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A Level H1 Chemistry Periodic Table Quiz
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Questions
A-Level Chemistry H1 Quiz - Periodic Table
Name: ________________________
Class: ________________________
Date: ________________________
Score: ______ / 40
Duration: 45 minutes
Total Marks: 40
Instructions:
- This quiz contains 20 questions on the topic of the Periodic Table.
- Answer ALL questions in the spaces provided.
- Show all working for calculation questions.
- A Data Booklet is provided for reference.
- Marks for each question are indicated in brackets.
Section A: Short Answer (10 marks)
Answer all questions in this section.
1. State the trend in atomic radius across Period 3 from sodium to chlorine. Explain this trend in terms of nuclear charge and shielding effect. [2]
2. The first ionisation energy of aluminium (578 kJ mol⁻¹) is lower than that of magnesium (738 kJ mol⁻¹), despite aluminium having a greater nuclear charge. Explain this anomaly. [2]
3. Define the term electronegativity. State how electronegativity changes across a period and down a group. [2]
4. Identify the Period 3 element that forms an oxide which is amphoteric. Write equations to show the amphoteric nature of this oxide. [2]
5. Explain why the atomic radius of a chloride ion (Cl⁻) is larger than that of a chlorine atom (Cl). [2]
Section B: Structured Questions (10 marks)
Answer all questions in this section.
6. The table below shows some properties of Period 3 elements.
| Element | Na | Mg | Al | Si | P | S | Cl |
|---|---|---|---|---|---|---|---|
| Atomic radius / pm | 186 | 160 | 143 | 117 | 110 | 104 | 99 |
| Melting point / °C | 98 | 650 | 660 | 1410 | 44 | 119 | –101 |
(a) Explain the general trend in atomic radius across Period 3. [2]
(b) Explain why silicon has a much higher melting point than phosphorus. [3]
7. The graph below shows the first ionisation energies of the Period 3 elements.
[A graph is provided showing first ionisation energy (kJ mol⁻¹) on the y-axis against elements Na to Ar on the x-axis. The graph shows a general increase from Na to Ar, with dips at Al and S.]
(a) Explain the general increase in first ionisation energy across Period 3. [2]
(b) Explain why the first ionisation energy of sulfur is lower than that of phosphorus. [2]
(c) Predict, with a reason, whether the first ionisation energy of potassium is higher or lower than that of sodium. [1]
Section C: Data-Based Questions (10 marks)
Answer all questions in this section.
8. Chlorine exists as two isotopes, ³⁵Cl and ³⁷Cl, with relative abundances of 75.5% and 24.5% respectively.
(a) Define the term isotope. [1]
(b) Calculate the relative atomic mass of chlorine. Give your answer to one decimal place. [2]
9. The oxides of Period 3 elements show a trend from basic to acidic character across the period.
(a) Classify each of the following oxides as basic, amphoteric, or acidic:
- (i) Na₂O [1]
- (ii) Al₂O₃ [1]
- (iii) SO₂ [1]
(b) Write a balanced equation, including state symbols, for the reaction of sodium oxide with water. [2]
(c) Write a balanced equation, including state symbols, for the reaction of sulfur dioxide with water. Name the acid formed. [2]
Section D: Application and Analysis (10 marks)
Answer all questions in this section.
10. The chlorides of Period 3 elements also show a trend in structure and bonding.
(a) State the type of structure and bonding in:
- (i) Sodium chloride [1]
- (ii) Silicon tetrachloride [1]
(b) Explain why sodium chloride has a high melting point while silicon tetrachloride has a low melting point. [3]
11. Magnesium oxide is used to line the inside of furnaces, but silicon dioxide cannot be used for this purpose. Explain why. [2]
12. Silicon tetrachloride reacts vigorously with water to produce silicon dioxide and hydrogen chloride. Write a balanced equation for this reaction. [1]
13. State the type of structure and bonding present in:
- (i) Sodium [1]
- (ii) Silicon [1]
14. Explain why the first ionisation energy generally increases across Period 3, but there is a slight decrease from phosphorus to sulfur. [2]
15. Describe the trend in the acid-base character of the oxides of Period 3 elements, from sodium to chlorine. Provide one example of a basic oxide and one example of an acidic oxide. [2]
Section E: Further Analysis (10 marks)
Answer all questions in this section.
16. The melting points of Period 3 elements show a general increase from sodium to silicon, then a sharp decrease to phosphorus. Explain this trend in terms of structure and bonding. [3]
17. Explain why the ionic radius of Na⁺ is smaller than the atomic radius of Na, while the ionic radius of Cl⁻ is larger than the atomic radius of Cl. [2]
18. The second ionisation energy of sodium is much higher than its first ionisation energy. Explain why. [2]
19. Predict, with a reason, the relative electronegativity values of magnesium and chlorine. [1]
20. Describe and explain the trend in electrical conductivity across Period 3 elements, from sodium to argon. [2]
END OF QUIZ
Check your answers carefully before submitting.
Answers
A-Level Chemistry H1 Quiz - Periodic Table - Answer Key
Total Marks: 40
Section A: Short Answer (10 marks)
1. State the trend in atomic radius across Period 3 from sodium to chlorine. Explain this trend in terms of nuclear charge and shielding effect. [2]
Answer:
- Atomic radius decreases across Period 3 from Na to Cl. [1]
- Across the period, the number of protons (nuclear charge) increases, while the shielding effect remains approximately constant as electrons are added to the same outer shell. The increased nuclear charge attracts the outer electrons more strongly, pulling them closer to the nucleus and decreasing atomic radius. [1]
2. The first ionisation energy of aluminium (578 kJ mol⁻¹) is lower than that of magnesium (738 kJ mol⁻¹), despite aluminium having a greater nuclear charge. Explain this anomaly. [2]
Answer:
- Magnesium has electronic configuration 1s²2s²2p⁶3s²; aluminium has 1s²2s²2p⁶3s²3p¹. [1]
- In aluminium, the outermost electron is in a 3p orbital, which is of higher energy and further from the nucleus than the 3s orbital of magnesium. The 3p electron in Al is also shielded by the 3s electrons, making it easier to remove despite the higher nuclear charge. [1]
3. Define the term electronegativity. State how electronegativity changes across a period and down a group. [2]
Answer:
- Electronegativity is the ability of an atom to attract bonding electrons in a covalent bond. [1]
- Electronegativity increases across a period (due to increasing nuclear charge and decreasing atomic radius) and decreases down a group (due to increasing atomic radius and increased shielding). [1]
4. Identify the Period 3 element that forms an oxide which is amphoteric. Write equations to show the amphoteric nature of this oxide. [2]
Answer:
- Aluminium (Al) forms the amphoteric oxide Al₂O₃. [1]
- Reaction with acid: Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l)
- Reaction with base: Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄(aq) [1 for both equations]
5. Explain why the atomic radius of a chloride ion (Cl⁻) is larger than that of a chlorine atom (Cl). [2]
Answer:
- A chloride ion has gained one electron compared to a chlorine atom, resulting in 18 electrons but still 17 protons. [1]
- The increased electron-electron repulsion in the outer shell, combined with the same nuclear charge, causes the electron cloud to expand, resulting in a larger ionic radius. [1]
Section B: Structured Questions (10 marks)
6. (a) Explain the general trend in atomic radius across Period 3. [2]
Answer:
- Atomic radius decreases from Na to Cl. [1]
- Across the period, nuclear charge increases while electrons are added to the same principal quantum shell (n=3). The shielding effect remains approximately constant, so the increased nuclear attraction pulls electrons closer, decreasing atomic radius. [1]
(b) Explain why silicon has a much higher melting point than phosphorus. [3]
Answer:
- Silicon has a giant covalent (macromolecular) structure with strong covalent bonds between Si atoms throughout the lattice. [1]
- Phosphorus exists as discrete P₄ molecules with weak van der Waals' forces between molecules. [1]
- Melting silicon requires breaking many strong covalent bonds, requiring a large amount of energy. Melting phosphorus only requires overcoming weak intermolecular forces, requiring much less energy. [1]
7. (a) Explain the general increase in first ionisation energy across Period 3. [2]
Answer:
- First ionisation energy generally increases from Na to Ar. [1]
- Nuclear charge increases across the period while shielding remains approximately constant (electrons added to same shell). Atomic radius decreases, so outer electrons are held more tightly and require more energy to remove. [1]
(b) Explain why the first ionisation energy of sulfur is lower than that of phosphorus. [2]
Answer:
- P: [Ne]3s²3p³ (half-filled 3p subshell); S: [Ne]3s²3p⁴. [1]
- In sulfur, one 3p orbital contains a pair of electrons. The repulsion between these paired electrons makes it easier to remove one of them compared to removing an unpaired electron from the half-filled, more stable 3p subshell of phosphorus. [1]
(c) Predict, with a reason, whether the first ionisation energy of potassium is higher or lower than that of sodium. [1]
Answer:
- Potassium has a lower first ionisation energy than sodium. [0.5]
- Potassium's outer electron is in the 4s orbital, which is further from the nucleus than sodium's 3s orbital. There is also increased shielding from inner electron shells, so the outer electron is less strongly attracted and easier to remove. [0.5]
Section C: Data-Based Questions (10 marks)
8. (a) Define the term isotope. [1]
Answer:
- Isotopes are atoms of the same element with the same number of protons (atomic number) but different numbers of neutrons (different mass numbers). [1]
(b) Calculate the relative atomic mass of chlorine. Give your answer to one decimal place. [2]
Answer:
- Ar(Cl) = (35 × 75.5/100) + (37 × 24.5/100) [1]
- = 26.425 + 9.065 = 35.49 ≈ 35.5 [1]
9. (a) Classify each oxide: [3]
- (i) Na₂O: Basic [1]
- (ii) Al₂O₃: Amphoteric [1]
- (iii) SO₂: Acidic [1]
(b) Write a balanced equation for the reaction of sodium oxide with water. [2]
Answer:
- Na₂O(s) + H₂O(l) → 2NaOH(aq) [1 for correct formulae, 1 for balancing and state symbols]
(c) Write a balanced equation for the reaction of sulfur dioxide with water. Name the acid formed. [2]
Answer:
- SO₂(g) + H₂O(l) → H₂SO₃(aq) [1]
- Acid formed: Sulfurous acid [1]
Section D: Application and Analysis (10 marks)
10. (a) State the type of structure and bonding: [2]
- (i) Sodium chloride: Giant ionic lattice with ionic bonding between Na⁺ and Cl⁻ ions. [1]
- (ii) Silicon tetrachloride: Simple molecular structure with covalent bonds within SiCl₄ molecules and weak van der Waals' forces between molecules. [1]
(b) Explain the difference in melting points. [3]
Answer:
- NaCl has a giant ionic structure with strong electrostatic forces of attraction between oppositely charged Na⁺ and Cl⁻ ions throughout the lattice. [1]
- SiCl₄ has a simple molecular structure with weak van der Waals' forces between discrete SiCl₄ molecules. [1]
- Melting NaCl requires overcoming many strong ionic bonds, requiring a large amount of energy (high melting point). Melting SiCl₄ only requires overcoming weak intermolecular forces, requiring little energy (low melting point). [1]
11. Explain why magnesium oxide is used to line furnaces but silicon dioxide cannot be used. [2]
Answer:
- MgO is basic and has a very high melting point (≈2852°C), making it suitable as a refractory lining that can withstand high temperatures without reacting with basic materials. [1]
- SiO₂ is acidic and would react with basic slags or materials in the furnace. It also has a lower melting point than MgO and may not withstand the extreme temperatures. [1]
12. Write a balanced equation for the reaction of SiCl₄ with water. [1]
Answer:
- SiCl₄(l) + 2H₂O(l) → SiO₂(s) + 4HCl(g) [1]
13. State the type of structure and bonding present in: [2]
- (i) Sodium: Giant metallic structure with metallic bonding (delocalised electrons and Na⁺ cations). [1]
- (ii) Silicon: Giant covalent (macromolecular) structure with strong covalent bonds between Si atoms. [1]
14. Explain why the first ionisation energy generally increases across Period 3, but there is a slight decrease from phosphorus to sulfur. [2]
Answer:
- General increase: Nuclear charge increases while shielding remains approximately constant, so outer electrons are held more tightly. [1]
- Decrease P to S: Phosphorus has a half-filled 3p subshell (3p³) which is more stable. Sulfur has a paired electron in one 3p orbital (3p⁴); the repulsion between paired electrons makes it easier to remove one, lowering the ionisation energy. [1]
15. Describe the trend in the acid-base character of the oxides of Period 3 elements, from sodium to chlorine. Provide one example of a basic oxide and one example of an acidic oxide. [2]
Answer:
- The oxides change from basic (left) to amphoteric (middle) to acidic (right) across the period. [1]
- Basic oxide example: Na₂O (sodium oxide). Acidic oxide example: SO₂ (sulfur dioxide) or SO₃. [1]
Section E: Further Analysis (10 marks)
16. The melting points of Period 3 elements show a general increase from sodium to silicon, then a sharp decrease to phosphorus. Explain this trend in terms of structure and bonding. [3]
Answer:
- Na, Mg, Al: Giant metallic structures. Melting point increases due to increasing number of delocalised electrons and smaller ionic radius, leading to stronger metallic bonding. [1]
- Si: Giant covalent structure with strong covalent bonds throughout the lattice, giving it the highest melting point. [1]
- P, S, Cl, Ar: Simple molecular structures with weak van der Waals' forces between molecules. Melting points are low because only weak intermolecular forces need to be overcome. [1]
17. Explain why the ionic radius of Na⁺ is smaller than the atomic radius of Na, while the ionic radius of Cl⁻ is larger than the atomic radius of Cl. [2]
Answer:
- Na⁺: Loss of the 3s electron means the outer shell is now n=2. The remaining electrons experience a greater effective nuclear charge and are pulled closer, reducing the radius. [1]
- Cl⁻: Gain of one electron increases electron-electron repulsion in the n=3 shell. With the same nuclear charge, the electron cloud expands, increasing the radius. [1]
18. The second ionisation energy of sodium is much higher than its first ionisation energy. Explain why. [2]
Answer:
- First IE: Removal of the 3s¹ electron from a neutral Na atom. [0.5]
- Second IE: Removal of an electron from the Na⁺ ion, which has the stable noble gas configuration of neon (2p⁶). [0.5]
- The second electron is removed from a lower energy level (n=2), closer to the nucleus, with much less shielding, and from a positively charged ion. This requires significantly more energy. [1]
19. Predict, with a reason, the relative electronegativity values of magnesium and chlorine. [1]
Answer:
- Chlorine has a higher electronegativity than magnesium. [0.5]
- Chlorine is further to the right in Period 3, with a greater nuclear charge and smaller atomic radius, so it attracts bonding electrons more strongly. [0.5]
20. Describe and explain the trend in electrical conductivity across Period 3 elements, from sodium to argon. [2]
Answer:
- Na, Mg, Al: Good conductors due to metallic bonding with delocalised electrons. Conductivity increases from Na to Al as the number of delocalised electrons per atom increases. [1]
- Si: A semiconductor with low conductivity due to its giant covalent structure.
- P, S, Cl, Ar: Non-conductors (insulators) as they have simple molecular structures with no mobile charged particles. [1]
END OF ANSWER KEY