Exothermic & Endothermic

• Enthalpy ($H$): This is the total heat content of a system. While we cannot measure absolute enthalpy, we can measure the change in enthalpy ($\Delta H$) during a chemical reaction.

• Enthalpy Change ($\Delta H$): Calculated as $H_{products} - H_{reactants}$. In exothermic reactions, $\Delta H$ is negative because products have less heat content than reactants; in endothermic reactions, it is positive because the system has gained heat energy.

• Thermodynamic Stability: Substances with lower enthalpy are generally more stable. Therefore, products of exothermic reactions are more energetically stable than their reactants under standard conditions.

• Calorimetry Setup: To measure heat changes, reactions are performed in insulated containers like polystyrene cups to minimize heat exchange with the environment. Stirring ensures an even distribution of heat during the process.

• The Energy Equation: The heat energy change ($Q$) is calculated using the formula $Q = m \times c \times \Delta T$, where $m$ is the mass of the substance being heated, $c$ is the specific heat capacity (typically $4.18 \, J/g/^\circ C$ for water), and $\Delta T$ is the temperature change.

• Molar Enthalpy Calculation: To find the energy per mole ($\Delta H$), divide the calculated heat energy $Q$ (in kJ) by the number of moles of the limiting reactant. The sign is then adjusted based on whether the temperature rose (negative) or fell (positive).

Comparison of Reaction Types

• Reaction Direction: In an exothermic reaction, the system loses energy to the environment, whereas in an endothermic reaction, the system acts as an energy sink, drawing heat inward.

• Check the Sign: Always remember that a temperature increase corresponds to an exothermic reaction, which requires a negative $\Delta H$ sign in your final answer. Forgetting this sign is a common way to lose marks.

• Mass Selection: In $Q = mc\Delta T$ calculations for solutions, the mass ($m$) is the mass of the liquid being heated (usually $1 \, g = 1 \, cm^3$), not the mass of the solid reactant added or the mass of the fuel burned.

• Unit Conversion: Calculations of $Q$ usually result in Joules ($J$). Exam questions typically ask for Enthalpy in $kJ/mol$, so you must divide by $1000$ to convert to $kJ$ before dividing by moles.

• Standard Assumptions: In calorimetry, we assume the density and specific heat capacity of the solution are the same as water ($1.0 \, g/cm^3$ and $4.18 \, J/g/^\circ C$) unless stated otherwise.

• Temperature vs. Energy: Students often confuse the temperature of the surroundings with the energy of the system. In an exothermic reaction, the temperature of the water/mixture goes up, but the energy of the chemical products has gone down.

• Bond Energy Confusion: A common mistake is using (Products - Reactants) for bond energy calculations. The correct approach is (Energy to Break Reactant Bonds - Energy Released Forming Product Bonds).

• Mass in Combustion: In spirit burner experiments, ensure you use the mass of the water in the calorimeter for 'm' in the energy equation, not the mass of the fuel that was burned.