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

Section 1 of 32Rate of reaction

Definition

Definition the change in concentration of a reactant or product per unit time, in mol dm⁻³ s⁻¹.

Formula

• Formula rate = change in concentration ÷ time.

Graph reading

• Graph reading the steeper the gradient of a concentration-time graph, the faster the rate.

Rate at a time

• Rate at a time draw a tangent to the curve and calculate its gradient, Δy/Δx, using a large triangle that meets gridlines.

Initial rate

• Initial rate the gradient of the tangent at t = 0.

Sign

• Sign rate is always positive, so a negative reactant gradient is quoted as a positive rate.

Equilibrium

• Equilibrium in reversible reactions the concentrations stay constant once dynamic equilibrium is reached.

Section 2 of 32Measuring the rate of reaction

• Principle
  • follow a property that changes with time and is proportional to the concentration of a reactant or product.

Method	Property followed	Limitation
Mass loss	Gas escaping, e.g. CO2, with cotton wool in the neck	Gas must be dense; CO2 (M = 44) works but H2 (M = 2) does not
Gas volume	Gas syringe or inverted measuring cylinder over water	Gas must have low solubility, e.g. H2
Colorimetry	Light passing through a coloured solution	Cannot follow precipitates, which scatter light
Conductivity or pH	Change in ion concentration	Needs ionic species that change
Titration of quenched samples	Concentration at fixed times	Sampling affects the mixture and is not continuous
Clock reaction	Time to a fixed visible point, e.g. disappearing cross	Gives only one value, an average rate (rate ∝ 1/t)

• Plateau
  • the reaction has finished when the gas volume levels off.

Section 3 of 32Sampling, quenching and clock reactions

1Quenching

• Quenching stopping the reaction in each sample by cooling or dilution so its concentration can be found by titration.

2Disappearing cross

• Na₂S₂O₃(aq) + 2HCl(aq) → 2NaCl(aq) + SO₂(g) + H₂O(l) + S(s), where yellow sulfur obscures a cross.

3Dissolving metal

• Dissolving metal timing how long a piece of magnesium takes to dissolve in acid is another clock method.

4Method steps

• Method steps start the timer as the reactants mix, keep other variables constant and record at regular intervals.

Section 4 of 32Collision theory

Requirements

• Requirements particles must collide with energy at least equal to the activation energy and in the correct orientation.

Four factors

• Four factors collision frequency, collision energy, activation energy and collision geometry.

Successful collision

• Successful collision has sufficient energy and the correct orientation; otherwise the particles bounce apart.

Rate

• Rate depends on the number of successful collisions per unit time.

Why most collisions fail

• Why most collisions fail mainly because they lack the activation energy, not because of orientation.

Orientation

• Orientation matters most for large molecules such as proteins with active sites.

Section 5 of 32Factors affecting rate

Factor	Change	Explanation
Concentration	Increase	More particles per volume, so more frequent collisions
Pressure of gases	Increase	Same particles in a smaller volume, so more frequent collisions
Surface area	Increase (powder)	More exposed particles, so more collisions
Temperature	Increase	More frequent collisions and a greater proportion with energy ≥ Ea; the energy effect dominates
Catalyst	Added	Alternative route with lower Ea, so more collisions succeed

Rule of thumb

• Rule of thumb raising the temperature by 10 K often roughly doubles the rate.

Kinetic energy

• KE = ½mv²,
• so at the same temperature lighter particles move faster and diffuse more quickly.
  What the symbols mean: E_k = kinetic energy (J) · m = mass (kg) · v = speed (m s⁻¹)

Section 6 of 32Activation energy and energy profiles

Activation energy Ea

Activation energy Ea the minimum energy colliding particles need for a successful reaction.

Exothermic

• Exothermic reactants are higher in energy than products; endothermic — reactants are lower.

Reverse reaction

• E_a(reverse) = E_a(forward) − ΔH, so it is larger for an exothermic reaction.

Calculation from data

• Calculation from data not required at SL.

Kinetic stability

• Kinetic stability a negative ΔH shows a reaction is favourable, but a high E_a can make it too slow to observe.

Section 7 of 32Catalysts

Action

• Action provide an alternative pathway with lower E_a and are unchanged at the end.

Effect on thermodynamics

• Effect on thermodynamics they do not change ΔH or the position of equilibrium.

Homogeneous

• Homogeneous in the same phase as the reactants.

Heterogeneous

• Heterogeneous in a different phase, e.g. solid iron in the Haber process.

Enzymes

• Enzymes biological catalysts that allow reactions at lower temperature and pressure.

Transition metals

• Transition metals often good catalysts because of variable oxidation states.

Sustainability

• Sustainability catalysts cut energy use, are reused in small amounts so give less waste, and increase selectivity.

Section 8 of 32Maxwell-Boltzmann distribution

Definition

Definition a curve showing the spread of kinetic energies of particles at a given temperature.

Shape

• Shape starts at the origin, peaks at the most probable energy and approaches but never touches the x-axis.

Area

• Area equals the total number of particles and stays constant when T changes.

Higher temperature

• Higher temperature the peak is lower and shifts right, the curves cross once and the high-energy tail is higher.

Catalyst

• Catalyst moves the E_a line to the left on the same curve, so the whole area to the right of the new E_a can react.

Section 9 of 32Why temperature raises the rate

1

Temperature increases.

2

Average kinetic energy increases.

3

The Maxwell-Boltzmann curve flattens and shifts right.

4

More particles have energy ≥ E_a.

5

More successful collisions occur per second.

6

The rate increases.

Section 10 of 32More from the specification

More from the specification

Nuclear magnetic resonance identification

• Nuclear magnetic resonance identification measuring chemical shifts and proton spin-spin splitting provides a definitive means to determine organic molecular connectivity.

Reaction rate definition

• The change in concentration of a reactant or product per unit time, measured in mol dm^-3 s^-1.

Maxwell-Boltzmann distribution curve

• A probability density distribution displaying the range of molecular kinetic energies present at a given absolute temperature.

Collision theory requirements

• Successful chemical reactions occur only when collisions result in particle kinetic energy exceeding activation energy (E >= Ea) with suitable geometry.

Reaction energy profiles

• Graphical diagrams showing potential energy versus reaction progress, illustrating activation energy barriers with and without catalysts.

Maxwell-Boltzmann distribution

• A graph showing how particle energies are spread at a given temperature.

Energy profile

• Diagram showing energy during a reaction

Section 11 of 32Following the progress of a reaction

Principle

• Principle measure a quantity that changes steadily and is proportional to the concentration of a reactant or product.

Conductivity

• Conductivity a probe follows reactions that make or remove ions, because ion concentration affects electrical conductivity.

Other options

• Other options gas volume with a syringe, loss of mass on a balance, colorimetry for coloured reagents and a pH meter for acid–base changes.

1

Mix the reactants and start the timer.

2

Record the property at regular intervals.

3

Plot the property against time.

4

Draw tangents to find the rate at chosen times.

Section 12 of 32Clock reactions

Definition

• Definition a non-continuous method in which the time to reach a fixed visible point is measured,
  • such as a cross being hidden by sulfur or a colour appearing.

Rate

• Rate taken as proportional to 1/time.

Advantages

• Advantages quick and uses simple equipment.

Limitations

• Limitations gives only an average rate over the interval, and judging the end point is subjective.

Iodine clock

• Iodine clock a small fixed amount of thiosulfate reacts with iodine as it forms,
  • and when the thiosulfate is used up, iodine gives a sudden blue-black colour with starch.

Section 13 of 32Kinetic theory and collision theory

Kinetic energy

• Kinetic energy the energy of motion, E_k = ½mv²,
  • and the average value is proportional to absolute temperature.
    What the symbols mean: E_k = kinetic energy (J) · m = mass (kg) · v = speed (m s⁻¹)

Collision requirement

• Collision requirement particles must collide, but not all collisions lead to reaction.

Successful collision

• Successful collision particles have at least the activation energy and the correct orientation, especially for large molecules of complex shape.

Catalyst

• Catalyst gives an alternative route with a lower activation energy and does not alter the collision frequency.

Section 14 of 32Changing collision frequency

Concentration

• Concentration more particles in the same volume collide more often.

Gas pressure

• Gas pressure a higher pressure has the same effect as higher concentration.

Surface area

• Surface area only particles on the surface of a solid can collide with the other reactant, so a powder reacts faster than a lump.

Temperature

• Temperature particles move faster, so they collide more often.

Section 15 of 32Planning a rate experiment

• Aim
  • investigate the effect of thiosulfate concentration on its reaction with hydrochloric acid.

1

Vary the thiosulfate concentration by dilution while keeping the total volume constant.

2

Time how long a cross under the flask takes to disappear.

3

Repeat each concentration.

4

Plot 1/time against concentration.

Controlled variables

• Controlled variables temperature, volumes, acid concentration, the same cross and the same observer.

Interpretation

• Interpretation a straight line through the origin suggests the rate is proportional to concentration.

Improvement

• Improvement a colorimeter removes the subjective judgement of the end point.

Section 16 of 32Choosing the best method

Gas syringe

• Gas syringe accurate for a released gas but must be airtight.

Balance

• Balance simple but only suits a dense gas such as CO₂.

Colorimeter

• Colorimeter objective and works with any coloured species.

Clock reaction

• Clock reaction quick but gives only an average rate.

Conductivity probe

• Conductivity probe useful when ions are used or produced, and a data logger gives a continuous record.

Section 17 of 32More from the specification

1Collision theory and reaction rates

• According to collision theory, the rate of a chemical reaction increases when the frequency of collisions between reactant particles increases and when a greater proportion of collisions possess energy exceeding activation energy.

2Maxwell-Boltzmann distribution curve

• A Maxwell-Boltzmann distribution graph shows the spread of molecular kinetic energies in a gas sample;
• increasing temperature flattens and shifts the peak to the right,
• meaning a much larger fraction of molecules exceeds activation energy.

3Reaction progress graphs and gradient

• When reaction progress is monitored over time, the gradient of the curve on a concentration-time graph represents the instantaneous reaction rate,
• which is highest initially and decreases as reactants are consumed.

4Catalytic lowering of activation energy

• Adding a catalyst provides an alternative pathway with reduced activation energy,
• which means an increased number of particles have sufficient energy to react at a given temperature.

5Collision theory and reaction rates

• According to collision theory, the rate of a chemical reaction increases when the frequency of collisions between reactant particles increases and when a greater proportion of collisions possess energy exceeding activation energy.

6Maxwell-Boltzmann distribution curve

• A Maxwell-Boltzmann distribution graph shows the spread of molecular kinetic energies in a gas sample;
• increasing temperature flattens and shifts the peak to the right,
• meaning a much larger fraction of molecules exceeds activation energy.

7Reaction progress graphs and gradient

• When reaction progress is monitored over time, the gradient of the curve on a concentration-time graph represents the instantaneous reaction rate,
• which is highest initially and decreases as reactants are consumed.

8Catalytic lowering of activation energy

• Adding a catalyst provides an alternative pathway with reduced activation energy,
• which means an increased number of particles have sufficient energy to react at a given temperature.

9Energy profile

• Diagram showing energy during a reaction

10Theory

• A tested hypothesis that is widely accepted.

11Gradient

• The change in y divided by the change in x on a graph.

12Maxwell-Boltzmann distribution

• A graph showing how particle energies are spread at a given temperature.

Section 18 of 32Temperature

Temperature

Kinetic energy

• Kinetic energy a rise in temperature makes particles move faster.

Frequency effect

• Frequency effect the number of collisions per second increases.

Energy effect

• Energy effect a greater proportion of collisions has energy ≥ E_a,
• so the proportion of successful collisions increases, and this is the main reason for the large rate increase.

Rule of thumb

• Rule of thumb a rise of about 10 K roughly doubles the rate of many reactions.

Answer structure

• Answer structure state both parts: more frequent collisions and a greater proportion of particles with E ≥ E_a.

Section 19 of 32Surface area

Surface collisions

• Surface collisions only particles on the surface of a solid can collide with the other reactant.

Grinding

• Grinding a lump ground into a powder has more particles on the surface, so collision frequency and rate rise.

Energy

• Energy particle energy is unchanged, so the proportion of successful collisions does not change.

Section 20 of 32Catalysts

Action

• Action an alternative pathway with a lower activation energy, so more collisions are successful.

Chemical state

• Chemical state unchanged at the end, though it may take part in intermediate steps.

Thermodynamics

• Thermodynamics it does not change ΔH or the position of equilibrium.

Examples

• Examples iron in the Haber process, platinum or rhodium in catalytic converters, manganese(IV) oxide in the decomposition of hydrogen peroxide and enzymes in living organisms.

Heterogeneous catalyst

• Heterogeneous catalyst in a different phase from the reactants.

Section 21 of 32Activation energy and energy profiles

Activation energy Ea

Activation energy Ea the minimum energy reactant particles must have to react.

Exothermic

• Exothermic reactants are higher in energy than products.

Endothermic

• Endothermic reactants are lower in energy than products.

Profile shape

• Profile shape rises to a peak then falls, and a catalysed pathway is a lower peak between the same reactant and product levels.

Orientation

• Orientation even with the correct orientation, particles will not react without at least E_a.

Section 22 of 32Comparing the factors

Factor	Collision frequency	Proportion with E ≥ Ea	Effect on rate
Temperature up	Up	Up (large)	Increases a lot
Surface area up	Up	Same	Increases
Catalyst	Same	Up (lower Ea)	Increases

Section 23 of 32Explaining a rate graph

Gas volume against time

• Gas volume against time compares conditions, and a steeper initial gradient means a higher initial rate.

Same final volume

• Same final volume the amount of limiting reagent is unchanged and only the speed changed,
  • as with a higher temperature, a catalyst or a larger surface area.

Different final volume

• Different final volume the quantity of limiting reagent has changed.

Answer pattern

• Answer pattern state the change, relate it to the particles and link it to the graph.

Section 24 of 32More from the specification

More from the specification

Maxwell-Boltzmann Distribution

• A distribution curve displaying the spread of molecular kinetic energies among gas particles at a given temperature, starting at the origin.

Temperature and Distribution Curves

• Raising the temperature flattens the Maxwell-Boltzmann curve and shifts its peak to higher energy, significantly increasing the proportion of particles with energy ≥ Ea.

Collision Theory and Pressure

• Chemical reactions require reactant particles to collide with sufficient energy and correct orientation;
• increasing gas pressure or concentration increases collision frequency.

Measuring Reaction Rates

• Rates can be determined by continuously measuring physical properties such as absorbance of transmitted light, pressure,
• or mass loss over time.

Maxwell-Boltzmann Distribution

• A distribution curve displaying the spread of molecular kinetic energies among gas particles at a given temperature, starting at the origin.

Temperature and Distribution Curves

• Raising the temperature flattens the Maxwell-Boltzmann curve and shifts its peak to higher energy, significantly increasing the proportion of particles with energy ≥ Ea.

Collision Theory and Pressure

• Chemical reactions require reactant particles to collide with sufficient energy and correct orientation;
• increasing gas pressure or concentration increases collision frequency.

Measuring Reaction Rates

• Rates can be determined by continuously measuring physical properties such as absorbance of transmitted light, pressure,
• or mass loss over time.

Concentration

• Moles ÷ volume in dm³, or g/dm³.

Pressure

• Force of gas particles colliding with a surface.

Theory

• A tested hypothesis that is widely accepted.

Maxwell-Boltzmann distribution

• A graph showing how particle energies are spread at a given temperature.

Collision theory

• Particles react only if they collide with at least the activation energy and the correct orientation.

Section 25 of 32What the curve shows

Meaning

• Meaning the spread of kinetic energies of particles in a gas or liquid at a given temperature.

Axes

• Axes kinetic energy on the x-axis and the number (or fraction) of particles on the y-axis.

Start

• Start the curve begins at the origin.

Peak

• Peak it rises to the most probable energy and then falls with a long tail to the right.

Tail

• Tail it approaches the energy axis but never touches it, since there is no maximum energy.

Area

• Area the total area equals the total number of particles and is the same at every temperature.

Activation energy

• Activation energy shown as a vertical line, and the area to its right represents the particles able to react.

Section 26 of 32Effect of temperature

Effect of temperature

Energy gain

• Energy gain particles gain kinetic energy and move faster.

Curve change

• Curve change the peak shifts to the right and becomes lower and broader, because the area must stay constant.

Reacting particles

• Reacting particles a larger proportion has energy greater than E_a, so the shaded area increases.

Rate

• Rate successful collisions per second rise, and collisions are also more frequent.

Drawing

• Drawing the higher-temperature curve has its peak lower and to the right, crosses the first curve once and ends above it in the tail, with curves labelled T₁ and T₂.

Examination style

• Examination style questions often ask about decreasing the temperature, which applies the same theory in reverse.

Section 27 of 32Effect of a catalyst

Effect of a catalyst

Distribution

• Distribution a catalyst does not change the energy distribution, so the curve is unchanged.

Activation energy

• Activation energy it lowers E_a, so a second E_a line is drawn to the left of the original.

Reacting particles

• Reacting particles the area to the right of the new line is much larger, so more particles can react.

Shading

• Shading shade the whole area to the right of the new E_a, including the area already beyond the old line, not just the extra strip.

Section 28 of 32Drawing a good sketch

1

Label the axes: kinetic energy (x) and number of particles (y).

2

Start both curves at the origin and keep them off the x-axis.

3

Draw the higher-temperature curve with a lower, broader peak.

4

Mark E_a (and E_a with catalyst) with vertical lines.

5

Shade the areas representing particles with E ≥ E_a.

Section 29 of 32Linking the distribution to rate

1Strong temperature dependence

• Strong temperature dependence the fraction of particles with energy at least E_a increases roughly exponentially with temperature,
  • so small rises cause large rate increases.

2Catalyst effect

• Catalyst effect lowering E_a enlarges the area beyond E_a without any temperature change.

3Wording

• Wording use the phrase "proportion of particles with energy greater than or equal to the activation energy".

4Care

• Care never say all particles gain energy equally.

Section 30 of 32More from the specification

More from the specification

Spectrometric molecular weight determination

• Spectrometric molecular weight determination measuring mass-to-charge ratios of parent ions provides a definitive analytical means to establish exact molecular formulas.

Measuring the rate of reaction

• The rate of a chemical reaction is defined as the change in concentration of reactants or products per unit of time.
  • (rate = delta_[product] / delta_t = -delta_[reactant] / delta_t)

Experimental methods for measuring rates

• Techniques to measure reaction rate include tracking the volume of gas evolved using a gas syringe, measuring change in mass of the system on a balance,
• or monitoring absorbance with a colorimeter.

Rate graphs and gradient determination

• On a graph of gas volume or product concentration plotted against time, the rate at any given instant is equal to the gradient (slope) of the tangent drawn to the curve.

Effect of surface area and pressure

• Increasing the surface area of solid reactants or increasing the pressure of reacting gases increases collision frequency between particles, thereby increasing the rate of reaction.

Concentration

• Moles ÷ volume in dm³, or g/dm³.

Reactants

• The substances that react together at the start of a reaction.

Product

• A substance formed in a reaction.

Volume

• The space taken up by an object.

Pressure

• Force of gas particles colliding with a surface.

Gradient

• The change in y divided by the change in x on a graph.

Orientation

• The alignment of colliding particles needed for reaction.

Energy profile

• Diagram showing energy during a reaction

Section 31 of 32Key terms

Term	Meaning
Rate of reaction	The change in concentration of a reactant or product per unit time.
Collision theory	Reactions occur when particles collide with energy ≥ Ea and the correct orientation.
Activation energy	The minimum energy for a collision to lead to reaction.
Catalyst	A substance that speeds a reaction by giving a lower-Ea route and is unchanged at the end.
Maxwell-Boltzmann distribution	The spread of kinetic energies of particles at a given temperature.
Quenching	Stopping a reaction in a sample so it can be analysed.

Term	Meaning
Clock reaction	A method that measures the time to reach a fixed point.
Collision frequency	The number of collisions per unit time.
Kinetic theory	A model describing matter as moving particles.
Orientation	The alignment of colliding particles needed for reaction.
Conductivity	The ability of a solution to carry current via ions.

Term	Meaning
Activation energy	The minimum energy needed for a successful collision.
Catalyst	A substance that speeds up a reaction by providing a lower-Ea pathway.
Surface area	The total area of solid exposed to reactants.
Successful collision	A collision with sufficient energy and the correct orientation.
Heterogeneous catalyst	A catalyst in a different phase from the reactants.

Term	Meaning
Maxwell-Boltzmann distribution	The spread of particle energies at a given temperature.
Most probable energy	The energy at the peak of the curve.
Activation energy	The minimum energy for reaction, marked on the energy axis.
Kinetic energy	The energy of motion of a particle.

Section 32 of 32Exam technique

1

Giving a negative rate — quote the rate as a positive value.

2

Saying a catalyst changes ΔH or the equilibrium position — it lowers E_a only.

3

Explaining the effect of temperature by collision frequency alone — the energy effect dominates.

4

Drawing Maxwell-Boltzmann curves that touch the x-axis or cross twice — they approach the axis and cross only once.

1

Using the clock time as the rate — the rate is proportional to 1/t.

2

Saying increased surface area raises the energy of the particles — it only increases collisions.

3

Saying more concentrated reactants make all collisions successful — only the frequency increases.

4

Changing the volume of reagent between runs — keep the total volume constant.

1

Saying a catalyst gives particles more energy — it lowers the activation energy.

2

Saying surface area changes the activation energy — it changes only collision frequency.

3

Saying temperature only changes collision frequency — it also raises the proportion with E ≥ E_a.

4

Omitting the orientation requirement — a successful collision needs the right orientation and enough energy.

1

Drawing a curve that touches the x-axis at high energy — it only approaches the axis.

2

Shifting the peak upward at higher temperature — the peak is lower.

3

Moving the curve when a catalyst is added instead of moving the E_a line.

4

Shading only the strip between the two E_a lines for the catalysed total — shade everything beyond the new E_a.