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Predict Before You Read
In 1884, Henry Le Chatelier at the École des Mines published his principle using iron(III) thiocyanate as a demonstration reaction: Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq). He showed that adding more Fe³⁺ deepened the red colour — the first quantitative demonstration that adding a reactant shifts equilibrium toward products. A test tube contains a deep red solution of iron(III) thiocyanate. A student adds a few drops of concentrated iron(III) nitrate solution to the test tube. Before reading on — predict what happens to the colour. Does it get darker red, lighter, or stay the same? Then predict what would happen if instead the student added a few drops of sodium hydroxide solution, which reacts with Fe³⁺ ions to form a precipitate (removing Fe³⁺ from solution). Write both predictions with your reasoning before reading on.
Key Formulas & Rules
⚠ Temperature is the ONLY variable that changes Keq; concentration shifts change equilibrium POSITION but NOT Keq
By the end of this lesson
Know
- State Le Chatelier’s Principle using precise scientific language
- Distinguish between changes that shift equilibrium position and changes that alter Keq
Understand
- Predict and explain the direction of equilibrium shift for concentration changes
- Predict and explain the direction of equilibrium shift and Keq change for temperature disturbances
Can Do
- Describe the iron(III) thiocyanate and cobalt(II) chloride investigations and explain observations using LCP
Scan these before reading
️ Core Content
Le Chatelier’s Principle — The Statement and the Logic
Le Chatelier’s Principle is chemistry’s most powerful prediction tool for equilibrium — one sentence that lets you predict the direction of every disturbance without any calculation.
Le Chatelier’s Principle states: when a closed system at dynamic equilibrium is disturbed by a change in conditions, the system shifts in the direction that minimises the effect of the disturbance and restores a new equilibrium. This principle was formulated by Henri Le Chatelier in 1884 and applies to any equilibrium system — chemical, physical, or biological.
The key word is “minimise” — the system does not eliminate the disturbance, it partially counteracts it. For example, if you add more reactant to an equilibrium system, the system shifts forward to consume some of the added reactant — but not all of it. The new equilibrium has more product and more reactant than the original, not the same amount of reactant as before.
What counts as a “disturbance”:
- Concentration changes (adding or removing a species)
- Temperature changes
- Pressure/volume changes (for gases)
What does NOT disturb equilibrium: catalysts (they affect both directions equally).
Le Chatelier's Principle: when a closed equilibrium system is disturbed, it shifts to minimise (not eliminate) the disturbance; it predicts direction only — disturbances include concentration, temperature, and pressure/volume changes, but NOT catalysts.
Pause — copy the highlighted LCP statement into your book before moving on.
A catalyst is added to an equilibrium mixture. According to Le Chatelier’s Principle, the equilibrium:
Concentration Changes — Predicting Direction of Shift
We just saw Le Chatelier's Principle as a qualitative prediction tool — disturb a system and it shifts to minimise the effect. That raises a question: how exactly does changing concentration shift the equilibrium, and why does Keq stay the same? This card answers it → with four concentration-change cases and the collision-theory mechanism behind each.
Concentration disturbances are the most straightforward LCP predictions — adding a species pushes the equilibrium away from it; removing a species pulls the equilibrium toward it.
For any reversible reaction at equilibrium:
- Adding a reactant: increases reactant concentration → forward collision frequency increases → forward rate > reverse rate → system shifts RIGHT → more product formed
- Adding a product: increases product concentration → reverse collision frequency increases → reverse rate > forward rate → system shifts LEFT → more reactant formed
- Removing a reactant: forward rate drops → system shifts LEFT to replace some of the removed reactant
- Removing a product: reverse rate drops → forward rate > reverse rate → system shifts RIGHT to replace some of the removed product
The iron(III) thiocyanate equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) demonstrates this visually. Adding Fe³⁺ (reactant) → shift right → more FeSCN²⁺ → darker red. Removing Fe³⁺ by precipitation with NaOH → shift left → less FeSCN²⁺ → lighter colour.
Add reactant
Remove reactant
Add product
Remove product
Concentration changes shift the equilibrium position but never change Keq: adding a species drives the equilibrium away from it (forward collision frequency increases); removing a species pulls it toward it (that direction's rate drops) — in every case Keq is constant.
Add the highlighted rule — shift direction AND Keq unchanged — to your notes before continuing.
LCP concentration rules — all four cases; note Keq is unchanged in every case
Adding more product to an equilibrium mixture increases Keq because it drives the reverse reaction.
The Iron(III) Thiocyanate Investigation
We just saw how concentration changes shift equilibrium position without changing Keq. That raises a question: what is the NESA-specified experiment that demonstrates concentration LCP shifts — and what does each addition do to the colour? This card answers it → the Fe³⁺/SCN⁻/FeSCN²⁺ investigation with a six-row results table.
The iron(III) thiocyanate equilibrium is chemistry’s most useful visual demonstration of concentration effects — every addition you make changes the colour of the solution in a predictable, vivid, and immediately visible way.
The equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) is ideal because FeSCN²⁺ is intensely deep red while Fe³⁺ and SCN⁻ are virtually colourless. Any shift in equilibrium position is immediately visible as a colour change.
| Addition | Species Affected | Shift Direction | Colour Change |
|---|---|---|---|
| Fe(NO₃)₃ added | Fe³⁺ (reactant) added | Right → | Darker red |
| KSCN added | SCN⁻ (reactant) added | Right → | Darker red |
| AgNO₃ added | SCN⁻ precipitated as AgSCN (removed) | Left ← | Paler/lighter |
| NaF added | Fe³⁺ forms FeF²⁺ complex (removed) | Left ← | Paler/lighter |
| Solution heated | Temperature increased (exothermic forward) | Left ← | Paler |
| Solution cooled | Temperature decreased | Right → | Darker red |
Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq): adding a reactant (Fe³⁺ or SCN⁻) shifts right → deeper red; precipitating or complexing a reactant (AgNO₃ removes SCN⁻; NaF removes Fe³⁺) shifts left → paler; Keq unchanged for all concentration disturbances.
Pause — write the highlighted equation and colour-change rules into your book before the check below.
When AgNO₃ is added to an iron(III) thiocyanate equilibrium, Ag⁺ ions precipitate SCN⁻ as AgSCN(s). This __________ a reactant from the equilibrium, causing a shift __________, producing __________ FeSCN²⁺ and a __________ colour.
Show answer
removes / left / less / paler/lighter. Removing a reactant (SCN⁻) decreases the forward rate → reverse rate > forward rate → system shifts left → FeSCN²⁺ decomposes → paler colour. Keq unchanged.
Temperature Changes — Predicting Direction and Effect on Keq
We just saw that concentration changes shift equilibrium position but leave Keq unchanged. That raises a question: is there any disturbance that actually changes the value of Keq itself — and if so, how does the shift direction depend on whether the reaction is exothermic or endothermic? This card answers it → temperature is the only variable that changes Keq, via a 2×2 ΔH × temperature matrix.
Temperature changes are qualitatively different from concentration changes — they don’t just shift the equilibrium position, they change the value of Keq itself, because they change the thermodynamic landscape of the reaction.
When temperature is increased, the system shifts in the direction of the endothermic reaction — the direction that absorbs the added heat and partially counteracts the temperature increase (Le Chatelier). When temperature is decreased, the system shifts in the direction of the exothermic reaction.
Example 1 — exothermic forward reaction (ΔH < 0): e.g. N₂ + 3H₂ ⇌ 2NH₃
- Increase T → shift LEFT → more N₂ and H₂, less NH₃ → Keq decreases
- Decrease T → shift RIGHT → more NH₃ → Keq increases
Example 2 — endothermic forward reaction (ΔH > 0): e.g. N₂O₄ ⇌ 2NO₂
- Increase T → shift RIGHT → more NO₂ → Keq increases
- Decrease T → shift LEFT → more N₂O₄ → Keq decreases
| Forward Reaction | Temperature Change | Direction of Shift | Effect on Keq |
|---|---|---|---|
| Exothermic (ΔH < 0) | Increase T | Left ← | Decreases |
| Exothermic (ΔH < 0) | Decrease T | Right → | Increases |
| Endothermic (ΔH > 0) | Increase T | Right → | Increases |
| Endothermic (ΔH > 0) | Decrease T | Left ← | Decreases |
Temperature × ΔH sign matrix — direction of shift and Keq change for all four cases
Temperature is the ONLY variable that changes Keq: increase T shifts toward the endothermic direction; decrease T shifts toward the exothermic direction; for exothermic forward (ΔH < 0): ↑T → shift left, Keq decreases; for endothermic forward (ΔH > 0): ↑T → shift right, Keq increases.
Add the highlighted temperature rules and Keq effect to your notes before the check below.
For the reaction N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92 kJ/mol), increasing temperature will:
Cobalt(II) Chloride Humidity Indicator — LCP in Everyday Life
We just saw how temperature shifts equilibrium direction and changes Keq. That raises a question: where does LCP appear in everyday life — and what is the NESA-specified cobalt(II) chloride example? This card answers it → CoCl₂ paper as a humidity indicator, with pink/blue colour changes driven by LCP.
Cobalt(II) chloride paper is in every silica gel packet in a new shoe box, every camera bag, and every pharmaceutical package — and its colour change is Le Chatelier’s Principle operating every time humidity changes.
The equilibrium is:
Pink (hexahydrate) ⇌ Blue (anhydrous) + water vapour
- In humid conditions: water vapour concentration is high → adding H₂O(g) to the system → reverse reaction favoured (Le Chatelier shifts left) → CoCl₂ absorbs water to form hexahydrate → paper turns PINK
- In dry conditions: water vapour concentration is low → H₂O(g) effectively removed → forward reaction favoured (Le Chatelier shifts right) → hexahydrate loses water to form anhydrous CoCl₂ → paper turns BLUE
- When heated: forward reaction is endothermic → heat shifts right → paper turns blue (used to regenerate desiccant indicators)
CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g): humid → H₂O(g) added → shift left → pink (hexahydrate); dry → H₂O(g) removed → shift right → blue (anhydrous); heated (endothermic forward) → shift right → blue. Mnemonic: "Pink in the rain, blue in the desert."
Pause — copy the highlighted equation and colour-change rules into your book before the check below.
Cobalt(II) chloride paper turns blue when humidity is high (wet conditions).
Put the Le Chatelier's Principle reasoning method in the correct order for a concentration change.
- Identify which direction of shift counteracts (opposes) the disturbance.
- Identify the disturbance (e.g. “concentration of reactant increased”).
- State the new equilibrium position (more products / more reactants).
- State LCP: the system will shift to minimise the effect of the disturbance.
- State the direction of shift (forward = right; reverse = left).
Worked Examples
The esterification equilibrium CH₃COOH(aq) + C₂H₅OH(aq) ⇌ CH₃COOC₂H₅(aq) + H₂O(l) is at equilibrium. Predict the direction of shift for: (a) more acetic acid added; (b) ethyl acetate removed by distillation; (c) water added.
Step 1 (a): Adding acetic acid (reactant) → forward collision frequency increases → forward rate > reverse → system shifts RIGHT. Effect: ethyl acetate concentration increases; ethanol and acetic acid concentrations decrease (partially).
Step 2 (b): Removing ethyl acetate (product) → reverse collision frequency decreases → forward rate > reverse → system shifts RIGHT. Reactants consumed to produce more ethyl acetate.
Step 3 (c): Adding water (product) → reverse collision frequency increases → reverse rate > forward → system shifts LEFT. Note: in this esterification reaction, water is a product, not the solvent — adding water shifts equilibrium left (hydrolysis direction). Acetic acid and ethanol concentrations increase.
Answer: (a) Shift right — adding reactant increases forward rate. (b) Shift right — removing product decreases reverse rate. (c) Shift left — adding product (water) increases reverse rate.
The reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), ΔH = −196 kJ/mol, is at equilibrium at 450°C with Keq = 1.7 × 10⁵. (a) Predict the direction of shift when temperature is increased to 600°C. (b) Will Keq at 600°C be greater than, equal to, or less than 1.7 × 10⁵? (c) Predict the direction of shift when temperature is decreased to 300°C.
Forward reaction is exothermic (ΔH = −196 kJ/mol). Increasing temperature adds thermal energy. Le Chatelier shifts in the endothermic direction (reverse) to absorb some of the added heat. Equilibrium shifts LEFT. [SO₃] decreases; [SO₂] and [O₂] increase.
Shift left at higher temperature → more reactants, fewer products → the ratio [SO₃]²/([SO₂]²[O₂]) is smaller at the new equilibrium. Therefore Keq at 600°C < 1.7 × 10⁵. Keq decreases when temperature increases for an exothermic forward reaction.
Decreasing temperature to 300°C — the system shifts in the exothermic direction (forward) to release heat and counteract the temperature decrease. Equilibrium shifts RIGHT. [SO₃] increases; [SO₂] and [O₂] decrease. Keq at 300°C > 1.7 × 10⁵ (increases).
Answer: (a) Shift left — exothermic forward; increase T favours endothermic reverse. (b) Keq decreases below 1.7 × 10⁵ — shift left means smaller Keq ratio. (c) Shift right — decrease T favours exothermic forward; Keq increases above 1.7 × 10⁵.
✏️ Multiple Choice
1. The equilibrium N₂(g) + O₂(g) ⇌ 2NO(g) is endothermic in the forward direction (ΔH = +180 kJ/mol). Which correctly predicts the effect of increasing temperature?
2. A student adds sodium chloride to the equilibrium AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq). The NaCl dissolves and adds Cl⁻ ions. Which prediction is correct?
3. The cobalt chloride equilibrium CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g) produces pink (hydrate) or blue (anhydrous). A chemist heats the pink paper in an oven at 110°C. Which observation and explanation is correct?
4. The equilibrium 2CrO₄²⁻(aq) + 2H⁺(aq) ⇌ Cr₂O₇²⁻(aq) + H₂O(l) produces yellow (chromate) or orange (dichromate). NaOH is added to an orange solution. What happens?
5. A student claims: “Removing a product from an equilibrium mixture increases Keq because more product must form to restore equilibrium.” Evaluate this claim.
Short Answer
Q4. The equilibrium 2CrO₄²⁻(aq) + 2H⁺(aq) ⇌ Cr₂O₇²⁻(aq) + H₂O(l) produces a yellow (CrO₄²⁻) to orange (Cr₂O₇²⁻) colour change. A student adds a few drops of concentrated hydrochloric acid to a yellow solution of chromate ions. (a) Predict the colour change. (b) Identify which species is disturbed. (c) Explain using Le Chatelier’s Principle.
Q5. The Contact Process reaction 2SO₂(g) + O₂(g) ⇌ 2SO₃(g), ΔH = −196 kJ/mol, produces sulfur trioxide for sulfuric acid manufacture. Explain why industrial chemists use a high-temperature reactor despite this reducing the equilibrium yield of SO₃.
Q6. A student is investigating the iron(III) thiocyanate equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq). They observe that adding AgNO₃ causes the solution to become much paler. Explain this observation fully using Le Chatelier’s Principle and identify what Ag⁺ ions are doing to the equilibrium system. Would you expect Keq to change? Justify your answer.
Model Answers
✏️ Multiple Choice
1. B — Forward reaction is endothermic. Increasing temperature adds heat — Le Chatelier shifts in the endothermic direction (forward, right). More NO is produced. Temperature changes Keq, and the shift is toward products → Keq increases. Option C is wrong because temperature does change Keq.
2. C — Adding Cl⁻ (a product) → system shifts LEFT → Ag⁺ and Cl⁻ consumed → more AgCl(s). Ag⁺ concentration decreases. Keq is UNCHANGED — concentration changes never change Keq. Option B is wrong because Keq does not change.
3. B — The forward reaction (hexahydrate → anhydrous + water vapour) is endothermic. Heating adds thermal energy — Le Chatelier shifts in the endothermic direction (forward, right) to absorb the heat. Water is driven off and anhydrous CoCl₂ (blue) is formed.
4. C — NaOH neutralises H⁺ (OH⁻ + H⁺ → H₂O), effectively removing H⁺ from the equilibrium. Removing a reactant causes the equilibrium to shift LEFT to replace some of the removed H⁺. More CrO₄²⁻ (yellow) forms. Keq unchanged.
5. C — Keq is unchanged. Keq depends only on temperature. Removing a product shifts the equilibrium position right (more product forms to partially replace what was removed), but the value of Keq is the same — only temperature can change it.
Short Answer Model Answers
Q4 (3 marks): (a) The solution changes from yellow to orange — the colour of Cr₂O₇²⁻ [1]. (b) H⁺ ions (a reactant) are added by the HCl [1]. (c) Adding H⁺ increases the concentration of a reactant. Le Chatelier’s Principle: the system shifts to the right to minimise the disturbance by consuming some of the added H⁺. The forward reaction proceeds at a faster rate → more Cr₂O₇²⁻ (orange) and H₂O produced. The solution turns orange. Keq is unchanged — only temperature changes Keq [1].
Q5 (3 marks): Although high temperature shifts the Contact Process equilibrium to the LEFT (forward reaction is exothermic, ΔH = −196 kJ/mol) and decreases Keq — reducing the equilibrium yield of SO₃ — higher temperatures are used because the rate of reaction is much faster at higher temperatures [1]. At low temperatures, the reaction rate is too slow for industrial production (insufficient collisions with enough energy to overcome the activation energy), even with a catalyst [1]. The higher temperature provides a commercially acceptable rate of SO₃ production, and the lower per-pass yield can be compensated by recycling unreacted SO₂ and O₂. This is the classic rate–yield trade-off in industrial chemistry [1].
Q6 (4 marks): Silver ions (Ag⁺) react with SCN⁻ to form a white precipitate of AgSCN(s) [1]. This effectively removes SCN⁻ ions from the equilibrium system. Le Chatelier’s Principle: removing a reactant (SCN⁻) from the equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) causes the equilibrium to shift LEFT → FeSCN²⁺ (deep red) decomposes to reform Fe³⁺ and SCN⁻ — but the SCN⁻ is immediately precipitated again [1]. The equilibrium continues to shift left → less FeSCN²⁺ → paler colour [1]. Keq does NOT change — concentration changes (including precipitation of a species) do not alter Keq. Only temperature changes Keq [1].
Le Chatelier’s Principle
Put your knowledge of Le Chatelier’s Principle to the test. Answer correctly to deal damage — get it wrong and the boss hits back. Pool: lessons 1–5.
Return to your Think First predictions
Return to your Think First predictions about Le Chatelier's 1884 iron(III) thiocyanate demonstration. This is the exact reaction he used to first demonstrate his principle. Using what you have now learned:
- Scenario 1 (adding Fe(NO₃)₃): Adding Fe³⁺ (reactant) → shift right → more FeSCN²⁺ formed → solution becomes darker red. This is exactly what Le Chatelier observed in 1884 — the result that led him to formulate his principle.
- Scenario 2 (adding NaOH): NaOH reacts with Fe³⁺ to form Fe(OH)₃ precipitate, removing Fe³⁺ from solution. Removing a reactant → shift left → FeSCN²⁺ decomposes → solution becomes paler/lighter. This is less intuitive — you need to recognise that precipitation effectively removes a species from the equilibrium.