the third shell: 11–20 electrons (Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca)
Examples:
• Hydrogen (H): 1 electron → 1
• Carbon (C, Z=6): 6 electrons → 2,4 (first shell 2, second shell 4)
• Oxygen (O, Z=8): 8 electrons → 2,6
• Sodium (Na, Z=11): 11 electrons → 2,8,1
• Sulphur (S, Z=16): 16 electrons → 2,8,6
• Calcium (Ca, Z=20): 20 electrons → 2,8,8
The periodic table is arranged so that atoms with the same number of outer electrons are in the same group (vertical column). This is why they have similar chemical properties.
Group number = number of electrons in the outer shell (for main group elements). Group 1 atoms (alkali metals like Na) have 1 outer electron. Group 2 atoms (alkaline earth metals like Ca) have 2 outer electrons. Group 15 atoms (like N, P) have 5 outer electrons. Group 17 atoms (halogens like Cl) have 7 outer electrons.
Group 8 (or Group 0 in old notation) = the noble gases (He, Ne, Ar) have full outer shells. This is why they're so unreactive—their outer shell is complete, so they don't need to gain, lose, or share electrons.
Period number = number of shells. Period 1 has 1 shell (H, He). Period 2 has 2 shells (Li, Be, B, C, N, O, F, Ne). Period 3 has 3 shells (Na, Mg, Al, Si, P, S, Cl, Ar).
Element: A substance made of only one type of atom. All atoms have the same atomic number. Examples: pure oxygen gas (O₂), pure iron (Fe), pure gold (Au). Elements appear on the periodic table.
Compound: A substance made of two or more elements chemically bonded together in fixed ratios. Examples: water (H₂O = 2 hydrogen atoms + 1 oxygen atom), carbon dioxide (CO₂ = 1 carbon + 2 oxygen atoms), sodium chloride (NaCl = 1 sodium + 1 chlorine atom). Compounds have completely different properties from their elements. For example, hydrogen (flammable gas) + oxygen (supports combustion) = water (extinguishes fire). Compounds are pure substances.
Mixture: Two or more substances physically combined but NOT chemically bonded. Examples: salt water (salt dissolved in water), brass (copper and zinc mixed), air (nitrogen, oxygen, argon, and other gases). Mixtures have variable composition—you can add more salt to salt water and it's still salt water. Mixtures can be separated by physical methods (evaporation, filtering, chromatography) because no chemical bonds need to be broken.
Isotopes are atoms of the SAME ELEMENT with the SAME number of protons but DIFFERENT numbers of neutrons. Since they have the same atomic number, they're the same element. Since they have different mass numbers, they have different masses.
Examples:
• Chlorine has two main isotopes: 35Cl (17 protons, 18 neutrons) and 37Cl (17 protons, 20 neutrons).
• Carbon has 12C and 13C and the radioactive 14C.
• Hydrogen has 1H (protium), 2H or D (deuterium), and 3H or T (tritium).
Why chemical properties are the same: Chemical reactions involve electrons, not the nucleus. Isotopes have the same number of electrons (same atomic number), so they behave the same way chemically. 35Cl and 37Cl form the same compounds (both form HCl, both form NaCl) and react the same way.
Why physical properties differ slightly: Mass affects some physical properties. 35Cl and 37Cl have slightly different boiling points, melting points, and densities because they have different masses. But the difference is usually small.
An ion is an atom (or group of atoms) that has gained or lost electrons.
If an atom loses electrons, it becomes positively charged (cation). For example, Na loses 1 electron to become Na⁺ (sodium ion). The Na atom originally had 11 electrons (same as protons), but after losing 1, it has only 10 electrons while still having 11 protons, so the net charge is +1.
If an atom gains electrons, it becomes negatively charged (anion). For example, Cl gains 1 electron to become Cl⁻ (chloride ion). The Cl atom originally had 17 electrons and 17 protons (neutral). After gaining 1, it has 18 electrons and 17 protons, so the net charge is −1.
Notation: We write Na⁺ and Cl⁻. The superscript shows the charge. For ions with more than one charge, we write 2+ or 2−, like O²⁻ (oxygen ion with 2 negative charges).
Electronic configuration of ions: It's the same as the atom, but with fewer or more electrons. For Na⁺: originally 11 electrons (2,8,1), but loses 1, so now 10 electrons (2,8). For Cl⁻: originally 17 electrons (2,8,7), but gains 1, so now 18 electrons (2,8,8).
Since isotopes have different masses, the atomic mass on the periodic table is an average. It's called the relative atomic mass (Ar).
If chlorine has two isotopes with the abundance:
• 35Cl: 75% abundance, mass = 35
• 37Cl: 25% abundance, mass = 37
Then: Ar = (35 × 75 + 37 × 25) / 100 = (2625 + 925) / 100 = 3550 / 100 = 35.5
This is why the periodic table lists chlorine as 35.5, not 35 or 37.
Formula: Ar = Σ(isotope mass × % abundance) / 100
These definitions use the exact wording expected in IGCSE mark schemes. Click each term to reveal.
An atom has the symbol 19F. State the number of protons and the number of neutrons in an atom of fluorine.
✓ Number of protons = 9 [1 mark — this is the atomic number]
✓ Number of neutrons = 19 − 9 = 10 [1 mark for calculation]
Note: Fluorine always has 9 protons (that's what makes it fluorine). The mass number 19 = 9 + neutrons, so neutrons = 10.
Write the electronic configuration of an atom of nitrogen (N, atomic number 7).
✓ Atomic number 7 means nitrogen has 7 electrons [1 mark for understanding]
First shell (max 2): 2 electrons
Second shell (remaining): 7 − 2 = 5 electrons
✓ Electronic configuration = 2,5 [1 mark for correct answer]
Explanation: We fill the inner shells first. The first shell holds max 2, second shell holds max 8. Nitrogen has only 7 electrons total, so first shell gets 2, and the second shell gets the remaining 5.
Chlorine has two isotopes: 35Cl (75% abundance) and 37Cl (25% abundance). Calculate the relative atomic mass of chlorine.
✓ Ar = (35 × 75 + 37 × 25) / 100 [1 mark for correct formula and substitution]
Ar = (2625 + 925) / 100
Ar = 3550 / 100
✓ Ar = 35.5 [1 mark for correct calculation]
✓ This matches the value on the periodic table for chlorine. [1 mark for recognition]
Explain why isotopes of an element have different physical properties but the same chemical properties.
✓ Chemical properties depend on the number of electrons (electronic configuration), which is the same for all isotopes of an element because they have the same atomic number. [1 mark]
✓ Physical properties like density and boiling point depend on mass, which differs between isotopes. [1 mark]
Complete the table below for the ion Na⁺:
Number of protons:___ | Number of electrons:___ | Number of neutrons:___
(Relative atomic mass of sodium = 23, atomic number = 11)
✓ Number of protons = 11 [1 mark — atomic number is always number of protons, unchanged by ionization]
✓ Number of electrons = 10 [1 mark — Na⁺ has lost 1 electron, so 11 − 1 = 10]
✓ Number of neutrons = 23 − 11 = 12 [1 mark — mass number minus atomic number]
Note: Ionization doesn't change the nucleus, only the electrons.
Mistake 1: Confusing atomic number with mass number. Remember: Atomic number is on the BOTTOM LEFT (tells you the element), mass number is on the TOP LEFT (total protons + neutrons).
Mistake 2: Forgetting that in a neutral atom, electrons = protons. Don't assume an atom is charged unless told otherwise.
Mistake 3: Writing electronic configuration wrongly. Remember: Fill inner shells first (1st shell max 2, 2nd shell max 8, 3rd shell max 8 for first 20 elements).
Mistake 4: Confusing "number of neutrons" with "mass number." Remember: Neutrons = mass number − atomic number, NOT the mass number itself.
Mistake 5: When calculating Ar from isotopes, forgetting to divide by 100 (or not using percentages correctly). The formula requires: (mass₁ × %₁ + mass₂ × %₂) / 100.
Mistake 6: Saying isotopes have different numbers of protons. Remember: Isotopes have the SAME number of protons (same element) but DIFFERENT numbers of neutrons.
Mistake 7: Forgetting the charge notation on ions. Write Na⁺ and Cl⁻, not just Na and Cl, when the question asks about ions.
Always draw or describe: nucleus (with protons and neutrons), surrounded by electrons in shells. Label each part. This shows the examiner you understand the model. Use the diagram to help answer numerical questions.
Write out the shells step-by-step. First shell gets max 2, second shell gets max 8, remainder in third shell. Write as numbers separated by commas (e.g., 2,8,6). No spaces or dashes—examiners look for exact format.
Always show the formula: Ar = (mass₁ × abundance₁ + mass₂ × abundance₂) / 100. Show substitution of numbers. Calculate step-by-step. Round to 1 decimal place unless told otherwise. Check your answer against the periodic table—it should match (or be very close).
Protons never change (they define the element). Only electrons change. If told an atom gains or loses electrons, calculate the new electron number and write the symbol with the charge (e.g., Na⁺, O²⁻). Electronic configuration changes, but the number of protons/neutrons stays fixed.
Tick off each objective as you master it. These are the exact learning objectives from the Cambridge 0620 syllabus (2026-2028).
Bonding is the heart of chemistry. It explains why NaCl is a brittle solid that melts at 801°C, while Cl₂ is a yellow-green gas that liquefies at −34°C. Both are made of non-metals, but their bonding is completely different. Understanding bonding is not just about knowing facts — it's about predicting properties from structure. This is tested heavily in IGCSE: you may be given a structure and asked to predict melting point, conductivity, or solubility.
Before any bond forms, atoms must gain or lose electrons to become ions.
Cations are positively charged ions formed when atoms lose electrons. Group I metals (Li, Na, K) lose 1 electron to form 1+ cations. Group II metals (Mg, Ca) lose 2 electrons to form 2+ cations. Aluminium (Al) loses 3 electrons to form Al³⁺.
Anions are negatively charged ions formed when atoms gain electrons. Non-metals gain electrons:
• Group VII (halogens: Cl, Br, I) gain 1 electron to form 1− anions (Cl⁻, Br⁻, I⁻)
• Group VI (O, S) gain 2 electrons to form 2− anions (O²⁻, S²⁻)
• Group V (N, P) gain 3 electrons to form 3− anions (N³⁻, P³⁻)
Why do they do this? All atoms "want" to achieve the electron configuration of a noble gas (8 electrons in the outer shell, except H which wants 2). Losing or gaining electrons to reach this stable configuration is the driving force behind bonding.
Definition: An ionic bond is a strong electrostatic attraction between oppositely charged ions.
How it forms: One atom (a metal) transfers one or more electrons to another atom (a non-metal), creating ions that attract each other.
Example: Sodium Chloride (NaCl)
Sodium has 1 electron in its outer shell. Chlorine has 7 electrons in its outer shell.
Other ionic compounds:
• MgO: Mg (2 outer) transfers 2 electrons to O (6 outer) → Mg²⁺ + O²⁻
• CaF₂: Ca (2 outer) transfers 2 electrons, and two F atoms each gain 1 → Ca²⁺ + 2F⁻
• Al₂O₃: Two Al atoms (3 outer each) transfer 6 electrons to three O atoms (6 outer each) → 2Al³⁺ + 3O²⁻
Ionic compounds form a giant lattice structure — a repeating 3D arrangement of alternating cations and anions, extending throughout the entire crystal.
Why high melting and boiling points?
The electrostatic forces between ions are very strong and work in all directions (each ion attracts multiple ions around it). Breaking these bonds requires enormous energy. Ionic compounds like NaCl melt at 801°C and boil at 1413°C.
Why do they conduct electricity when molten or aqueous, but not when solid?
In a solid, ions are locked in fixed positions — they can't move, so no current flows. When heated above the melting point (molten) or dissolved in water (aqueous), ions become mobile and can carry electrical charge. This is the key difference: ionic compounds only conduct when the ions are free to move.
Why are they often soluble in water?
Water molecules are polar (oxygen pulls electrons toward itself). Water's negative end (oxygen) attracts the cations, and water's positive end (hydrogen) attracts the anions, pulling apart the ionic lattice.
Definition: A covalent bond is formed when a pair of electrons is shared between two atoms, allowing each atom to achieve a noble gas configuration.
Unlike ionic bonding (electron transfer), covalent bonding involves sharing. Neither atom completely loses or gains electrons; they pool their outer electrons.
Example: Hydrogen Molecule (H₂)
Each hydrogen atom has 1 electron. When two H atoms approach, their electrons overlap (share). Both atoms now "feel" 2 electrons in their sphere, achieving the stable He configuration.
Other common covalent molecules:
• Cl₂: Each Cl has 7 outer electrons. Sharing 1 pair gives each Cl 8 electrons (Ar configuration)
• H₂O: Oxygen (6 outer) bonds with two hydrogens. Each H gets 2, O gets 8
• CH₄: Carbon (4 outer) forms 4 bonds with 4 hydrogens. C gets 8, each H gets 2
• NH₃: Nitrogen (5 outer) bonds with 3 hydrogens, leaving 1 lone pair
• HCl: H and Cl each share a pair, both achieving stable configurations
Multiple bonds:
Some atoms need to share more than one pair:
• O₂: A double bond (2 shared pairs) between the oxygens
• N₂: A triple bond (3 shared pairs) — very strong, hard to break
• C₂H₄: Ethene has a double bond between the carbons
Molecules like H₂, Cl₂, H₂O, CO₂ are held together by strong covalent bonds within the molecule, but the molecules themselves are held together by weak intermolecular forces (Van der Waals forces, hydrogen bonds).
Why low melting and boiling points?
When you heat a molecular substance like ice, you don't break the covalent bonds (O−H bonds in water). Instead, you break the intermolecular forces between separate H₂O molecules. These forces are weak, so little energy is needed. Ice melts at 0°C, water boils at 100°C — much lower than ionic compounds.
Why poor electrical conductivity?
Molecular substances have no ions (or very, very few). Electricity is carried by moving charges (ions or free electrons). Since they have neither, they don't conduct. Even when molten, most simple molecules don't conduct because they don't split into ions.
Intermolecular forces explained:
Molecules are electrically neutral overall, but the electrons within them are unevenly distributed. A slightly positive end of one molecule attracts the slightly negative end of another — these weak attractions are intermolecular forces. They're much weaker than covalent or ionic bonds.
Some elements form giant covalent structures where every atom is bonded to every neighbouring atom by strong covalent bonds throughout the entire crystal. There are no molecules — just one enormous network.
Diamond:
Each carbon atom forms 4 covalent bonds to 4 other carbon atoms in a tetrahedral arrangement. This continues in all directions, creating an extremely rigid, hard structure. Diamond is one of the hardest known substances — used for cutting, grinding, and polishing. It melts only at 3823°C.
Graphite:
Carbon atoms form a layered structure. Within each layer, each carbon bonds to 3 others in a hexagonal pattern. The bonds within layers are strong covalent bonds, but between layers, only weak intermolecular forces hold the layers together. This is why graphite:
• Is slippery (layers slide easily past each other) — used as a lubricant
• Conducts electricity (delocalized electrons in the layers can move freely) — used in electrodes
• Has a lower melting point than diamond (though still very high, ~3650°C) because weak forces between layers break before the covalent bonds do
Silicon(IV) Oxide (SiO₂):
Similar to diamond, but with alternating Si and O atoms. Each Si bonds to 4 oxygen atoms, and each O bonds to 2 Si atoms. It's a giant 3D covalent network. Sand, quartz, and glass are all forms of SiO₂. It's hard (used in abrasives) and has a very high melting point (~1710°C).
Diamond vs. SiO₂: Both are giant covalent structures with very high melting points and hardness. The key similarity is the 3D network of strong covalent bonds with no free electrons, which is why neither conducts electricity.
Definition: Metallic bonding consists of metal cations held together by a "sea" of delocalised electrons that move freely throughout the structure.
In a metal, the outermost electrons are not fixed to individual atoms. Instead, they form a mobile "sea" that flows throughout the crystal. Metal cations (the positively charged cores) are held in place by attraction to this electron sea.
Metal properties explained by this model:
• Good electrical conductivity: The free electrons can move easily, carrying electrical current
• Good thermal conductivity: Free electrons transfer heat energy rapidly
• Malleability: Layers of metal atoms can slide past each other (the electron sea maintains bonding even as positions shift)
• Ductility: Metals can be drawn into wires without breaking — the electron sea holds the structure together even under stress
• Shiny appearance: Electrons can absorb and re-emit light easily
• High melting points: Breaking apart a metal requires disrupting the strong attraction between cations and the electron sea
| Property | Ionic | Simple Molecular | Giant Covalent | Metallic |
|---|---|---|---|---|
| Structure | 3D lattice of ions | Discrete molecules | 3D network of atoms | Cations + electron sea |
| Melting point | Very high | Very low | Very high | High |
| Boiling point | Very high | Very low | Very high | High |
| Solid conductivity | No | No | No (except graphite) | Yes |
| Molten conductivity | Yes | No | No | Yes |
| Solubility (water) | Often soluble | Variable | Insoluble | Insoluble |
| Solidity | Brittle | Liquid/gas | Hard, rigid | Malleable, ductile |
These definitions use the exact wording expected in IGCSE mark schemes. Click each term to reveal.
Describe the ionic bonding in sodium chloride (NaCl) using a dot-and-cross diagram. Explain why sodium chloride has a high melting point.
Dot-and-cross diagram:
✓ Sodium loses 1 electron (dot-and-cross diagram showing Na⁺ with 0 outer electrons, Cl⁻ with 8). [1 mark]
✓ Chlorine gains 1 electron from sodium. [1 mark]
Explanation of high melting point:
✓ Ionic compounds have a giant 3D lattice structure where each ion is surrounded by multiple oppositely charged ions. [1 mark]
✓ The strong electrostatic forces between ions extend throughout the entire structure, requiring large amounts of energy to break. [1 mark]
Compare the electrical conductivity of solid NaCl and molten NaCl. Explain your answer in terms of ion movement.
✓ Solid NaCl does not conduct electricity. [1 mark]
✓ Molten NaCl conducts electricity. [1 mark]
✓ In solid NaCl, ions are fixed in the lattice and cannot move. In molten NaCl, ions are free to move and carry electrical charge (current). [1 mark]
Draw dot-and-cross diagrams for H₂ and Cl₂. Explain why these molecules have low melting points.
Dot-and-cross diagrams:
✓ H₂: Two H atoms sharing one electron pair (both with 2 electrons). [1 mark]
✓ Cl₂: Two Cl atoms sharing one electron pair (each with 8 electrons). [1 mark]
Why low melting points:
✓ H₂ and Cl₂ are simple molecular substances with weak intermolecular forces (Van der Waals forces) between molecules. [1 mark]
✓ When melting, these weak intermolecular forces are overcome, not the strong covalent bonds within molecules. Therefore, little energy is needed. [1 mark]
Explain why graphite conducts electricity but diamond does not, although both are forms of carbon.
✓ Both graphite and diamond have giant covalent structures with strong C−C bonds. [1 mark]
✓ Graphite has layers with delocalised electrons within each layer that are free to move and carry electrical charge. [1 mark]
✓ Diamond has all electrons localized in C−C bonds with no free electrons, so it cannot conduct electricity. [1 mark]
Explain why metals such as copper are malleable (can be bent without breaking).
✓ Metals consist of cations held together by a sea of delocalised electrons. [1 mark]
✓ When a metal is bent, layers of cations can slide past each other. [1 mark]
✓ The electron sea continues to hold the cations together even after they move, so the metal doesn't break. [1 mark]
State the number of electrons lost by a sodium atom to form a Na⁺ cation, and explain why sodium is willing to lose this electron.
Which of the following is the best description of an ionic bond?
Magnesium oxide (MgO) is an ionic compound. Draw a dot-and-cross diagram showing the outer electrons in Mg²⁺ and O²⁻.
Why does solid NaCl not conduct electricity, but aqueous NaCl (solution) does?
Draw a dot-and-cross diagram for a water molecule (H₂O), showing the covalent bonds.
Compare the melting points of water (100°C at boiling), methane (−162°C), and sodium chloride (801°C). Explain these differences in terms of bonding and structure.
State two properties of giant covalent structures like diamond. Explain each property in terms of structure.
Explain why graphite is used as a lubricant and as an electrode in electrolysis, linking to its structure.
Describe metallic bonding and explain why metals are both good electrical conductors and malleable.
Explain why SiO₂ has properties similar to diamond (both hard, high melting points, poor conductivity) but is made of different elements.
Don't just list facts. Link structure to property. Always follow this format:
1. Describe the structure (e.g., "giant 3D lattice of ions")
2. Identify what holds it together (e.g., "strong electrostatic forces")
3. Explain the property (e.g., "high melting point because lots of energy is needed to break these strong forces")
Example: If asked why diamond is hard: "Diamond has a giant covalent structure where each carbon atom is bonded to 4 other carbons in a tetrahedral arrangement. These strong C−C covalent bonds extend throughout the entire 3D crystal. To break the substance, you must break these strong bonds, which requires a lot of energy. Therefore, diamond is very hard." ✓
• Clearly show which atoms the electrons come from (usually dots from one atom, crosses from the other)
• For ionic compounds, show the charge (e.g., Na⁺ and Cl⁻)
• For covalent bonds, show the shared pair clearly in the overlap region
• Show all outer shell electrons, including lone pairs
• Practice reading the element's position in the periodic table to determine outer electrons (Group I = 1, Group II = 2, Group VI = 6, Group VII = 7)
Always ask: "What carries the charge?" and "Can it move?"
• Ionic solids: No moving charges (ions locked in place) → No conduction
• Molten/aqueous ionic: Ions can move → Conducts
• Simple molecular: No ions, no free electrons → No conduction
• Metals: Free electrons → Conducts (even in solid)
• Graphite: Free electrons in layers → Conducts
• Diamond: No free electrons → No conduction
Use a table to organize your answer:
Diamond: Giant covalent, all electrons in C−C bonds, no delocalised electrons, very hard, very high mp
Graphite: Giant covalent, delocalised electrons in layers, conducts, slippery, high mp but lower than diamond
Similarity: Both are giant covalent structures made of C
Difference: Arrangement and electron arrangement (delocalised vs. localized)
Question: "Why does water have a low boiling point compared to sodium chloride?"
Don't say: "Because covalent bonds are weaker."
Do say: "Water's molecules are held together by weak intermolecular forces (hydrogen bonds/Van der Waals forces), while NaCl has a giant ionic lattice held together by strong electrostatic forces. Melting/boiling water requires breaking intermolecular forces, not covalent bonds, so less energy is needed."
Tick off each objective as you master it. These are the exact learning objectives from the Cambridge 0620 syllabus (2026-2028).
Chemistry is a language. Formulae and equations are how we communicate exactly what substances are involved, in what proportions, and what conditions they require. Get the formula wrong, and everything that follows — calculations, predictions, understanding — falls apart. IGCSE tests this heavily because it's fundamental to everything you do in chemistry.
Molecular formula: The actual number of atoms of each element in a molecule. Examples: H₂O (water), CO₂ (carbon dioxide), C₆H₁₂O₆ (glucose).
Empirical formula: The simplest whole number ratio of atoms of each element in a compound. For glucose, the empirical formula is CH₂O (the ratio is 1:2:1), but the molecular formula is C₆H₁₂O₆ (which is 6 × CH₂O).
For ionic compounds, use the valency (oxidation state) method. The positive charge of the cation must balance the negative charge of the anion.
Example: Calcium Fluoride
• Ca is a Group II metal → forms Ca²⁺
• F is a Group VII non-metal → forms F⁻
• To balance: 1 Ca²⁺ needs 2 F⁻ ions
• Formula: CaF₂
Example: Aluminum Oxide
• Al is a Group III metal → forms Al³⁺
• O is a Group VI non-metal → forms O²⁻
• To balance: 2 Al³⁺ ions (total +6) with 3 O²⁻ ions (total −6)
• Formula: Al₂O₃
A balanced chemical equation shows the correct proportions of reactants and products. Atoms are neither created nor destroyed (Law of Conservation of Mass).
Process:
1. Write the unbalanced equation with correct formulae
2. Count atoms of each element on both sides
3. Add coefficients (numbers in front) to balance
4. Never change subscripts (the numbers in formulae) — only add coefficients
5. Check: each element has the same count on both sides
Example: Burning magnesium
Unbalanced: Mg + O₂ → MgO
• Left: 1 Mg, 2 O
• Right: 1 Mg, 1 O (not balanced)
Add coefficient 2 to MgO:
2Mg + O₂ → 2MgO
• Left: 2 Mg, 2 O
• Right: 2 Mg, 2 O ✓ Balanced
Example: Combustion of ethane
Unbalanced: C₂H₆ + O₂ → CO₂ + H₂O
• C: 2 on left, 1 on right → need 2 CO₂
• H: 6 on left, 2 on right → need 3 H₂O
• O: need to check with new coefficients
C₂H₆ + O₂ → 2CO₂ + 3H₂O
• O on right: (2×2) + (3×1) = 7. Need 7/2 = 3.5 O₂
2C₂H₆ + 7O₂ → 4CO₂ + 6H₂O (multiply by 2 to avoid fractions)
Always include state symbols in equations:
• (s) = solid
• (l) = liquid
• (g) = gas
• (aq) = aqueous (dissolved in water)
Example: Magnesium burns in oxygen:
2Mg(s) + O₂(g) → 2MgO(s)
In IGCSE, equations without state symbols often lose marks. Always include them.
Relative atomic mass (Ar): The mass of an atom relative to 1/12 of a carbon-12 atom. It's a dimensionless number, but numerically equal to the atomic mass in g/mol. Common values you should know:
• H = 1, C = 12, N = 14, O = 16, Na = 23, Mg = 24, S = 32, Cl = 35.5, K = 39, Ca = 40
Relative molecular mass (Mr): The sum of the relative atomic masses of all atoms in a molecule.
Example: Calculate Mr of H₂SO₄
Mr = (2 × 1) + 32 + (4 × 16) = 2 + 32 + 64 = 98
Definition: A mole is a quantity equal to Avogadro's constant, 6.02 × 10²³ particles (atoms, molecules, electrons, ions, etc.).
Why is this useful? It converts between the particle level (tiny, uncountable) and the lab level (grams, volumes). One mole of any substance has a mass in grams numerically equal to its Mr:
• 1 mole of H₂ (Mr = 2) = 2 g
• 1 mole of O₂ (Mr = 32) = 32 g
• 1 mole of H₂SO₄ (Mr = 98) = 98 g
Key formula: n = m / Mr
where n = number of moles, m = mass in grams, Mr = relative molecular mass
Example: How many moles are in 24 g of carbon? (Ar of C = 12)
n = 24 / 12 = 2 moles
Once you know the mole ratio from a balanced equation, you can calculate how much of each reactant or product is involved.
Example: 2Mg + O₂ → 2MgO
The equation tells us: 2 moles of Mg react with 1 mole of O₂ to produce 2 moles of MgO.
• If you have 3 moles of Mg, you'll need 1.5 moles of O₂
• If 2 moles of O₂ react, then 4 moles of Mg reacted and 4 moles of MgO formed
Reacting mass calculation:
If 4.8 g of Mg reacts with excess O₂, how much MgO is produced?
• First, find moles of Mg: n = 4.8 / 24 = 0.2 moles
• Use equation ratio: 2 Mg → 2 MgO, so 0.2 moles Mg → 0.2 moles MgO
• Find mass of MgO: Mr of MgO = 24 + 16 = 40. Mass = 0.2 × 40 = 8 g
Often one reactant runs out first. The one that runs out first is the limiting reactant — it limits how much product can form.
Example: 2H₂ + O₂ → 2H₂O
If you have 5 g H₂ and 16 g O₂, which runs out first?
• Moles of H₂: 5 / 2 = 2.5 moles
• Moles of O₂: 16 / 32 = 0.5 moles
• Equation ratio: 2 H₂ : 1 O₂. If we use all 0.5 moles O₂, we need 1 mole H₂
• We have 2.5 moles H₂, but only need 1, so O₂ is the limiting reactant
• Water produced: 0.5 × 2 = 1 mole = 18 g (Mr of H₂O = 18)
At room temperature and pressure (rtp): 1 mole of any gas occupies 24 dm³ (or 24 L).
Example: How much space does 2 moles of CO₂ occupy at rtp?
Volume = 2 × 24 = 48 dm³
Reverse calculation: What volume of oxygen gas is needed to react with 1.2 g of carbon?
C + O₂ → CO₂
• Moles of C: 1.2 / 12 = 0.1 moles
• Moles of O₂ needed: 0.1 moles (1:1 ratio)
• Volume of O₂: 0.1 × 24 = 2.4 dm³
Concentration in g/dm³: grams of solute per dm³ of solution
Example: "This solution contains 5 g/dm³ of salt" means 5 g of salt dissolved in enough water to make 1 dm³ total.
Concentration in mol/dm³ (molarity): moles of solute per dm³
Formula: concentration (mol/dm³) = n / V
where n = moles and V = volume in dm³
Converting between them:
If a solution has concentration 2 mol/dm³ of NaCl (Mr = 58.5), what's the concentration in g/dm³?
• In 1 dm³, there are 2 moles of NaCl
• Mass = 2 × 58.5 = 117 g
• Concentration = 117 g/dm³
Titrations measure unknown concentrations by reacting a known volume/concentration of one substance with an unknown quantity of another.
Example: HCl vs. NaOH
25 cm³ (0.025 dm³) of NaOH solution was neutralized by 20 cm³ of 0.5 mol/dm³ HCl.
Find the concentration of NaOH.
• Equation: HCl + NaOH → NaCl + H₂O (1:1 ratio)
• Moles of HCl: 0.5 × 0.020 = 0.01 moles
• Moles of NaOH: 0.01 moles (1:1)
• Concentration of NaOH: 0.01 / 0.025 = 0.4 mol/dm³
In real reactions, you never get the theoretical amount. Percentage yield tells you how efficient a reaction was.
Formula: % yield = (actual mass / theoretical mass) × 100
Example: Reaction between Cu and O₂ produces CuO. Theory predicts 8 g of CuO, but you only get 6.4 g.
% yield = (6.4 / 8) × 100 = 80%
What percentage of a compound's mass is due to a particular element?
Formula: % by mass = (Ar × number of atoms) / Mr × 100
Example: H₂SO₄
• Mr = 98
• % S = (32 / 98) × 100 = 32.7%
• % H = (2 / 98) × 100 = 2.0%
• % O = (64 / 98) × 100 = 65.3%
Real samples are often impure. Percentage purity tells you what fraction is the desired substance.
Formula: % purity = (mass of pure substance / total mass) × 100
Example: A 50 g sample of iron ore contains 40 g of pure iron.
% purity = (40 / 50) × 100 = 80%
These definitions use the exact wording expected in IGCSE mark schemes. Click each term to reveal.
Balance this equation and include state symbols:
C + O₂ → CO₂
✓ Balanced equation: C(s) + O₂(g) → CO₂(g) [1 mark for balance]
✓ State symbols included: (s), (g), (g) [1 mark]
✓ Each element has the same number of atoms on both sides: C: 1 on each; O: 2 on each [1 mark for explanation]
Calculate the relative molecular mass (Mr) of sulfuric acid, H₂SO₄.
(Use: H = 1, S = 32, O = 16)
Mr = (2 × 1) + 32 + (4 × 16)
✓ Mr = 2 + 32 + 64 [1 mark for substitution]
✓ Mr = 98 [1 mark for correct answer]
✓ Units: Mr is relative (no units) [1 mark for understanding]
Calculate how many moles are in 12 g of carbon dioxide (CO₂).
(Use: C = 12, O = 16)
Step 1: Calculate Mr of CO₂
Mr = 12 + (2 × 16) = 12 + 32 = 44
✓ Mr = 44 [1 mark]
Step 2: Use n = m / Mr
n = 12 / 44 = 0.273 moles (or 3/11 moles)
✓ Correct formula and substitution [1 mark]
✓ n = 0.27 moles (to 2 sig figs) [1 mark for correct answer]
✓ Appropriate significant figures [1 mark]
In a reaction, 4.8 g of magnesium burns in oxygen to form magnesium oxide. Calculate the mass of magnesium oxide produced.
(Use: Mg = 24, O = 16)
Step 1: Write and balance the equation
✓ 2Mg + O₂ → 2MgO [1 mark]
Step 2: Calculate moles of Mg
n = 4.8 / 24 = 0.2 moles
✓ Correct calculation [1 mark]
Step 3: Use stoichiometry
From equation: 2 moles Mg → 2 moles MgO
0.2 moles Mg → 0.2 moles MgO
✓ Correct mole ratio [1 mark]
Step 4: Calculate mass of MgO
Mr of MgO = 24 + 16 = 40
Mass = 0.2 × 40 = 8 g
✓ Correct Mr and mass calculation [1 mark]
✓ Final answer: 8 g [1 mark]
Calculate the percentage composition of carbon in CO₂.
(Use: C = 12, O = 16)
Step 1: Calculate Mr of CO₂
Mr = 12 + (2 × 16) = 44
✓ Mr = 44 [1 mark]
Step 2: Apply percentage formula
% C = (Ar × atoms) / Mr × 100 = (12 × 1) / 44 × 100
✓ Correct formula and substitution [1 mark]
% C = 12/44 × 100 = 27.3%
✓ Correct calculation [1 mark]
✓ Final answer: 27.3% (or 27%) [1 mark]
Define the term "relative molecular mass" and state its units.
What is the empirical formula of a compound with molecular formula C₆H₁₂?
Balance this equation and include state symbols: Fe + O₂ → Fe₂O₃
Write the formula of calcium carbonate using the ion charges method. (Ca²⁺ and CO₃²⁻)
Calculate the Mr of sodium nitrate (NaNO₃). (Use: Na = 23, N = 14, O = 16)
Calculate how many moles are in 6.4 g of oxygen gas (O₂). (O = 16)
What volume (in dm³) would 0.5 moles of hydrogen gas occupy at room temperature and pressure?
A solution contains 0.1 mol/dm³ of hydrochloric acid. Calculate the concentration in g/dm³. (H = 1, Cl = 35.5)
A reaction theoretically produces 20 g of product, but actually produces 16 g. Calculate the percentage yield.
In the reaction 2H₂ + O₂ → 2H₂O, what is the limiting reactant if 6 g of H₂ and 20 g of O₂ are mixed? Show your calculations.
1. Start with the most complex compound (usually the one with the most atoms)
2. Balance that first
3. Move to the next element
4. Balance oxygen last (if it appears in multiple compounds)
5. Never use fractions — multiply the entire equation to make all coefficients whole numbers
Example: P + O₂ → P₂O₅
• Start with P₂O₅: need 2 P, so 2P on left
• That needs (2×5)/2 = 5 O₂ on left
• 4P + 5O₂ → 2P₂O₅ ✓
1. Write the balanced equation
2. Calculate Mr (if not given)
3. Convert mass to moles using n = m/Mr
4. Use the equation ratios to find moles of product/other reactant
5. Convert moles back to mass or volume
6. Always show units and significant figures
Always work in moles as an intermediate step — never try to go directly from mass to mass.
1. Calculate moles of EACH reactant
2. Divide each by its coefficient in the equation
3. The smallest answer is the limiting reactant
4. Use only that reactant to calculate the product
Example: 3H₂ + N₂ → 2NH₃. If 6 mol H₂ and 2 mol N₂:
• H₂: 6/3 = 2
• N₂: 2/1 = 2
• Both are equal — neither is limiting. Both react completely.
Always write the balanced equation first to see the mole ratio. Then use:
n₁V₁ = n₂V₂ (for equal ratios)
or multiply through by the ratio if it's not 1:1
Remember: 1 dm³ = 1000 cm³, so convert volumes carefully.
Never confuse "actual" with "theoretical":
• Theoretical: calculated from the stoichiometry, assuming reaction goes to completion
• Actual: measured in the lab (often less because reactions don't always go to completion)
% yield is always ≤ 100%. If you get >100%, you've made an error.
Tick off each objective as you master it. These are the exact learning objectives from the Cambridge 0620 Chemistry syllabus (2026-2028).
Every chemical reaction involves energy. Some reactions release energy (warming the surroundings), while others absorb energy (cooling the surroundings). This is the core idea of chemical energetics.
Exothermic reactions release thermal energy. Think of combustion: when you burn petrol or wood, heat is released. The surroundings warm up. Everyday exothermic reactions include rusting of iron, neutralisation of acids with bases, and explosions.
Endothermic reactions absorb thermal energy from the surroundings. Examples: melting ice (requires heat), evaporation, photosynthesis, and instant cold packs used in first aid. The surroundings cool down.
A reaction pathway diagram shows the energy of reactants, products, and the transition state. The vertical axis is energy (or enthalpy), and the horizontal axis represents the reaction progress.
For an exothermic reaction: The products are at a lower energy level than the reactants. The reaction pathway goes down. ΔH is negative (energy is released).
For an endothermic reaction: The products are at a higher energy level than the reactants. The reaction pathway goes up. ΔH is positive (energy is absorbed).
The activation energy (Ea) is the energy "hump" that must be overcome for the reaction to proceed. It's the minimum collision energy needed. Even if a reaction is exothermic overall, it still needs activation energy to get started.
Enthalpy (H) is the total heat content of a substance. Enthalpy change (ΔH) is the difference between the enthalpy of products and enthalpy of reactants.
ΔH = H(products) − H(reactants)
If ΔH is negative (exothermic), energy has been released — the products have less energy stored than the reactants. If ΔH is positive (endothermic), energy has been absorbed — the products have more energy than the reactants.
Every chemical bond contains energy. Breaking a bond requires energy (endothermic process). Making a bond releases energy (exothermic process).
The net energy change in a reaction depends on which wins: the energy needed to break bonds in reactants, or the energy released when making new bonds in products.
ΔH = Energy to break bonds − Energy to make bonds
If more energy is released in making bonds than required to break bonds, the reaction is exothermic (ΔH negative). If more energy is needed to break bonds than released in making bonds, the reaction is endothermic (ΔH positive).
Combustion: Burning fuel (exothermic). The fuel molecules break apart (requires energy), but the products (CO₂ and H₂O) form with strong bonds, releasing much more energy than was needed. Overall: highly exothermic.
Neutralisation: Acid + base → salt + water (exothermic). Even though it seems "small", neutralisation always releases a measurable amount of heat (about 57 kJ per mole of water formed in dilute solutions).
Evaporation: Liquid → gas (endothermic). Breaking intermolecular forces requires energy. This is why sweating cools you down — the energy to evaporate sweat comes from your skin.
Calculate the enthalpy change for the reaction: H₂ + Cl₂ → 2HCl
Bond energies: H–H = 436 kJ/mol, Cl–Cl = 244 kJ/mol, H–Cl = 432 kJ/mol
Step 1: Identify bonds broken (energy required)
H–H: 436 kJ/mol
Cl–Cl: 244 kJ/mol
Total energy to break bonds = 436 + 244 = 680 kJ/mol [1 mark]
Step 2: Identify bonds made (energy released)
2 × H–Cl = 2 × 432 = 864 kJ/mol [1 mark]
Step 3: Calculate ΔH
ΔH = 680 − 864 = −184 kJ/mol [1 mark]
The reaction is exothermic (ΔH negative). More energy is released making HCl bonds than required to break H₂ and Cl₂ bonds.
Explain why combustion reactions are always exothermic, even though breaking bonds requires energy.
✓ Breaking bonds in fuel molecules requires energy (endothermic). [1 mark]
✓ Making bonds in products (CO₂ and H₂O) releases more energy than required to break reactant bonds. [1 mark]
The net result is that energy is released, making the reaction exothermic (ΔH negative).
A student plots a reaction pathway diagram for an exothermic reaction with a catalyst.
(a) How does the diagram show the reaction is exothermic? [1]
(b) How does the catalyst appear on the diagram? [1]
(c) What does the activation energy represent? [1]
(a) The products are at a lower energy than the reactants / the pathway goes downward. [1 mark]
(b) The catalyst lowers the activation energy / the peak of the pathway is lower when catalyst is present. [1 mark]
(c) The activation energy is the minimum energy required for particles to collide and react / the energy difference between reactants and the transition state. [1 mark]
Describe an exothermic reaction and an endothermic reaction, using temperature change as evidence.
✓ Exothermic reaction releases energy to the surroundings, causing the temperature to increase. Example: combustion of methane or neutralisation of acid with alkali. [1 mark]
✓ Endothermic reaction absorbs energy from the surroundings, causing the temperature to decrease. Example: evaporation, dissolving ammonium nitrate in water, or thermal decomposition. [1 mark]
Which of the following is an exothermic reaction?
Calculate ΔH for: N₂ + 3H₂ → 2NH₃. Bond energies: N≡N = 945, H–H = 436, N–H = 392 kJ/mol
What is meant by activation energy?
Which sign indicates an endothermic reaction?
Explain why a catalyst increases the rate of reaction but does not affect ΔH.
In a reaction pathway diagram for an endothermic reaction, where are the products positioned relative to the reactants?
Calculate ΔH for: CH₄ + 2O₂ → CO₂ + 2H₂O. Bond energies: C–H = 413, O=O = 498, C=O = 799, O–H = 467 kJ/mol
Explain why a cold pack containing ammonium nitrate gets cold when activated.
Which process involves breaking chemical bonds and is endothermic?
Compare the effect of a catalyst on activation energy and on enthalpy change. Which one does it affect?
When calculating ΔH from bond energies, explicitly write out which bonds you're counting. This earns marks even if your final answer is wrong.
If asked "explain why this is exothermic", sketch a diagram showing products lower than reactants. A diagram is worth explanation time — examiners award marks for correct diagrams.
If the reaction releases more energy than it uses, the overall change is negative. This is counterintuitive to some — but the mnemonic helps: exothermic = "exit" = "−" (negative).
Exothermic = surroundings get hotter (temperature increases). Endothermic = surroundings get cooler (temperature decreases). Use these exact phrases in exam answers.
Tick off each objective as you master it.
At the most fundamental level: acids donate protons (H⁺), bases accept protons. This is the Brønsted-Lowry definition.
When hydrochloric acid (HCl) dissolves in water, it completely breaks apart: HCl → H⁺ + Cl⁻. The H⁺ ion (a proton) is what makes it acidic. When sodium hydroxide (NaOH) dissolves, it releases OH⁻ ions, which accept protons: OH⁻ + H⁺ → H₂O. The OH⁻ is what makes it basic (alkaline).
Strong acids completely dissociate in water. Examples: HCl, HBr, HI, HNO₃, H₂SO₄, HClO₄.
When you dissolve HCl in water, 100% of the HCl molecules break into H⁺ and Cl⁻. There are no HCl molecules left in solution.
Weak acids partially dissociate. Example: acetic acid (CH₃COOH, found in vinegar).
When you dissolve CH₃COOH in water, only a tiny fraction (maybe 1%) breaks into H⁺ and CH₃COO⁻. Most of it remains as intact CH₃COOH molecules. This is shown as a reversible reaction with ⇌:
CH₃COOH ⇌ H⁺ + CH₃COO⁻
pH measures the concentration of H⁺ ions. The scale runs from 0 (very acidic) to 14 (very alkaline), with 7 being neutral (pure water).
| pH | Description |
|---|---|
| 0–3 | Strong acid |
| 4–6 | Weak acid |
| 7 | Neutral |
| 8–10 | Weak alkali |
| 11–14 | Strong alkali |
A universal indicator changes colour gradually across the pH range, allowing you to estimate pH. Litmus is simpler: red in acid, blue in alkali.
With metals (less reactive than hydrogen):
Acid + metal → salt + hydrogen gas
2HCl + Mg → MgCl₂ + H₂↑
With bases (metal oxides/hydroxides):
Acid + base → salt + water (neutralisation)
2HCl + CuO → CuCl₂ + H₂O
With carbonates:
Acid + carbonate → salt + water + carbon dioxide
2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂↑
With acids:
This is neutralisation. The ionic equation is always:
H⁺(aq) + OH⁻(aq) → H₂O(l)
With ammonium salts (containing NH₄⁺):
Warm alkali + ammonium salt → ammonia gas + water + salt
2NaOH + 2NH₄Cl → 2NH₃↑ + 2NaCl + 2H₂O
Or in ionic form: OH⁻(aq) + NH₄⁺(aq) → NH₃(g) + H₂O(l)
| Indicator | In acid | In alkali |
|---|---|---|
| Litmus | Red | Blue |
| Methyl orange | Red | Yellow |
| Thymolphthalein | Colourless | Blue |
Write the balanced equation: hydrochloric acid + calcium carbonate
✓ 2HCl + CaCO₃ → CaCl₂ + H₂O + CO₂ [3 marks]
Explain the difference between a strong acid and a weak acid.
✓ Strong acid completely dissociates; weak acid only partially dissociates (reversible). [2 marks]
Write the balanced equation and ionic equation for: nitric acid + copper oxide
✓ Balanced equation: 2HNO₃ + CuO → Cu(NO₃)₂ + H₂O [1 mark]
✓ Ionic equation: H⁺(aq) + OH⁻(aq) → H₂O(l) [2 marks for correct simplified ionic equation]
Warm sodium hydroxide solution is added to a solution of ammonium chloride. Identify the gas produced and state how it can be detected.
✓ Gas produced is ammonia (NH₃). [1 mark]
✓ It can be detected by its pungent smell, or by holding damp red litmus paper above the solution which turns blue. [1 mark]
A student adds universal indicator to three solutions (A, B, C) with pH values 2, 7, and 12 respectively. Describe the colour change observed.
✓ Solution A (pH 2): Red/pink colour - strongly acidic. [1 mark]
✓ Solution B (pH 7): Green colour - neutral. [1 mark]
✓ Solution C (pH 12): Blue/purple colour - strongly alkaline. [1 mark]
What is an alkali?
What colour does litmus paper turn in the presence of an acid?
Write the balanced equation for: Sulfuric acid + Magnesium metal
Which of the following is a proton donor?
What products are formed when hydrochloric acid reacts with calcium carbonate?
Which colour does methyl orange turn in an alkaline solution?
Explain the difference between a strong acid and a weak acid in terms of dissociation.
What is the ionic equation for all acid-base neutralisation reactions?
Write the balanced equation: Sodium hydroxide + Hydrochloric acid
Describe what happens when a base reacts with an ammonium salt and explain the gas produced.
Method 1: Titration – Acid + alkali. Add acid dropwise until indicator changes colour. Evaporate filtrate.
Method 2: Excess Metal – Acid + excess Zn powder. Filter out excess metal. Evaporate.
Method 3: Excess Insoluble Base – Acid + excess CuO powder. Filter out excess base. Evaporate.
Method 4: Excess Insoluble Carbonate – Acid + excess CaCO₃ powder. Filter out excess. Evaporate.
| Compound | Solubility | Exceptions |
|---|---|---|
| All Na⁺, K⁺, NH₄⁺ salts | Soluble | None |
| All nitrates | Soluble | None |
| Chlorides | Soluble | AgCl, PbCl₂ |
| Sulfates | Soluble | BaSO₄, CaSO₄, PbSO₄ |
| Carbonates | Insoluble | Na₂CO₃, K₂CO₃, (NH₄)₂CO₃ |
| Hydroxides | Insoluble | NaOH, KOH, Ca(OH)₂ |
Hydrated salts contain water bonded in the crystal (e.g., CuSO₄·5H₂O, blue).
Anhydrous salts contain no water (e.g., CuSO₄, white).
Water of crystallisation = water molecules bonded within the crystal structure.
Mix two soluble salts whose ions form an insoluble product. Filter, wash with distilled water, dry in oven.
Example: AgNO₃ + NaCl → AgCl↓ + NaNO₃
Describe preparation of pure NaCl by titrating HCl with NaOH.
✓ Measure NaOH into conical flask with indicator. [1 mark]
✓ Add HCl dropwise from burette until indicator changes. [1 mark]
✓ Evaporate solution to obtain NaCl crystals. [1 mark]
Explain why copper sulfate crystals are blue but anhydrous copper sulfate is white.
✓ Blue crystals contain water of crystallisation (CuSO₄·5H₂O). Anhydrous CuSO₄ has no water. [2 marks]
Is BaSO₄ soluble or insoluble?
Which salt preparation method would you use to prepare an insoluble salt?
Describe the steps for preparing crystals of sodium chloride from sodium hydroxide and hydrochloric acid.
What technique is used to separate insoluble crystals from a liquid?
State what is meant by the water of crystallisation in a salt crystal.
What is the difference between a hydrated and an anhydrous salt?
Why should crystals be washed with a small amount of cold distilled water during salt preparation?
Which solubility rule is correct?
When preparing a salt by titration, how is the volume of acid required for neutralisation determined?
Explain why a salt prepared from an excess of acid would contain impurities and how this is avoided.
Periods: Horizontal rows. Increasing proton number, new electron shell is filled.
Groups: Vertical columns. Same number of valence electrons = similar properties.
Left to right: metallic → non-metallic character. Left = metal, middle = transition, right = non-metal, far right = noble gas (unreactive).
Properties: Soft, low density, shiny (tarnish quickly), 1 valence electron, form 1+ ions.
Trends down group: mp DECREASES, density varies, reactivity INCREASES.
Reaction: 2Na + 2H₂O → 2NaOH + H₂↑
Properties: Diatomic molecules (F₂, Cl₂, Br₂, I₂), 7 valence electrons, form 1− ions.
Appearance: Cl₂ = pale yellow-GREEN gas, Br₂ = RED-BROWN liquid, I₂ = GREY-BLACK solid.
Trends down group: density ↑, reactivity ↓, boiling point ↑
Displacement: Cl₂ + 2NaBr → 2NaCl + Br₂ (more reactive displaces less reactive)
Characteristics: High density, high mp, form coloured compounds, catalysts, variable oxidation numbers (e.g., Fe²⁺/Fe³⁺).
Characteristics: Monatomic, full valence shell (8 electrons, except He = 2), completely unreactive under normal conditions. Explain: full shell = no tendency to gain/lose electrons.
Explain why bromine is less reactive than chlorine.
✓ Both in Group 17; Cl above Br in group. [1 mark]
✓ Down group, atomic size increases, valence electrons further from nucleus, weaker attraction, harder to gain electron. Cl (smaller) is more reactive. [1 mark]
Chlorine can displace bromine from potassium bromide solution. Write the balanced equation and explain why this happens.
✓ Balanced equation: Cl₂ + 2KBr → 2KCl + Br₂ [1 mark]
✓ Chlorine is more reactive than bromine (higher up Group VII) [1 mark].
✓ Chlorine readily gains electrons (is a strong oxidising agent) and displaces less reactive halogens [1 mark].
Explain why sodium has a lower melting point (97.8°C) compared to lithium (180°C) in Group I.
✓ The metallic bonding in the alkali metals becomes weaker down the group [1 mark].
✓ Atoms are larger; valence electrons are further from the nucleus and less strongly attracted, resulting in weaker metallic bonding and lower melting point [1 mark].
Write the equation: potassium + water
✓ 2K + 2H₂O → 2KOH + H₂↑ [2 marks]
Chlorine displaces bromine from KBr solution: (a) Write equation [1] (b) Explain why [1] (c) State colour of Br₂ [1]
(a) Cl₂ + 2KBr → 2KCl + Br₂ [1 mark]
(b) Cl more reactive than Br (higher in Group 17). [1 mark]
(c) Red-brown liquid. [1 mark]
Chlorine can displace bromine from potassium bromide solution. Write the balanced equation and explain why this happens.
✓ Balanced equation: Cl₂ + 2KBr → 2KCl + Br₂ [1 mark]
✓ Chlorine is more reactive than bromine (higher up Group VII) [1 mark].
✓ Chlorine readily gains electrons (is a strong oxidising agent) and displaces less reactive halogens [1 mark].
Explain why sodium has a lower melting point (97.8°C) compared to lithium (180°C) in Group I.
✓ The metallic bonding in the alkali metals becomes weaker down the group [1 mark].
✓ Atoms are larger; valence electrons are further from the nucleus and less strongly attracted, resulting in weaker metallic bonding and lower melting point [1 mark].
Which element is a transition metal?
State the colour of Cl₂, Br₂, and I₂ at room temperature.
Explain why noble gases are unreactive.
As you go down Group I (alkali metals), what happens to the reactivity?
Which element has the highest melting point in Group I?
Compare the physical properties of fluorine and iodine (Group VII elements).
Explain the trend in reactivity of Group VII elements (halogens) going down the group.
Which halogen is most likely to displace the other halogens from solutions of their salts?
Define the term transition element and give one property that distinguishes transition elements from other elements.
Predict the properties of element X, which is in Period 3 and Group V of the periodic table. Explain your reasoning.