Rates of Reaction
• Rate of reaction is a measure of the progress of a reaction in a certain amount of
time.
• Rate of reaction = Quantity of reactant used up, product formed, change in mass,
volume of gas produced, 1 / time taken.
• We can measure the change in mass of a reaction to see rate. For example, solid
marble chips reacting with hydrochloric acid. They produce a gas which, if it can
escape, the mass decreases as the reaction takes place. Using this we can work out
the mass of gas escaped.
• So in this case, change in mass (g) / time (s). Always include the units.
• Another way to show rate is to draw a graph of its progress as it happens (time on x
axis, mass on y). We can see fast and slower parts of the reaction. At the end it
plateaus as all the reactants are used up.
• We can use the volume of gas produced (product). If we attach a delivery tube with a
bung (connected to reaction, conical flask) to a measuring cylinder upside down filled
with water inside a trough, the gas will displace the water as it progresses allowing us to
measure volume. Always measure volumes from the bottom of the meniscus (the
bottom of a curve). cm^2 for volume. Volume increases fast, slows down and plateaus
when all reactants are used up.
• This could also be done with a gas syringe. It is useful for more accurate
measurements OR if the gas is soluble (so it doesn’t dissolve in water). This also works
with a delivery tube and bung. HOWEVER it is quite prone to leaking which impacts
results.
• Another way to measure rate of reaction is through a colour change. Draw a cross on
a piece of paper below your conical flask. A precipitate (solid formed) makes the
reaction go cloudy, so as it progresses eventually we can no longer see the cross. We
time how long it takes for us to no longer be able to see the cross. As we don’t have a
particular measurement, we use 1 for product formed and 0 for not formed. 1 / time. Unit
is /s, for example 0.033/s.
• If asked to describe a rate of reaction graph, explain how the curve is steeper to show
faster rate of reaction.
• Rate of reaction decreases as there are less reacting particles in a given volume,
meaning the particles are less likely to collide (with enough energy) /// less successful
collisions.
,
Interpreting rate graphs
• Time (s) is on the x axis. Reactant used up / product produced is on the y axis.
• Always start at the origin (zero). The rate of reaction is fastest (steepest) at the
beginning. As the reactants are used up it gets less steep and it flattens out (plateaus)
as all the product is used up / no more is produced.
• Compare: Time for reaction to stop, rate (gradient), amount of product produced.
• To calculate the rate of reaction at a specific time, draw a tangent. The only point of
the curve it’s touching is the specific time. Aim for what looks accurate and convenient
points to work with. Change in y / change in x. Broad range of what answer could be.
Factors affecting rates of reaction
• Temperature, concentration (particles present in a volume of a solution), pressure
(particles present in a volume of a gas), surface area (to volume ratio), catalysts.
• Temperature – As you increase the temperature, the rate of reaction increases. There
won’t be any additional product.
• Concentration is a measurement of the number of particles dissolved in a given
volume of solvent.
• Pressure is a measurement of the number of particles in a given volume of gas.
• For concentration and pressure, if you increase the number of particles the rate
increases.
• Surface area is the total area of a solid that is exposed to a reaction at any one time.
It involves solids reacting with solutions or gases. As a large block only the outside can
react first. If it is split into smallest pieces there is a larger surface area to volume ratio
(more surface area, same mass). For an even larger surface area use a powder. Larger
SA = higher rate of reaction.
• A catalyst is a substance that speeds up a chemical reaction without itself being used
up. Adding a catalyst increases the rate of reaction. It reduces the activation energy
(same for exothermic and endothermic).
REQUIRED PRACTICAL – Investigating the effect of concentration on rate of reaction -
• For example, sodium thiosulfate and hydrochloric acid. This produces a white
precipitate which makes the solution cloudy and the cross disappear. We need to
change the concentration of sodium thiosulfate (independent variable). We can do this
by increasing the ratio of it to water, (e.g. in 50cm^3, 1:4 of sodium thiosulfate to water,
then 2:3, then 3:2, then 4:1, then 5:0).
, • Time how long it takes for the cross to disappear using a stopclock. Do 2-3 trials to
identify any anomalies and find a mean to draw your graph from. Mix up your
appropriate concentrations of sodium thiosulfate and then a specific amount of
hydrochloric acid, for example 20cm^3. Add the same amount each time. Start the stop
clock as soon as it’s added. When the cross can’t be seen stop the stopclock.
• Concentration (g/dm^3) = Mass (g) / Volume (dm^3)
• As we increase concentration of sodium thiosulfate, time taken decreases. This
means this is inversely proportional having a negative correlation graph.
• Always put independent on the x axis and dependant on the y axis. Use logical to put
on each axis. Plot your graph, label axis appropriately, sensible/accurate scale, use
units, correctly plotted points, draw a line of best fit (often straight, but can also be
curved if appropriate. Just don’t join the dots, one pencil line. Doesn’t need to start at
0.).
Collision theory and activation energy (including catalysts)
• Collision theory states that chemical reactions can only occur when reacting particles
collide with each other and with sufficient energy.
• The minimum amount of energy that particles must have to react is called the
activation energy.
• By increasing the concentration of reactants in solution, the pressure of reacting
gases, and the surface are of solid reactants increases the frequency of collisions
(number per second) and so increases the rate of reaction.
• Increasing the temperature increases the frequency of collisions and makes the
collisions more energetic, so increases rate of reaction.
- The particles have more energy à so move faster à so there’s more collisions PER
SECOND à so collide with more energy à so more likely to have enough energy to
react.
• If we increase concentration or pressure there is more particles in the same space à
so more likely to collide with enough energy for a successful collision à so more
collisions. Resultingly the reactant is used up in half the time
• Higher surface area à more particles exposed to the other reactants à more frequent
collisions à so faster rate of reactions.
• Catalysts increase the rate of a reaction by providing a different pathway for the
reaction that has a lower activation energy. This means less energy is required by
particles for a successful collision.
, Reversible reactions and equilibrium
• A reversible reaction is one in which the products of the reaction can react to produce
the original reactions again (forwards and backwards). ⇌
• e.g. NH4Cl ßà NH3 + HCl. Changed by changing the conditions. Heat goes forward,
cooling down forms back into ammonium chloride
• One direction is always exothermic (negative energy change) and the other direction
is always endothermic (positive energy change).
• The same amount of energy is transferred in each direction/case.
• Hydrated copper sulfate heated strongly breaks down the forces that hold it to water
molecules and water is released as water vapour, leaving us with anhydrous copper
sulfate (a white powder). We can reverse this by adding water. This can be done over
and over again. This can be used as a test for water.
• Equilibrium is when the forward and reverse reactions occur at exactly the same rate
(IN A CLOSED SYSTEM, no substances can be added/removed). One reaction may be
faster than the other, however eventually they will occur at the same rate so there is no
overall change in the amount present (looks as though nothing is changing).
• It doesn’t mean there will always be an even amount of reactants/products or
concentration.
Factors affecting equilibrium
• Le Chantelier’s Principle states that if a system is at equilibrium and a change is
made to any of the conditions (such as temperature, concentration, etc), then the
system responds to counteract the change. Only applies in a closed system.
• For example, in A + B ßà C + D, if the concentration of A is increased then the
reaction forward will be favoured and creates more products of C and D, using more A
and B. Whilst this occurs the equilibrium temporarily shifts towards the right making
more products of C and D. The same applies in the reverse.
• When they make ammonium chloride it is exothermic (releases energy to
surroundings). If we heat it, it returns and is endothermic.
• If we increase the temperature, the reaction that takes IN the heat is favoured
(endothermic). The system reacts to cool it down.
• This creates more of the reactants and less of the product (NH4Cl) as it is being used
to do the backwards reaction and return to equilibrium.
• In a closed system the pressure the gases are directly related to the number of moles
of gas contained in it. If we increase number of moles we increase pressure.
• To figure out what happens see how many moles of gas you have on the LEFT and
• Rate of reaction is a measure of the progress of a reaction in a certain amount of
time.
• Rate of reaction = Quantity of reactant used up, product formed, change in mass,
volume of gas produced, 1 / time taken.
• We can measure the change in mass of a reaction to see rate. For example, solid
marble chips reacting with hydrochloric acid. They produce a gas which, if it can
escape, the mass decreases as the reaction takes place. Using this we can work out
the mass of gas escaped.
• So in this case, change in mass (g) / time (s). Always include the units.
• Another way to show rate is to draw a graph of its progress as it happens (time on x
axis, mass on y). We can see fast and slower parts of the reaction. At the end it
plateaus as all the reactants are used up.
• We can use the volume of gas produced (product). If we attach a delivery tube with a
bung (connected to reaction, conical flask) to a measuring cylinder upside down filled
with water inside a trough, the gas will displace the water as it progresses allowing us to
measure volume. Always measure volumes from the bottom of the meniscus (the
bottom of a curve). cm^2 for volume. Volume increases fast, slows down and plateaus
when all reactants are used up.
• This could also be done with a gas syringe. It is useful for more accurate
measurements OR if the gas is soluble (so it doesn’t dissolve in water). This also works
with a delivery tube and bung. HOWEVER it is quite prone to leaking which impacts
results.
• Another way to measure rate of reaction is through a colour change. Draw a cross on
a piece of paper below your conical flask. A precipitate (solid formed) makes the
reaction go cloudy, so as it progresses eventually we can no longer see the cross. We
time how long it takes for us to no longer be able to see the cross. As we don’t have a
particular measurement, we use 1 for product formed and 0 for not formed. 1 / time. Unit
is /s, for example 0.033/s.
• If asked to describe a rate of reaction graph, explain how the curve is steeper to show
faster rate of reaction.
• Rate of reaction decreases as there are less reacting particles in a given volume,
meaning the particles are less likely to collide (with enough energy) /// less successful
collisions.
,
Interpreting rate graphs
• Time (s) is on the x axis. Reactant used up / product produced is on the y axis.
• Always start at the origin (zero). The rate of reaction is fastest (steepest) at the
beginning. As the reactants are used up it gets less steep and it flattens out (plateaus)
as all the product is used up / no more is produced.
• Compare: Time for reaction to stop, rate (gradient), amount of product produced.
• To calculate the rate of reaction at a specific time, draw a tangent. The only point of
the curve it’s touching is the specific time. Aim for what looks accurate and convenient
points to work with. Change in y / change in x. Broad range of what answer could be.
Factors affecting rates of reaction
• Temperature, concentration (particles present in a volume of a solution), pressure
(particles present in a volume of a gas), surface area (to volume ratio), catalysts.
• Temperature – As you increase the temperature, the rate of reaction increases. There
won’t be any additional product.
• Concentration is a measurement of the number of particles dissolved in a given
volume of solvent.
• Pressure is a measurement of the number of particles in a given volume of gas.
• For concentration and pressure, if you increase the number of particles the rate
increases.
• Surface area is the total area of a solid that is exposed to a reaction at any one time.
It involves solids reacting with solutions or gases. As a large block only the outside can
react first. If it is split into smallest pieces there is a larger surface area to volume ratio
(more surface area, same mass). For an even larger surface area use a powder. Larger
SA = higher rate of reaction.
• A catalyst is a substance that speeds up a chemical reaction without itself being used
up. Adding a catalyst increases the rate of reaction. It reduces the activation energy
(same for exothermic and endothermic).
REQUIRED PRACTICAL – Investigating the effect of concentration on rate of reaction -
• For example, sodium thiosulfate and hydrochloric acid. This produces a white
precipitate which makes the solution cloudy and the cross disappear. We need to
change the concentration of sodium thiosulfate (independent variable). We can do this
by increasing the ratio of it to water, (e.g. in 50cm^3, 1:4 of sodium thiosulfate to water,
then 2:3, then 3:2, then 4:1, then 5:0).
, • Time how long it takes for the cross to disappear using a stopclock. Do 2-3 trials to
identify any anomalies and find a mean to draw your graph from. Mix up your
appropriate concentrations of sodium thiosulfate and then a specific amount of
hydrochloric acid, for example 20cm^3. Add the same amount each time. Start the stop
clock as soon as it’s added. When the cross can’t be seen stop the stopclock.
• Concentration (g/dm^3) = Mass (g) / Volume (dm^3)
• As we increase concentration of sodium thiosulfate, time taken decreases. This
means this is inversely proportional having a negative correlation graph.
• Always put independent on the x axis and dependant on the y axis. Use logical to put
on each axis. Plot your graph, label axis appropriately, sensible/accurate scale, use
units, correctly plotted points, draw a line of best fit (often straight, but can also be
curved if appropriate. Just don’t join the dots, one pencil line. Doesn’t need to start at
0.).
Collision theory and activation energy (including catalysts)
• Collision theory states that chemical reactions can only occur when reacting particles
collide with each other and with sufficient energy.
• The minimum amount of energy that particles must have to react is called the
activation energy.
• By increasing the concentration of reactants in solution, the pressure of reacting
gases, and the surface are of solid reactants increases the frequency of collisions
(number per second) and so increases the rate of reaction.
• Increasing the temperature increases the frequency of collisions and makes the
collisions more energetic, so increases rate of reaction.
- The particles have more energy à so move faster à so there’s more collisions PER
SECOND à so collide with more energy à so more likely to have enough energy to
react.
• If we increase concentration or pressure there is more particles in the same space à
so more likely to collide with enough energy for a successful collision à so more
collisions. Resultingly the reactant is used up in half the time
• Higher surface area à more particles exposed to the other reactants à more frequent
collisions à so faster rate of reactions.
• Catalysts increase the rate of a reaction by providing a different pathway for the
reaction that has a lower activation energy. This means less energy is required by
particles for a successful collision.
, Reversible reactions and equilibrium
• A reversible reaction is one in which the products of the reaction can react to produce
the original reactions again (forwards and backwards). ⇌
• e.g. NH4Cl ßà NH3 + HCl. Changed by changing the conditions. Heat goes forward,
cooling down forms back into ammonium chloride
• One direction is always exothermic (negative energy change) and the other direction
is always endothermic (positive energy change).
• The same amount of energy is transferred in each direction/case.
• Hydrated copper sulfate heated strongly breaks down the forces that hold it to water
molecules and water is released as water vapour, leaving us with anhydrous copper
sulfate (a white powder). We can reverse this by adding water. This can be done over
and over again. This can be used as a test for water.
• Equilibrium is when the forward and reverse reactions occur at exactly the same rate
(IN A CLOSED SYSTEM, no substances can be added/removed). One reaction may be
faster than the other, however eventually they will occur at the same rate so there is no
overall change in the amount present (looks as though nothing is changing).
• It doesn’t mean there will always be an even amount of reactants/products or
concentration.
Factors affecting equilibrium
• Le Chantelier’s Principle states that if a system is at equilibrium and a change is
made to any of the conditions (such as temperature, concentration, etc), then the
system responds to counteract the change. Only applies in a closed system.
• For example, in A + B ßà C + D, if the concentration of A is increased then the
reaction forward will be favoured and creates more products of C and D, using more A
and B. Whilst this occurs the equilibrium temporarily shifts towards the right making
more products of C and D. The same applies in the reverse.
• When they make ammonium chloride it is exothermic (releases energy to
surroundings). If we heat it, it returns and is endothermic.
• If we increase the temperature, the reaction that takes IN the heat is favoured
(endothermic). The system reacts to cool it down.
• This creates more of the reactants and less of the product (NH4Cl) as it is being used
to do the backwards reaction and return to equilibrium.
• In a closed system the pressure the gases are directly related to the number of moles
of gas contained in it. If we increase number of moles we increase pressure.
• To figure out what happens see how many moles of gas you have on the LEFT and
