Practical: Investigating Temperature Changes

• The law of Conservation of Energy dictates that energy cannot be created or destroyed, only transferred. Therefore, any heat released by an exothermic reaction is absorbed by the surrounding solution, causing a measurable temperature rise.

• The quantity of heat energy ($Q$) can be calculated using the fundamental thermochemical equation:

Heat Equation: $Q = m \times c \times \Delta T$

• In this equation, $m$ represents the total mass of the solution (calculated from total volume and density), $c$ is the specific heat capacity, and $\Delta T$ is the change in temperature (final $-$ initial).

• To simplify calculations, chemists assume the density of the solution is $1.0$ g/cm$^3$ and that the heat capacity is identical to that of pure water. These assumptions allow volume measurements to be converted directly into mass for use in the heat equation.

• Preparation: Measure a fixed volume of the first reactant (e.g., an alkali) into the calorimeter and record its initial stable temperature. This establishes the baseline for energy calculations.

• Execution: Add the second reactant (e.g., an acid) in controlled increments or as a single excess amount, depending on the goal. Continuous stirring is essential to ensure that the heat generated by the reaction is evenly distributed throughout the solution.

• Recording: Observe the thermometer closely to identify the maximum temperature reached (for exothermic reactions) or the minimum temperature (for endothermic reactions). The difference between this value and the initial baseline is used to calculate $Q$.

• Graphical Analysis: For titration-style thermochemical investigations, plot the temperature reached against the volume of reactant added. The point where the temperature peak occurs indicates the stoichiometric neutralization point.

• The choice of container is critical: Polystyrene is used for solutions because its low thermal conductivity prevents heat from leaking out. In contrast, Copper is used for combustion to efficiently transfer heat from the flame to the water inside.

• Total Mass Calculation: A common mistake is using only the mass of one reactant. Always add the volumes of all liquids combined in the calorimeter to find the total mass ($m$) for the formula $Q=mc\Delta T$.

• Unit Precision: The energy $Q$ is calculated in Joules (J). Most exam questions require the answer in kilojoules (kJ), so remember to divide the result by $1000$ before finalizing molar enthalpy calculations.

• Sign Convention: For exothermic reactions (where temperature increases), the enthalpy change ($\Delta H$) must be recorded as a negative value. Conversely, endothermic reactions have a positive $\Delta H$.

• Assumption Verification: Be prepared to list the assumptions made during the experiment, such as assuming the density is exactly $1$ g/cm$^3$ and ignoring the heat absorbed by the calorimeter itself.

• Heat Loss to Surroundings: Despite insulation, some heat inevitably escapes. This leads to an underestimate of the temperature change, resulting in a calculated enthalpy value that is lower than the true theoretical value.

• Incomplete Reaction: If the reactants are not properly stirred or if the concentration is lower than expected, the reaction may not reach completion. This suppresses the maximum temperature reached and distorts the data.

• Thermal Inertia: Students often forget that the thermometer and the cup itself absorb some heat energy. In precise calorimetry, the heat capacity of the calorimeter would be included, but it is often ignored in introductory practicals.