Cambridge O Level Chemistry 5070 Syllabus

Chemistry 5070 syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Chemistry 5070 in Cambridge O Level, not a summary of it.

1. 1

   States of Matter and Separation

   3 topics

   • Kinetic Particle Theory
   • Diffusion
   • Methods of Purification and Analysis
2. 2

   Atoms, Elements and Compounds

   3 topics

   • Atomic Structure
   • Bonding
   • Structure and Properties of Materials
3. 3

   Stoichiometry

   3 topics

   • Formulae and Equations
   • The Mole Concept
   • Reacting Masses and Concentrations
4. 4

   Electrochemistry

   2 topics

   • Electrolysis
   • Applications of Electrolysis
5. 5

   Chemical Energetics

   3 topics

   • Exothermic and Endothermic Reactions
   • Energy Profile Diagrams
   • Fuels and Energy
6. 6

   Rate of Reaction

   2 topics

   • Factors Affecting Rate
   • Measuring Reaction Rates
7. 7

   Reversible Reactions and Equilibrium

   2 topics

   • Reversible Reactions
   • Dynamic Equilibrium and the Haber Process
8. 8

   Acids, Bases and Salts

   4 topics

   • Properties of Acids and Bases
   • Oxides and Neutralisation
   • Preparation of Salts
   • Identification of Ions and Gases
9. 9

   The Periodic Table

   3 topics

   • Periodic Trends
   • Group Properties
   • Transition Elements
10. 10

    Metals

    3 topics

    • Properties and Reactivity Series
    • Extraction of Metals
    • Alloys and Corrosion
11. 11

    Chemistry of the Environment

    2 topics

    • Air and Water
    • Sulfur and Carbonates
12. 12

    Organic Chemistry

    4 topics

    • Fuels and Petroleum
    • Alkanes and Alkenes
    • Alcohols and Carboxylic Acids
    • Macromolecules and Polymers

Chemistry 5070 flashcards for Cambridge O Level

24 of 51 cards from the Chemistry 5070 deck — real questions with worked answers.

1. State the main points of the kinetic particle theory of matter.

   All matter is made of tiny particles that are in constant motion; the particles have spaces between them and attractive forces between them; the higher the temperature, the greater the average kinetic energy and speed of the particles.

2. Compare the arrangement, movement and forces between particles in solids, liquids and gases.

   Solid: regular close-packed arrangement, vibrate about fixed positions, strong forces. Liquid: irregular close arrangement, move/slide around each other, moderate forces. Gas: far apart and random, move fast in all directions, negligible forces.

3. Define melting, boiling, evaporation, condensation, freezing and sublimation in terms of state changes.

   Melting: solid to liquid. Boiling/evaporation: liquid to gas. Condensation: gas to liquid. Freezing/solidification: liquid to solid. Sublimation: solid directly to gas (and vice versa).

4. Why does a gas exert pressure on the walls of its container?

   Gas particles are in constant random motion and collide with the container walls; these collisions exert a force per unit area, which is the gas pressure.

5. Using kinetic theory, explain why a gas expands when heated at constant pressure.

   Heating increases the average kinetic energy and speed of the particles, so they collide with the walls more often and harder; to keep pressure constant the volume must increase.

6. Define diffusion.

   Diffusion is the net movement of particles from a region of higher concentration to a region of lower concentration due to the random motion of particles.

7. How does the relative molecular mass of a gas affect its rate of diffusion?

   Lighter gases (lower relative molecular mass) diffuse faster than heavier gases; rate of diffusion is inversely related to the square root of the relative molecular mass.

8. In the ammonia and hydrogen chloride diffusion experiment, why does the white ring of ammonium chloride form nearer the HCl end of the tube?

   Ammonia (Mr 17) is lighter and diffuses faster than hydrogen chloride (Mr 36.5), so they meet closer to the HCl-soaked end.

9. How does temperature affect the rate of diffusion?

   Increasing temperature increases the kinetic energy and speed of particles, so the rate of diffusion increases.

10. What is the difference between a pure substance and a mixture, and how do their melting/boiling points differ?

    A pure substance contains only one substance and has a sharp, fixed melting/boiling point. A mixture contains more than one substance and melts/boils over a range of temperatures.

11. Describe how filtration separates a mixture and what it separates.

    Filtration separates an insoluble solid from a liquid; the mixture is poured through filter paper, the solid (residue) stays on the paper and the liquid (filtrate) passes through.

12. What does crystallisation separate and how is it carried out?

    Crystallisation separates a dissolved solid (solute) from its solution by evaporating some solvent to form a saturated solution, then cooling so crystals form, which are filtered and dried.

13. How does simple distillation work and what does it separate?

    Simple distillation separates a solvent from a solution (e.g. water from salt water); the solution is heated, the solvent evaporates, then passes through a condenser where it cools and condenses back to liquid, collected as distillate.

14. How does fractional distillation separate a mixture of miscible liquids?

    It separates liquids with different boiling points using a fractionating column; the liquid with the lowest boiling point distils over first, the column provides a temperature gradient so each component condenses and re-evaporates until separated.

15. Explain how paper chromatography separates a mixture of soluble coloured substances.

    A spot of the mixture is placed on paper and a solvent moves up; components that are more soluble and less strongly attracted to the paper travel further, separating the mixture into spots.

16. Define the Rf value in chromatography and give its formula.

    Rf = distance moved by the substance / distance moved by the solvent front. It is a ratio with no units, used to identify components.

17. State the relative charge, relative mass and location of a proton, neutron and electron.

    Proton: charge +1, mass 1, in nucleus. Neutron: charge 0, mass 1, in nucleus. Electron: charge -1, mass ~1/1840, in shells around the nucleus.

18. Define proton (atomic) number and nucleon (mass) number.

    Proton number is the number of protons in an atom's nucleus. Nucleon number is the total number of protons and neutrons in the nucleus.

19. Define isotopes.

    Isotopes are atoms of the same element with the same number of protons (same proton number) but different numbers of neutrons (different nucleon numbers).

20. State the maximum number of electrons that can occupy the first three electron shells.

    First shell: 2 electrons; second shell: 8 electrons; third shell: 8 electrons (up to calcium).

21. How is an ion formed, and how do metals and non-metals form ions?

    An ion forms when an atom loses or gains electrons. Metals lose electrons to form positive ions (cations); non-metals gain electrons to form negative ions (anions).

22. Define ionic bonding.

    Ionic bonding is the strong electrostatic attraction between oppositely charged ions, formed by the transfer of electrons from a metal to a non-metal.

23. Define covalent bonding.

    Covalent bonding is the sharing of one or more pairs of electrons between two non-metal atoms, so each atom achieves a stable (noble gas) electron configuration.

24. Describe metallic bonding.

    Metallic bonding is the electrostatic attraction between a lattice of positive metal ions and a 'sea' of delocalised (mobile) electrons.

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