The shape of the lesson
By the end of the lesson, learners can
- Acid rain: Burn a high-sulfur fuel in the simulator and follow the sulfur through the atmosphere to the rain, then watch it react with limestone.
- The Contact process: Run the Contact process converter in the simulator and vary the conditions.
- Source: sulfur in fossil fuels.
- Acid rain: SO2 → SO3 → H2SO4.
- Remove: CaCO3 + SO2 → CaSO3 + CO2.
- Contact: V2O5, 450 °C, 1 to 2 atm.
One picture every idea returns to
Fossil fuels contain sulfur, which burns to sulfur dioxide (S +4). In the air SO2 is oxidised to SO3 (S +6), helped by the NO2 catalyst, and dissolves in rain to form sulfuric acid, a major cause of acid rain. Acid rain damages limestone buildings, lakes, fish and trees. SO2 can be removed from power-station flue gases by reacting it with calcium carbonate or calcium oxide. In small amounts SO2 is a food preservative, because it is an antioxidant and kills bacteria. Industrially, SO2 is converted to SO3 in the Contact process to make sulfuric acid.
Forty-five minutes, phase by phase
| Time | Phase | What happens in the room | Grouping |
|---|---|---|---|
| 0 to 5 min | Hook: predict | Open with the question in the title, Sulfur up the chimney.. Coal and oil contain traces of sulfur. Put the figure on the board: it is the one picture every idea in the lesson returns to. | Whole class, then pairs |
| 5 to 17 min | Teach and check: Acid rain | Burn a high-sulfur fuel in the simulator and follow the sulfur through the atmosphere to the rain, then watch it react with limestone. Project tab 4 of the lab (Nitrogen and sulfur lab, acid rain) and let a student drive it while the class predicts. | Whole class, then pairs |
| 17 to 30 min | Teach and check: The Contact process | Run the Contact process converter in the simulator and vary the conditions. Project tab 5 of the lab (Nitrogen and sulfur lab, Contact process) and let a student drive it while the class predicts. | Anticipation / Reaction Guide |
| 30 to 36 min | Check | The four Stage 1 questions from the student lesson, answered on mini whiteboards or by show of fingers. Re-teach the one most get wrong. | Individual, then whole class |
| 36 to 43 min | Plenary | The exam-style questions from Stage 2, one at a time: thirty seconds alone, then a random call. The feedback lines in the key give the examiner's wording. | Individual, random call |
| 43 to 45 min | Exit and homework | Set the worksheet below (6 questions, 14 marks). Learners self-assess against the objectives before they leave. | Individual |
Protect the parts that carry the learning
The sequence carries about fifty minutes of material, so the lab is driven by one student while the class predicts rather than explored freely. Times are a guide.
Protected: the check at 30 minutes. It is the only individual read on whether the key idea landed, and it is not cut.
Cut first if the build overruns: the last check in The Contact process, never the check or the plenary.
In a 60 minute block: run worksheet questions 1 to 3 in class as a Rally Coach set and model both the hardest worked example and its check live.
Anticipation / Reaction Guide
Before the lesson, students mark five statements true or false. After the teaching they mark them again and justify any change of mind in one sentence each. The statements come from the examiner traps below.
The full structure, with its variations and the other twelve, is on the cooperative-learning page.
What to head off, and how
| Trap learners fall into | Teaching move that pre-empts it |
|---|---|
| Acid rain is caused by sulfur and nitrogen oxides, not by carbon dioxide; rain is naturally slightly acidic (about pH 5.6) from dissolved CO 2, but SO 2 and NO 2 push it much lower. | Remember the oxidation number of sulfur: +4 in SO 2 and H 2 SO 3, +6 in SO 3 and H 2 SO 4. |
Sulfur dioxide and acid rain: worksheet
6 questions, 14 marks. Written in the style of the papers; no past paper question is reproduced.
Write equations to show how sulfur in coal leads to sulfuric acid in rain.
Give two harmful effects of acid rain.
Calculate the mass of SO2 produced when 500 tonnes of oil containing 3.00% sulfur by mass is burned. (Ar S 32.1; Mr SO2 64.1)
Calculate the mass of calcium carbonate needed to absorb 32.0 kg of SO2 in a desulfurisation unit. (Mr: CaCO3 100.1, SO2 64.1)
Explain why sulfur dioxide is used in small amounts as a preservative in dried fruit.
State the conditions of the Contact process and explain why a low pressure is used even though high pressure would increase the yield of SO3.
Answer key and marking notes
- S + O2 → SO2.
- 2SO2 + O2 → 2SO3 (catalysed by NO2).
- SO3 + H2O → H2SO4.
- Damages limestone and marble buildings and statues.
- Lowers the pH of lakes, killing fish; damages trees by leaching nutrients from soil.
- Mass of S = 0.0300 × 500 = 15.0 tonnes.
- Mass of SO2 = 15.0 × 64.1 ÷ 32.1 = 30.0 tonnes.
- n(SO2) = 32 000 ÷ 64.1 = 499 mol = n(CaCO3).
- Mass = 499 × 100.1 = 50 000 g = 50.0 kg.
- It is an antioxidant, preventing the food from being oxidised and discoloured.
- It inhibits the growth of bacteria and fungi.
- V2O5 catalyst, about 450 °C, 1 to 2 atm.
- The equilibrium already lies far to the right (about 99% conversion), so extra pressure gives little benefit.
- High-pressure equipment would be expensive.
8 multiple-choice questions on this lesson
From the topic question bank. Assign them in the app for a marked score in your gradebook, or run the topic as a class game.
What is the main source of the sulfur dioxide that causes acid rain?
- Aburning fossil fuels such as coal and oil that contain sulfur compounds
- Bcarbon dioxide from burning fuels, which dissolves in rain to form a strong acid
- Cnitrogen and oxygen reacting in car engines
- Dsulfur released from fertilisers containing ammonium sulfate
The diagram shows how sulfur in fuel becomes acid rain. In which step is sulfur oxidised from +4 to +6?
- AH2SO3 is oxidised (for example by O2) to H2SO4
- BSO2 + H2O → H2SO3
- Csulfur in the fuel burns to SO2
- Dthe acid rain reacts with limestone
Acid rain slowly dissolves buildings made of limestone. Which equation shows this?
- ACaCO3 + H2SO4 → CaSO4 + H2O + CO2
- BCaCO3 + H2SO4 → CaSO4 + H2
- CCaCO3 + SO2 → CaSO3 + CO2, because limestone reacts mainly with CO2
- DCaO + H2O → Ca(OH)2
Sulfur dioxide is added in small amounts to some foods and drinks, such as dried fruit and wine. Why?
- Ait is a preservative: it acts as an antioxidant and inhibits the growth of bacteria and fungi
- Bit adds sulfur as an essential nutrient
- Cit acts as an oxidising agent, speeding up the ripening of the fruit
- Dit neutralises acids in the food, making it taste less sour
What are the oxidation numbers of sulfur in SO2, SO3, H2SO4 and H2S, in that order?
- A+4, +6, +6, −2
- B+2, +3, +4, −1
- C+4, +6, +8, +2
- D+4, +6, +4, −2
In the Contact process, 2SO2 + O2 ⇌ 2SO3 (forward reaction exothermic). Which set of conditions is used, and why is the pressure only 1 to 2 atm?
- AV2O5 catalyst, about 450 °C, 1 to 2 atm; the conversion is already about 99% at low pressure, so high pressure is not worth the extra cost
- Biron catalyst, about 450 °C, 200 atm; high pressure is essential because there are fewer gas moles on the right
- CV2O5 catalyst, about 1000 °C, 1 to 2 atm; a high temperature gives the highest yield
- Dno catalyst, room temperature, 1 atm; SO2 reacts rapidly with oxygen on its own
Flue gas desulfurisation removes SO2 from power-station waste gases. Which statement is correct?
- Apowdered CaCO3 (or CaO) reacts with the acidic SO2: CaCO3 + SO2 → CaSO3 + CO2
- Bthe waste gas is passed through acid, which neutralises the SO2
- CCaCO3 + SO2 → CaSO3 + CO2, which also removes all the CO2 from the gas
- Dthe SO2 is oxidised to sulfur, which is then released harmlessly into the air
A power station produces 1.00 tonne of SO2. What mass of calcium carbonate is needed to remove it all? CaCO3 + SO2 → CaSO3 + CO2 (Mr SO2 = 64.1, CaCO3 = 100.1)
- A1.56 tonne
- B0.640 tonne
- C3.12 tonne
- D1.00 tonne
Quiz key
Coal and crude oil contain sulfur compounds; when the fuel burns, the sulfur is oxidised: S + O2 → SO2. Power stations burning coal are a major source. Blaming CO2 is the classic error: dissolved CO2 makes rain only weakly acidic (about pH 5.6), and it is SO2 and nitrogen oxides that make acid rain. Engines form NO, a nitrogen oxide, not sulfur dioxide. The sulfate ion in ammonium sulfate is stable and is not converted into SO2 in the soil.
Sulfur is +4 in both SO2 and H2SO3 and +6 in H2SO4, so the +4 to +6 oxidation is the step H2SO3 → H2SO4 (2H2SO3 + O2 → 2H2SO4). Dissolving SO2 in water is an acid-base reaction with no change in oxidation number (+4 to +4). Burning the fuel is an oxidation, but sulfur goes from a low state (for example 0 or −2) to +4, not from +4 to +6. Attack on limestone is an acid-carbonate neutralisation; the sulfur stays +6 as SO42−.
Limestone is calcium carbonate; acid rain contains sulfuric acid, so CaCO3 + H2SO4 → CaSO4 + H2O + CO2. The calcium sulfate is washed away or flakes off, so the stone erodes. Carbonates give CO2 with acids, not hydrogen; hydrogen comes from reactive metals. CaCO3 + SO2 is the flue-gas desulfurisation reaction at a power station, not the action of rain, and the reason given is wrong. Calcium oxide reacting with water describes making slaked lime, not the effect of acid rain.
SO2 (and sulfites) preserve food: as a reducing agent it is oxidised in preference to the food, acting as an antioxidant that stops browning and spoilage, and it inhibits the growth of microorganisms. It is not added as a nutrient; some people are sensitive to it, so its use is controlled. In this role it acts as a reducing agent, not an oxidising agent, and it is used to slow down changes, not speed up ripening. SO2 is an acidic oxide, so it cannot neutralise acids.
In SO2: S + 2(−2) = 0, S = +4. In SO3: S = +6. In H2SO4: 2(+1) + S + 4(−2) = 0, S = +6. In H2S: 2(+1) + S = 0, S = −2. The second row treats oxygen as −1 and gives hydrogen sulfide the wrong value. +8 ignores the +1 hydrogens in H2SO4 (and is above the maximum of +6 for sulfur), and +2 for H2S has the wrong sign. +4 for H2SO4 confuses it with sulfurous acid, H2SO3.
The Contact process uses a vanadium(V) oxide catalyst at about 450 °C and 1 to 2 atm. Fewer gas moles on the right means high pressure would favour SO3, but the equilibrium yield at low pressure is already about 99%, so the cost of high-pressure plant is not justified. An iron catalyst and 200 atm are the Haber process conditions. Raising the temperature shifts an exothermic equilibrium to the left, so 1000 °C would lower the yield; 450 °C is a compromise between rate and yield. Without a catalyst SO2 is oxidised only very slowly; in the atmosphere NO2 acts as the catalyst.
SO2 is an acidic oxide, so it reacts with a base: CaCO3 + SO2 → CaSO3 + CO2 (or CaO + SO2 → CaSO3). The CaSO3 can be oxidised to CaSO4 (gypsum) and sold for plasterboard. An acid cannot neutralise another acidic gas. Using limestone actually releases CO2, it does not remove it; removing CO2 is the classic misconception. Going from SO2 (+4) to sulfur (0) is reduction, not oxidation, and sulfur is not released into the air.
The ratio CaCO3 : SO2 is 1 : 1, so mass of CaCO3 = 1.00 × 100.1 / 64.1 = 1.56 tonne (1.00 tonne = 1.00 × 106 g, so n(SO2) = 1.56 × 104 mol, needing 1.56 × 104 mol of CaCO3). 0.640 tonne inverts the ratio of Mr values (64.1 / 100.1). 3.12 tonne doubles the answer, as if two moles of CaCO3 were needed per mole of SO2. 1.00 tonne assumes equal masses react, ignoring the different molar masses.
What to project
The student lesson embeds these tabs. Open them full screen on the projector; every slider has a drag handle on the stage.
Slides for this lesson
A projectable deck for every pack is on the way. Until then, present the student lesson full screen: its sections, figure and lab tabs follow the same order as the plan above.