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Chemistry Y12 · Module 5 · Lesson 5

IQ2 — Le Chatelier's Principle

Le Chatelier's Principle — Concentration & Temperature

Cobalt(II) chloride paper turns pink in humid conditions and blue when dry — a reversible colour change used in humidity indicators for decades, and a live demonstration of Le Chatelier's Principle every time the weather changes.

Understand Predict Apply

Misconceptions to Fix

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Wrong: Le Chatelier's Principle says the system opposes the disturbance by returning to the original concentrations.

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Right: Le Chatelier's Principle states the system shifts to minimise the disturbance. After adding a reactant, the system consumes some of it — but the final concentration is still higher than before. The system minimises, not eliminates, the disturbance. Original concentrations are not restored.

Learning Intentions

State Le Chatelier's Principle using precise scientific language

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

Describe the iron(III) thiocyanate and cobalt(II) chloride investigations and explain observations using LCP

Distinguish between changes that shift equilibrium position and changes that alter Keq

Printable worksheet

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🤔 Think First — Before You Read

A test tube contains a deep red solution of iron(III) thiocyanate — the equilibrium Fe³⁺(aq) + SCN⁻(aq) ⇌ FeSCN²⁺(aq) (colourless reactants → deep red product). Here is the case: 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.

Your prediction:

Key Formulas & Rules

Le Chatelier's Principle: when a closed system at dynamic equilibrium is disturbed, it shifts to minimise the effect of the disturbance and restore equilibrium

Concentration disturbance rule:

Add reactant OR remove product → shift RIGHT (forward)
Add product OR remove reactant → shift LEFT (reverse)

Temperature disturbance rule:

Forward reaction exothermic (ΔH < 0): increase T → shift LEFT; decrease T → shift RIGHT
Forward reaction endothermic (ΔH > 0): increase T → shift RIGHT; decrease T → shift LEFT

⚠ Temperature is the ONLY variable that changes Keq; concentration shifts change equilibrium POSITION but NOT Keq

Understand

1. 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).

Must know: Le Chatelier's Principle predicts direction of shift only — not the magnitude or new equilibrium concentrations. For quantitative predictions you need Keq and ICE tables (L09–L11). In HSC questions asking you to "predict the effect," give the direction (left or right) AND justify using LCP language.

Common error: "Adding more reactant shifts the equilibrium — therefore the concentration of reactant at the new equilibrium is the same as before." Wrong. The new equilibrium has MORE of both reactant and product than the original equilibrium. The reactant was not completely consumed. LCP says the system partially counteracts the disturbance, not completely eliminates it.

Predict

2. Concentration Changes — Predicting Direction of Shift

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

Rate Effect: Forward rate increases

Direction of Shift: Right →

Effect on Products: Products increase

Remove reactant

Rate Effect: Forward rate decreases

Direction of Shift: Left ←

Effect on Products: Products decrease

Add product

Rate Effect: Reverse rate increases

Direction of Shift: Left ←

Effect on Products: Reactants increase

Remove product

Rate Effect: Reverse rate decreases

Direction of Shift: Right →

Effect on Products: Products replaced

Must know: Concentration changes shift the position of equilibrium but do NOT change the value of Keq. After the system reaches its new equilibrium following a concentration disturbance, the ratio of product to reactant concentrations equals the same Keq as before. This is a critical HSC point.

Common error: "Adding more reactant increases Keq." Wrong — Keq is unchanged by concentration changes. Only temperature changes Keq. The new equilibrium has different concentrations, but the same Keq.

LCP concentration rules — all four cases; note Keq is unchanged in every case

Apply

3. The Iron(III) Thiocyanate Investigation

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

HSC exam format: "Describe the observation and explain using Le Chatelier's Principle when [substance] is added to an iron(III) thiocyanate equilibrium mixture." Your answer must include: (1) the observation (colour change); (2) which species is affected; (3) direction of shift; (4) reason using LCP.

Insight: The iron(III) thiocyanate equilibrium is also used to measure Keq experimentally using colourimetry — because the concentration of FeSCN²⁺ can be determined from the absorbance of the red colour using Beer's Law. You will encounter this quantitative application in L13.

Predict

4. Temperature Changes — Predicting Direction and Effect on Keq

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

Full marks requires 3 components: (1) whether the forward reaction is exo or endothermic; (2) direction of shift; (3) whether Keq increases or decreases. "Increasing temperature shifts the equilibrium left for an exothermic forward reaction, decreasing the value of Keq" is the minimum complete answer.

Common error: "Increasing temperature always shifts equilibrium to the right because higher temperature means more energy and faster reactions." Wrong. For an exothermic forward reaction, the reverse (endothermic) reaction has higher Eₐ — a greater proportion of particles exceed this higher energy barrier — so the reverse rate increases MORE → shift LEFT. Never use "more energy = more products" reasoning.

Temperature × ΔH sign matrix — direction of shift and Keq change for all four cases

Key Terms — scan these before reading

Le Chatelier's PrincipleIf a system at equilibrium is disturbed, it will shift to minimise the disturbance.

Concentration disturbance ruleThe amount of solute present in a given quantity of solution or solvent.

Le Chatelier's Principle statesIf a system at equilibrium is disturbed, it will shift to minimise the disturbance.

Dynamic equilibriumA state where forward and reverse reaction rates are equal.

Equilibrium constant (Keq)The ratio of product to reactant concentrations at equilibrium.

Reaction quotient (Q)The ratio of product to reactant concentrations at any instant.

Apply

5. Cobalt(II) Chloride Humidity Indicator — LCP in Everyday Life

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:

CoCl₂·6H₂O(s) ⇌ CoCl₂(s) + 6H₂O(g)

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)

Memory aid: "Pink in the rain, blue in the desert." Hexahydrate (wet, 6 water molecules) = pink; anhydrous (dry, no water) = blue.

NESA-specified investigation: Know the colours (pink = humid, blue = dry), the direction of each LCP shift, and the effect of heating (endothermic forward → heat shifts right → blue). This experiment appears in HSC exam questions regularly.

Common error: Students sometimes remember the colours backwards — blue for wet, pink for dry. Use the mnemonic above. The hexahydrate (6 water molecules attached) is the wet form = pink.

Interactive — Le Chatelier's Principle Simulator

Revisit Your Thinking

At the start of this lesson you predicted what would happen when concentrated Fe³⁺ was added to the iron(III) thiocyanate equilibrium. The solution should become darker red. According to Le Chatelier's principle, adding a reactant (Fe³⁺) shifts the equilibrium to the right, producing more coloured FeSCN²⁺. For temperature: because the forward reaction is endothermic, increasing temperature also shifts right (darker), while decreasing temperature shifts left (paler).

Simulator: Le Chatelier's Principle Simulator

Revisit Your Initial Thinking

Look back at what you wrote in the Think First section. What has changed? What did you get right? What surprised you?

Worked Examples

Apply

Example 1 — Predicting Concentration Effects with LCP

Problem: 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.

Band 5–6

Example 2 — Predicting Temperature Effects and Keq Changes

Problem: 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.

Step 1 (a): 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.

Step 2 (b): 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.

Step 3 (c): 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.

Checkpoint Questions

1 mark

Q1: 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?

A Shift left; Keq decreases

B Shift right; Keq increases

C Shift right; Keq is unchanged

D Shift left; Keq is unchanged

1 mark

Q2: 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?

A Equilibrium shifts right; Ag⁺ concentration increases

B Equilibrium shifts left; Ag⁺ concentration decreases; Keq decreases

C Equilibrium shifts left; Ag⁺ concentration decreases; Keq unchanged

D No shift occurs because NaCl is not part of the equilibrium expression

1 mark

Q3: 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?

A The paper becomes darker pink because heating shifts equilibrium left, forming more hydrate

B The paper turns blue because heating shifts equilibrium right (forward endothermic direction), removing water and forming anhydrous CoCl₂

C The paper turns blue because heating breaks all chemical bonds in the hydrate simultaneously

D The paper remains pink because the pink hexahydrate is thermally stable at 110°C

Short Answer Practice

3 marks

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.

3 marks

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.

4 marks

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.

Q4 Model Answer:

(a) The solution changes from yellow to orange — the colour of Cr₂O₇²⁻.

(b) H⁺ ions (a reactant) are added by the HCl.

(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.

Q5 Model Answer:

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. 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. 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.

Q6 Model Answer:

Silver ions (Ag⁺) react with SCN⁻ to form a white precipitate of AgSCN(s). 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) decreases the reverse rate less than it decreases the forward rate. The reverse rate now exceeds the forward rate, and the equilibrium shifts LEFT. The FeSCN²⁺ (deep red) decomposes to reform Fe³⁺ and SCN⁻ — but the SCN⁻ is immediately precipitated again. The equilibrium continues to shift left → less FeSCN²⁺ → paler colour. Keq does NOT change — concentration changes (including precipitation of a species) do not alter Keq. Only temperature changes Keq.

Revisit: Think First Review

Return to your Think First predictions. Using what you have now learned:

Scenario 1 (adding Fe(NO₃)₃): Adding Fe³⁺ (reactant) → shift right → more FeSCN²⁺ formed → solution becomes darker red. You likely predicted this correctly if you understood that adding a reactant shifts the equilibrium toward products.
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.