HL IB Chemistry Study Notes on Rates of Reaction

Rate of Reaction

• Definition of Rate of Reaction:
  • Indicates speed at which a chemical reaction occurs.
  • Measured as change in concentration of reactants or products per unit time.
  • Units: $ext{mol} ext{dm}^{-3} ext{s}^{-1}$ .
• Calculating Rate of Reaction:
  • Formula:
    $ext{Rate of Reaction} = \frac{\text{change in concentration (mol dm}^{-3}\text{)}}{\text{time (s)}}$
• Rate of Reaction Graphs:
  • Concentrations of reactants/products change over time.
  • Steeper gradient indicates a faster reaction.
  • Find the rate at a specific time by calculating the gradient of the curve at that point.

Gradient Calculation

• To find the gradient of a curve:
  • Draw a tangent to the curve.
  • Use gradient formula:
    $ext{Gradient} = \frac{\Delta y}{\Delta x}$
• Graph of Reactants vs. Time:
  • Shows negative gradient.
  • Example: $ext{gradient} = -20.6 ext{ mol dm}^{-3} ext{s}^{-1}$ converts to $ext{rate} = 20.6 ext{ mol dm}^{-3} ext{s}^{-1}$ .
• Graph of Products vs. Time:
  • Shows positive gradient and does not require conversion.

Worked Example - Iodine and Methanoic Acid Reaction

• Reaction: $I<em>{(aq)} + HCOOH</em>{(aq)} → 2I<em>{(aq)} + 2H^{+}</em>{(aq)} + CO_{(g)}$
• Calculate rate at 20 seconds:
  • Draw tangent at this point.
  • Complete triangle to find $x$ and $y$ values.
  • Gradient calculation example:
    $ext{Rate of Reaction} = \frac{24}{40} ext{ cm s}^{-1} = 0.60 ext{ cm s}^{-1}$.

Measuring Rates of Reaction

• Methods to measure:
  • Mass Loss:
    • Weight of vessel decreases as gas escapes.
  • Gas Production:
    • Measure volume of gas produced over time.
  • Colorimetry:
    • Light intensity changes as reaction proceeds.

Colorimetry Measurement

• Setup:
  • Use colorimeter/spectrophotometer to monitor light through solution.
• Data Collection:
  • Measure light intensity periodically; plot data to show concentration changes.
  • Example of limitations: Colorimetry unable to monitor colored precipitate formation.

Mass Change Measurement

• Setup:
  • Use a balance to measure mass loss in reactions producing gas (e.g., $ext{CaCO}<em>3 + ext{HCl} → ext{CO}</em>2$).
• Considerations:
  • Gas must be dense enough to measure.

Volume of Gas Measurement

• Measure gas volume produced over time using gas syringe or water displacement method.
• Plot volume vs. time to find reaction rate.

Measuring Concentration Changes

• Use titration while stopping the reaction to analyze the concentration.
• Quenching: To stabilize a sample for titration analysis.

Conductivity Measurements

• Monitor changes in electrical conductivity as ions react in solutions.

Clock Reactions

• Measurement when a specific visual point is reached (e.g., precipitate formation).

Collision Theory

• Explanation:
  • Reactions occur when particles collide with sufficient energy and correct orientation.
• Important Factors
  1. Collision Frequency:
     • Number of collisions per unit time influenced by concentration, pressure, temperature, surface area.
  2. Collision Energy:
     • Energy during collision; not all collisions are successful.
  3. Activation Energy $(E_a)$:
     • Minimum energy required to initiate a reaction.

Activation Energy

• Defined as the minimum energy required for reaction to occur.
• For a successful reaction, particles must collide with energy $\geq E_a$ .
• Impact of Catalysts: Reduces $E_a$.

Factors Affecting Rates of Reaction

• Concentration: Increased concentration raises collision frequency.
• Pressure: Increased pressure in gas reactions increases collision frequency.
• Temperature: Higher temperatures result in faster motion and greater energy, increasing reaction rates.
• Surface Area: More exposed surfaces increase collision opportunities.
• Catalysts: Provide alternative pathways with lower $E_a$.

Rate Equations (HL)

• Defined as expression showing relationship between concentration of reactants and rate of reaction.
• General Formula: $ext{Rate} = k[A]^m[B]^n$ where $m$ and $n$ are reaction orders.
• Common expression:
  • $ext{Rate} \propto [D]$ or $ext{Rate} = k[D]$ for some reactions.
• Orders of Reaction:
  • Zero Order: No effect from concentration.
  • First Order: Directly proportional to concentration.
  • Second Order: Proportional to the square of concentration.

Worked Example - Determining Reaction Orders

• Given thermal decomposition reaction: $2N<em>2O</em>5(g) → 4NO(g) + O_2(g)$
  • Order with respect to $N<em>2O</em>5$ is 1.
  • Doubling concentration triples the rate.

The Rate Constant (HL)

• Rate constant $k$ is influenced by temperature and activation energy.
• Formula for rate constant:
  1. Experimental data: $k = \frac{\text{rate}}{[A]^m[B]^n}$

Calculating Rate Constant

• Units of $k$ depend on reaction order:
  • Zero order: $ext{mol dm}^{-3} ext{s}^{-1}$
  • First order: $ext{s}^{-1}$
  • Second order: $ext{mol}^{-1} ext{dm}^3 ext{s}^{-1}$
• How temperature affects $k$: Raises rate and $k$ value.

Reaction Mechanisms

• Description: Series of elementary steps leading to the overall reaction, including intermediates that do not appear in the overall equation.

Energy Profiles

• Show energy changes throughout a reaction; include activation energies of steps.

Molecularity (HL)

• Defined as number of reactants in an elementary step:
  • Unimolecular: one reactant.
  • Bimolecular: two reactants.
  • Termolecular: three reactants (rare).

The Arrhenius Equation (HL)

• Formula: $k = Ae^{-\frac{E_a}{RT}}$ where:
  • $k$ is the rate constant.
  • $A$ is the Arrhenius factor.
  • $E_a$ is activation energy.
  • $R$ is gas constant.
  • $T$ is temperature.

Logs in Arrhenius Equation

• Taking natural logarithm helps simplify the equation to:
  $ext{ln} k = ext{ln} A - \frac{E_a}{RT}$ .

Worked Example - Activation Energy

• Calculation process to find activation energy based on provided values and experiments.

Graphing the Arrhenius Equation

• Graph of $ext{ln} k$ vs. $1/T$ provides gradient related to activation energy.
• You can determine $A$ from the y-intercept and subsequent calculations.