If a student performs an endothermic reaction in a calorimeter, how does the calculated value of \(\Delta H\) differ from the actual value if the heat exchanged with the calorimeter is not taken into account?

An endothermic reaction is a chemical reaction that absorbs heat from its surroundings. During this reaction, energy is taken in by the system, which causes an increase in temperature. A calorimeter is a device used for measuring the heat absorbed or released during a chemical reaction. It helps in calculating ΔH for a given reaction.

In ideal circumstances, when a reaction is performed in a calorimeter, the energy absorbed or released is only associated with the chemicals involved in the reaction. To calculate ΔH, we use the formula: ΔH = q / n where q is the heat exchanged (absorbed or released) during the reaction and n is the amount of substance (in moles) involved in the reaction.

When a reaction is performed in a calorimeter, some heat may be exchanged between the reaction and the calorimeter itself. By not taking this heat exchange into account, we are making the assumption that the calorimeter is a perfect insulator – meaning it does not absorb or lose any heat energy. In reality, no calorimeter is perfectly insulated, so the calculated ΔH will deviate from the actual value. In an endothermic reaction, since the reaction is absorbing heat from its surroundings, if the calorimeter also absorbs some of this heat, the actual value of ΔH will be greater than the calculated value. This is because the calorimeter has absorbed some of the heat that was initially supposed to be absorbed by the reaction.

If the heat exchanged with the calorimeter is not taken into account during the calculation of ΔH for an endothermic reaction, the calculated value will be less than the actual value. The actual value of ΔH in an endothermic reaction would be greater because some heat is absorbed by the calorimeter, in addition to the heat absorbed by the reaction itself.