Activation Energy

- For a reaction to occur, the particles must collide and possess enough energy to allow it to proceed.

- Activation Energy is the minimum amount of energy particles need in order to react with each other. It acts as a barrier to the reaction.
- Some reactions require high amounts of energy, and some require less.

General Rule:

- The bigger the activation energy required, the slower the reaction will be at a certain temperature. E.g., the combustion of ethanol would be slower on a cold day than when they collide explosively in a fast reaction.
- If particles possess energy equal to or greater than the activation energy, then when they collide, they can react because they will have sufficient energy to break the bonds within the reactants; a successful collision can occur.
- If particles possess less energy than the activation energy, then when they collide, they don't react, and they will just bounce off each other: an unsuccessful collision will occur.

- Activation energy can be represented on a reaction profile diagram:

- The higher the activation energy, the slower the reaction is.

Faster reactions will require less energy to initiate, resulting in a lower activation energy and a smaller graph.

Rate of Reaction

- The rate of a reaction is a measure of how quickly a reactant is being used up or a product is formed. A graph can illustrate this.

There are numerous ways to track the rate of product formation:

1) When a gas is produced, we can measure the loss of mass.
2) Measuring the volume of gas produced.
→ Gas can be collected over water and be measured using an upturned cylinder.
→ A gas syringe is a tool for measuring gases.
3) When a precipitate is produced, we can measure the decreasing light passing through.
→ Eg. Add dilute acid to a beaker (clear) placed over a cross and start timing. Then time: how long does it take for the cross to disappear?
→ You could also use a light sensor to check when the cross has disappeared—this would be more accurate. However, light sensors are expensive.
4) For reactions that use up or produce an acid/alkali, the change in pH may be followed.
5) If the reaction is fairly exothermic or endothermic, you can follow the temperature change.
6) For some reactions, the change in colour of a solution may be monitored using a colourimeter.

It is always essential to time the progress of the reaction in conjunction with one of these methods, as the rate represents a change per time unit.

Data Interpretation

Graph to show the volume of gas produced over time when hydrochloric acid reacted with CaCO₃.

A: The process is fastest at the beginning because there are many reactant particles that can collide with each other.

B: Slows down because there are fewer reactant particles, as some have formed products. This means that there are fewer collisions in a certain time.

C: Stops when at least one reactant is used completely, preventing product formation and reaction-causing collisions.

Calculating Rate

Units:

- Mass = g/s
- Volume = cm³/s
- Amount = mol/s

Calculating Rate at a Specific Time:

- You can use the rate of reaction graph to find the rate of reaction at a specific time using a tangent.
- Draw a tangent to the curve at that time; a straight line that just touches the curve at that point.
- Then construct a right-angle triangle using the tangent as the hypotenuse.
- Then work out the gradient—rise/run—and you have your rate.

Example:

Surface Area

- Marble chips: Calcium carbonate ranges from large to medium to small in size. They can react with hydrochloric acid.

Equation:

- CaCO_{3 (s)} + 2HCl_(aq) → CaCl_{2 (aq)} + CO_{2 (g)} + H₂O_(l)

Observations (small chips and large chips had the same mass):

- Small chips: More effervescence (disappeared quicker)
- Large chips: Slower/some effervescence (disappeared slower)

- The rate was quicker when using a larger surface area (small marble chips). The smaller surface area (large chips) has a slower rate of reaction.
- The larger the pieces of CaCO₃ the smaller the surface area available for reactions to take place on. Therefore, there are less surfaces available at a certain time, so there are fewer collisions that can take place - slowing the rate of reaction.
- This is the opposite for the larger surface area - smaller chips.
- The collision frequency is higher for a larger surface area.

- The surface area doesn't affect the total amount of CO₂ gas formed, but the smaller the particle size, the steeper the initial gradient of the curve.

Temperature

- Reaction between sodium thiosulfate and hydrochloric acid.
2HCl_(aq) + Na₂S₂O_{3 (aq)} → 2NaCl_(aq) + SO_{2 (aq)} + S_(s) + H₂O_(l)
- When this reacts, the clear solution turns cloudy as a precipitate → Sulfur is formed.
- Therefore, the rate of reaction can be followed by timing how long it takes for the solution to become cloudy/to see an 'X' on a paper beneath the reaction.

Observations:

- The 'X' disappears quicker at a higher temperature. This is because the reaction rate is faster, so the solution turns cloudy quickly.
- As temperature increases, the energies of reactant particles increase, and so they move around faster and collide more often; the frequency of collisions increases.
- The collisions will be more energetic, and so the proportion of particles with energy greater than the activation energy (or equal) will increase, and the proportion of successful collisions will increase.
- Increasing the temperature by 20 degrees usually has a dramatic effect on the rate of a reaction. An increase as small as usual has a dramatic effect on the rate of a reaction. An increase as small as 10ºC will cause the rate of many reactions to double. The rate of many reactions varies rapidly and exponentially with temperature - They are not directly proportional.

- The total amount of CO₂ gas formed, but the higher the temperature, the steeper the initial gradient of the curve.

Example:

- Copper oxide reacts much faster with acid at 40ºC than at 20ºC as particles gain kinetic energy and move faster, leading to more frequent collisions. There are also more energetic collisions as more particles have E ≥ Ea. This means there's a higher proportion of successful collisions.

Concentration

- Concentration is a measure of the number of particles per unit volume.
- In the case of gases, an increase in pressure compresses the gas into a smaller volume, resulting in a decrease in particle proximity.
- As the concentration increases, the number of reactant particles per unit volume and its proximity increase.
- The particles will have the same amount of energy, but they are closer, and so there will be more frequent collisions; the frequency of collisions increases.
- Since the energy of the particles is the same, the proportion of successful collisions will be the same. However, the increased frequency of collisions means the rate of reaction will increase.
- Simple version: Increasing the concentration increases the rate because there are more particles in a certain area. This means the particles are closer together, and so there are more frequent collisions, leading to an increased rate of reaction.

Graphs:

- If you increase the concentration of the reactant in excess, the rate increases, but the amount of product doesn't. If you increase the concentration of the limiting reactant, the rate increases, and more product is made.

Graphs

Graph 1:

a) 50cm³ of 0.1 moldm^-3 acid with excess marble chips.
b) 50cm³ of 0.2 moldm^-3 acid with excess marble chips.

Graph 2:

- Doubling the concentration of a particular reactant will double the number of particles per unit volume. So, this will generally result in a doubling of the reaction rate. This means the reaction rate is often directly proportional to the concentration of a reactant.

Graph 3:

- Shows the reaction of excess acid and marble chips for two different acid concentrations. The concentration of acid is double the rate of response in excess, so the mass of CO₂ is produced as acid is present in excess. A less steep gradient shows the rate of reaction with more dilute acid.

Graph 4:

- Shows the reaction of 2 different volumes of acid with excess marble chips. Doubling the volume as acid doubles the mass of CO₂ is produced initially. (When the concentration of acid is identical in both experiments. The gradient indicates a similar rate of reaction at each concentration.

Catalysts

How Do Catalysts Work?

- A catalyst alters the route by which the chemical process takes place. The alternative, catalysed route of reaction will have a lower Ea than the uncatalysed process.
- In the presence of a catalyst, the reactant particles will have the same amount of energy. However, by decreasing the Ea, the proportioning of particles with E ≥ Ea will increase, so the proportion of successful collisions increases.
- A catalyst is a substance that increases the reaction rate but is not used up. It provides an alternative pathway with lower Ea. The proportion of particles with E ≥ Ea increases.
- When the reaction stops, the catalyst is still present. Catalysts contain transition metals.

Decomposition of H₂O₂ (hydrogen peroxide)

- Hydrogen peroxide decomposes slowly at room temperature into water and oxygen.

Order of catalysts - rate of reaction:

Fastest

1) Liver
2) Manganese (IV) oxide
3) Peeled potatoes
4) Magnesium oxide - very slow
5) Sodium chloride - very slow (slowest)