Key Takeaways for GCSE Chemistry

1. The Periodic Table

  • Periods vs. Groups
    • Periods: Horizontal rows. Elements in the same period have the same number of electron shells.
      Example: In the table [N, K, L], N and L are in the same period.
    • Groups: Vertical columns. Elements in the same group have similar chemical properties and the same number of outer electrons.
      Common mistake: Confusing groups with periods (e.g., Student 2 listed L and J as same period).
  • Transition Metals
    • Located in the middle block of the Periodic Table.
      Example: K is a transition metal.
  • Ion Charges
    • Metals lose electrons → positive ions.
    • Non-metals gain electrons → negative ions.
      Example: Element L (non-metal) forms a 1−1− ion.
      Tip: Group 7 elements gain 1 electron → 1−1−; Group 6 → 2−2−.
  • Noble Gases
    • Unreactive due to full outer electron shells.
      Example: M (noble gas) has a full outer shell.

2. Electronic Configurations

  • Structure
    • Written as numbers separated by commas, e.g., sodium: 2,8,12,8,1.
      Mistake: Student 1 wrote 2,82,8 for sodium (ion configuration, not atom).
  • Reactivity Trends
    • Group 1 (Alkali Metals): Reactivity increases down the group.
      Reason: Outer electron is farther from the nucleus → weaker attraction.
      Example: Potassium (K) reacts more violently with water than sodium (Na).
    • Group 7 (Halogens): Reactivity decreases down the group.
      Example: Fluorine (F) is more reactive than chlorine (Cl).

3. Chemical Equations & Reactions

  • Balancing Equations
    • Example:
      2Na(s)+2H2O(l)→2NaOH(aq)+H2(g)2Na(s)+2H2​O(l)→2NaOH(aq)+H2​(g)
      Tip: Balance metals first, then non-metals, and leave H/O last.
  • Oxidation & Reduction
    • Oxidation: Gain of oxygen (e.g., Mg+O2→MgOMg+O2​→MgO).
    • Reduction: Loss of oxygen (e.g., Fe2O3+Al→Fe+Al2O3Fe2​O3​+Al→Fe+Al2​O3​).
  • Percentage of Oxygen in Air
    • Calculation:
      Percentage=Volume lostInitial volume×100Percentage=Initial volumeVolume lost​×100
      Example: 21100×100=21%10021​×100=21%.

4. Tests for Ions & Gases

  • Chloride Ions:
    • Add dilute nitric acid, then silver nitrate → white precipitate.
  • Sulfate Ions:
    • Add dilute hydrochloric acid, then barium chloride → white precipitate.
  • Carbon Dioxide:
    • Bubble through limewater → cloudy/milky precipitate.
  • Hydrogen:
    • Lit splint → “pop” sound.

5. Reactivity Series

  • Displacement Reactions
    • More reactive metals displace less reactive ones from compounds.
      Example: Magnesium displaces iron from FeSO4FeSO4​:
      Mg(s)+FeSO4(aq)→MgSO4(aq)+Fe(s)Mg(s)+FeSO4​(aq)→MgSO4​(aq)+Fe(s)
  • Order of Reactivity:
    Metals: K > Na > Li > Ca > Mg > Al > Zn > Fe > Cu.

6. Practical Tips

  • Control Variables:
    • In experiments (e.g., testing stone corrosion), ensure mass, acid volume, and time are constant.
  • Avoiding Suck-Back:
    • Remove the delivery tube from water before stopping heating.

7. Common Mistakes & Fixes

  • Flame Tests:
    • Calcium → brick red; Sodium → yellow; Potassium → lilac.
  • Ion Charges:
    • Iron forms Fe2+Fe2+ (iron(II)) and Fe3+Fe3+ (iron(III)).
  • Electron Configuration:
    • Atoms ≠ ions! Sodium atom: 2,8,12,8,1; Sodium ion: 2,82,8.

50 GCSE Chemistry Questions

Section 1: Periodic Table

  1. Which elements are in the same period if the table has N, K, L?
  2. Identify the transition metal from N, K, L.
  3. What is the charge on an ion of element L?
  4. Explain why element M is unreactive.
  5. Name the block where transition metals are located.

Section 2: Electronic Configurations

  1. Write the electronic configuration of lithium (atomic number 3).
  2. Why does reactivity increase down Group 1?
  3. Fluorine has the electronic configuration 2,72,7. Draw its electron arrangement.
  4. What is the electronic configuration of a chloride ion?
  5. Why is potassium more reactive than sodium?

Section 3: Chemical Equations

  1. Balance: Na(s)+H2O(l)→NaOH(aq)+H2(g)Na(s)+H2​O(l)→NaOH(aq)+H2​(g)
  2. Write the equation for magnesium burning in oxygen.
  3. Name the gas released when zinc carbonate decomposes.
  4. What type of reaction is Fe2O3+Al→Fe+Al2O3Fe2​O3​+Al→Fe+Al2​O3​?
  5. Calculate the percentage of oxygen in air if 21 cm³ reacts with copper.

Section 4: Reactivity Series

  1. List Mg, Cu, Fe, Zn in order of reactivity (most to least).
  2. Which metal displaces iron from FeSO4FeSO4​?
  3. Why does sodium react violently with water?
  4. What is the test for hydrogen gas?
  5. Explain why copper does not react with dilute HCl.

Section 5: Ion Tests

  1. Describe the test for sulfate ions.
  2. How would you test for chloride ions in gypsum?
  3. What flame colour indicates calcium ions?
  4. Explain why barium chloride is used in sulfate tests.
  5. How do you test for carbonate ions?

Section 6: Oxidation & Reduction

  1. Define oxidation in terms of oxygen.
  2. What is reduced in Al+Fe2O3→Fe+Al2O3Al+Fe2​O3​→Fe+Al2​O3​?
  3. Name the gas produced when magnesium burns in CO₂.
  4. Write the ionic equation for chlorine reacting with sodium bromide.
  5. Why is sulfur dioxide an atmospheric pollutant?

Section 7: Practical Work

  1. Describe an experiment to test if marble corrodes faster in acid rain.
  2. How would you test rainwater for purity?
  3. Explain how to avoid “suck-back” in gas collection.
  4. What safety precaution is needed when testing for hydrogen?
  5. Describe a test to confirm carbon dioxide.

Section 8: Data Analysis

  1. Use the reactivity table to order Al, Cu, Na, Zn.
  2. Which halogen is most reactive: Cl₂, Br₂, or I₂?
  3. Interpret the results of chlorine water reacting with sodium iodide.
  4. Why does iron rust in water but not in dry air?
  5. Predict the pH of nitrogen oxide’s aqueous solution.

Section 9: Formulae & Calculations

  1. Calculate the relative formula mass of (NH4)2CO3(NH4​)2​CO3​.
  2. Write the formula for gypsum.
  3. What is the charge on the iron ion in FeOFeO?
  4. Balance: Mg(s)+CO2(g)→MgO(s)+C(s)Mg(s)+CO2​(g)→MgO(s)+C(s)
  5. Name the gas produced when HCl dissolves in water.

Section 10: Bonding & Structure

  1. Explain why HCl conducts electricity when dissolved in water.
  2. Draw a dot-and-cross diagram for HCl.
  3. Why is argon unreactive?
  4. Compare the bonding in NaCl and CO₂.
  5. Why does potassium sulfate dissolve in water?

Detailed Answers

  1. N and L (same period = same row).
  2. K (transition metals are in the middle block).
  3. 1−1− (Group 7 gains 1 electron).
  4. Full outer electron shell (no tendency to react).
  5. Middle block (between Groups 2 and 3).
  6. 2,12,1 (Li has atomic number 3).
  7. Outer electron is farther from nucleus → weaker attraction.
  8. Fluorine diagram
  9. 2,8,82,8,8 (gains 1 electron).
  10. K has more electron shells → outer electron less attracted.
  11. 2Na+2H2O→2NaOH+H22Na+2H2​O→2NaOH+H2​
  12. 2Mg+O2→2MgO2Mg+O2​→2MgO
  13. Carbon dioxide (CO2CO2​).
  14. Redox reaction (Al is oxidised, Fe is reduced).
  15. 21100×100=21%10021​×100=21%
  16. Mg > Zn > Fe > Cu (displacement reactions).
  17. Magnesium (more reactive than iron).
  18. Outer electron easily lost → rapid reaction.
  19. Lit splint produces a “pop”.
  20. Copper is below H in reactivity series → no reaction.
  21. Add HCl + BaCl₂ → white precipitate.
  22. Add HNO₃ + AgNO₃ → white precipitate.
  23. Brick red flame.
  24. Ba²⁺ reacts with SO₄²⁻ → BaSO₄ precipitate.
  25. Add acid → effervescence; gas turns limewater cloudy.
  26. Gain of oxygen (e.g., Mg → MgO).
  27. Iron oxide (loses oxygen).
  28. Carbon (black smoke).
  29. Cl2+2Br−→2Cl−+Br2Cl2​+2Br−→2Cl−+Br2​
  30. Forms acid rain (reacts with water → H₂SO₃).
  31. Weigh stone samples → immerse in acid → measure mass loss.
  32. Test boiling point (pure water boils at 100°C).
  33. Remove delivery tube before cooling.
  34. Use a safety screen (H₂ is explosive).
  35. Bubble through limewater → milky precipitate.
  36. Na > Al > Zn > Cu (displacement data).
  37. Cl₂ (displaces Br₂ and I₂).
  38. Orange → brown (Cl₂ oxidises I⁻ to I₂).
  39. Rusting requires oxygen and water.
  40. pH < 7 (non-metal oxide = acidic).
  41. N: 14, H: 1, C: 12, O: 16 → 2(14+4)+12+3(16)=962(14+4)+12+3(16)=96.
  42. CaSO4⋅2H2OCaSO4​⋅2H2​O.
  43. 2+2+ (O²⁻ requires Fe²⁺ for neutrality).
  44. 2Mg+CO2→2MgO+C2Mg+CO2​→2MgO+C
  45. H⁺ and Cl⁻ ions (dissociation).
  46. Ionises into H⁺ and Cl⁻ → conducts electricity.
  47. HCl dot-and-cross
  48. Full outer shell (no electron transfer).
  49. NaCl: ionic; CO₂: covalent.
  50. Ionic compound → dissolves into ions in water.