To test the properties of a fertilizer, $\pu{15.0g}$ of urea, $\ce{NH2CONH2_{(s)}}$, is dissolved in $\pu{150 mL}$ of water in a simple calorimeter. A temperature change from $\pu{20.6^\circ C}$ to $\pu{17.8^\circ C}$ is measured. Calculate the molar enthalpy of solution for the fertilizer urea

I worked through this question by finding $Q = mc\Delta T$, and then dividing $Q$ by the moles of urea present. I can tell the process is endothermic because $\Delta T$ is negative, however my answer for $\Delta H$ comes out as negative, which would only make sense if this was an exothermic reaction. I'm not sure where I am wrong to be honest.

Here is my work:

$$ \begin{align} \Delta H &= \frac{(\pu{150ml}) \times (\pu{1g mL^{-1}}) \times (\pu{4.18J g^{-1} K ^{-1}}) \times (\pu{-2.8 K})} {(\pu{15g}/\pu{60.07g})}\\ &= \pu{-7030.59J/mol}\\ &= \pu{-7.03kJ/mol} \end{align} $$

TL;DR - question asks for $\Delta H$ of an endothermic process, not sure if my answer should be positive or negative


1 Answer 1


The sign of Q depends on the perspective. The water temperature decreased because it "lost" heat. The process of dissolving urea required energy, it "gained" energy. If I give you a penny, should that be +1 or -1 penny? Well, it depends who you ask.

In your answer, you are missing a negative sign in $\Delta H=−Q$ the way you start out with $Q$ from the perspective of the water.


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