Explain why the absolute values of the standard enthalpy change in the reaction and heat measured by the calorimeter are not equal. Assume that there are no measurement errors of any kind.
In the previous problem I solved that Delta H rxn was -1500 and q = 1.519 x 10^6 J and I said that it dealt with internal energy and got half credit. Can someone try to explain the theory to me?
The absolute values of the standard enthalpy change in the reaction and the heat measured by the calorimeter are not equal because they represent different quantities and measure slightly different aspects of the reaction.
The standard enthalpy change (ΔH) of a reaction is a thermodynamic property that determines the heat released or absorbed during a chemical reaction under standard conditions (usually at a temperature of 25°C and pressure of 1 bar). It is calculated based on the difference in enthalpy (heat content) between the products and the reactants. The standard enthalpy change is often given in units of kilojoules per mole (kJ/mol), and it represents the change in energy associated with the reaction itself.
On the other hand, the heat measured by the calorimeter (q) represents the heat released or absorbed by the reaction mixture (including the reactants and any solvent or solution present) as it is placed in a calorimeter and undergoes the reaction. The calorimeter is an apparatus used to measure the heat exchange of a system, and it typically consists of a vessel surrounded by a heat-insulating material.
While both ΔH and q represent energy changes, they are not expected to be exactly equal due to a couple of factors. First, the calorimeter itself may have some heat capacity and absorb or release a certain amount of heat during the measurement process, which adds to or subtracts from the heat associated with the reaction. Second, there can be heat losses to the surroundings (such as through conduction, radiation, or evaporation) during the measurement, leading to a discrepancy between ΔH and q.
In your previous problem, you mentioned that ΔHrxn was -1500 kJ/mol and q was 1.519 x 10^6 J. The negative sign of ΔHrxn indicates an exothermic reaction, where heat is released. It is a measure of the energy change associated with the reaction taking place. The q value you obtained represents the amount of heat actually measured by the calorimeter. The discrepancy between the two values can be attributed to the factors mentioned earlier, such as the heat capacity of the calorimeter and the heat losses to the surroundings.
In summary, while ΔH and q are related to energy changes in a reaction, they measure different aspects and are not expected to be exactly equal due to experimental considerations such as the heat capacity of the calorimeter and heat losses to the surroundings.