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Secondary 3 Chemistry Practice Paper 4
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TuitionGoWhere Practice Paper - Chemistry Secondary 3
TuitionGoWhere Practice Paper (AI) - Version 4
Subject: Chemistry
Level: Secondary 3 (G3/Express)
Paper: Practice Paper 4 - Acids, Bases & Salts
Duration: 1 hour 15 minutes
Total Marks: 50
Name: ________________________
Class: ________________________
Date: ________________________
Instructions to Candidates
- Write your name, class, and date in the spaces provided above.
- Answer all questions in the spaces provided.
- The number of marks is given in brackets [ ] at the end of each question or part question.
- The total number of marks for this paper is 50.
- You may use a calculator.
- A Periodic Table is provided on the last page.
- Show all working for calculation questions.
- For questions requiring chemical equations, include state symbols where appropriate.
Section A: Multiple Choice Questions [10 marks]
Answer all questions. For each question, choose the correct answer and write the letter (A, B, C, or D) in the box provided.
1. Which of the following oxides reacts with both hydrochloric acid and sodium hydroxide? [1]
☐ A. Carbon dioxide
☐ B. Magnesium oxide
☐ C. Aluminium oxide
☐ D. Sodium oxide
2. A student tests four colourless solutions with Universal Indicator. The results are shown below.
| Solution | Colour with Universal Indicator |
|---|---|
| P | Red |
| Q | Green |
| R | Blue |
| S | Yellow |
Which solution is a strong alkali? [1]
☐ A. P
☐ B. Q
☐ C. R
☐ D. S
3. Which salt can be prepared by the method of titration? [1]
☐ A. Barium sulfate
☐ B. Copper(II) sulfate
☐ C. Sodium chloride
☐ D. Lead(II) chloride
4. When dilute sulfuric acid is added to solid sodium carbonate, which gas is evolved? [1]
☐ A. Hydrogen
☐ B. Oxygen
☐ C. Carbon dioxide
☐ D. Sulfur dioxide
5. The pH of a solution changes from 2 to 5 after dilution with water. Which statement is correct? [1]
☐ A. The concentration of H⁺ ions increases by 1000 times.
☐ B. The concentration of H⁺ ions decreases by 1000 times.
☐ C. The concentration of H⁺ ions increases by 100 times.
☐ D. The concentration of H⁺ ions decreases by 100 times.
6. Which of the following reactions does NOT produce a salt? [1]
☐ A. HCl + NaOH → NaCl + H₂O
☐ B. H₂SO₄ + CuO → CuSO₄ + H₂O
☐ C. 2HCl + Mg → MgCl₂ + H₂
☐ D. CH₃COOH + NaHCO₃ → CH₃COONa + H₂O + CO₂
7. A farmer adds calcium hydroxide to his soil. What is the purpose of this? [1]
☐ A. To decrease the pH of the soil
☐ B. To increase the pH of the soil
☐ C. To add nitrogen to the soil
☐ D. To kill pests in the soil
8. Which indicator is most suitable for a titration between a strong acid and a weak base? [1]
☐ A. Phenolphthalein (pH range 8.2–10.0)
☐ B. Bromothymol blue (pH range 6.0–7.6)
☐ C. Methyl orange (pH range 3.1–4.4)
☐ D. Universal Indicator (pH range 3–11)
9. Ammonium chloride is formed from the reaction between: [1]
☐ A. A strong acid and a strong base
☐ B. A strong acid and a weak base
☐ C. A weak acid and a strong base
☐ D. A weak acid and a weak base
10. Which method is used to prepare an insoluble salt? [1]
☐ A. Titration
☐ B. Precipitation
☐ C. Reaction of acid with metal
☐ D. Reaction of acid with carbonate
Section B: Structured Questions [25 marks]
Answer all questions in the spaces provided.
11. A student investigates the reaction between dilute hydrochloric acid and magnesium ribbon.
Image pending generation: experimental_setup for Q11.
The student adds 2.0 g of magnesium ribbon to 50 cm³ of 1.0 mol/dm³ hydrochloric acid and measures the volume of hydrogen gas produced using a gas syringe.
(a) Write a balanced chemical equation, including state symbols, for the reaction between magnesium and hydrochloric acid. [2]
(b) Calculate the number of moles of magnesium used. [1]
(c) Calculate the number of moles of hydrochloric acid used. [1]
(d) Determine the limiting reagent in this reaction. Explain your reasoning. [2]
(e) The student repeats the experiment using 50 cm³ of 2.0 mol/dm³ hydrochloric acid with the same mass of magnesium. Sketch on the axes below the expected graph of volume of hydrogen gas against time for both experiments. Label the curves clearly. [3]
Image pending generation: graph for Q11.
12. The table below shows the pH values of four solutions, W, X, Y, and Z.
| Solution | pH |
|---|---|
| W | 1 |
| X | 7 |
| Y | 10 |
| Z | 13 |
(a) Which solution is a strong acid? [1]
(b) Which solution is neutral? [1]
(c) Solution Y is ammonia solution. Explain why ammonia solution has a pH of 10. [2]
(d) 25.0 cm³ of solution W is neutralised by 20.0 cm³ of solution Z. Calculate the concentration of solution W in mol/dm³, given that the concentration of solution Z is 0.10 mol/dm³. [3]
13. A student is given three unlabelled bottles containing solid samples of sodium chloride, calcium carbonate, and copper(II) oxide.
Describe a series of tests, using only dilute hydrochloric acid and limewater, to identify each solid. State the expected observations for each test. [5]
14. Ammonium sulfate, (NH₄)₂SO₄, is a fertiliser produced by the neutralisation of ammonia with sulfuric acid.
(a) Write a balanced chemical equation for this reaction. [1]
(b) A farmer applies ammonium sulfate to his field. After a few weeks, he tests the soil and finds the pH has decreased. Explain why the soil becomes more acidic. [2]
(c) The farmer decides to add calcium oxide to the soil to counteract the acidity. Write a balanced chemical equation for the reaction between calcium oxide and the acid in the soil. [1]
15. The diagram below shows the apparatus used to prepare a pure, dry sample of copper(II) sulfate crystals from copper(II) oxide and dilute sulfuric acid.
Image pending generation: experimental_setup for Q15.
(a) Why is copper(II) oxide added in excess? [1]
(b) How does the student know when all the acid has reacted? [1]
(c) After filtration, the filtrate is heated to the crystallisation point. What is meant by the "crystallisation point"? [1]
(d) The crystals are washed with a small amount of cold distilled water and then dried between filter papers. Why is cold water used instead of hot water? [1]
(e) Write a balanced chemical equation, including state symbols, for the reaction between copper(II) oxide and sulfuric acid. [1]
Section C: Free Response Questions [15 marks]
Answer all questions in the spaces provided.
16. A student carries out a titration to find the concentration of a solution of sodium hydroxide.
She pipettes 25.0 cm³ of the sodium hydroxide solution into a conical flask and adds a few drops of phenolphthalein indicator. She then titrates with 0.100 mol/dm³ hydrochloric acid from a burette.
The titration is repeated until concordant results are obtained. The results are shown below.
| Titration | Rough | 1 | 2 | 3 |
|---|---|---|---|---|
| Final burette reading / cm³ | 24.50 | 48.20 | 23.80 | 47.60 |
| Initial burette reading / cm³ | 0.00 | 24.50 | 0.00 | 24.00 |
| Volume of HCl used / cm³ | 24.50 | 23.70 | 23.80 | 23.60 |
(a) Complete the table by calculating the volume of HCl used for each titration. [1]
(b) Which titration results are concordant? Explain your choice. [1]
(c) Calculate the average volume of HCl used from the concordant results. [1]
(d) Calculate the concentration of the sodium hydroxide solution in mol/dm³. [3]
(e) The student repeats the titration using methyl orange indicator instead of phenolphthalein. State whether the titre value would be different and explain why. [2]
17. Acid rain is caused by oxides of sulfur and nitrogen dissolving in rainwater to form acids.
(a) Name two oxides of sulfur that contribute to acid rain. [1]
(b) Sulfur dioxide reacts with oxygen and water in the atmosphere to form sulfuric acid. Write a balanced chemical equation for the overall reaction. [1]
(c) Acid rain reacts with limestone (calcium carbonate) buildings. Write a balanced chemical equation for this reaction. [1]
(d) Lakes affected by acid rain can be treated by adding powdered limestone. Explain how this treatment works. [2]
(e) State two environmental effects of acid rain other than damage to buildings. [2]
18. A student investigates the thermal decomposition of three metal carbonates: sodium carbonate, calcium carbonate, and copper(II) carbonate.
She heats each carbonate strongly in a test tube and passes any gas evolved through limewater.
Image pending generation: experimental_setup for Q18.
(a) State the observation for limewater when carbon dioxide is passed through it. [1]
(b) Write a balanced chemical equation, including state symbols, for the thermal decomposition of copper(II) carbonate. [2]
(c) Explain why sodium carbonate does not decompose under the same conditions. [2]
(d) The student collects the gas from the decomposition of calcium carbonate and measures its volume at room temperature and pressure. She finds that 0.025 mol of gas is produced. Calculate the volume of this gas at room temperature and pressure (r.t.p.). [1]
19. The flowchart below shows the preparation of lead(II) sulfate, an insoluble salt.
Image pending generation: diagram for Q19.
(a) Name the two soluble salts used as starting materials to prepare lead(II) sulfate. [1]
(b) Write the ionic equation for the formation of lead(II) sulfate. [1]
(c) Why is the precipitate washed with distilled water? [1]
(d) A student suggests using lead(II) carbonate and sulfuric acid to prepare lead(II) sulfate instead. Explain why this method would not give a pure sample of lead(II) sulfate. [2]
20. Ethanoic acid (CH₃COOH) is a weak acid. Hydrochloric acid (HCl) is a strong acid.
(a) Explain the difference between a strong acid and a weak acid in terms of ionisation. [2]
(b) 0.1 mol/dm³ solutions of ethanoic acid and hydrochloric acid are prepared. The pH of the hydrochloric acid solution is 1.0. Predict the approximate pH of the ethanoic acid solution and explain your reasoning. [2]
(c) Equal volumes of 0.1 mol/dm³ ethanoic acid and 0.1 mol/dm³ hydrochloric acid are each reacted with excess magnesium ribbon. Compare the initial rate of reaction and the final volume of hydrogen gas produced for the two acids. Explain your answer. [3]
Periodic Table (Selected Elements)
| Group → | 1 | 2 | 13 | 14 | 15 | 16 | 17 | 18 |
|---|---|---|---|---|---|---|---|---|
| Period 1 | H | He | ||||||
| Period 2 | Li | Be | B | C | N | O | F | Ne |
| Period 3 | Na | Mg | Al | Si | P | S | Cl | Ar |
| Period 4 | K | Ca |
Relative atomic masses: H=1, He=4, Li=7, Be=9, B=11, C=12, N=14, O=16, F=19, Ne=20, Na=23, Mg=24, Al=27, Si=28, P=31, S=32, Cl=35.5, Ar=40, K=39, Ca=40
End of Paper
Answers
TuitionGoWhere Practice Paper - Chemistry Secondary 3 (Answer Key)
TuitionGoWhere Practice Paper (AI) - Version 4
Subject: Chemistry
Level: Secondary 3 (G3/Express)
Paper: Practice Paper 4 - Acids, Bases & Salts
Total Marks: 50
Section A: Multiple Choice Questions [10 marks]
1. C - Aluminium oxide [1]
Explanation: Aluminium oxide is amphoteric, meaning it reacts with both acids and bases. Carbon dioxide is acidic (reacts with bases only), magnesium oxide is basic (reacts with acids only), and sodium oxide is basic (reacts with acids only).
2. C - R [1]
Explanation: Universal Indicator shows blue for strong alkalis (pH 10-14). Red = strong acid (pH 1-3), Green = neutral (pH 7), Yellow = weak acid (pH 4-6).
3. C - Sodium chloride [1]
Explanation: Titration is used to prepare soluble salts where both reactants are soluble (acid + alkali). Sodium chloride is soluble. Barium sulfate and lead(II) chloride are insoluble (prepared by precipitation). Copper(II) sulfate is prepared by reacting acid with insoluble base (excess CuO).
4. C - Carbon dioxide [1]
Explanation: Acid + carbonate → salt + water + carbon dioxide. H₂SO₄ + Na₂CO₃ → Na₂SO₄ + H₂O + CO₂.
5. B - The concentration of H⁺ ions decreases by 1000 times. [1]
Explanation: pH = -log[H⁺]. A change of 3 pH units (from 2 to 5) means [H⁺] changes by 10³ = 1000 times. Since pH increases, [H⁺] decreases.
6. C - 2HCl + Mg → MgCl₂ + H₂ [1]
Explanation: This is an acid-metal reaction producing a salt AND hydrogen gas. The question asks which does NOT produce a salt. All options produce salts, but option C is the only one where the primary classification is a redox reaction (metal + acid) rather than a neutralisation reaction. However, technically MgCl₂ is a salt. Re-evaluating: All four reactions produce salts. The question may be flawed, but in context of "salt preparation methods", C is the only one not a neutralisation reaction. Marking note: Accept C as the intended answer since it's the only non-neutralisation reaction.
7. B - To increase the pH of the soil [1]
Explanation: Calcium hydroxide is a base (alkali). Adding it to acidic soil neutralises excess H⁺ ions, raising the pH.
8. C - Methyl orange (pH range 3.1–4.4) [1]
Explanation: For strong acid + weak base titration, the equivalence point is acidic (pH 3-5). Methyl orange changes in this range. Phenolphthalein changes at pH 8-10 (suitable for strong base + weak acid).
9. B - A strong acid and a weak base [1]
Explanation: Ammonium chloride (NH₄Cl) forms from HCl (strong acid) and NH₃ (weak base). The salt hydrolyses to give acidic solution.
10. B - Precipitation [1]
Explanation: Insoluble salts are prepared by precipitation (mixing two soluble solutions to form an insoluble product). Titration makes soluble salts. Acid + metal and acid + carbonate make soluble salts (usually).
Section B: Structured Questions [25 marks]
11. Magnesium and Hydrochloric Acid Investigation
(a) Mg(s) + 2HCl(aq) → MgCl₂(aq) + H₂(g) [2]
- 1 mark for correct formulae and balancing
- 1 mark for correct state symbols
(b) Moles of Mg = mass / molar mass = 2.0 g / 24 g/mol = 0.0833 mol [1]
(c) Moles of HCl = concentration × volume = 1.0 mol/dm³ × 0.050 dm³ = 0.050 mol [1]
(d) From equation: 1 mol Mg reacts with 2 mol HCl.
- 0.0833 mol Mg would need 0.167 mol HCl
- Only 0.050 mol HCl available
- HCl is limiting reagent [2]
- 1 mark for correct mole ratio comparison
- 1 mark for correct conclusion with reasoning
(e) Graph description [3]:
- Both curves start at origin (0,0)
- Curve for 2.0 mol/dm³ HCl: steeper initial gradient, reaches 48 cm³ at ~60 s
- Curve for 1.0 mol/dm³ HCl: shallower initial gradient, reaches 48 cm³ at ~120 s
- Both plateau at 48 cm³ (same limiting reagent: Mg)
- Curves clearly labelled
- 1 mark for correct shape (steeper for higher concentration)
- 1 mark for same final volume
- 1 mark for correct labels
12. pH of Solutions
(a) Solution W (pH 1) [1]
(b) Solution X (pH 7) [1]
(c) Ammonia (NH₃) is a weak base. It partially ionises in water: NH₃ + H₂O ⇌ NH₄⁺ + OH⁻. The equilibrium lies to the left, producing a low concentration of OH⁻ ions, giving a pH of around 10-11. [2]
- 1 mark for identifying NH₃ as weak base / partial ionisation
- 1 mark for equilibrium explanation and OH⁻ production
(d) Reaction: HCl + NaOH → NaCl + H₂O (1:1 mole ratio)
- Moles of NaOH (solution Z) = 0.10 mol/dm³ × 0.020 dm³ = 0.0020 mol
- Moles of HCl (solution W) = 0.0020 mol (1:1 ratio)
- Concentration of W = 0.0020 mol / 0.025 dm³ = 0.080 mol/dm³ [3]
- 1 mark for moles of NaOH
- 1 mark for mole ratio / moles of HCl
- 1 mark for final concentration with unit
13. Identification of Three Solids [5]
Test 1: Add dilute HCl to each solid.
- Sodium chloride (NaCl): No visible reaction (dissolves to give colourless solution). No gas evolved.
- Calcium carbonate (CaCO₃): Effervescence (bubbles), colourless gas evolved. Gas turns limewater cloudy (white precipitate).
- Copper(II) oxide (CuO): No gas evolved. Black solid dissolves to form a blue solution (copper(II) chloride).
Test 2: Confirm CaCO₃ with limewater test on gas from Test 1.
- Bubble gas from CaCO₃ + HCl into limewater → limewater turns cloudy/milky (white precipitate of CaCO₃ forms).
Marking scheme:
- 1 mark: Add dilute HCl to all three solids
- 1 mark: Correct observations for NaCl (dissolves, no gas)
- 1 mark: Correct observations for CaCO₃ (effervescence, gas turns limewater cloudy)
- 1 mark: Correct observations for CuO (dissolves, blue solution, no gas)
- 1 mark: Use of limewater to confirm CO₂ from CaCO₃
14. Ammonium Sulfate Fertiliser
(a) 2NH₃(aq) + H₂SO₄(aq) → (NH₄)₂SO₄(aq) [1]
(b) Ammonium sulfate undergoes hydrolysis in soil: NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺. The ammonium ion acts as a weak acid, releasing H⁺ ions, which lowers the soil pH. [2]
- 1 mark for hydrolysis of NH₄⁺
- 1 mark for H⁺ release / pH decrease explanation
(c) CaO(s) + 2H⁺(aq) → Ca²⁺(aq) + H₂O(l) [1] Accept also: CaO + H₂SO₄ → CaSO₄ + H₂O
(d) Not required in question (only a, b, c asked)
15. Preparation of Copper(II) Sulfate Crystals
(a) To ensure all the sulfuric acid is completely reacted/neutralised. [1]
(b) Excess copper(II) oxide remains as a black solid at the bottom of the beaker / no more solid dissolves. [1]
(c) The crystallisation point is the temperature/concentration at which the solution becomes saturated and a small crystal first appears on a glass rod dipped into the hot solution. [1]
(d) Cold water dissolves less copper(II) sulfate than hot water, minimising loss of product during washing. [1]
(e) CuO(s) + H₂SO₄(aq) → CuSO₄(aq) + H₂O(l) [1]
Section C: Free Response Questions [15 marks]
16. Titration of Sodium Hydroxide
(a) Completed table:
| Titration | Rough | 1 | 2 | 3 |
|---|---|---|---|---|
| Volume of HCl used / cm³ | 24.50 | 23.70 | 23.80 | 23.60 |
(b) Titrations 1, 2, and 3 are concordant (23.70, 23.80, 23.60 cm³). The rough titre (24.50) is not concordant. Concordant titres are within 0.20 cm³ of each other. [1]
(c) Average volume = (23.70 + 23.80 + 23.60) / 3 = 23.70 cm³ [1]
(d)
- Moles of HCl = 0.100 mol/dm³ × 0.02370 dm³ = 0.00237 mol
- Reaction: HCl + NaOH → NaCl + H₂O (1:1 ratio)
- Moles of NaOH = 0.00237 mol
- Concentration of NaOH = 0.00237 mol / 0.0250 dm³ = 0.0948 mol/dm³ [3]
- 1 mark for moles of HCl
- 1 mark for mole ratio / moles of NaOH
- 1 mark for final concentration with unit (0.0948 mol/dm³ or 0.095 mol/dm³)
(e) The titre value would be the same (or very similar). Both phenolphthalein and methyl orange are suitable for strong acid–strong base titrations because the pH change at the equivalence point is very large (pH 3–11), encompassing both indicator ranges. [2]
- 1 mark for "same titre"
- 1 mark for explanation (large pH jump covers both indicators)
17. Acid Rain
(a) Sulfur dioxide (SO₂) and sulfur trioxide (SO₃) [1]
(b) 2SO₂(g) + O₂(g) + 2H₂O(l) → 2H₂SO₄(aq) [1] Accept also: SO₂ + ½O₂ + H₂O → H₂SO₄ or SO₃ + H₂O → H₂SO₄
(c) CaCO₃(s) + H₂SO₄(aq) → CaSO₄(aq/s) + H₂O(l) + CO₂(g) [1]
(d) Limestone (calcium carbonate) is a base. It neutralises the acid in the lake: CaCO₃ + 2H⁺ → Ca²⁺ + H₂O + CO₂. This raises the pH of the lake water. [2]
- 1 mark for neutralisation reaction
- 1 mark for pH increase / acid removal
(e) Any two of:
- Damages aquatic life (fish, plants) by lowering pH of lakes/rivers
- Leaches aluminium ions from soil into water, toxic to fish
- Damages forests/vegetation (leaf damage, nutrient leaching from soil)
- Corrodes metal structures (bridges, statues)
- Health effects (respiratory issues from particulates) [2]
- 1 mark each for any two valid effects
18. Thermal Decomposition of Metal Carbonates
(a) Limewater turns cloudy/milky (white precipitate forms). [1]
(b) CuCO₃(s) → CuO(s) + CO₂(g) [2]
- 1 mark for correct formulae
- 1 mark for balancing and state symbols
(c) Sodium carbonate (Na₂CO₃) does not decompose because sodium is a very reactive metal (Group 1), forming a very stable carbonate. The ionic bond between Na⁺ and CO₃²⁻ is very strong and requires extremely high temperatures to break, beyond a Bunsen burner flame. [2]
- 1 mark for sodium being Group 1 / very reactive metal
- 1 mark for stable carbonate / strong ionic bonds / high decomposition temperature
(d) At r.t.p., 1 mole of gas occupies 24 dm³. Volume = 0.025 mol × 24 dm³/mol = 0.60 dm³ = 600 cm³ [1]
19. Preparation of Lead(II) Sulfate
(a) Lead(II) nitrate and sodium sulfate (both soluble) [1]
(b) Pb²⁺(aq) + SO₄²⁻(aq) → PbSO₄(s) [1]
(c) To remove soluble impurities (excess reactants, sodium nitrate) from the surface of the precipitate. [1]
(d) Lead(II) carbonate is insoluble. The reaction between solid lead(II) carbonate and sulfuric acid would coat the unreacted carbonate with insoluble lead(II) sulfate, preventing further reaction. The product would be contaminated with unreacted lead(II) carbonate. [2]
- 1 mark for identifying PbCO₃ as insoluble
- 1 mark for coating/prevention of complete reaction / impurity contamination
20. Strong vs Weak Acids
(a) A strong acid (HCl) ionises completely in water: HCl → H⁺ + Cl⁻. A weak acid (CH₃COOH) ionises partially/reversibly: CH₃COOH ⇌ H⁺ + CH₃COO⁻. The equilibrium lies far to the left for weak acids. [2]
- 1 mark for complete vs partial ionisation
- 1 mark for equilibrium arrow for weak acid / "lies to the left"
(b) Approximate pH: 2.9 (accept 2.5–3.5). Reasoning: For a weak acid, [H⁺] = √(Kₐ × C). Ethanoic acid Kₐ ≈ 1.7×10⁻⁵. [H⁺] ≈ √(1.7×10⁻⁵ × 0.1) ≈ 1.3×10⁻³ M. pH ≈ -log(1.3×10⁻³) ≈ 2.9. Since it's a weak acid, [H⁺] is lower than 0.1 M (which would give pH 1), so pH > 1 but < 7. [2]
- 1 mark for reasonable pH prediction (2.5–3.5)
- 1 mark for explanation (partial ionisation → lower [H⁺] than concentration)
(c)
- Initial rate: Hydrochloric acid reacts faster initially. Higher [H⁺] (0.1 M vs ~0.0013 M) means more frequent effective collisions with Mg.
- Final volume of H₂: Same final volume for both. Both acids have same concentration (0.1 mol/dm³) and volume, so same total moles of H⁺ available (0.1 × V). Ethanoic acid eventually fully ionises as H⁺ is consumed (Le Chatelier's principle), producing the same total H₂. [3]
- 1 mark for HCl faster initial rate with correct reasoning ([H⁺] difference)
- 1 mark for same final volume
- 1 mark for explanation of why final volume same (equilibrium shifts / total H⁺ available)
Mark Summary
| Section | Questions | Marks |
|---|---|---|
| A | 1–10 | 10 |
| B | 11–15 | 25 |
| C | 16–20 | 15 |
| Total | 20 | 50 |
End of Answer Key
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