Reactivity How fast? The rate of chemical change : R2.2.3 Factors affecting the rate of reaction IB DP Chemistry Study Notes - New Syllabus 2025
Reactivity How fast? The rate of chemical change – IB DP Chemistry- Study Notes
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Reactivity 2.2.5 – Effect of Changing Conditions on Rate of Reaction
Effect of Temperature on the Rate of Reaction
Temperature is a measure of the average kinetic energy of particles in a system. An increase in temperature increases the speed and energy of particles, making them collide more frequently and with greater energy.
Collision Theory Explanation:
According to collision theory, a reaction occurs when particles collide with sufficient energy and correct orientation. The minimum energy required for a reaction to occur is called the activation energy \( (E_a) \).
As temperature increases:
- The average kinetic energy of particles increases.
- More particles have energy \( \geq E_a \).
- The frequency and effectiveness of collisions increase.
- This leads to a higher rate of reaction.
Maxwell-Boltzmann Distribution (Effect of Temperature)
An increase in temperature shifts the curve to the right and flattens it, increasing the proportion of particles with energy greater than \( E_a \):
\( \text{Fraction of particles with } E \geq E_a \uparrow \Rightarrow \text{Reaction rate} \uparrow \)
Quantitative Insight (Arrhenius Equation):
The Arrhenius equation gives the relationship between temperature and rate constant:
\( k = A e^{\frac{-E_a}{RT}} \)
- \( k \) = rate constant
- \( A \) = frequency factor
- \( E_a \) = activation energy (J/mol)
- \( R \) = gas constant = 8.314 J·mol−1·K−1
- \( T \) = temperature (K)
$lnk=lnA-\frac{E_a}{RT}$ : Another Form
As temperature increases, \( T \uparrow \Rightarrow \frac{-E_a}{RT} \uparrow \Rightarrow e^{\frac{-E_a}{RT}} \uparrow \Rightarrow k \uparrow \)
Example:
Decomposition of Hydrogen Peroxide , At higher temperature how it will react.
\( 2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2 \)
▶️Answer/Explanation
At a higher temperature, more H2O2 molecules can overcome the activation energy. Oxygen gas is produced faster, making the reaction noticeably quicker.
Example:
Reaction of Magnesium with Hydrochloric Acid , on Increasing the temperature how it will react.
\( \text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g) \)
▶️Answer/Explanation
Increasing the temperature speeds up the collision frequency between H+ and Mg atoms. Hydrogen gas is released more rapidly due to the increased reaction rate.
Temperature and Rate of Reaction
| Temperature Change | Effect on Rate | Reason |
|---|---|---|
| Increase in temperature | Rate increases | More particles have energy \( \geq E_a \); more frequent and effective collisions |
| Decrease in temperature | Rate decreases | Fewer particles have energy \( \geq E_a \); collisions become less frequent and less energetic |
Effect of Concentration and Pressure on the Rate of Reaction
Concentration refers to the number of moles of solute per unit volume of solution (mol/dm3).
Pressure refers to the force exerted by gas particles when they collide with the walls of a container. In gaseous reactions, pressure is directly related to the concentration of gas molecules.
Collision Theory Explanation:
According to collision theory, increasing the concentration of reactants in solution or the pressure of gaseous reactants leads to more particles per unit volume. This increases the frequency of collisions between reactant particles and, hence, the rate of reaction.
- Higher concentration or pressure → more particles in a given volume.
- More frequent collisions per second.
- Greater chance of collisions with energy \( \geq E_a \).
- Thus, rate of reaction increases.
Concentration Formula:
\( C = \frac{n}{V} \)
- \( C \) = concentration (mol/dm3)
- \( n \) = number of moles of solute
- \( V \) = volume of solution (dm3)
Example:
Reaction Between Hydrochloric Acid and Magnesium At higher concentration how it will react.
\( \text{Mg}(s) + 2\text{HCl}(aq) \rightarrow \text{MgCl}_2(aq) + \text{H}_2(g) \)
▶️Answer/Explanation
At higher HCl concentrations, more H+ ions are available to collide with the magnesium surface. This increases the frequency of effective collisions, producing hydrogen gas more rapidly.
Example:
Combustion of Propane at Different Pressures what will happen on Increasing the pressure of the gaseous reactants.
\( \text{C}_3\text{H}_8(g) + 5\text{O}_2(g) \rightarrow 3\text{CO}_2(g) + 4\text{H}_2\text{O}(g) \)
▶️Answer/Explanation
Increasing the pressure of the gaseous reactants increases the number of molecules per unit volume. This results in more frequent collisions and a faster combustion reaction.
Concentration/Pressure and Rate of Reaction
| Condition Changed | Effect on Rate | Reason |
|---|---|---|
| Increase concentration | Rate increases | More particles in solution → more frequent collisions |
| Decrease concentration | Rate decreases | Fewer particles in solution → fewer collisions |
| Increase pressure (gases) | Rate increases | More gas particles per volume → more collisions |
| Decrease pressure (gases) | Rate decreases | Fewer gas particles → fewer collisions |
Effect of Surface Area (Particle Size) on the Rate of Reaction
Surface area refers to the total area of the exposed surface of a solid reactant. Smaller particles have a larger surface area relative to their volume, exposing more reactant particles to potential collisions.
Collision Theory Explanation:
Increasing the surface area of a solid reactant increases the frequency of collisions between particles in contact with the solid surface. This leads to a faster reaction rate.
- Smaller particles → larger surface area-to-volume ratio
- More surface sites available for collisions
- More frequent successful collisions → faster reaction
Note: This only applies to reactions involving solids.
Example :
Reaction of Calcium Carbonate with Hydrochloric Acid
\( \text{CaCO}_3(s) + 2\text{HCl}(aq) \rightarrow \text{CaCl}_2(aq) + \text{H}_2\text{O}(l) + \text{CO}_2(g) \)
How does the rate of reaction differ when using powdered CaCO3 compared to large chips? Explain why.
▶️Answer/Explanation
Powdered calcium carbonate has a much larger surface area than large chips. This increases the frequency of collisions between HCl particles and CaCO3, resulting in faster CO2 evolution and an increased rate of reaction.
Example:
Why does powdered sawdust combust faster than a block of wood under the same conditions?
▶️Answer/Explanation
Sawdust exposes more surface area to oxygen than a solid wood block. This results in more frequent collisions between oxygen molecules and cellulose particles, increasing the combustion rate.
Surface Area and Rate of Reaction
| Change in Particle Size | Effect on Rate | Reason |
|---|---|---|
| Decrease in particle size | Rate increases | More surface area exposed → more frequent collisions |
| Increase in particle size | Rate decreases | Less surface area exposed → less frequent collisions |
Effect of Catalysts on the Rate of Reaction
A catalyst is a substance that increases the rate of a chemical reaction without itself being consumed or permanently changed in the process.
How Catalysts Work (Mechanism):
Catalysts provide an alternative reaction pathway with a lower activation energy \( (E_a) \), meaning a greater proportion of particles have sufficient energy to react. This increases the frequency of successful collisions.
Collision Theory Explanation:
- A catalyst does not increase the kinetic energy of particles.
- Instead, it lowers the activation energy.
- As a result, more collisions become effective (energy ≥ \( E_a \)).
- The overall rate of reaction increases.
Energy Profile Diagram:
In a reaction profile, the catalysed pathway has a lower peak than the uncatalysed one.
Note: Catalysts are reaction-specific. A catalyst for one reaction may have no effect on another.
Example:
Decomposition of Hydrogen Peroxide
\( 2\text{H}_2\text{O}_2(aq) \rightarrow 2\text{H}_2\text{O}(l) + \text{O}_2(g) \)
How does the addition of manganese(IV) oxide affect this reaction? Explain why.
▶️Answer/Explanation
Manganese(IV) oxide (MnO2) acts as a catalyst, providing a lower-energy reaction pathway. It increases the rate of oxygen gas production without being consumed.
Example:
Haber Process
\( \text{N}_2(g) + 3\text{H}_2(g) \leftrightarrow 2\text{NH}_3(g) \)
What is the role of iron in the Haber Process, and how does it affect the rate of ammonia production?
▶️Answer/Explanation
Iron acts as a heterogeneous catalyst. It speeds up the formation of ammonia by allowing nitrogen and hydrogen to react on its surface with lower activation energy.
Catalyst and Rate of Reaction
| Change | Effect on Rate | Reason |
|---|---|---|
| Add a catalyst | Rate increases | Lowers activation energy → more effective collisions |
| Remove catalyst | Rate returns to original (slower) | Higher activation energy → fewer effective collisions |