I know I pored this already, but I for it to mention the assumption. See below.
A calorimeter contains 30.0 mL of water at 15.0 C. When 1.50 g of X (a substance with a molar mas of 46.0g/ mol is added, it dissolves via the reaction
X (s) + H2O (l) ----> X (aq)
and the temperature of the solution increases to 26.5 C.
Calculate the enthalpy change for this reaction.
ASSUMPTION:
Assume that the specific heat an density of the resulting solution are equal to those of water and that no heat is lost to the calorimeter itself, not to the surroundings.
Is the delta H negative or positive?
I'm not sure I understand the question. The assumptions don't change the way you work the problem; i.e., no heat is lost is obvious and it means density of the solution is 1.00 g/mL which means mass of the solution is 30.0 grams.
To determine whether the enthalpy change (ΔH) for this reaction is negative or positive, we can use the information given in the question and apply the principle of energy conservation.
The question states that when 1.50 g of X is added to the water in the calorimeter, the temperature of the solution increases from 15.0°C to 26.5°C. This increase in temperature indicates that energy is being absorbed by the system, which means that the reaction is endothermic.
For an endothermic reaction, the enthalpy change (ΔH) is positive. So, in this case, ΔH is positive.
However, it is important to note that the assumption given in the question states that no heat is lost to the calorimeter or the surroundings, which means that all the energy absorbed by the system (water and X) is reflected in the increase in temperature of the solution. This assumption simplifies the calculation of ΔH by allowing us to consider the temperature change as the only indicator of energy transfer.
In summary, the enthalpy change (ΔH) for this reaction is positive, indicating that it is an endothermic process.