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How do you find the experimental molar enthalpy?

Published in Experimental Enthalpy 4 mins read

The experimental molar enthalpy change can be found by measuring the heat transferred during a reaction and relating it to the moles of the limiting reactant using a specific formula.

Understanding Molar Enthalpy Change

Molar enthalpy change, denoted as ΔH, represents the heat absorbed or released when one mole of a substance reacts. This value is crucial in understanding the energetics of chemical reactions. Experimentally, we calculate it using calorimetric data and a specific equation.

The Key Equation

The core of finding experimental molar enthalpy is the following equation, derived from the reference:

  • ΔH = (m c ΔT) / n
Where:
 *   **ΔH** is the molar enthalpy change, typically in kJ/mol.
 *   **m** is the mass of the substance that absorbs or releases the heat. In calorimetry, this usually represents the mass of the solution or water in the calorimeter (in grams).
 *   **c** is the specific heat capacity of the substance (usually the solution) that absorbs or releases the heat (in J/g°C). For water, the specific heat capacity is approximately 4.184 J/g°C.
 *  **ΔT** is the temperature change observed during the reaction (in °C or K). It is calculated as the final temperature minus the initial temperature.
 *  **n** is the number of moles of the limiting reactant that reacted in the chemical reaction. This can be derived from the mass of the limiting reactant that was used.

Steps to Calculate Experimental Molar Enthalpy

Here are the steps to determine the experimental molar enthalpy, using the information from the provided reference and incorporating further detail:

  1. Perform the reaction in a calorimeter: Conduct the reaction in a calorimeter, a device designed to measure heat transfer.

  2. Record initial and final temperature:

    • Measure the initial temperature of the calorimeter's contents before the reaction.
    • Measure the final temperature after the reaction is complete.
    • Calculate the change in temperature using the following equation: ΔT = final temperature - initial temperature.

  3. Determine the mass and specific heat capacity: Determine the mass of the solution within the calorimeter (m), and the substance's specific heat capacity (c). For dilute aqueous solutions, the specific heat capacity of water can be assumed.

  4. Calculate the heat transferred (q): Calculate the heat released or absorbed by the calorimeter contents, using q = m c ΔT .

  5. Determine the number of moles (n) of the limiting reactant. In a reaction, if one of the reactants runs out before the other, it is known as the limiting reactant. We must know how many moles of the limiting reactant were used in the reaction.

  6. Calculate the molar enthalpy change (ΔH): Divide the heat transferred (q) by the number of moles (n) of the limiting reactant: ΔH = q / n. This will give you the molar enthalpy change, typically in J/mol, which may be converted to kJ/mol by dividing by 1000.

  • Remember to pay attention to the sign of the enthalpy change. A negative ΔH indicates an exothermic reaction (heat is released), and a positive ΔH indicates an endothermic reaction (heat is absorbed).

Practical Insights

  • Calorimeter Calibration: For accurate results, the calorimeter should be calibrated to determine its heat capacity, which may be required for specific experiments.

  • Assumptions: Certain assumptions are often made, such as neglecting heat loss to the surroundings and assuming that the solution has the properties of pure water. These assumptions can introduce minor errors.

  • Limiting Reactant: The number of moles of the limiting reactant, n, is crucial for an accurate result. Always verify that you are using the number of moles of the limiting reactant.

Example:

Suppose you perform a reaction where:

  • Mass of solution (m) = 100 g of water
  • Specific heat capacity of water (c) = 4.184 J/g°C
  • Temperature change (ΔT) = +5.0 °C (an increase, indicating the reaction is exothermic)
  • Moles of limiting reactant (n) = 0.02 moles

  1. Calculate q: q = (100 g) (4.184 J/g°C) (5.0 °C) = 2092 J

  2. Calculate ΔH: ΔH = 2092 J / 0.02 mol = 104600 J/mol or 104.6 kJ/mol. Since the temperature increased, it is an exothermic process, therefore ΔH = - 104.6 kJ/mol.

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