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