A-level Chemistry: 3.1.11 Rate Equations & 3.1.12 Kp

Change in amount of reactant or product per unit time

Mol dm-3 s-1

rate = k[A]m[B]n

orders of the reaction

k = rate constant

Tell you how reactant concentrations affect the rate

(e.g. m tells you how concentration of reactant A affects rate and n tells you same for reactant B)

m + n

From experiments

The order of reaction with respect to A is 0

The order of reaction with respect to A is 1

The order of reaction with respect to A is 2

reactant concentrations to rate at a particular temperature

faster the reaction

the same

• ∵ ↑ temp. = rate of reaction ↑

• ∵ increasing no. of collisions between reactant molecules + energy of each collisions

• But conc. of reactants and orders of reaction stay the same

• So k must increase for rate equation to balance

Rate right at the start of the reaction

Find it from a concentration-time graph by calculating the gradient of the tangent at time = 0

Used to create rate equations

1. Repeat an experiment several times using different initial concentrations of the reactants
   

   • Usually only change 1 conc. of at time

2. Calculate initial rate for each experiment

3. See how initial concentration affects initial rates and figure out the order for each reactant

• Iodine clock reaction

• Reactions that produce precipitates
  

  • Measure time it takes for mark underneath reaction vessel to disappear

• Other reactions
  

  • Measure time taken for small amount of product to be formed

1. Add dilute sulfuric acid and starch solution to beaker

2. Add sodium thiosulfate to reaction mixture
   

   1. Add potassium iodide solution

   2. Add hydrogen peroxide solution

3. Sodium thiosulfate reacts with iodine being formed

4. Once all sodium thiosulfate is used up = any more iodine formed remains in solution

5. Turns starch indicator blue-black

6. Varying conc. of iodide or hydrogen peroxide while keeping everything constant = different times for colour change
   

   • Used to work out reaction order

H₂O₂ + 2I⁻ + 2H⁺ → I₂ + 2H₂O

I₂ + 2S₂O₃²⁻ → 2I⁻ + S₄O₆²⁻

• Measuring Initial Reaction Rate

• Continuous Monitoring

• Can follow reaction all way to its end by recording amount of product (or reactant) you have at regular time intervals

• Use results to work out how rate changes over time

• Loss of Mass

• Colour Change

• Gas Volume

• Change in pH

Measures absorbance

More concentrated the colour of the solution is

Can measure change in absorbance:

1. Plot calibration curve

   • Graph of known concentrations of coloured solution (e.g. I₂) plotted against absorbance

2. During experiment, take small samples from your reaction solution at regular intervals and read the absorbance

3. Use calibration curve to convert absorbance at each time point into a concentration

1. If gas is given off, system will lose mass

2. Can measure this at regular intervals with a balance

3. Use mole calculations to work out how much gas you've lost

4. & thus how many moles of reactants are left

1. If gas given off, could collect it in gas syringe & record how you've got at regular time intervals (e.g. every 15s)

   • e.g. acid + carbonate = CO₂

   • e.g. magnesium ribbon + HCl

2. Find conc. of reactant at each time point

3. Use ideal gas equation to work how many mole of gas

4. Then use molar ratio to work out conc. of reactant

If reaction produces or uses H+ ions, can measure pH of solution at regular intervals & calculate the conc. of H+

Can use data from concentration-time graph to construct it

Reaction order

1. Find gradient at various points on graph

   1. Gives you rate at that particular concentration

   2. For curve, need to draw tangents

2. Plot each point on new graph with axes rate and concentration

   1. Then draw smooth line or curve through points

   2. Shape of line will tell you the order of the reaction with respect to that reactant

different rate

Slowest step in a multi-step reaction

By the step with the slowest rate = rate determining step (aka rate-limiting step)

If a reactant appears in the rate equation, it must affect the rate. ∴ this reactant, or something derived from it, must be in the rate determining step.

first

chemical equation

Shows the number of molecules of that reactant that are involved in the rate determining step

intermediate

How rate constant (k) varies with temperature (T) and activation energy (Ea)

• k = rate constant

• Ea = activation energy (J)

• T = temperature (K)

• R = gas constant (8.31 J K-1 mol-1)

• A = the Arrhenius constant

• Large Ea = slow rate

• If reactions has high Ea, not many reactant particles have enough energy to react

• ∴ few of collisions will result in reaction occurring & rate will be slow

• Higher temperatures mean reactant particles move around faster with more energy

• More likely to collide and more likely to collide with E ≥ Ea

• So reaction rate ↑

Ea or Arrhenius constant

In a mixture of gases, each gas exerts its own pressure = partial pressure

sum of all partial pressure of the individual gases

Proportion of a gas mixture that is made up a particular gas

Mole fraction of a gas = (Number of moles of the gas) ÷ (Total number of moles of all gases in the mixture)

Partial pressure of a gas = (Mole fraction of the gas) × (Total pressure of the gas mixture)

Kp

Temperature

| (Kp is only valid for a given temp.)

• ∵ changing temp. changes how much product is formed at equilibrium

• Changes mole fractions of gases present which changes their partial pressures

pressure

Equilibrium will shift to keep it the same

won't affect

• Reaction occurs when molecules have E ≥ Ea

• Doubling T causes many more molecules to have this E

• Whereas doubling [E] only doubles the number with this E

Kp increases for endothermic reactions and decreases for exothermic reactions.

At time 0, the conc. are known

• Calculate gradient of the tangent

• At t = 0 OR at start of graph

• By measuring the time it takes for the colour change to happen

• Shorter the time = faster rate