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R2.2Reactivity 2 — How much, how fast, how far?

How fast? The rate of chemical change — IB Chemistry HL Revision Notes & Topical Questions

40 min read 66 parts Lesson 2 of 3 1 quiz

📚 IB Chemistry HL Hub / Topical Revision
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Curriculum Equation Key قوانين الرقم الهيدروجيني والاتزان الأيوني - pH & Acid-Base Equilibrium FormulasFast Exam Solution
pH = -log[H⁺] | pOH = -log[OH⁻] | pH + pOH = 14 | [H⁺][OH⁻] = 10⁻¹⁴
Word Equation & Conditions: الرقم الهيدروجيني = -لوغاريتم تركيز أيون الهيدروجين | pH + pOH = 14
Location & Specifications: الاتزان الأيوني، كيمياء 3 ثانوي، التحصيلي، و AP Chemistry.
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Lesson summary

R2.2 How fast? The rate of chemical change — IB CHEMISTRY HL notes: Measure and explain reaction rates with collision theory, and use rate equations, mechanisms and the Arrhenius equation (HL).Definition — the change in concentration of a reactant or product per unit time, in mol…

  • Rate at a time
  • Gas volume
  • Colorimetry
  • Conductivity
  • Clock reaction
  • Requirement

Part 1 of 10Rate of reaction

Definition
  • Definition the change in concentration of a reactant or product per unit time, in mol dm−3 s−1.
Graph
  • Graph the steeper the concentration–time graph, the faster the rate.
Rate at a time
  • Rate at a time the gradient of the tangent to the curve at that time.
Sign
  • Sign the rate is positive, so the negative gradient of a reactant graph is changed to positive.
Tangent
  • Tangent draw a large gradient triangle using gridline intersections to reduce error.

Part 2 of 10Measuring rates

Principle
  • Principle follow a property proportional to concentration, such as mass, volume or colour.
Mass loss
  • Mass loss a balance records escaping gas; the gas must be dense enough, so CO2 works but H2 does not.
Gas volume
  • Gas volume a gas syringe or an inverted measuring cylinder over water, the latter only for gases of low solubility.
Colorimetry
  • Colorimetry measures light passing through a coloured solution, but not for coloured precipitates.
Titration
  • Titration samples are taken and quenched, then titrated.
    • (the reaction is stopped)
Conductivity
  • Conductivity follows changing ion concentration.
Clock reaction
  • Clock reaction times a fixed visible point, as in the disappearing cross with sodium thiosulfate and acid, giving one data point per run.

Part 3 of 10Collision theory

Requirement
  • Requirement particles must collide with enough energy and the correct orientation.
Collision frequency
  • Collision frequency how often particles collide, changed by concentration, pressure, temperature and surface area.
Collision energy
  • Collision energy the combined energy of the colliding particles.
Activation energy
  • Activation energy the minimum energy for a successful collision, fixed for a given route.
Collision geometry
  • Collision geometry orientation matters most for large molecules such as proteins.
Success
  • Success most collisions fail because of insufficient energy rather than wrong orientation.
Rate
  • Rate depends on the number of successful collisions per unit time.

Part 4 of 10Factors affecting rate

Concentration
  • Concentration more particles per unit volume give more collisions and more successful collisions.
Pressure
  • Pressure in gases, it acts like concentration.
Temperature
  • Temperature particles collide more often and a larger proportion have energy ≥ Ea;
    • the second effect is more important.
Surface area
  • Surface area more surface particles of a solid can collide, so powders react faster than lumps.
Catalyst
  • Catalyst provides a lower-Ea pathway, so a greater proportion of collisions succeed.

Part 5 of 10Energy profiles and catalysts

Energy profiles and catalysts
Exothermic profile
  • Exothermic profile reactants are higher than products.
Endothermic profile
  • Endothermic profile reactants are lower than products.
Reverse Ea
  • Reverse Ea exothermic Ea(reverse) = Ea(forward) + |ΔH|; endothermic Ea(reverse) = Ea(forward) − ΔH.
Catalyst
  • Catalyst lowers the peak but leaves ΔH and the start and end levels unchanged, and is chemically unchanged.
Homogeneous
  • Homogeneous same phase as the reactants; heterogeneous — different phase.
Industry
  • Industry catalysts lower energy needs, reduce waste and increase selectivity.
Enzymes and transition metals
  • Enzymes and transition metals enzymes are biological catalysts;
    • transition metals are catalysts because of variable oxidation states.

Part 6 of 10Maxwell–Boltzmann curves

Curve
  • Curve shows how particle energies are distributed at a temperature, starting at the origin and never touching the x-axis.
Peak
  • Peak the most probable energy.
Area
  • Area the area under the curve to the right of Ea is the number of particles able to react;
    • the total area stays constant.
Higher temperature
  • Higher temperature the curve flattens, the peak lowers and shifts right, and more particles exceed Ea;
    • the curves cross once.
Catalyst
  • Ea moves left on the same curve, so more particles exceed it.

Part 7 of 10Rate equations (HL)

Rate equation
  • Rate equation rate = k[A]m[B]n, found experimentally and not from the stoichiometric equation.
    What the symbols mean: rate (mol dm−3 s−1) · k = rate constant · [ ] = concentration (mol dm−3) · each power is the order with respect to that reactant
    • (units depend on the overall order)
Order
  • Order the power of a concentration;
    • zero order has no effect, first order doubles the rate when concentration doubles, second order quadruples it.
Overall order
  • Overall order m + n.
Initial rates
  • Initial rates compare experiments in which only one concentration changes: order = rate factor.
    • (concentration factor)
Rate constant
  • Rate constant k = rate ÷ concentration terms, with units depending on overall order (s−1 for first order, mol−1 dm3 s−1 for second).
Temperature
  • Temperature k increases with temperature; the rate roughly doubles for each 10 K rise.

Part 8 of 10Mechanisms (HL)

1Mechanism
  • Mechanism a sequence of elementary steps summing to the overall equation.
2Intermediate
  • Intermediate made in one step and used in another, so it does not appear in the overall equation or the rate equation.
3Rate-determining step
  • Rate-determining step the slowest step; its reactants appear in the rate equation.
4Molecularity
  • Molecularity the number of particles in an elementary step: unimolecular, bimolecular or termolecular (rare).
5Energy profile
  • Energy profile each step has a peak; the step with the greatest activation energy is rate-determining.
6Kinetics
  • Kinetics it can disprove a mechanism but not prove it.

Part 9 of 10The Arrhenius equation (HL)

Equation
  • Equation k = Ae−Ea/RT.
    What the symbols mean: k = rate constant · A = pre-exponential (frequency) factor · Ea = activation energy (J mol−1) · R = 8.31 J mol−1 K−1 · T = temperature (K)
Terms
  • A the frequency factor, Ea the activation energy, R = 8.31 J K−1 mol−1 and T in kelvin.
Logarithmic form
  • Logarithmic form ln k = −Ea/RT + ln A.
Graph
  • Graph ln k against 1/T is a straight line with gradient −Ea/R and intercept ln A.
Calculation
  • Ea = −gradient × R, and A = eintercept.
Effects
  • Effects higher Ea gives a lower k, and higher T gives a higher k.

Part 10 of 10More from the specification

1Reaction order with respect to reactant
  • The power to which a reactant concentration is raised in the experimental rate equation.
    • (rate = k [A]^m [B]^n)
2Maxwell-Boltzmann distribution curves
  • Energy distribution graphs showing the fraction of molecules possessing sufficient kinetic energy (>= Ea) at a given temperature.
3Effect of temperature on distribution
  • Higher temperature flattens and broadens the curve to the right, dramatically increasing the proportion of particles with energy exceeding activation energy.
4Catalyst mechanism on distribution
  • Catalysts provide an alternative pathway with lower activation energy, shifting Ea to the left without altering curve shape.
5Hydrogen
  • A gas that gives a squeaky pop with a lighted splint.
6Maxwell-Boltzmann distribution
  • A graph showing how particle energies are spread at a given temperature.

Key terms

TermMeaning
Rate of reactionChange in concentration per unit time.
Activation energyMinimum energy for a successful collision.
Order of reactionThe power of a concentration in the rate equation.
Rate constantThe constant k in the rate equation, dependent on temperature.
Rate-determining stepThe slowest elementary step in a mechanism.
IntermediateA species formed and used up within a mechanism.
Arrhenius factorThe pre-exponential factor A.

Avoid these mistakes

1
Giving the rate of a reactant graph as negative.
2
Saying a catalyst changes ΔH or increases the particles' energy — it lowers Ea.
3
Reading orders from the stoichiometric coefficients.
4
Using temperature in °C in the Arrhenius equation.
5
Saying the rate-determining step has the highest peak rather than the greatest activation energy.

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