This practical guide outlines the investigation of the effect of temperature on the rate of reaction between sodium thiosulfate and hydrochloric acid. The rate is monitored by measuring the time taken for a fixed amount of sulfur precipitate to form, rendering a marked black cross invisible.
Core Specification Link
This practical connects directly to topic 3.1.5 Kinetics, exploring collision theory, activation energy (\(E_a\)), and Maxwell-Boltzmann distributions.
Aim
To investigate the effect of temperature on the rate of reaction between sodium thiosulfate and hydrochloric acid, and to show that the reaction rate increases exponentially with temperature.
Equipment List
- Sodium thiosulfate solution (0.10 mol dm⁻³)
- Dilute hydrochloric acid (1.0 mol dm⁻³)
- Conical flask (250 cm³)
- Measuring cylinders (50 cm³ and 10 cm³)
- Thermometer (−10 to 110 °C)
- Stopwatch
- White card marked with a bold black cross
- Water bath or Bunsen burner, tripod, gauze, and heatproof mat
Experimental Method
- Measure 40 cm³ of 0.10 mol dm⁻³ sodium thiosulfate solution using a measuring cylinder and pour it into a clean 250 cm³ conical flask.
- Place the conical flask directly on top of the white card marked with a black cross.
- Record the initial temperature of the sodium thiosulfate solution.
- Using a smaller measuring cylinder, measure 5 cm³ of 1.0 mol dm⁻³ hydrochloric acid.
- Add the acid to the conical flask, start the stopwatch immediately, and swirl the flask once to mix.
- Observe the black cross by looking vertically downwards through the mouth of the flask.
- Stop the stopwatch at the exact moment the cross becomes completely obscured by the cloudy yellow-white sulfur precipitate . Record the time in seconds.
- Dispose of the toxic reaction mixture immediately in a sodium carbonate stop bath to neutralise dissolved sulfur dioxide gas . Wash the flask thoroughly.
- Repeat the experiment at four higher temperatures (approximately 30 °C, 40 °C, 50 °C, and 60 °C) . Gently heat the sodium thiosulfate solution in a water bath prior to adding acid, measuring the temperature immediately before acid addition.
- Calculate the initial rate for each temperature, which is proportional to \(1 / t\) (where \(t\) is the time taken in seconds).
Chemical Equation
\[ \text{Na}_2\text{S}_2\text{O}_3(\text{aq}) + 2\text{HCl}(\text{aq}) \rightarrow 2\text{NaCl}(\text{aq}) + \text{S}(\text{s}) + \text{SO}_2(\text{g}) + \text{H}_2\text{O}(\text{l}) \]The turbidity is caused by the slow precipitation of colloidal solid sulfur, S(s).
- Temperature (\(T\)): 55 °C
- Time (\(t\)): 18.5 s
Step 1: Convert temperature to Kelvin (K) and calculate \(1/T\)
\[ T = 55 + 273.15 = 328.15\text{ K} \] \[ \frac{1}{T} = \frac{1}{328.15} = 3.05 \times 10^{-3}\text{ K}^{-1} \]Step 2: Calculate the relative rate (\(1/t\)) and its natural logarithm (\(\ln(1/t)\))
\[ \text{Rate} = \frac{1}{t} = \frac{1}{18.5} = 0.0541\text{ s}^{-1} \] \[ \ln(\text{Rate}) = \ln(0.0541) = -2.92 \]Step 3: Relate the Arrhenius equation to the gradient of the graph
The Arrhenius equation is: \( \ln(k) = -\frac{E_a}{RT} + \ln(A) \). Plotting \(\ln(\text{rate})\) against \(1/T\) yields a straight line with gradient: \[ \text{Gradient} = -\frac{E_a}{R} \]
Step 4: Calculate the gradient using the two data points
\[ \text{Gradient} = \frac{y_2 - y_1}{x_2 - x_1} = \frac{-2.12 - (-2.92)}{(3.03 \times 10^{-3}) - (3.05 \times 10^{-3})} = \frac{0.80}{-2.00 \times 10^{-5}} = -40,000\text{ K} \]Step 5: Calculate the activation energy (\(E_a\))
\[ E_a = -\text{Gradient} \times R = -(-40,000) \times 8.31 = 332,400\text{ J mol}^{-1} = 332\text{ kJ mol}^{-1} \]The activation energy for the reaction is 332 kJ mol⁻¹ (3 significant figures).
Safety & Risk Assessment
| Hazard | Risk | Precaution |
|---|---|---|
| Sulfur dioxide (SO₂) gas | Toxic gas that causes respiratory irritation, especially in asthmatics. | Carry out the reaction in a well-ventilated laboratory. Place the flask in a neutralising stop bath (sodium carbonate) immediately after recording the time. |
| Dilute hydrochloric acid | Skin and eye irritation. | Wear safety goggles and a lab coat. Wash any splashes from skin immediately. |
| Hot water baths | Thermal burns. | Handle hot conical flasks with care using tongs or heat-resistant gloves. |
Results & Data Analysis
Since the reaction is stopped when a fixed amount of sulfur precipitate has formed, the concentration of sulfur is constant at the endpoint. We can assume that the average rate of reaction is proportional to 1/time.
Theory: Why rate increases with temperature
According to collision theory, a reaction only occurs when reactant particles collide with sufficient energy (greater than or equal to the activation energy, \(E_a\)) in the correct orientation. Increasing the temperature increases the kinetic energy of the particles. Consequently:
- A much larger fraction of particles possess energy exceeding the activation energy.
- Particles move faster, increasing collision frequency.
The first factor is primary: the Maxwell-Boltzmann distribution shifts to the right, significantly increasing the area under the curve beyond the activation energy barrier. As a result, the frequency of successful collisions increases exponentially.
Sources of Error & Improvements
| Error Source | Classification | Consequence & Mitigation |
|---|---|---|
| Subjective endpoint judgement | Random | Different observers judge when the cross disappears at slightly different times, introducing variability. Mitigation: Use a colorimeter or light sensor connected to a datalogger to record a set level of light transmittance. |
| Temperature cooling during reaction | Systematic | Warmed thiosulfate cools down while reacting in the flask, causing a discrepancy between measured temperature and real reaction temperature. Mitigation: Perform the reaction inside a temperature-controlled water bath. |
| Human reaction time error | Random | A delay of about 0.2 seconds in starting or stopping the stopwatch. Mitigation: Minimise relative error by using concentrations that yield reaction times between 10 and 100 seconds. |
Exam questions, marked instantly
Write your answer and the AI examiner checks it against the mark scheme, showing exactly which marks you earned and what to add.
Exam-style question
3 marksA 10 °C rise in temperature roughly doubles the rate of many reactions, even though the collision frequency increases only slightly. Explain why.
Show Mark Scheme
- At a higher temperature, many more particles have energy greater than or equal to the activation energy (the Maxwell–Boltzmann distribution shifts to higher energies) [1]
- So a much larger proportion of collisions are successful [1]
- This effect is far larger than the small increase in collision frequency [1]
Exam-style question
2 marksIn the thiosulfate "disappearing cross" experiment, explain why 1/t can be used as a measure of the rate of reaction.
Show Mark Scheme
- The same amount of sulfur (product) must form to hide the cross in every trial [1]
- Rate = amount of product ÷ time, so the rate is proportional to 1/t [1]
CPAC Skills Assessed
- CPAC 1: Follows written procedures to heat reactants and run kinetics experiments.
- CPAC 3: Safely sets up heating equipment and manages toxic sulfur dioxide gas emissions.
- CPAC 4: Accurately measures and records temperature values and reaction times.
When drawing a Maxwell-Boltzmann curve for two temperatures, remember that the peak for the higher temperature (\(T_2\)) must be lower and shifted to the right compared to the peak for the lower temperature (\(T_1\)). The total area under both curves must remain equal, as the total number of particles is constant.