The shape of the lesson
By the end of the lesson, learners can
- Heating carbonates and nitrates: Heat each carbonate and nitrate in the simulator and compare the temperature at which it starts to decompose.
- Solubility trends: Compare how much of each hydroxide and sulfate dissolves in the solubility simulator.
- Carbonate → oxide + CO2; nitrate → oxide + NO2 + O2.
- Stability increases down: lower charge density, less polarisation.
- Hydroxides more soluble down; sulfates less.
- Sulfate test: HCl then BaCl2, white precipitate.
One picture every idea returns to
Group 2 carbonates decompose on heating to the oxide and carbon dioxide; nitrates decompose to the oxide, nitrogen dioxide and oxygen. Thermal stability increases down the group: the M2+ ion gets larger with the same charge, so its charge density falls and it polarises (distorts) the large carbonate or nitrate anion less. The hydroxides become more soluble down the group, so the pH of their saturated solutions rises; the sulfates become less soluble, so barium sulfate is insoluble, which is the basis of the test for sulfate ions.
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, Stability up, solubility both ways.. Heat magnesium carbonate and it breaks down easily; barium carbonate barely notices. 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: Heating carbonates and nitrates | Heat each carbonate and nitrate in the simulator and compare the temperature at which it starts to decompose. Project tab 3 of the lab (Group 2 lab, thermal decomposition) and let a student drive it while the class predicts. | Whole class, then pairs |
| 17 to 30 min | Teach and check: Solubility trends | Compare how much of each hydroxide and sulfate dissolves in the solubility simulator. Project tab 4 of the lab (Group 2 lab, solubility trends) and let a student drive it while the class predicts. | Corners |
| 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 Solubility trends, 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.
Corners
Put one answer or viewpoint in each corner of the room. Students move to the corner they agree with, justify the choice to a partner there, then hear one voice from each corner before the reveal.
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 |
|---|---|
| The two solubility trends go in opposite directions: hydroxides become more soluble down the group, sulfates less. | When explaining thermal stability, talk about the cation polarising the anion; the anion is the same throughout, so the change must come from the size of M 2+. |
Group 2 compounds and their trends: worksheet
6 questions, 14 marks. Written in the style of the papers; no past paper question is reproduced.
Write equations for the thermal decomposition of (a) strontium carbonate and (b) calcium nitrate.
Explain why magnesium carbonate decomposes at a lower temperature than barium carbonate.
State the trends in solubility of the Group 2 hydroxides and sulfates down the group.
Describe how to test a solution for sulfate ions, giving the reagents, the result and an ionic equation.
Calculate the loss in mass when 2.50 g of calcium carbonate is heated to constant mass. (Mr: CaCO3 100.1, CO2 44.0)
Saturated solutions of Mg(OH)2 and Ba(OH)2 are tested with a pH meter. Predict and explain the difference.
Answer key and marking notes
- (a) SrCO3 → SrO + CO2.
- (b) 2Ca(NO3)2 → 2CaO + 4NO2 + O2.
- Mg2+ is smaller than Ba2+ but has the same charge, so it has a higher charge density.
- It polarises (distorts) the electron cloud of the carbonate ion more, weakening the C–O bonds.
- So less energy is needed to break off CO2.
- Hydroxides become more soluble.
- Sulfates become less soluble.
- Add dilute hydrochloric acid (to remove carbonate ions), then aqueous barium chloride.
- A white precipitate shows sulfate ions.
- Ba2+(aq) + SO42−(aq) → BaSO4(s).
- n(CaCO3) = 2.50 ÷ 100.1 = 0.0250 mol = n(CO2) lost.
- Mass lost = 0.0250 × 44.0 = 1.10 g.
- Ba(OH)2 gives a higher pH (about 13 to 14) than Mg(OH)2 (about 10).
- Hydroxide solubility increases down the group, so more OH− ions are in solution.
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.
Which equation shows the thermal decomposition of calcium carbonate?
- ACaCO3 → CaO + CO2
- BCaCO3 → Ca + CO2 + ½O2
- CCaCO3 + H2O → Ca(OH)2 + CO2
- D2CaCO3 → 2CaO + 2CO + O2
Samples of MgCO3, CaCO3, SrCO3 and BaCO3 are heated in turn. Which carbonate needs the highest temperature to decompose?
- ABaCO3
- BMgCO3
- CCaCO3
- Dthey all decompose at the same temperature, because each contains the CO32− ion
What are the products when magnesium nitrate is heated strongly?
- Amagnesium oxide, nitrogen dioxide and oxygen
- Bmagnesium nitrite and oxygen only
- Cmagnesium oxide and nitrogen dioxide only
- Dmagnesium, nitrogen and oxygen
The chart shows how the solubilities of the Group 2 hydroxides and sulfates change from Mg to Ba. Which pair contains the MORE soluble compound of each type?
- ABa(OH)2 and MgSO4
- BMg(OH)2 and BaSO4
- CBa(OH)2 and BaSO4
- DMg(OH)2 and MgSO4
Acidified BaCl2(aq) is added to a solution suspected to contain sulfate ions. What is seen if sulfate is present, and what is the ionic equation?
- Aa white precipitate; Ba2+(aq) + SO42−(aq) → BaSO4(s)
- Ba white precipitate; Ba2+(aq) + 2Cl−(aq) → BaCl2(s)
- Cno change, because BaSO4 is soluble
- Da yellow precipitate; Ba2+(aq) + SO42−(aq) → BaSO4(s)
1.483 g of magnesium nitrate, Mg(NO3)2, is heated until decomposition is complete. What total volume of gas is produced, measured at room conditions? (Mr Mg(NO3)2 = 148.3; molar gas volume = 24.0 dm3 mol−1)
- A0.600 dm3
- B0.480 dm3
- C1.20 dm3
- D0.120 dm3
Saturated solutions of magnesium hydroxide and barium hydroxide are made at the same temperature. Which statement about their pH values is correct?
- Athe Ba(OH)2 solution has the higher pH, because barium hydroxide is much more soluble, giving a higher [OH−]
- Bthe Mg(OH)2 solution has the higher pH, because hydroxide solubility decreases down the group
- Cboth have pH 14, because all Group 2 hydroxides are strong bases
- Dboth have pH 7, because Group 2 hydroxides are insoluble in water
Which explanation best accounts for the increase in thermal stability of the Group 2 carbonates and nitrates down the group?
- Athe cation gets larger with the same 2+ charge, so its charge density falls and it polarises (distorts) the large anion less, making the anion harder to break up
- Bthe cation gets larger, so it polarises the anion more, which strengthens the C-O bonds
- Cthe lattice energy increases down the group, because larger ions pack more closely
- Dthe anion gets smaller down the group, so it is held more strongly by the cation
Quiz key
Group 2 carbonates decompose on strong heating to the metal oxide and carbon dioxide: CaCO3 → CaO + CO2. Heating does not reduce the compound to the metal; calcium stays in the +2 state as Ca2+ in CaO. The third equation involves water, which is not a reactant in thermal decomposition, and calcium carbonate does not react with water in this way. Carbon monoxide and oxygen are not produced: the carbonate ion simply loses CO2, leaving O2−.
Thermal stability of the carbonates increases down Group 2, so BaCO3 is the most stable and needs the highest temperature (well above 1000 °C). MgCO3 is the least stable and decomposes most easily; choosing it reverses the trend, the classic error of thinking stability decreases down the group. CaCO3 decomposes at a lower temperature than BaCO3 (around 900 °C in a lime kiln). The anion is the same, but the cation differs in size, and the cation's polarising power decides how easily the anion breaks up, so the temperatures are not equal.
Group 2 nitrates decompose to the oxide, nitrogen dioxide and oxygen: 2Mg(NO3)2 → 2MgO + 4NO2 + O2. Brown fumes of NO2 are seen, and the O2 relights a glowing splint. Nitrite plus oxygen is the decomposition of Group 1 nitrates such as NaNO3, not Group 2. Oxide and NO2 only cannot balance: oxygen atoms are left over, and they form O2. The metal is not produced; magnesium stays as Mg2+ in MgO, and the nitrogen ends up in NO2, not as N2.
Hydroxide solubility increases down the group, so Ba(OH)2 is the most soluble hydroxide (Mg(OH)2 is only sparingly soluble). Sulfate solubility decreases down the group, so MgSO4 is the most soluble sulfate (BaSO4 is insoluble). Mg(OH)2 and BaSO4 are the LEAST soluble of each type, the reverse of the question. Choosing BaSO4 assumes sulfate solubility increases down the group like the hydroxides, which is the classic error. Choosing Mg(OH)2 assumes hydroxide solubility decreases down the group.
Barium sulfate is insoluble, so Ba2+ and SO42− combine to give a dense white precipitate: Ba2+(aq) + SO42−(aq) → BaSO4(s). The acid removes carbonate ions, which would otherwise also precipitate. BaCl2 is the reagent and is soluble; it does not precipitate from its own ions. BaSO4 is insoluble, not soluble; sulfate solubility decreases down the group. The precipitate is white, not yellow (a yellow precipitate is characteristic of silver iodide, for example).
n(Mg(NO3)2) = 1.483 / 148.3 = 0.0100 mol. From the equation, 2 mol of nitrate gives 4 mol NO2 + 1 mol O2 = 5 mol gas, so 1 mol gives 2.5 mol. n(gas) = 0.0250 mol and V = 0.0250 × 24.0 = 0.600 dm3. 0.480 dm3 counts only the NO2 (0.0200 mol) and forgets the oxygen. 1.20 dm3 uses 5 mol gas per mole of nitrate, ignoring the 2 in front of Mg(NO3)2. 0.120 dm3 counts only the oxygen (0.00500 mol).
Hydroxide solubility increases down Group 2. Mg(OH)2 is only sparingly soluble, so its saturated solution contains very few OH− ions (pH about 10), whereas Ba(OH)2 is quite soluble, giving a high OH− concentration (pH about 13 or more). Saying Mg(OH)2 is higher reverses the solubility trend. A strong base is fully dissociated, but the pH still depends on how much dissolves, so they are not both pH 14. They are not insoluble: both dissolve enough to make the solution alkaline, so pH 7 is wrong.
Mg2+ is small with a 2+ charge, so it has a high charge density and strongly polarises the electron cloud of the nearby CO32− or NO3− ion, weakening the bonds inside the anion so that CO2 (or NO2 and O2) is released at a lower temperature. Ba2+ has the same charge but is much larger, so its charge density and polarising power are lower; the anion is distorted less and the compound is more stable. Larger cations polarise LESS, not more, and polarisation weakens the bonds in the anion rather than strengthening them. Lattice energy becomes less exothermic, not more, as the ions get larger. The anion is the same throughout; only the cation changes.
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.