Welcome to Thermochemical and Kinetics Experiments!
In this chapter, we explore the "how-to" of measuring energy changes and reaction speeds in the lab. These skills are essential for Unit 3 (Practical Skills in Chemistry I), where you will be tested on your ability to plan, carry out, and evaluate experiments based on the theory you learned in Units 1 and 2. Don't worry if the calculations look scary at first—we will break them down step-by-step!
1. Thermochemical Experiments: Measuring Heat
Thermochemistry is all about measuring the heat energy released or absorbed during a chemical reaction. In the lab, we call this the enthalpy change, represented by the symbol \( \Delta H \).
The Basic Setup: The Calorimeter
To measure heat change, we use a calorimeter. For most AS Level experiments, this is simply a polystyrene cup.
Why polystyrene? It is an excellent thermal insulator, which means it reduces heat loss to the surroundings, making our measurements more accurate.
The Key Formula
To calculate the heat energy transferred (\( Q \)), we use:
\( Q = mc\Delta T \)
Where:
\( Q \) = energy transferred in Joules (\( J \))
\( m \) = mass of the substance being heated (usually the solution in the cup, assuming \( 1 \text{ cm}^3 = 1 \text{ g} \))
\( c \) = specific heat capacity (for water/dilute solutions, this is \( 4.18 \text{ J g}^{-1} \text{ }^\circ\text{C}^{-1} \))
\( \Delta T \) = change in temperature (\( T_{final} - T_{initial} \))
Converting \( Q \) to Enthalpy Change (\( \Delta H \))
Once you have \( Q \), you need to find the molar enthalpy change in \( \text{kJ mol}^{-1} \).
1. Convert \( Q \) from Joules to kiloJoules: \( Q / 1000 \).
2. Calculate the number of moles (\( n \)) of the limiting reactant.
3. Use the formula: \( \Delta H = -Q / n \).
Quick Tip: If the temperature goes up (exothermic), \( \Delta H \) must be negative. If the temperature goes down (endothermic), \( \Delta H \) must be positive.
Key Takeaway:
Always use the mass of the liquid in the cup for \( m \), not the mass of the solid added!
2. Core Practical 2: Enthalpy Change and Hess’s Law
Sometimes we can't measure an enthalpy change directly (e.g., the reaction is too slow or dangerous). Instead, we use Hess's Law, which states that the total enthalpy change is the same regardless of the route taken.
Example Experiment: Finding the enthalpy change for the decomposition of potassium hydrogencarbonate.
Since we can't easily heat the solid and measure the energy change, we react both the reactant and the product with hydrochloric acid and use an enthalpy cycle to calculate the "missing link."
Reducing Errors in Thermochemistry
In Unit 3, you are often asked how to improve these experiments. Common methods include:
1. Cooling Curve Corrections: If a reaction is slow, heat is lost while it's happening. We plot a graph of temperature against time, then extrapolate the cooling curve back to the time of mixing to find a more accurate "theoretical" maximum temperature.
2. Lids and Insulation: Use a lid on the polystyrene cup to prevent heat loss by evaporation and convection.
3. Measurement Uncertainty: Use a thermometer with smaller graduations (e.g., \( 0.1 \text{ }^\circ\text{C} \) instead of \( 1 \text{ }^\circ\text{C} \)) to reduce the percentage error.
3. Kinetics Experiments: Measuring Rate
Kinetics is the study of reaction rates—how fast a reaction happens. In Unit 3, you need to know how to measure this rate experimentally.
The General Method
Rate is defined as the change in concentration of a reactant or product per unit time. However, in simple lab experiments, we often use:
\( \text{Rate} = 1 / \text{time} \)
This is used when we measure how long it takes for a visible change to occur (like a color change or a precipitate forming).
Core Practical 5: Hydrolysis of Halogenoalkanes
In this experiment, we compare how quickly different halogenoalkanes react with water.
The Setup:
1. Mix the halogenoalkane with ethanol (acting as a solvent so the reactants mix) and aqueous silver nitrate.
2. Start a stopwatch and measure how long it takes for a silver halide precipitate to form.
3. The faster the precipitate forms (shorter time), the higher the rate.
Did you know? This experiment proves that iodoalkanes react faster than chloroalkanes because the \( \text{C–I} \) bond is weaker (lower bond enthalpy) than the \( \text{C–Cl} \) bond.
4. Factors Affecting Rate (Experimental Context)
When performing these experiments, you must control variables to make the test valid:
1. Temperature: Use a water bath to keep the temperature constant.
2. Concentration: Ensure the volumes and concentrations of reactants are measured precisely using a pipette or burette.
3. Surface Area: If using solids, ensure the particle size is consistent.
Quick Review:
To find the rate at a specific point on a concentration-time graph, draw a tangent to the curve and calculate its gradient (\( \text{rise} / \text{run} \)).
5. Safety and Risk Assessment
In Unit 3, you must distinguish between a hazard (the potential to cause harm) and a risk (the likelihood of that harm occurring).
Common Precautions:
1. Corrosive acids/alkalis: Wear safety goggles and gloves.
2. Flammable liquids (like ethanol): Use a water bath for heating instead of a Bunsen burner (no naked flames).
3. Toxic gases: Perform the experiment in a fume cupboard.
Summary Checklist for the Exam
- Can you calculate \( Q \) and \( \Delta H \) from experimental data?
- Do you know why we use polystyrene cups and cooling curves?
- Can you explain how to measure the rate of hydrolysis in halogenoalkanes?
- Are you comfortable identifying hazards and suggesting safety precautions?
Remember: Unit 3 is about being a scientist in the lab. Think about what could go wrong (errors) and how to make your results more reliable!