Core practicals 8, 9, 13, 14: enthalpy, Ka, rates and activation energy

Introduction to Core Practicals 8, 9, 13, and 14

Welcome to your study guide for some of the most important practicals in the Pearson Edexcel Chemistry A Level! These four practicals cover Energetics, Acid-Base Equilibria, and Kinetics. In Paper 3, you won't just be asked what you did; you'll be asked why you did it, how to improve the method, and how to process the data. Don't worry if these seem complex—we will break them down step-by-step.

Core Practical 8: To determine the enthalpy change of a reaction using Hess’s Law

Sometimes, we want to find the enthalpy change (\(\Delta H\)) of a reaction that is impossible to measure directly. For example, if a reaction is too slow or requires intense heating, we use Hess's Law. This law states that the total enthalpy change is the same regardless of the route taken.

The Setup

A common exam example is finding the enthalpy change for the thermal decomposition of potassium hydrogencarbonate (\(2KHCO_3 \rightarrow K_2CO_3 + CO_2 + H_2O\)). Since we can't easily measure the heat of "breaking down" a solid, we react both the reactant (\(KHCO_3\)) and the product (\(K_2CO_3\)) with excess hydrochloric acid (\(HCl\)).

Step-by-Step Procedure

1. Weigh a known mass of the solid (e.g., \(K_2CO_3\)).
2. Place a known volume of \(HCl\) (in excess) into a polystyrene cup.
3. Measure the initial temperature of the acid for a few minutes to ensure it is stable.
4. Add the solid, stir continuously, and record the maximum (or minimum) temperature reached.

Calculating the Result

First, calculate the energy change (\(Q\)) using:
\(Q = mc\Delta T\)
Where \(m\) is the mass of the solution (usually taken as the volume of the acid), \(c\) is the specific heat capacity (\(4.18 \, J \, g^{-1} \, K^{-1}\)), and \(\Delta T\) is the temperature change.

Then, convert this to \(\Delta H\) in \(kJ \, mol^{-1}\) by dividing by the number of moles and adding the correct sign (\(-\) for exothermic, \(+\) for endothermic).

Common Errors and Improvements

Heat loss is the biggest problem. To fix this, we use a polystyrene cup (an insulator) and a lid. You can also use a "cooling curve" graph to extrapolate the temperature back to the exact moment of mixing to find a more accurate \(\Delta T\).

Quick Review: Hess's Law allows us to calculate \(\Delta H\) for "impossible" reactions by using an indirect route through a common intermediate (like a salt solution).

Core Practical 9: Finding the \(K_a\) value for a weak acid

The acid dissociation constant (\(K_a\)) tells us how much a weak acid dissociates (splits into ions) in water. In this practical, we use a pH meter and a titration curve to find it.

The "Half-Neutralisation" Trick

The easiest way to find \(K_a\) is to look at the half-neutralisation point.
The expression for \(K_a\) is: \(K_a = \frac{[H^+][A^-]}{[HA]}\)
When we have neutralised exactly half of the acid, the concentration of the remaining acid \([HA]\) equals the concentration of the salt produced \([A^-]\).
These two terms cancel out, leaving: \(K_a = [H^+]\).
This also means that \(pK_a = pH\) at the half-neutralisation point!

Step-by-Step Procedure

1. Pipette a fixed volume of a weak acid (e.g., ethanoic acid) into a beaker.
2. Calibrate a pH meter using buffer solutions (this is vital for accuracy!).
3. Add a strong base (e.g., \(NaOH\)) from a burette in small increments.
4. Record the pH after each addition and plot a pH curve (pH vs volume of base added).

Data Handling

1. From your graph, find the equivalence point (the vertical section).
2. Note the volume of base added at this point (e.g., \(20 \, cm^3\)).
3. Divide this volume by 2 (e.g., \(10 \, cm^3\)).
4. Look up the pH on your graph at that half-volume. This pH is your \(pK_a\).
5. Calculate \(K_a\) using: \(K_a = 10^{-pH}\).

Did you know? We calibrate pH meters because they can "drift" over time. Using buffers of known pH (usually pH 4, 7, and 10) ensures the meter's scale is reading correctly.

Core Practical 13: Rates of Reaction

This practical is split into two parts: following a reaction over time (continuous monitoring) and measuring the initial rate (the "clock" reaction).

13a: Following the rate of the iodine-propanone reaction

The reaction is: \(CH_3COCH_3 + I_2 \xrightarrow{H^+} CH_3COCH_2I + H^+ + I^-\)
We follow this by taking samples (aliquots) at regular intervals.

1. Mix propanone, iodine, and sulfuric acid catalyst.
2. Every few minutes, remove a sample and add sodium hydrogencarbonate. This "quenches" the reaction by neutralising the acid catalyst, effectively stopping the reaction in that sample.
3. Titrate the remaining iodine with sodium thiosulfate (\(Na_2S_2O_3\)) using starch as an indicator.
4. Plot a graph of iodine concentration against time. A straight line indicates the reaction is zero order with respect to iodine.

13b: Investigating a clock reaction (The Iodine Clock)

In a "clock" reaction, we measure the time (\(t\)) it takes for a distinct color change to occur. We assume the initial rate is proportional to \(1/t\).

1. Mix hydrogen peroxide, iodide ions, and a small, fixed amount of thiosulfate and starch.
2. The peroxide produces iodine, which is immediately used up by the thiosulfate.
3. Once the thiosulfate is all gone, the iodine starts to build up and turns the starch blue-black.
4. By varying the concentration of one reactant and keeping the others constant, you can determine the order of reaction.

Core Practical 14: Finding the activation energy of a reaction

The activation energy (\(E_a\)) is the minimum energy required for a collision to result in a reaction. We find this by seeing how the rate constant (\(k\)) changes with temperature.

The Arrhenius Equation

You will be given this in the exam: \(k = Ae^{-E_a/RT}\).
To make it useful for a graph, we use the logarithmic form:
\(\ln k = -\frac{E_a}{R} \cdot \frac{1}{T} + \ln A\)
This matches the straight-line equation \(y = mx + c\).

Step-by-Step Procedure

1. Perform a clock reaction (like the one in CP13b) at several different temperatures (e.g., \(20^\circ C, 30^\circ C, 40^\circ C, 50^\circ C\)).
2. Calculate the rate for each (Rate \(\approx 1/t\)).
3. Since the rate is proportional to \(k\), we can plot \(\ln(\text{rate})\) on the y-axis against \(1/T\) on the x-axis (where \(T\) is in Kelvin).

Calculating \(E_a\)

1. Calculate the gradient of your line of best fit (\(\frac{\Delta y}{\Delta x}\)). The gradient will be negative.
2. Set the gradient equal to \(-E_a/R\).
3. \(E_a = -\text{gradient} \times R\) (where \(R = 8.31 \, J \, mol^{-1} \, K^{-1}\)).
4. Your answer will be in \(J \, mol^{-1}\). Usually, you should divide by 1000 to give the final answer in \(kJ \, mol^{-1}\).

Common Mistake: Forgetting to convert Celsius to Kelvin! Always add 273 to your Celsius temperatures before doing the calculation.

Summary of Key Skills for Paper 3

• Safety: Always mention wearing eye protection. Many of these chemicals (like \(HCl\) or propanone) are irritants or flammable.
• Accuracy: Use a volumetric pipette instead of a measuring cylinder for small volumes. Use a white tile to see color changes clearly in titrations.
• Calculations: Be comfortable with \(Q = mc\Delta T\), the Arrhenius graph, and \(pH = -\log[H^+]\).
• Variables: In kinetics, always state that you keep concentrations constant when changing temperature, and vice versa.

Don't worry if the math in the Arrhenius equation feels heavy! Just remember: plot \(\ln(\text{rate})\) against \(1/T\), find the gradient, and multiply by \(-8.31\). You've got this!