Chemistry » Energy and Chemical Reactions » Rates Of Reaction And Factors Affecting Rate

# Reaction Rate Factors Continued

## Optional Experiment: Temperature, concentration and reaction rate

### Aim

To determine the effect of temperature and concentration on the average reaction rate of the iodine clock experiment. This experiment is best done in groups.

### Apparatus

• Potassium iodide ($$\text{KI}$$), soluble starch, sodium thiosulfate solution ($$\text{Na}_{2}\text{S}_{2}\text{O}_{3}$$), dilute (around $$\text{0.2}$$ $$\text{mol.dm^{-3}}$$) sulfuric acid ($$\text{H}_{2}\text{SO}_{4}$$), 3% hydrogen peroxide ($$\text{H}_{2}\text{O}_{2}$$) solution

• Five beakers, a measuring cylinder, a hotplate, an ice bath, a glass stirring rod, a stop-watch

### Method

• Preheat the hotplate to $$\text{40}$$ $$\text{℃}$$

• Label a beaker solution 1. Measure $$\text{75}$$ $$\text{ml}$$ $$\text{H}_{2}\text{SO}_{4}$$ into the beaker. Add $$\text{25}$$ $$\text{ml}$$ 3% $$\text{H}_{2}\text{O}_{2}$$. Remember to use dilute ($$\text{0.2}$$ $$\text{mol.dm^{-3}}$$) sulfuric acid.

The equations for what is occuring in this reaction are given below:

$$\text{H}_{2}\text{O}_{2}(\text{l}) + 2\text{KI}(\text{s}) + \text{H}_{2}\text{SO}_{4}(\text{l})$$ $$\to$$ $$\color{red}{\text{I}_{2}{\text{(s)}}}$$ $$+ \text{K}_{2}\text{SO}_{4}(\text{aq}) + 2\text{H}_{2}\text{O}(\text{l})$$

$$\color{red}{\text{I}_{2}{\text{(s)}}}$$ $$+ 2\text{Na}_{2}\text{S}_{2}\text{O}_{3}(\text{aq})$$ $$\to$$ $$\text{Na}_{2}\text{S}_{4}\text{O}_{6}(\text{aq}) + 2\text{NaI}(\text{aq})$$

It is good scientific practice to vary only one factor at a time during an experiment. Therefore, this experiment has two parts. First we will vary the concentration of $$\text{KI}$$, then we will vary the temperature:

• Varying the concentration

1. Weigh out $$\text{0.5}$$ $$\text{g}$$ of $$\text{KI}$$ into a beaker and label it A.

2. Weigh out $$\text{1}$$ $$\text{g}$$ of $$\text{KI}$$ into a different beaker and label it B.

3. Add $$\text{20}$$ $$\text{ml}$$ $$\text{Na}_{2}\text{S}_{2}\text{O}_{3}$$ to both beaker A and beaker B.

4. Add a spatula of soluble starch to both beaker A and beaker B and stir with a glass rod.

5. Measure $$\text{15}$$ $$\text{ml}$$ of solution 1 with the measuring cylinder. Get your stopwatch ready. Pour the $$\text{15}$$ $$\text{ml}$$ of solution 1 into beaker A and start timing.

Stop timing when the solution starts to change colour. Write down your time in the table below.

6. Repeat step $$\text{5}$$ with beaker B.

 Beaker Concentration (M) Temperature (℃) Time (s) A approx. 0.15 room temperature B approx. 0.3 room temperature
• Varying the temperature

1. Weigh out $$\text{0.5}$$ $$\text{g}$$ of $$\text{KI}$$ into a new beaker and label it C.

2. Add $$\text{20}$$ $$\text{ml}$$ $$\text{Na}_{2}\text{S}_{2}\text{O}_{3}$$ to beaker C.

3. Add a spatula of soluble starch to beaker C and stir with a glass rod.

4. Measure $$\text{15}$$ $$\text{ml}$$ of solution 1 with the measuring cylinder.

5. Place beaker C in the ice bath.

6. Get your stopwatch ready. Pour the $$\text{15}$$ $$\text{ml}$$ of solution 1 into beaker C and start timing. Stop timing when the solution starts to change colour. Write down your time in the table below.

7. Repeat steps 1 – 4 (label the beaker D).

Place beaker D on the hotplate. Then repeat step 6

 Beaker Concentration (M) Temperature (℃) Time (s) A approx. 0.15 room temperature C approx. 0.15 0 D approx. 0.15 40

Beaker A has been included here because it has the same concentration as beakers C and D, but is at a different temperature.

### Results

Make a table with the information for all the beakers. Include columns for concentration, temperature, time, and reaction rate.

### Questions and discussion

• Did beaker A or B have the faster reaction rate?

• Why did it have a faster reaction rate?

• Did beaker A, C or D have the fastest reaction rate? Why?

• Did beaker A, C or D have the slowest reaction rate? Why?

### Conclusions

You will notice that the faster reaction rate occurs in the beaker with the higher concentration of $$\text{KI}$$. You should also see that the higher the temperature, the faster the reaction rate.

### Optional Video: Iodine Clock Reaction

This video shows how this experiment can be used as a clock with the concentration chosen so that the experiment changes colour at a specific time (or with a particular part of a song). This is why this experiment is known as the iodine clock reaction.

## Example: Reaction Rates

### Question

Write a balanced equation for the exothermic reaction between $$\text{Zn}(\text{s})$$ and $$\text{HCl}(\text{l})$$. Also name three ways to increase the rate of this reaction.

### Step 1: Write the equation for zinc and hydrochloric acid

The products must be a salt and hydrogen gas. Zinc ions have a charge of 2+ while chloride ions have a charge of 1-. Therefore the salt must be $$\text{ZnCl}_{2}$$.

$$\text{Zn}(\text{s}) + \text{HCl}(\text{aq})$$ $$\to$$ $$\text{ZnCl}_{2}(\text{aq}) + \text{H}_{2}(\text{g})$$

### Step 2: Balance the equation if necessary

There are more chloride ions and hydrogen atoms on the right side of the equation. Therefore there must be $$\text{2}$$ $$\text{HCl}$$ on the left side of the equation.

$$\text{Zn}(\text{s}) + 2\text{HCl}(\text{aq})$$ $$\to$$ $$\text{ZnCl}_{2}(\text{aq}) + \text{H}_{2}(\text{g})$$

### Step 3: Think about the methods mentioned in this section that would increase reaction rate

• A catalyst could be added

• The zinc solid could be ground into a fine powder to increase its surface area

• The $$\text{HCl}$$ concentration could be increased