Chemistry · Simulation and virtual lab Free
Chemical equilibrium: Le Chatelier’s principle with colour changes
Add reagents, heat or cool, and watch the colour show which way the equilibrium moves.
- You will learn to
- Observe how concentration and temperature move the position of an equilibrium.
- Use Le Chatelier’s principle and Q against K to explain each colour change.
- Decide from a temperature change whether the forward reaction is endothermic or exothermic.
- You will measure
- Colour and the direction of shift for each change; the % of the product form; K at each temperature.
- Time needed
- About 25–35 min
- Equipment
- Nothing extra. Real version: cobalt(II) chloride solution, concentrated HCl, silver nitrate, water baths and ice, test tubes (teacher demonstration for chromate).
- Price
- Free: no payment, no account.
Virtual lab
Follow the steps with the simulation above (go back to the simulation ↑). Each time you tap Record, a row is added to the table and a point to the graph.
Aim
To investigate how changes in concentration and temperature affect the position of an equilibrium, and to relate the observations to Le Chatelier’s principle and the equilibrium constant K.
What to do
- Choose the cobalt chloride system. Note the starting colour (purple), the bars and the Q / K gauge. Tap Record.
- Add concentrated HCl. Watch the colour and the bars. Which way did the equilibrium move? Record.
- Add water several times, then silver nitrate. Explain each change. Record.
- Start again. Heat the tube to 80 °C, then cool to 0 °C. Which direction is endothermic? Does K change?
- Choose the chromate system. Add sulfuric acid, then sodium hydroxide. Record the colours.
- Write a rule: what does an equilibrium do when you change a condition?
Results
Questions
State Le Chatelier’s principle.
If a system at equilibrium is disturbed by a change in conditions, the position of equilibrium moves in the direction that tends to oppose the change.
Why does adding concentrated HCl turn pink cobalt chloride blue?
It adds Cl⁻ ions; the equilibrium moves to the right to use them up, forming more blue [CoCl₄]²⁻.
Why does adding silver nitrate turn the solution pink?
Ag⁺ removes Cl⁻ as a white precipitate of AgCl, so the equilibrium moves left to replace Cl⁻, forming pink [Co(H₂O)₆]²⁺.
Heating turns the solution blue. Is the forward reaction exothermic or endothermic?
Endothermic: heating favours the endothermic direction, and the forward (blue) direction is favoured, so ΔH is positive.
Which change alters the value of K: concentration or temperature?
Only temperature. Concentration changes move the position of equilibrium but K stays the same.
Why does adding NaOH turn orange dichromate yellow?
OH⁻ removes H⁺ (forming water), so 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O moves left, forming yellow chromate.
For teachers: expected results and how the model works
Cobalt: [Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O, ΔH = +50 kJ/mol (model value), K(25 °C) = 0.006, start [Cl⁻] = 3 mol/dm³, giving about 33 % blue form (purple). K at other temperatures from the van ’t Hoff equation: about 0.001 at 0 °C and 0.14 at 80 °C, so heating gives blue and cooling gives pink. Concentrated HCl raises [Cl⁻] (blue); water dilutes Cl⁻, which appears to the fourth power in Q = [CoCl₄²⁻] ÷ ([Co(H₂O)₆²⁺][Cl⁻]⁴), so Q rises above K and the equilibrium moves left (pink); AgNO₃ removes Cl⁻ as white AgCl (pink). Chromate: 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O, K about 3 × 10¹³, almost thermoneutral, so temperature has little effect; acid gives orange, alkali yellow. Learners record changes rather than numbers; the Q / K gauge shows why the system moves: if Q < K it moves right, if Q > K it moves left. Safety for the real version: concentrated HCl and chromates are hazardous; chromates are toxic and must be used only by the teacher.
Exam practice Exam set
National-exam style questions on this simulation (12 marks). Try each one on paper first, then open the marking scheme.
For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ/mol, state and explain the effect on the yield of ammonia of (a) increasing the pressure, (b) increasing the temperature. (4 marks)
Marking scheme: (a) Yield increases (1): the equilibrium moves to the side with fewer gas molecules, 2 instead of 4 (1). (b) Yield decreases (1): the forward reaction is exothermic, so heating moves the equilibrium to the left (endothermic direction) (1).
Write the expression for Kc for 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) and give its units. (3 marks)
Marking scheme: Kc = [SO₃]² ÷ ([SO₂]² [O₂]) (2); units dm³ mol⁻¹ (1).
In the cobalt chloride equilibrium, describe what you would see when the blue solution is placed in ice water, and explain. (3 marks)
Marking scheme: It turns pink (1); cooling favours the exothermic direction (1), which is the backward reaction forming [Co(H₂O)₆]²⁺ (1).
Explain why a catalyst does not change the position of equilibrium. (2 marks)
Marking scheme: It speeds up the forward and backward reactions equally (1), so equilibrium is reached faster but the amounts at equilibrium are the same (1).
Donat Sciences and MathsChemistry · Virtual lab worksheet
School: ........................................ Name: ................................ Class: ........ Date: ............
Chemical equilibrium: Le Chatelier’s principle with colour changes
Aim: To investigate how changes in concentration and temperature affect the position of an equilibrium, and to relate the observations to Le Chatelier’s principle and the equilibrium constant K.
Apparatus (real laboratory): cobalt(II) chloride solution, concentrated HCl, silver nitrate, water baths and ice, test tubes (teacher demonstration for chromate).
What to do (with the simulation on www.donatstem.co.rw)
- Choose the cobalt chloride system. Note the starting colour (purple), the bars and the Q / K gauge. Tap Record.
- Add concentrated HCl. Watch the colour and the bars. Which way did the equilibrium move? Record.
- Add water several times, then silver nitrate. Explain each change. Record.
- Start again. Heat the tube to 80 °C, then cool to 0 °C. Which direction is endothermic? Does K change?
- Choose the chromate system. Add sulfuric acid, then sodium hydroxide. Record the colours.
- Write a rule: what does an equilibrium do when you change a condition?
Results
| # | System | Change made | Colour before | Colour after | Direction of shift | Reason |
|---|
| 1 | | | | | | |
| 2 | | | | | | |
| 3 | | | | | | |
| 4 | | | | | | |
| 5 | | | | | | |
| 6 | | | | | | |
| 7 | | | | | | |
| 8 | | | | | | |
Graph
No graph: complete the table and write Le Chatelier’s principle in your own words.
Questions
- State Le Chatelier’s principle. [2]
- Why does adding concentrated HCl turn pink cobalt chloride blue? [2]
- Why does adding silver nitrate turn the solution pink? [2]
- Heating turns the solution blue. Is the forward reaction exothermic or endothermic? [2]
- Which change alters the value of K: concentration or temperature? [2]
- Why does adding NaOH turn orange dichromate yellow? [2]
- For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ/mol, state and explain the effect on the yield of ammonia of (a) increasing the pressure, (b) increasing the temperature. [4]
- Write the expression for Kc for 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) and give its units. [3]
- In the cobalt chloride equilibrium, describe what you would see when the blue solution is placed in ice water, and explain. [3]
- Explain why a catalyst does not change the position of equilibrium. [2]
Donat Sciences and MathsChemistry · Virtual lab worksheet
School: ........................................ Name: ................................ Class: ........ Date: ............
Marking scheme: Chemical equilibrium: Le Chatelier’s principle with colour changes
Expected results
Cobalt: [Co(H₂O)₆]²⁺ + 4Cl⁻ ⇌ [CoCl₄]²⁻ + 6H₂O, ΔH = +50 kJ/mol (model value), K(25 °C) = 0.006, start [Cl⁻] = 3 mol/dm³, giving about 33 % blue form (purple). K at other temperatures from the van ’t Hoff equation: about 0.001 at 0 °C and 0.14 at 80 °C, so heating gives blue and cooling gives pink. Concentrated HCl raises [Cl⁻] (blue); water dilutes Cl⁻, which appears to the fourth power in Q = [CoCl₄²⁻] ÷ ([Co(H₂O)₆²⁺][Cl⁻]⁴), so Q rises above K and the equilibrium moves left (pink); AgNO₃ removes Cl⁻ as white AgCl (pink). Chromate: 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O, K about 3 × 10¹³, almost thermoneutral, so temperature has little effect; acid gives orange, alkali yellow. Learners record changes rather than numbers; the Q / K gauge shows why the system moves: if Q < K it moves right, if Q > K it moves left. Safety for the real version: concentrated HCl and chromates are hazardous; chromates are toxic and must be used only by the teacher.
Answers
- State Le Chatelier’s principle. [2]
If a system at equilibrium is disturbed by a change in conditions, the position of equilibrium moves in the direction that tends to oppose the change.
- Why does adding concentrated HCl turn pink cobalt chloride blue? [2]
It adds Cl⁻ ions; the equilibrium moves to the right to use them up, forming more blue [CoCl₄]²⁻.
- Why does adding silver nitrate turn the solution pink? [2]
Ag⁺ removes Cl⁻ as a white precipitate of AgCl, so the equilibrium moves left to replace Cl⁻, forming pink [Co(H₂O)₆]²⁺.
- Heating turns the solution blue. Is the forward reaction exothermic or endothermic? [2]
Endothermic: heating favours the endothermic direction, and the forward (blue) direction is favoured, so ΔH is positive.
- Which change alters the value of K: concentration or temperature? [2]
Only temperature. Concentration changes move the position of equilibrium but K stays the same.
- Why does adding NaOH turn orange dichromate yellow? [2]
OH⁻ removes H⁺ (forming water), so 2CrO₄²⁻ + 2H⁺ ⇌ Cr₂O₇²⁻ + H₂O moves left, forming yellow chromate.
- For N₂(g) + 3H₂(g) ⇌ 2NH₃(g), ΔH = −92 kJ/mol, state and explain the effect on the yield of ammonia of (a) increasing the pressure, (b) increasing the temperature. [4]
(a) Yield increases (1): the equilibrium moves to the side with fewer gas molecules, 2 instead of 4 (1). (b) Yield decreases (1): the forward reaction is exothermic, so heating moves the equilibrium to the left (endothermic direction) (1).
- Write the expression for Kc for 2SO₂(g) + O₂(g) ⇌ 2SO₃(g) and give its units. [3]
Kc = [SO₃]² ÷ ([SO₂]² [O₂]) (2); units dm³ mol⁻¹ (1).
- In the cobalt chloride equilibrium, describe what you would see when the blue solution is placed in ice water, and explain. [3]
It turns pink (1); cooling favours the exothermic direction (1), which is the backward reaction forming [Co(H₂O)₆]²⁺ (1).
- Explain why a catalyst does not change the position of equilibrium. [2]
It speeds up the forward and backward reactions equally (1), so equilibrium is reached faster but the amounts at equilibrium are the same (1).