Practical: Investigating Temperature Changes

The Law of Conservation of Energy dictates that energy cannot be created or destroyed, only transferred. In a calorimeter, we assume the heat change of the reaction ($q_{rxn}$) is equal and opposite to the heat change of the solution ($q_{sol}$), provided insulation is perfect.

The heat energy change ($Q$) is calculated using the formula: $Q = m \times c \times \Delta T$ where $m$ is the mass of the solution (g), $c$ is the specific heat capacity ($J/g/^{\circ}C$), and $\Delta T$ is the change in temperature ($^{\circ}C$).

Specific Heat Capacity ($c$) is a physical property representing the energy required to raise the temperature of one gram of a substance by one degree Celsius. For aqueous solutions, this is typically assumed to be $4.18 J/g/^{\circ}C$, the value for pure water.

Molar Enthalpy Change ($\Delta H$) scales the heat energy change to a per-mole basis, allowing for the comparison of different reactions: $\Delta H = \frac{-Q}{n}$ where $n$ is the number of moles of the limiting reactant.

Step 1: Initial Measurement: Use a measuring cylinder to add a fixed volume of the first reactant (e.g., an alkali) into the calorimeter and record its stable initial temperature.

Step 2: Controlled Addition: Add the second reactant (e.g., an acid) in measured increments or as a single bulk addition, depending on the specific investigation goals.

Step 3: Agitation: Stir the mixture continuously using the thermometer or a stirring rod to ensure the heat is distributed evenly throughout the solution for an accurate reading.

Step 4: Data Recording: Monitor the thermometer closely and record the maximum (for exothermic) or minimum (for endothermic) temperature reached before the mixture begins returning to room temperature.

Step 5: Analysis: Calculate the temperature difference ($\Delta T = T_{final} - T_{initial}$) and use the mass of the total liquid volume to calculate the energy change.

Unit Conversion: Always check if the question asks for the answer in Joules ($J$) or Kilojoules ($kJ$). Remember to divide by $1000$ to convert $J$ to $kJ$.

Density Assumption: In most exam problems, assume the density of an aqueous solution is $1.0 g/cm^3$. Therefore, a volume of $50 cm^3$ of solution is treated as a mass of $50 g$.

Sign Convention: For enthalpy change ($\Delta H$), a negative sign ($-$) MUST be used for exothermic reactions (temperature increase), and a positive sign ($+$) for endothermic reactions (temperature decrease).

Significant Figures: Ensure your final answer matches the precision of the data provided in the question, usually 2 or 3 significant figures.

Ignoring the Lid: Students often forget that a lid is essential to prevent heat loss via convection to the air, which would lead to an underestimation of the temperature change.

Mass Confusion: A common error is using the mass of the solid reactant instead of the mass of the solution being heated in the $Q = mc\Delta T$ formula.

Incomplete Stirring: Failing to stir the mixture results in 'hot spots' or 'cold spots,' leading to inaccurate thermometer readings that do not represent the bulk solution.

Specific Heat Capacity: Students sometimes use the specific heat capacity of the calorimeter material instead of the solution (water).