Course Topics
Introduction to Reaction Kinetics
Reaction Kinetics: The study of rates of chemical reactions, factors affecting rates, and reaction mechanisms.
Chemical Reactions Classification by Rate:
| Type | Examples | Characteristics |
|---|---|---|
| Very Fast Reactions | Ionic reactions, neutralization | Completed in microseconds |
| Very Slow Reactions | Rusting of iron, weathering | Take years to complete |
| Moderately Slow Reactions | Esterification, hydrolysis | Take minutes to hours |
Kinetics = Rates + Mechanisms + Factors – Remember the three pillars of reaction kinetics!
Rate of Reaction
Definition: The change in concentration of a reactant or product per unit time.
Units: mol dm⁻³ s⁻¹ (moles per cubic decimeter per second)
Rate = Concentration change ÷ Time taken. Units always have “per second” (s⁻¹) or similar time unit!
Rate Constant (k)
Definition: The proportionality constant in the rate equation. Rate = k[Reactants]ⁿ
Key Properties:
- Constant at constant temperature
- Independent of concentration
- Depends on temperature (increases with temperature)
- Characteristic of a specific reaction
Rate of Reaction
- Changes with time
- Depends on concentration
- Units: mol dm⁻³ s⁻¹
- Measured experimentally
Rate Constant (k)
- Constant at constant T
- Independent of concentration
- Units vary with order
- Characteristic of reaction
k is like the “personality” of a reaction – unique, constant at same T, and tells how fast it inherently is!
Factors Affecting Reaction Rate
| Factor | Effect on Rate | Reason |
|---|---|---|
| Nature of Reactant | Ionic reactions faster than covalent | Ions react immediately on contact |
| Concentration | Rate ∝ Concentrationⁿ | More particles = more collisions |
| Surface Area | Rate increases with surface area | More exposed particles for reaction |
| Temperature | Rate doubles for every 10°C rise | More particles have E ≥ Eₐ |
| Catalyst | Increases rate without being consumed | Lowers activation energy |
| Light | Increases photochemical reactions | Provides energy for activation |
Remember: Concentration, Temperature, Catalyst, Surface area, Nature = CTCSN!
Arrhenius Equation
Equation: k = A e^(-Eₐ/RT)
Where:
- k = rate constant
- A = frequency factor (collision frequency)
- Eₐ = activation energy (J/mol)
- R = gas constant (8.314 J/mol·K)
- T = temperature (K)
Logarithmic Form: ln k = ln A – Eₐ/RT
Arrhenius Equation: k = A e^(-Eₐ/RT) – Think “k equals A times e to the negative Ea over RT”!
Half-Life and Order of Reaction
Half-Life (t₁/₂): Time required for concentration to reduce to half its initial value.
| Order | Half-Life Relation | Dependence on [A]₀ |
|---|---|---|
| Zero | t₁/₂ = [A]₀/2k | Directly proportional |
| First | t₁/₂ = 0.693/k | Independent |
| Second | t₁/₂ = 1/(k[A]₀) | Inversely proportional |
| Third | t₁/₂ = 3/(2k[A]₀²) | Inversely proportional to square |
First order: Constant half-life! Zero order: t₁/₂ increases as [A]₀ increases. Second order: t₁/₂ decreases as [A]₀ increases!
Order of Reaction
Definition: Sum of exponents in the rate equation.
Types of Reactions:
| Order | Rate Equation | Example |
|---|---|---|
| Zero | Rate = k | Photochemical reactions |
| First | Rate = k[A] | Radioactive decay, N₂O₅ decomposition |
| Second | Rate = k[A]² or k[A][B] | 2HI → H₂ + I₂ |
| Third | Rate = k[A]³ or k[A]²[B] | 2NO + Cl₂ → 2NOCl |
| Pseudo First | Rate = k[A][B] but [B] is constant | Ester hydrolysis in excess water |
Order = Sum of exponents in rate law. NOT from stoichiometry! Always determine experimentally!
Determination of Order
Methods to Determine Order:
- Initial Rate Method: Measure initial rates at different concentrations
- Half-Life Method: Study how t₁/₂ changes with [A]₀
- Integrated Rate Law Method: Plot appropriate graphs
- Differential Method: Use rate vs concentration data
| Order | Integrated Rate Law | Linear Plot |
|---|---|---|
| Zero | [A] = [A]₀ – kt | [A] vs t (straight line, slope = -k) |
| First | ln[A] = ln[A]₀ – kt | ln[A] vs t (straight line, slope = -k) |
| Second | 1/[A] = 1/[A]₀ + kt | 1/[A] vs t (straight line, slope = k) |
Zero order: [A] vs t linear. First order: ln[A] vs t linear. Second order: 1/[A] vs t linear!
Activation Energy & Activated Complex
Activation Energy (Eₐ): Minimum energy required for effective collisions.
Activated Complex: High-energy transitional state between reactants and products.
- Exothermic: Eₐ(forward) < Eₐ(reverse)
- Endothermic: Eₐ(forward) > Eₐ(reverse)
- ΔH = Eₐ(forward) – Eₐ(reverse)
Key Points:
- Catalysts lower Eₐ
- Higher Eₐ means slower reaction
- Activated complex has partial bonds
- Cannot be isolated
Activation energy is the “energy hill” molecules must climb to react. Catalysts make this hill smaller!
Enzyme Catalysis
Enzymes: Biological catalysts (proteins) that speed up biochemical reactions.
Mechanism: Enzyme-Substrate Complex → Products + Enzyme
| Enzyme | Reaction Catalyzed | Substrate |
|---|---|---|
| Urease | Urea hydrolysis | Urea |
| Invertase | Sucrose hydrolysis | Sucrose |
| Zymase | Glucose fermentation | Glucose |
| Diastase | Starch hydrolysis | Starch |
| Maltase | Maltose hydrolysis | Maltose |
Enzyme names usually end in “-ase”! They’re like specialized tools – each enzyme fits only its specific substrate!
Units of Rate Constant
General Formula: Units of k = (concentration)¹⁻ⁿ (time)⁻¹
| Order (n) | Units of k | Example |
|---|---|---|
| 0 | mol dm⁻³ s⁻¹ | Same as rate |
| 1 | s⁻¹ | Radioactive decay |
| 2 | dm³ mol⁻¹ s⁻¹ | 2HI → H₂ + I₂ |
| 3 | dm⁶ mol⁻² s⁻¹ | 2NO + Cl₂ → 2NOCl |
Units of k: As order increases, power of concentration becomes more negative. 1st order: s⁻¹ (simplest)!
Reaction Mechanisms
Rate Determining Step (RDS): Slowest step in a multi-step reaction mechanism.
Reaction Intermediate: Species formed in one step and consumed in another, not in overall equation.
Rate = k[NO₂]² (independent of [CO])
Mechanism:
- NO₂ + NO₂ → NO₃ + NO (slow, RDS)
- NO₃ + CO → NO₂ + CO₂ (fast)
NO₃ is reaction intermediate
Rate law comes from RDS! Intermediates don’t appear in overall equation but appear in mechanism!