C5.2 Controlling Reactions Flashcards
Rate of reaction
Measure of how quickly reactants are used or products are formed
Rate of reaction = amount of reactant used / time taken
Rate of reaction = amount of product formed / time taken
Measuring volume of gas
Example:
Magnesium reacts with dilute hydrochloric acid to form magnesium chloride solution and hydrogen:
Mg(s) + 2HCI(aq) -> MgCl2(aq) + H2(g)
Measure how fast hydrogen is produced by carrying out the following steps:
Place dilute hydrochloric acid in a conical flask connected to a gas syringe
Add a piece of magnesium ribbon into the acid, stopper the flask, and start a stop clock immediately
Draw a results table and record the time and volume of hydrogen at regular intervals
Analyse your results by drawing a line graph (curved)
Before adding the stopper to the conical flask in this investigation, you should push the plunger all the way in, to make sure the reading starts at 0 cm³
You also need to make sure that you stopper the flask as soon as the reaction starts, to make sure all the hydrogen is collected
Calculating mean rate of reaction
Example: Calculate the mean rate of reaction between 60s and 90s
1) Calculate the changes in volume and time
Change in volume = (50 cm³ - 38 cm³) = 12 cm³
Change in time = (90s - 60s) = 30s
2) Calculate the mean gradient (equal to the mean rate of reaction)
Change in volume / change in time = 12cm³ / 30s = 0.4cm³/s
Instantaneous (specific) rate of reaction
Example: Find the rate of reaction at 75 seconds
In this example, a tangent to the curve has been drawn at 75s
Instantaneous rate of reaction at this time = gradient of the tangent
Why reactions go faster at higher temperatures
A reaction can only happen if:
Reactant particles collide with each other, and colliding particles have enough energy to react
Collision that leads to a reaction = successful collision
Collision will not be successful if particles have less energy than the activation energy
As temperature of a reaction mixture increases:
Particles move more quickly, so they collide more often
Greater proportion of the colliding particles have the activation energy or more
The greater the rate of successful collisions, the greater the rate of reaction
Investigating effect of temperature on a reaction
Example: Reaction between sodium thiosulfate solution and hydrochloric acid is often used to investigate the factors affecting rates of reaction:
Na2S2O3(aq) + 2HCI(aq) -+ 2NaCI(aq) + H2O(l) + SO2(g) + S(s)
The equation above looks complicated, but you can easily measure how quickly a pale yellow precipitate of sulfur appears:
1 In a beaker, mix the reactants
2 Look down through the mixture to a cross drawn on a piece of paper:
The longer it takes for the cross to disappear, the lower the rate of reaction
3 You can vary the temperature of the reaction mixture by warming up one of the solutions before mixing
Reaction will start as soon as the sodium thiosulfate solution and dilute hydrochloric acid are mixed together
This means that you should start the stop clock as soon as you mix the solutions
You should stop the clock as soon as the cross disappears, so that the timing will stop at the same point in all the reactions
How you can use reaction times
Rate of reaction is inversely proportional to the reaction time
This means that 1 / time is directly proportional to the rate:
1 / reaction time ∝ rate of reaction
You can plot a graph with rates of reaction instead of reaction times, if you calculate 1 / time for each result:
For example, a reaction time of 10s gives a rate of 1 / 10s = 0.1/s
Sometimes 1000/time is used instead of 1/time, as this can make the vertical axis scale easier to work out
Why reactions go faster at higher concentrations
Concentration of a solution is measure of how much solute is dissolved in the solvent
The more concentrated a solution is, the more solute is dissolved in the solvent
If a reaction involves one or more reactants in solution, rate of reaction increases as the concentration increases:
The particles become more crowded, so they collide more often
The energy stored in the particles does not change, but because the rate of collisions increases, rate of successful collisions increases
How to investigate effect of concentration on reaction rate
Reaction between magnesium ribbon and hydrochloric acid is often used to investigate rates of reaction
Mg(s) + 2HCI(aq) - MgCI2(aq) + H2(g)
You could use a gas syringe to measure the volume of hydrogen gas produced
However, it is easier to measure the time taken for a piece of magnesium ribbon to be used up in the reaction
You can add water to the hydrochloric acid to reduce its concentration
When you have recorded the reaction times at different concentrations of acid, you can calculate 1/time to obtain the reaction rates
When carrying out this investigation, the length of the magnesium ribbon needs to be controlled because the reaction time is affected by the mass (and therefore size) of the piece of magnesium, so this must be kept the same to make it a fair test
Limiting reactants
In a reaction involving 2 reactants, one of the reactants will be limiting and the other in excess
Amount of product will be proportional to amount of limiting reactant
Why reactions go faster at higher pressures
If a reaction involves 1 or more reactants in gas state:
Rate of reaction increases as the pressure of the gas increases:
Particles in the gas state become more crowded, so they collide more often
Energy stored in particles does not change but, because the rate of collisions increases, the rate of successful collisions increases
Industrial chemical processes often use high pressures to achieve high rates of reaction
Why reactions go faster with powders
Particles in a substance in the solid state can only vibrate about fixed positions: they cannot move from place to place
This means that only the particles at the surface can take part in collisions
Rate of reaction increases as surface area increases because:
More reactant particles are available for collisions
Collisions are more likely, so particles collide more often
Energy stored in the particles does not change but, because the rate of collisions increases, rate of successful collisions increases
Powders have an even larger surface area than lumps, so their reactions are very fast
What matters is surface area : volume ratio of the lumps:
As the size of the lumps decreases, surface area : volume ratio increases
Investigating effect of particle size
Reaction between calcium carbonate and dilute hydrochloric acid is often used to investigate rates of reaction:
CaCO3 (s) + 2HCl(aq) -> CaCl2(aq) + H2O(l) + CO2(g)
You could use a gas syringe to measure the volume of carbon dioxide gas produced, but another method involves weighing the reaction mixture:
Add dilute hydrochloric acid to a conical flask and stopper it with cotton wool
Place flask on a balance
Tare (counterweight) the balance, remove the cotton wool, drop the calcium carbonate into the acid, start a stop clock, and replace the cotton wool
The mass goes down as carbon dioxide escapes, so the balance reading must be recorded at regular intervals
You can repeat the experiment with different sized lumps of calcium carbonate or with calcium carbonate powder
Diagram:
. |*| — cotton wool
. / \ — conical flask
. /_____ \ — hydrochloric acid
. / ** \ —calcium carbonate
‾‾‾‾‾‾‾‾‾‾‾‾‾‾‾‾‾‾
| |-0.449| | — balance
______________
Catalysts
Substance which increases rate of reaction but is unchanged at the end of the reaction
Catalysts provide an alternate reaction pathway with a decreased activation energy
If you add 1g of catalyst to a reaction mixture, there will still be 1g left when the reaction has finished
Catalysts are specific to particular reactions: a substance that acts as a catalyst for one reaction may not catalyse a different reaction
Small amount of catalyst can catalyse reaction between large amounts of reactants
Helpful because catalytic converters use platinum, rhodium, and palladium
These metals very expensive so are coated onto an inert ceramic ‘honeycomb’
This uses only a few grams of catalyst, provides a large surface area for the reactions, and allows exhaust gases out
Investigating a catalyst
Example: Powdered manganese(IV) oxide, MnO,, catalyses the decomposition of hydrogen peroxide:
2H2O2 (aq) -> 2H2O(l) + O2(g)
Place hydrogen peroxide solution in a conical flask connected to a gas syringe
Put a little manganese(IV) oxide into the flask, stopper it, and start a stop clock
Only small amounts should be added, as gas can be produced very quickly and violently
Record the time and volume of oxygen at regular intervals
Draw a graph to analyse your results
You could vary the mass of catalyst or try other catalysts
When investigating a catalyst, you should keep volume, temperature, and concentration of hydrogen peroxide the same
This due to volume of oxygen produced is affected by these variables, so they must be kept the same to make it fair test
