Effect of Temperature Revision Notes for HSC SSCE Chemistry

Effect of Temperature

Understanding the equilibrium constant and temperature

The equilibrium constant ( $K$ ) gives us important information about the ratio of products to reactants in a chemical reaction when it reaches equilibrium. This ratio has a specific value at a particular temperature.

When you change the concentration of one substance in an equilibrium system, Le Chatelier's principle tells us that the system will respond by shifting to counteract that change. The concentrations of all substances will adjust, but importantly, the ratio of products to reactants returns to the same value - the equilibrium constant remains unchanged.

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Unlike concentration changes which don't affect $K$ , temperature changes actually alter the equilibrium constant itself. This is a crucial distinction that makes temperature a unique factor in equilibrium systems.

How temperature affects equilibrium constant

Connection to Le Chatelier's principle

Let's use an example to understand this concept. Consider the formation of the iron thiocyanate complex ion, which is an exothermic reaction:

$\text{Fe}^{3+}(aq) + \text{SCN}^{-}(aq) \rightleftharpoons \text{FeSCN}^{2+}(aq) + \text{heat}$

Because this reaction releases heat (negative enthalpy), we can think of heat as a product of the reaction.

What happens when temperature increases (exothermic reactions)

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Worked Example: Temperature Effect on an Exothermic Reaction

Consider the iron thiocyanate reaction above. When you increase the temperature of this system:

Step 1: According to Le Chatelier's principle, the system responds to minimize the change

Step 2: The reverse reaction is favoured because it absorbs (uses up) the added heat

Step 3: This means $[\text{Fe}^{3+}]$ and $[\text{SCN}^{-}]$ increase (more reactants)

Step 4: Meanwhile, $[\text{FeSCN}^{2+}]$ decreases (less product)

Step 5: This shift changes the equilibrium constant because:

$K = \frac{[\text{products}]}{[\text{reactants}]} = \frac{\text{smaller number (decreased concentration)}}{\text{bigger number (increased concentrations)}} = \text{smaller value for } K$

Conclusion: For exothermic reactions, the equilibrium constant decreases as temperature increases.

What happens for endothermic reactions

For endothermic reactions, the opposite pattern occurs:

Endothermic reactions absorb heat, so heat acts like a reactant
When temperature increases, the forward reaction is favoured
This produces more products and uses up reactants
Therefore, $K$ increases as temperature increases

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Key Principles to Remember:

Exothermic reactions: $K$ decreases when temperature increases (think: "EXO-down")
Endothermic reactions: $K$ increases when temperature increases (think: "ENDO-up")

Why temperature must always be specified

chatImportant

Because the equilibrium constant changes with temperature, you must always specify the temperature when stating a $K$ value. Without knowing the temperature, the equilibrium constant value has no meaning.

Examples of temperature dependence

The table above shows how the equilibrium constant changes with temperature for four different reactions. Let's analyze what these patterns tell us:

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By observing how $K$ changes with temperature, we can determine whether a reaction is exothermic or endothermic:

If $K$ decreases with increasing temperature → exothermic reaction
If $K$ increases with increasing temperature → endothermic reaction

Reaction (a): $\text{N}_2\text{O}_4(g) \rightleftharpoons 2\text{NO}_2(g)$

As temperature increases from 273 K to 500 K, $K$ increases dramatically from $5.7 \times 10^{-4}$ to $41.4$
This indicates an endothermic reaction (products favoured at higher temperatures)

Reaction (b): $\text{N}_2(g) + 3\text{H}_2(g) \rightleftharpoons 2\text{NH}_3(g)$ (ammonia synthesis)

As temperature increases from 298 K to 900 K, $K$ decreases from $4.0 \times 10^8$ to $5.4 \times 10^{-3}$
This indicates an exothermic reaction (reactants favoured at higher temperatures)
This has important implications for industrial ammonia production

Reaction (c):

Shows $K$ decreasing from 450.0 to 0.21 as temperature rises from 600 K to 1000 K
Another exothermic reaction example

Reaction (d):

Shows $K$ decreasing from 1360 to 190 as temperature increases from 273 K to 370 K
Also indicates an exothermic reaction

Practical implications

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These temperature effects on equilibrium constants are crucial for:

Industrial chemical processes where maximising product yield is important
Understanding reaction conditions and optimisation
Predicting how systems will respond to temperature changes
Designing efficient chemical reactors

Remember!

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Key Takeaways:

The equilibrium constant $K$ changes with temperature, unlike concentration changes which don't affect $K$
For exothermic reactions: increasing temperature causes $K$ to decrease (reverse reaction favoured)
For endothermic reactions: increasing temperature causes $K$ to increase (forward reaction favoured)
Temperature must always be specified when stating an equilibrium constant value
Table data helps predict whether reactions are exothermic or endothermic by observing how $K$ changes with temperature

Effect of Temperature (HSC SSCE Chemistry): Revision Notes