A calorimeter contains 25.0 mL of water at 12.0 C . When 1.20 g of X (a substanc
ID: 961844 • Letter: A
Question
A calorimeter contains 25.0 mL of water at 12.0 C . When 1.20 g of X (a substance with a molar mass of 44.0 g/mol ) is added, it dissolves via the reaction X(s)+H2O(l)X(aq) and the temperature of the solution increases to 26.0 C . Calculate the enthalpy change, H, for this reaction per mole of X. Assume that the specific heat of the resulting solution is equal to that of water [4.18 J/(gC)], that density of water is 1.00 g/mL, and that no heat is lost to the calorimeter itself, nor to the surroundings. Express the change in enthalpy in kilojoules per mole to three significant figures.
Explanation / Answer
q = mc (delta T)
= (25.0 g H2O + 1.20 g X) x (4.18 J/gC) (26.0 oC-12.0 oC)
= (26.2 g) x (4.18 J/goC) x (14.0 oC)
= 1533.224 J
Now,
q / [(1.20 g) / (44.0g/mole)] = - delta H
It should be noted that in this case delta H is negative because delta H is always the opposite of the q you calculated because the q calculated is the q for the surrounding and not system. The delta H that you want to find is q system that is why the delta H have an opposite sign of the q.
1533.224 J / 0.027 mole= - delta H
56786.07 J/mole = - delta H
56.8 kJ/mole = - delta H
delta H = - 56.8 kJ/mole.
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