10.05 Past Paper Practice
Questions A
1.The equation for two reactions are shown.
Equation 1: C + CO2 → xCO
Equation 2: 2C2H6 + 7O2 → 4CO2 + yH2O
Question: Which row shows the value of x, the value of y, and the equations that are for redox reactions?
| x | y | Redox Reactions | |
|---|---|---|---|
| A | 1 | 3 | Equation 1 only |
| B | 2 | 3 | Equations 1 and 2 |
| C | 2 | 6 | Equation 1 only |
| D | 2 | 6 | Equations 1 and 2 |
Answer: D
Correct Answer:
x=2, y=6 and both equations (Equations 1 and 2) are redox reactions.
Explanation:
Step 1: Identify the values of x and y
- For Equation 1: C + CO2→ x CO
In this reaction, one carbon atom reacts with carbon dioxide to produce carbon monoxide. To balance the equation:- Reactant side: 1 carbon from CCC and 1 carbon from CO2 → total of 2 carbons.
- Product side: Each CO contains 1 carbon → x=2
- For Equation 2: 2C2H6 + 7O2 → 4CO2 + y H2O
In this combustion reaction:- Reactant side: 2C2H6 contains 6 H2 molecules (total 12 hydrogen atoms).
- Product side: H2O contains 2 H atoms → y=6
Step 2: Check for redox reactions
- Redox reaction definition: A reaction in which oxidation and reduction occur simultaneously. Oxidation involves the loss of electrons, and reduction involves the gain of electrons.
- For Equation 1:
- Carbon (C) is oxidized from 0 (elemental state) to +2 in CO.
- Carbon in CO2 is reduced from +4 to +2 in CO.
- Therefore, Equation 1 is a redox reaction.
- For Equation 2:
- Carbon in C2H6 is oxidized from −3 to +4 in CO2.
- Oxygen (O2) is reduced from 0 to −2 in H2O.
- Therefore, Equation 2 is also a redox reaction.
.
2. Which row describes the changes that occur when metals burn in oxygen?
Answer: C
Explanation:
Step 1: What happens when metals burn in oxygen?
When metals burn in oxygen, they react with oxygen to form metal oxides. This process releases energy in the form of heat, which increases the surrounding temperature.
Step 2: Oxidation or Reduction?
- Oxidation: Metals lose electrons when they react with oxygen. This loss of electrons is called oxidation.
- Example: 2Mg+O2→2MgO2
Magnesium atoms (Mg) lose 2 electrons each to form Mg2+ in MgO.
- Example: 2Mg+O2→2MgO2
- Reduction: This occurs when a substance gains electrons. In this case, oxygen is reduced to O2− ions. However, the question specifically asks about the change to the metal, which is oxidized.
Step 3: Analyze the options
- Temperature: The burning process is highly exothermic, so the temperature increases.
- Metal: The metal undergoes oxidation.
Thus, the correct row is: Temperature increases, Metal oxidised (Option C).
3. Silver oxide reacts with magnesium to make silver and magnesium oxide:
Reaction: Ag2O + Mg →2 Ag + MgO
Question: Which substance is oxidised in this reaction?
| Options |
|---|
| A. Magnesium |
| B. Magnesium oxide |
| C. Silver |
| D. Silver oxide |
Answer: A (Magnesium)
Explanation:
Step 1: What is oxidation?
- Oxidation is the loss of electrons or an increase in oxidation state.
- Reduction is the gain of electrons or a decrease in oxidation state.
Step 2: Analyze the reaction
- Magnesium (Mg):
In the reactants, magnesium is in its elemental form (Mg0). In the products, it becomes part of magnesium oxide (MgO) where it is in the +2 oxidation state.- Conclusion: Magnesium is oxidized from 0 to +2.
- Silver oxide (Ag2O):
In the reactants, silver is in the +1 oxidation state within Ag2O. In the products, silver is in its elemental form (Ag0), meaning it is reduced.- Conclusion: Silver is reduced from +1 to 0.
Step 3: Eliminate incorrect options
- Silver oxide (D): Silver oxide is the compound reduced in the reaction.
- Magnesium oxide (B): This is a product and does not undergo oxidation or reduction.
- Silver (C): Silver is reduced, not oxidized.
4. Four redox equations and statements about the equations are shown:
- C + O2 → CO2 (carbon is oxidised)
- CO2 + C → 2CO (carbon dioxide is oxidised)
- CO2 + C → 2CO (carbon is oxidised)
- Fe2O3 + 3CO → 2Fe + 3CO2 (iron(III) oxide is oxidised)
Which statements about the equations are correct?
Answer: B (Reactions 1 and 3)
Correct Analysis:
Equation 1: C+O2→CO2 (carbon is oxidized)
- Carbon starts in the elemental state (C0) and forms CO2, where it is in the +4 oxidation state.
- Conclusion: The statement is correct, as carbon is oxidized.
Equation 2: CO2+C→2CO (carbon dioxide is oxidized)
- In CO2, carbon is in the +4 oxidation state.
- In CO, carbon is in the +2 oxidation state.
- Carbon in CO2 is reduced, not oxidized.
- Conclusion: The statement is incorrect, as carbon dioxide is reduced.
Equation 3: CO2+C→2CO (carbon is oxidized)
- The elemental carbon (C0) reacts to form CO, where it is in the +2 oxidation state.
- Carbon from C0 is indeed oxidized.
- Conclusion: The statement is correct.
Equation 4: Fe2O3 + 3CO → 2Fe + 3CO2 (iron(III) oxide is oxidized)
- In 3Fe2O3, iron is in the +3 oxidation state.
- In Fe, iron is in the 0 oxidation state.
- Iron in Fe2O3 is reduced, not oxidized.
- Conclusion: The statement is incorrect, as iron(III) oxide is reduced.
Final Answer:
- The correct statements are:
- 3 (carbon is oxidized)
- 1 (carbon is oxidized)
5. Which equation shows the reduction of copper?
- A) CuO + C → Cu + CO
- B) 2CuS + 3O2 → 2CuO + 2SO2
- C) Cu(g) → Cu(l)
- D) Cu(l) → Cu(s)
Answer: A (CuO + C → Cu + CO)
Explanation: Copper is reduced as it gains electrons, going from Cu2+ in CuO to Cu (metal).
6. Which equation shows an oxidation reaction?
Answer: A (C + O2 → CO2)
Explanation: In this reaction, carbon is oxidised as it combines with oxygen to form CO2.
7. The equation for the reaction between magnesium and copper(II) oxide is shown:
- Reaction: Mg + CuO → MgO + Cu
- Question: Which substance is oxidised?
- A) Cu
- B) CuO
- C) Mg
- D) MgO
Answer: C (Magnesium)
Explanation: Magnesium is oxidised as it forms Mg2+ in MgO, losing electrons in the process.
8. When magnesium is heated with zinc oxide, a reaction occurs:
- Reaction: Mg + ZnO → MgO + Zn
- Question: Which substance is oxidised?
- A) magnesium
- B) magnesium oxide
- C) zinc
- D) zinc oxide
Answer: A (Magnesium)
Explanation: Magnesium is oxidised as it forms Mg2+ in MgO. Zinc oxide is reduced as zinc gains electrons.
9. Hydrogen iodide is dissolved in water:
- Reaction: HI → H+ + I–
- Question: Which row describes the final colours seen when the solution is tested with damp red litmus paper and acidified aqueous potassium manganate(VII)?
Answer: D
Explanation: The solution is acidic, so red litmus paper remains red. Acidified potassium manganate(VII) is colourless in the presence of iodide ions.
10. The equation for the reaction of metal M with aqueous zinc sulfate is shown:
- Reaction:
M(s) + ZnSO4(aq) → MSO4(aq)+Zn(s) - Question: Which statement explains why metal M reacts with aqueous zinc sulfate?
- A) Zinc is less reactive than M because M is able to accept electrons from zinc ions.
- B) Zinc is a more powerful reducing agent than M.
- C) Zinc is more reactive than M because it can lose electrons more easily than M.
- D) Zinc ions can remove electrons from M.
D) Zinc ions can remove electrons from M.
Explanation:
In the reaction: M(s)+ZnSO4(aq)→MSO4(aq)+Zn(s)
- Metal M is oxidized: M(s)→ M2+ + 2e−
- Zinc ions (Zn²⁺) are reduced: Zn2++2e−→Zn(s)
This indicates that zinc ions (Zn2+) act as an oxidizing agent, meaning they remove electrons from metal M. Therefore, the reaction proceeds because zinc ions can accept electrons from metal M, facilitating the displacement.
Why the other options are incorrect:
- A) Misstates the electron transfer process. It’s not that M accepts electrons from zinc ions; rather, M loses electrons to zinc ions.
- B) Incorrect because zinc acts as an oxidizing agent in this reaction, not as a reducing agent.
- C) Incorrect because the reaction shows that metal M is more reactive than zinc, not the other way around.
Final Answer:
D) Zinc ions can remove electrons from M.
11. In which equation is the underlined substance acting as a reducing agent?
Explanation:
D (Incorrect):
The underlined CaO does not participate in a redox reaction. This is a neutralisation reaction where CaO reacts with H2O to form Ca(OH)2. No electron transfer occurs, so there is no reducing agent.
A (Correct):
The underlined CO is a reducing agent because it donates electrons to Fe2O3, reducing Fe3+ to metallic iron (Fe) while itself being oxidised to CO2.
B (Incorrect):
The underlined CO2 is not a reducing agent. Instead, carbon (C) in the reaction reduces CO2 to CO, making C the reducing agent, not CO2.
C (Incorrect):
The underlined CuO is not a reducing agent; it is an oxidising agent because it accepts electrons from H2, which is oxidised to H2O.
12. Ethanoic acid is made by reacting ethanol with acidified potassium manganate(VII).
Question: Which type of reaction occurs?
C (Correct):
In this reaction, ethanol (CH3CH2OH) is oxidised to ethanoic acid (CH3COOH) by the acidified potassium manganate(VII) (KMnO4). This involves the loss of hydrogen atoms from ethanol and the addition of oxygen to form ethanoic acid.
A (Incorrect):
A displacement reaction involves the replacement of one element by another in a compound. This does not occur in this reaction.
B (Incorrect):
Fermentation is a biological process where sugars are converted to ethanol and carbon dioxide by enzymes. This is unrelated to the chemical oxidation of ethanol to ethanoic acid.
D (Incorrect):
Neutralisation is a reaction between an acid and a base to form water and a salt. This reaction does not involve neutralisation but rather oxidation.
13. Sodium ions (Na+) and oxygen ions (O2−) combine with chromium ions to form Na2Cr2O7.
- Question: What is the oxidation state of chromium in this salt?
- A) +2
- B) +3
- C) +6
- D) +12
Explanation:
To determine the oxidation state of chromium (Cr) in Na2Cr2O7:
- Break down the formula:
- Sodium (Na+) contributes +1 for each ion, so +2 for two sodium ions.
- Oxygen (O2−) contributes −2 for each of the 7 oxygen atoms, totaling −14.
- Set up the charge balance: The overall charge of Na2Cr2O7 is neutral (0).
Let the oxidation state of each chromium ion be x: 2(+1) + 2x + 7(−2)=0 - Solve for x: 2 + 2x − 14 = 0
2x = 12x = +6 - Conclusion: Each chromium ion has an oxidation state of +6.
Why Other Options Are Incorrect:
- D (+12): Chromium cannot have an oxidation state this high in this compound.
- A (+2): This is too low and does not satisfy the charge balance.
- B (+3): Chromium in dichromate ions is not typically in the +3 state.
14. When chlorine gas dissolves in water:
- Reaction: Cl2 + H2O → HCl + HClO
- Question: Which row of the table identifies the oxidation number for chlorine in the chlorine-containing species?
| Cl2 | HCl | HClO |
|---|
| A: -1 | -1 | -1 |
| B: 0 | -1 | -1 |
| C: -1 | +1 | +1 |
| D: 0 | -1 | +1 |
Answer:
D: 0, -1, +1
Explanation:
- Chlorine gas (Cl2):
- Chlorine in molecular form (Cl2) has an oxidation number of 0 because it is in its elemental state.
- Hydrogen chloride (HCl):
- Chlorine in HCl has an oxidation number of -1, as it forms a chloride ion (Cl−) when bonded to hydrogen.
- Hypochlorous acid (HClOHClOHClO):
- In HClO, oxygen has an oxidation number of -2, hydrogen is +1, and the total charge of the molecule is 0. Therefore, chlorine has an oxidation number of +1 to balance the charges.
Why Other Options Are Incorrect:
- B (0, -1, -1): Incorrect because chlorine in HClO does not have an oxidation number of -1; it is +1.
- C (-1, +1, +1): Incorrect because chlorine in Cl2 is not -1; it is 0.
- A (-1, -1, -1): Incorrect because chlorine in Cl2 is in its elemental state with an oxidation number of 0.
15. Aqueous iron(III) chloride (FeCl3) reacts with aqueous potassium iodide (KIKIKI):
vFeCl3 + wKI → xFeCl2 + yKCl + I2
Which statements are correct?
- In the balanced equation, v, w, x, and y have the same value.
- Potassium iodide is an oxidising agent.
- A dark brown solution is produced in the reaction.
Options:
A: 1 and 2
B: 1 and 3
C: 2 only
D: 2 and 3
Answer:
B: 1 and 3
Explanation:
- Statement 1 (Correct):
In the balanced equation:FeCl3 + KI → FeCl2 + KCl + I2. The coefficients (v, w, x, and y) are all equal because for every one mole of FeCl3, one mole of KIKIKI is consumed, producing one mole each of FeCl2, KCl, and I2 . - Statement 2 (Incorrect):
Potassium iodide (KI) is a reducing agent, not an oxidising agent. It donates electrons to reduce Fe3+ in FeCl3 to Fe2+, while the iodide ions are oxidised to iodine (I2). - Statement 3 (Correct):
A dark brown solution is produced due to the formation of iodine (I2), which has a characteristic brown color in aqueous solution.
Why Other Options Are Incorrect:
- C (2 only): Incorrect because Statement 2 is wrong, and Statement 1 is correct.
- D (2 and 3): Incorrect because Statement 2 is wrong, though Statement 3 is correct.
- A (1 and 2): Incorrect because KIKIKI is not an oxidising agent.
Questions B
Question 1:
(a) Deduce the electronic configuration of magnesium. [1]
Answer:
(a) Electronic configuration of magnesium:
Magnesium has an atomic number of 12, so its electronic configuration is:
2,8,2
This represents 2 electrons in the first shell, 8 in the second shell, and 2 in the third shell.
(b) Magnesium can be produced by reducing magnesium oxide with barium:
MgO + Ba → Mg + BaO
Explain how this equation shows that magnesium oxide is reduced. [1]
(b) Explanation of reduction in the equation:
Reduction involves the gain of electrons or the loss of oxygen from a compound.
In the reaction:
MgO + Ba → Mg + BaO
- MgO: Magnesium is in the +2 oxidation state in magnesium oxide.
- Mg: Magnesium is in the elemental state (0 oxidation state).
- Conclusion: Magnesium is reduced because it goes from +2 in MgO to 0 in elemental magnesium.
Question 2:
The equation for the reaction of iron with steam is shown: 3Fe + 4H2O → Fe3O4 + 4H2
Describe how this equation shows that iron is oxidised.
Iron is oxidised in the reaction 3Fe+4H2O→Fe3O4+4H2 because it gains oxygen to form Fe3O4, an oxide of iron. This aligns with the classical definition of oxidation as the gain of oxygen. Additionally, the oxidation state of iron increases, indicating a loss of electrons, which also defines oxidation.
Question 3:
This question is about metals and metal compounds.
Iron is extracted in a blast furnace by reduction of iron(III) oxide (Fe2O3) with carbon monoxide.
Carbon monoxide is produced by the reaction of carbon with carbon dioxide:
C + CO2 → 2CO
(i) Explain how this equation shows that carbon dioxide is reduced. [1]
Additionally, the oxidation state of carbon in CO2 decreases from +4 to +2, confirming that carbon dioxide is reduced.
In the reaction C+CO2→2CO, carbon dioxide (CO2) gains oxygen from carbon to form carbon monoxide (CO).
Gaining oxygen is another way to define reduction.
(ii) Name the type of chemical reaction where oxidation and reduction take place simultaneously. [1]
Redox reaction.
Question 4:
Hydrogen sulfide burns in air to produce sulfur dioxide and water.
(i) Complete the chemical equation for this reaction:
….. H2S+…..O2→2H2O+2SO2 [2]
2H2S+3O2→2H2O+2SO2
(ii) Explain how this equation shows that hydrogen sulfide is oxidised. [1]
(ii) Explain how this equation shows that hydrogen sulfide is oxidised:
In H2S, sulfur has an oxidation state of −2.
In SO2, sulfur has an oxidation state of +4.
The increase in the oxidation state of sulfur from −2 to +4 indicates that hydrogen sulfide is oxidised.
Alternatively, oxidation is also defined as the gain of oxygen. In this reaction, sulfur in H2S gains oxygen to form SO2, confirming that it is oxidised.
Question 5:
Iron ore contains iron(III) oxide, Fe2O3. A blast furnace is used to extract iron from Fe2O3.
Equations for some reactions in the blast furnace are:
- C+O2→CO2
- CaCO3→CaO+CO2
- CaO+SiO2→CaSiO3
Iron(III) oxide in iron ore is converted to iron when it reacts with carbon monoxide (CO) in the blast furnace.
(i) Calculate the percentage by mass of iron in iron(III) oxide (Fe2O3):
- Percentage = …………………………% [2]
Molar mass of Fe2O3:
- Fe: 2×55.85=111.7 g/mol
- O: 3×16=48 g/mol
- Total = 111.7+48=159.7 g/mol
Percentage of iron:
Percentage = 70.0% (rounded to 1 decimal place)
(ii) State the type of reaction shown by equation 2 (CaCO3→CaO+CO2). [1]
Thermal decomposition.
(iii) State the name of the iron ore which consists mainly of iron(III) oxide. [1]
Haematite.
(iv) Describe how carbon monoxide is formed in the blast furnace. [1]
Carbon reacts with oxygen to form carbon dioxide:
C+O2→CO2 .
Carbon dioxide reacts with more carbon to form carbon monoxide:
CO2+C→2CO
(v) Write the symbol equation to show the reaction that occurs when iron(III) oxide is converted to iron in the blast furnace. [2]
Fe2O3+3CO→2Fe+3CO2
(vi) Name the chemical process which happens to iron when iron(III) oxide is converted to iron in the blast furnace. [1]
Reduction.
Question 6:
The symbols of the elements in Period 2 of the Periodic Table are shown:
Li, Be, B, C, N, O, F, Ne
Use the symbols of the elements in Period 2 to answer the questions that follow. Each symbol may be used once, more than once, or not at all.
(vi) Give the symbol of the element that only has an oxidation number of zero. [1]
The element that only has an oxidation number of zero is Ne (Neon).
Explanation:
It does not form compounds under normal conditions and therefore does not exhibit any oxidation states other than zero.
Neon is a noble gas and exists as an inert, monoatomic element.
Question 7:
This question is about sulfur and compounds of sulfur.
Sulfur is converted into sulfuric acid (H2SO4) by the Contact process. The process involves four stages:
- Stage 1: Molten sulfur is converted into sulfur dioxide.
- Stage 2: Sulfur dioxide reacts with oxygen to form sulfur trioxide.
- Stage 3: Sulfur trioxide combines with concentrated sulfuric acid to form oleum (H2S2O7).
- Stage 4: Oleum reacts to form concentrated sulfuric acid.
(a) (i) In stage 1, iron pyrites (FeS2) can be used instead of molten sulfur. The iron pyrites is heated strongly in air.
Balance the equation for the reaction occurring when iron pyrites reacts with oxygen in the air:
…..FeS2+…..O2→…..Fe2O3+…..SO2 [1]
Balance the equation: 4FeS2+11O2→2Fe2O3+8SO2
(ii) Name Fe2O3. Include the oxidation number of iron. [1]
- Iron(III) oxide
- The oxidation number of iron in Fe2O3 is +3.
Explanation:
- In Fe2O3, each oxygen has an oxidation state of −2.
- The total oxidation state for 3 O atoms is −6.
- To balance, the two iron atoms must have a total oxidation state of +6, making the oxidation state of each iron atom +3.
Question 8:
Ethanol is manufactured by two methods.
(c) The catalyst in method 2 is phosphoric acid, H3PO4. Dilute phosphoric acid is a weak acid that contains phosphate ions, PO43−.
(iii) Determine the oxidation number of phosphorus in the PO43−− ion. Show your working.
Oxidation number = ……………………. [2]
- Step 1: Write the formula and assign oxidation numbers
- The formula of the phosphate ion is PO43−.
- Oxygen (O) typically has an oxidation number of −2.
- Let the oxidation number of phosphorus (P) be xxx.
- Step 2: Set up the equation
- The total charge on the ion is −3.
- The contribution from 4 oxygen atoms is 4×−2=−8
- The equation for the total oxidation state is: x+(−8)=−3
- Step 3: Solve for x: x−8=−3
x=+5
Oxidation number = +5
Final Answer:
The oxidation nu mber of phosphorus in PO43− is +5.
Question 9:
Copper(II) oxide (CuO) undergoes the reaction: 4CuO→2Cu2
(i) Identify the changes in oxidation numbers of copper and oxygen in this reaction.
Explain in terms of changes in oxidation numbers why this is a redox reaction.
- Change in oxidation number of copper:
From +2 (in CuO) to +1 (in Cu2O). - Change in oxidation number of oxygen:
From -2 (in CuO) to 0 (in O2). - Explanation:
This is a redox reaction because:- Reduction: Copper is reduced as its oxidation number decreases from +2 (in CuO) to +1 (in Cu2O).
- Oxidation: Oxygen is oxidised as its oxidation number increases from −2 (in CuO) to 000 (in O2).
The simultaneous oxidation and reduction of different species confirm this is a redox reaction.