GCSE Chemistry Syllabus

Chemistry syllabus — full chapter and topic list

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

1. 1

   Atomic Structure and the Periodic Table

   3 topics

   • Atoms and Elements
   • The Periodic Table
   • Bonding and Structure
2. 2

   Chemical Changes

   3 topics

   • Reactivity of Metals
   • Electrolysis
   • Energy Changes
3. 3

   Organic Chemistry

   3 topics

   • Hydrocarbons
   • Polymers
   • Alcohols and Carboxylic Acids
4. 4

   Chemical Analysis

   3 topics

   • Purity and Formulations
   • Chromatography
   • Identification of Ions

Chemistry flashcards for GCSE

21 of 59 cards from the Chemistry deck — real questions with worked answers.

1. What is an element?

   A substance made of only one type of atom. There are about 100 different elements, each with its own chemical symbol, shown in the periodic table.

2. State the relative charges and relative masses of the three subatomic particles.

   Proton: charge $+1$, mass $1$. Neutron: charge $0$, mass $1$. Electron: charge $-1$, mass $\frac{1}{1840}$ (negligible).

3. What do the atomic (proton) number and mass number of an atom tell you?

   The atomic number is the number of protons (which equals the number of electrons in a neutral atom). The mass number is the total number of protons plus neutrons.

4. Define isotopes.

   Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, giving them the same atomic number but different mass numbers.

5. How is the relative atomic mass ($A_r$) of an element calculated from its isotopes?

   $$A_r = \frac{\sum (\text{isotope abundance} \times \text{isotope mass})}{\sum \text{abundances}}$$ It is a weighted mean of the isotope masses.

6. Describe the arrangement of electrons in the first three shells (energy levels).

   The first shell holds up to $2$ electrons, the second up to $8$, and the third up to $8$ (at GCSE). Shells fill from the innermost outwards, e.g. sodium is $2,8,1$.

7. How did the discovery of the electron and the nucleus change the model of the atom?

   J.J. Thomson's plum pudding model was replaced after Rutherford's alpha-scattering experiment showed atoms have a tiny, dense, positively charged nucleus. Bohr then proposed electrons orbit in fixed shells.

8. What is a compound and how does it differ from a mixture?

   A compound contains two or more elements chemically bonded together in fixed proportions and can only be separated chemically. A mixture contains substances not chemically joined and can be separated by physical means.

9. How is the modern periodic table arranged?

   Elements are arranged in order of increasing atomic (proton) number, in rows called periods and columns called groups, so elements with similar properties line up in the same group.

10. Why did Mendeleev leave gaps in his early periodic table?

    He left gaps for elements he believed had not yet been discovered, and used the gaps to predict the properties of those undiscovered elements, which later proved correct.

11. What does the group number tell you about an element's atoms?

    The group number equals the number of electrons in the outer shell (for groups 1-7/0), which determines the element's chemical properties.

12. Describe the trend in reactivity down Group 1 (alkali metals) and explain it.

    Reactivity increases down the group. The outer electron is further from the nucleus and more shielded, so it is lost more easily during reactions.

13. Describe the trend in reactivity down Group 7 (halogens) and explain it.

    Reactivity decreases down the group. Atoms get larger with more shielding, so it becomes harder to attract and gain an outer electron.

14. Why are the Group 0 elements (noble gases) unreactive?

    They have full outer electron shells (a stable electronic structure), so they have little tendency to gain, lose or share electrons.

15. Give the general properties of transition metals compared with Group 1 metals.

    Transition metals are harder, stronger, denser and have higher melting points; they are less reactive. They form coloured compounds, have variable oxidation states (ion charges) and are often used as catalysts.

16. What is ionic bonding and between which types of element does it occur?

    Ionic bonding is the electrostatic attraction between oppositely charged ions, formed when electrons transfer from a metal to a non-metal. E.g. $\ce{Na -> Na+ + e-}$ and $\ce{Cl + e- -> Cl-}$.

17. What is covalent bonding and where does it occur?

    Covalent bonding is a shared pair of electrons between two atoms. It occurs between non-metal atoms, e.g. in $\ce{H2O}$, $\ce{CH4}$ and $\ce{Cl2}$.

18. Describe metallic bonding.

    Metallic bonding is the electrostatic attraction between positive metal ions arranged in a lattice and a 'sea' of delocalised outer-shell electrons that are free to move.

19. Explain why ionic compounds have high melting points and conduct electricity only when molten or dissolved.

    Strong electrostatic forces in the giant ionic lattice need lots of energy to break, giving high melting points. Ions are fixed in the solid, but when molten or dissolved they are free to move and carry charge.

20. Why do simple molecular substances have low melting and boiling points?

    They have weak intermolecular forces between molecules that need little energy to overcome. The strong covalent bonds within molecules are not broken during melting or boiling.

21. Compare the structure and properties of diamond and graphite.

    Both are giant covalent carbon structures. In diamond each carbon forms 4 bonds, making it very hard with no free electrons (non-conductor). In graphite each carbon forms 3 bonds in layers that slide, and delocalised electrons let it conduct electricity.

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