Le Chatelier’s Principle: Concentration and Temperature
In 1884, Henri Le Chatelier at the École des Mines watched a blood-red iron(III) thiocyanate solution darken the instant extra iron(III) was tipped in, and lighten again when the iron was pulled out of solution as a precipitate. From that single test tube he wrote the one sentence that lets you predict the direction of every equilibrium shift without a calculation.
Today's hook, a deep red solution goes darker when you add one reagent and paler when you add another, and a scrap of cobalt chloride paper flips pink to blue as the air dries out. Two colours, one principle, zero calculations.
0/5QUESTS
1
You’re here
Get oriented
Warm up first
Three quick questions from earlier lessons. Pulling old material back to mind before you learn something new makes the new material stick better, so this is not busywork.
Worksheets
Practise this lesson
Four printable worksheets that build from the foundations up to exam-style questions, start at whatever level suits you.
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.
FORMULAS
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 rule: Add reactant OR remove product → shift RIGHT; Add product OR remove reactant → shift LEFT
Temperature rule: Exothermic fwd (ΔH < 0): ↑T → LEFT; Endothermic fwd (ΔH > 0): ↑T → RIGHT
⚠ Temperature is the ONLY variable that changes Keq; concentration shifts change equilibrium POSITION but NOT Keq
LEARNING INTENTIONS
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
KEY TERMS
Scan these before reading
Le Chatelier’s Principle
If a system at equilibrium is disturbed, it shifts to partially counteract the disturbance.
Concentration disturbance
Adding or removing a reactant or product shifts equilibrium towards the side that consumes/produces it.
Temperature disturbance
Increasing temperature shifts equilibrium in the endothermic direction; decreasing shifts it exothermically.
Exothermic reaction
A reaction that releases heat energy to the surroundings (ΔH < 0).
Endothermic reaction
A reaction that absorbs heat energy from the surroundings (ΔH > 0).
Position of equilibrium
A qualitative description of whether products or reactants are favoured at equilibrium.
2
The principle itself
️ Core Content
01UnderstandApply
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.
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.
+5 XP Quick Check
A catalyst is added to an equilibrium mixture. According to Le Chatelier’s Principle, the equilibrium:
3
Concentration shifts
02PredictApply
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
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.
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
+5 XP True or False
Adding more product to an equilibrium mixture increases Keq because it drives the reverse reaction.
4
The iron(III) thiocyanate investigation
03ApplyApply
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
Colour intensity is comparative evidence under matched conditions: connect each disturbance to the immediate rate imbalance, the net reaction direction, and the new equilibrium composition.
Trace the response: cover the final boxes, decide which rate becomes greater immediately after each disturbance, then predict whether the red mixture becomes darker or paler.
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.
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.
Safety (school-risk controls)
The iron(III) thiocyanate colour equilibrium is best shown as a teacher demonstration or with dilute solutions (about 0.002 mol/L) under supervision. Wear safety glasses and gloves; iron(III) nitrate and thiocyanate solutions are irritants. Do not acidify thiocyanate solutions, acidified thiocyanate can release toxic hydrogen cyanide gas. The silver nitrate variation produces toxic silver waste, collect all solutions as inorganic/heavy-metal waste and do not pour them down the sink.
+5 XP Fill the Gap
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.
3b
Investigation, Changing Concentration and Temperature on One Equilibrium
Syllabus: investigate the effects of temperature, concentration, volume and/or pressure on equilibrium systems
We just saw how Le Chatelier’s principle predicts the direction of a shift. That raises a question: does a real system actually behave that way, and can you tell a shift from a dilution? This card answers it → a controlled run where every change is made against a reference tube.
Inquiry question. How do changes in concentration and temperature affect the position of the iron(III) thiocyanate equilibrium, and are the observed shifts consistent with Le Chatelier’s principle?
Safety
Teacher supervision and a written risk assessment required. Safety glasses and gloves. Thiocyanate salts must never contact acid, which can release toxic hydrogen cyanide. Iron(III) salts stain. Use a water bath rather than a naked flame for heating, and tongs for hot tubes. Collect all iron/thiocyanate waste in the labelled residues container.
Equipment. 0.002 mol L⁻¹ Fe(NO₃)₃, 0.002 mol L⁻¹ KSCN, solid KSCN, solid Fe(NO₃)₃, 0.1 mol L⁻¹ NaOH, distilled water, 6 test tubes and rack, dropper pipettes, 2 beakers as hot and ice water baths, thermometer, white tile.
1
Prepare a stock by mixing equal volumes of the dilute Fe(NO₃)₃ and KSCN solutions. The pale red colour should let a change be seen in either direction, if it is too dark to see through, dilute it further and record that you did.
2
Divide equally between six tubes. Tube 1 is the reference and is never altered. Every observation below is a comparison against tube 1 on a white tile.
3
Tube 2: add solid KSCN. Tube 3: add solid Fe(NO₃)₃. Record each colour change.
4
Tube 4: add NaOH dropwise until a change is seen. Record the colour and any precipitate.
5
The dilution control. Tube 5: add the same number of drops of distilled water as you added of NaOH to tube 4. Record the colour. This separates a genuine equilibrium shift from simply making the solution more dilute.
6
Tube 6: stand in a hot water bath at about 60 °C for three minutes, record the colour and the temperature, then move it to an ice bath and record again once cooled.
7
Return tube 6 to room temperature and check whether the original colour is restored. Temperature shifts should be reversible, and this confirms nothing was decomposed.
Results, all colours judged against the untouched reference tube
Tube
Change
Colour vs reference
Shift direction
LCP explanation
1
reference
—
—
—
2
+ SCN⁻
3
+ Fe³⁺
4
+ OH⁻ (removes Fe³⁺)
5
+ water (dilution control)
6
heated, then cooled
Analysis.
Explain each concentration result using collision theory, in terms of the rate of the forward and reverse reactions immediately after the change, and how a new equilibrium is reached.
Compare tubes 4 and 5. What does the dilution control let you claim that tube 4 alone would not?
From the temperature result, deduce whether the forward reaction is exothermic or endothermic, and state your reasoning.
Which of your changes altered the value of Keq, and which only altered the position of the equilibrium? Justify the distinction.
Conclusion. State, for each factor tested, the direction of the shift and whether it matched the Le Chatelier prediction. Where an observation did not match, say whether the likely cause is the prediction or the technique, and what you would change to find out.
The control tube is the experiment: adding anything as a solution also dilutes. Without tube 5 you cannot distinguish “the equilibrium shifted left” from “there is now less of everything per mL”, and both look like a paler tube.
Common Error: Reporting that heating “shifted the equilibrium and increased Keq” for an exothermic forward reaction. Heating an exothermic forward reaction shifts it left and decreases Keq. Temperature is the only factor here that changes Keq at all.
Method: one untouched reference tube, one dilution control, then change one variable per tube. Adding a reactant shifts right, removing Fe³⁺ with OH⁻ shifts left, heating shifts an exothermic forward reaction left. Only temperature changes Keq; concentration changes move the position at constant Keq.
Pause, write the highlighted method and results into your book.
5
Temperature and Keq
04PredictAnalyse
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
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 the higher Ea, so raising the temperature lets a proportionally greater share of particles clear that higher barrier and the reverse rate increases MORE → shift LEFT. Never use “more energy = more products” reasoning.
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.
+5 XP Quick Check
For the reaction N₂ + 3H₂ ⇌ 2NH₃ (ΔH = −92 kJ/mol), increasing temperature will:
Cross-lesson links: Le Chatelier's 1884 iron(III) thiocyanate demonstration introduced here is extended with pressure and catalyst effects in L06. Temperature shifts introduced in Card 4 explain why Keq changes with temperature, a quantitative analysis you will complete in L13. The Cobalt(II) chloride system in Card 5 is a standard NESA-listed example; it reappears in L07 as an industrial analogy.
6
Cobalt(II) chloride in everyday life
05ApplyApply
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:
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.
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.
+5 XP True or False
Cobalt(II) chloride paper turns blue when humidity is high (wet conditions).
✓
Worked examples and activities
✏️ Worked Examples
Example 1, Predicting Concentration Effects with LCP
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 2 (b): Removing ethyl acetate (product) → reverse collision frequency decreases → forward rate > reverse → system shifts RIGHT. Reactants consumed to produce more ethyl acetate.
c
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.
Example 2, Predicting Temperature Effects and Keq Changes
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.
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.
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. Keqdecreases when temperature increases for an exothermic forward reaction.
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.
Activities
Sort the steps+7 XP
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).
Use the tool to warm the endothermic equilibrium N₂O₄(g) ⇌ 2NO₂(g) (ΔH > 0 in the forward direction). Raising the temperature will…
✓
Practice questions
01
Multiple Choice
+5 XP
A fresh set drawn from this lesson’s question bank, with feedback shown immediately. +5 XP per correct · +25 XP all correct
Short Answer
ApplyApply(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.
AnalyseAnalyse(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.
EvaluateAnalyse(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.
Model Answers
✏️ Multiple-choice traps worth knowing
Endothermic forward reaction, temperature raised. For N₂(g) + O₂(g) ⇌ 2NO(g) (ΔH = +180 kJ/mol), raising the temperature shifts the system in the endothermic direction (forward, right), so more NO forms. Because the shift is toward products and temperature has changed, Keq increases. Saying “Keq is unchanged” is the classic error here, temperature is the one variable that does change it.
Adding a common ion. Adding Cl⁻ to AgCl(s) ⇌ Ag⁺(aq) + Cl⁻(aq) adds a product, so the system shifts LEFT, Ag⁺ and Cl⁻ are consumed and more AgCl(s) forms. The Ag⁺ concentration falls but Keq is UNCHANGED, concentration changes never change Keq.
Heating cobalt(II) chloride paper. The forward reaction (hexahydrate → anhydrous + water vapour) is endothermic, so heating shifts the system in the endothermic direction (forward, right) to absorb the added heat. Water is driven off and blue anhydrous CoCl₂ is left behind.
Adding base to the chromate/dichromate demonstration. In 2CrO₄²⁻(aq) + 2H⁺(aq) ⇌ Cr₂O₇²⁻(aq) + H₂O(l), NaOH neutralises H⁺ (OH⁻ + H⁺ → H₂O), which effectively removes a reactant. The system shifts LEFT to replace some of the removed H⁺, so more yellow CrO₄²⁻ forms. Keq unchanged.
“Removing a product increases Keq.” Incorrect. Keq depends only on temperature. Removing a product shifts the equilibrium position to the right (more product forms to partially replace what was taken), but the value of Keq is the same.
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].
✓
Retrieve and reflect
Check what actually stuck
⚔️
Boss Challenge
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.
REVISIT YOUR THINKING
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.
