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Chemical Kinetics - Collision theory and rates of reaction

Grade 12IBChemistry

Review the key concepts, formulae, and examples before starting your quiz.

🔑Concepts

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Collision Theory: For a chemical reaction to occur, reactant particles must collide with a kinetic energy greater than or equal to the Activation Energy (EaE_a) and in the correct spatial orientation.

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Rate of Reaction: Defined as the change in the concentration of a reactant or product per unit time. The standard units are mol⋅dm−3⋅s−1mol \cdot dm^{-3} \cdot s^{-1}.

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Activation Energy (EaE_a): The minimum energy required for colliding particles to result in a successful reaction by reaching the transition state.

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Factors Affecting Rate: Concentration (increases collision frequency), Pressure (increases frequency for gases), Surface Area (increases frequency for solids), Temperature (increases both frequency and the fraction of particles with E≥EaE \geq E_a), and Catalysts.

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Maxwell-Boltzmann Distribution: A probability distribution curve showing the kinetic energies of particles at a specific temperature. Increasing the temperature flattens the curve and shifts it to the right, significantly increasing the area under the curve beyond EaE_a.

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Catalysts: Substances that increase the reaction rate by providing an alternative reaction pathway with a lower activation energy (EaE_a), without being consumed in the process.

📐Formulae

Rate=−Δ[Reactant]Δt=Δ[Product]ΔtRate = -\frac{\Delta[Reactant]}{\Delta t} = \frac{\Delta[Product]}{\Delta t}

k=Ae−EaRTk = A e^{-\frac{E_a}{RT}}

ln⁡k=−EaRT+ln⁡A\ln k = -\frac{E_a}{RT} + \ln A

ln⁡(k1k2)=EaR(1T2−1T1)\ln\left(\frac{k_1}{k_2}\right) = \frac{E_a}{R}\left(\frac{1}{T_2} - \frac{1}{T_1}\right)

💡Examples

Problem 1:

In a reaction between MgMg and HClHCl, the concentration of HClHCl drops from 2.00 mol⋅dm−32.00\ mol \cdot dm^{-3} to 1.50 mol⋅dm−31.50\ mol \cdot dm^{-3} in 25.025.0 seconds. Calculate the average rate of reaction with respect to HClHCl.

Solution:

Rate=−(1.50−2.00)25.0=0.020 mol⋅dm−3⋅s−1Rate = -\frac{(1.50 - 2.00)}{25.0} = 0.020\ mol \cdot dm^{-3} \cdot s^{-1}

Explanation:

The rate is calculated by dividing the change in concentration (Δ[HCl]\Delta [HCl]) by the time interval (Δt\Delta t). The negative sign in the formula accounts for the disappearance of the reactant to yield a positive rate value.

Problem 2:

Explain, using the Maxwell-Boltzmann distribution, why a 10 K10\ K increase in temperature typically doubles the reaction rate.

Solution:

An increase in temperature shifts the distribution curve to the right and lowers the peak. This causes a significant increase in the area under the curve to the right of the EaE_a line.

Explanation:

While the total number of collisions increases slightly, the primary reason for the rate doubling is the exponential increase in the fraction of particles that possess kinetic energy E≥EaE \geq E_a. More collisions become 'successful' per unit time.

Collision theory and rates of reaction Revision - Grade 12 Chemistry IB