A calorimeter reports joules. A data table reports kilojoules per mole. The bridge is $q = n\,\Delta H$ — q equals n delta-H. Read the three letters carefully, because two of them look alike: $q$ is the total heat for the sample actually burned or dissolved, in kilojoules; $n$ is the amount that reacted, in moles, which almost always arrives through $n = m/M$ first; and $\Delta H$ is the molar enthalpy change, in kilojoules per mole — the price for exactly one mole, and a property of the reaction rather than of your sample. Solve for whichever one the question leaves blank.

Now the sign, which is where marks go to die. Thermochemistry is written from the SYSTEM's point of view: the reaction is the accountant, and it reports its own balance. An exothermic reaction releases energy, so the system loses it and $\Delta H$ is negative; an endothermic reaction absorbs energy, so $\Delta H$ is positive. The surroundings — the water in the cup — see exactly the opposite, which is why a negative $\Delta H$ shows up as a thermometer going UP. And $q$ and $\Delta H$ always agree in sign, so if your rearrangement ever hands you a negative number of moles, the signs, not the arithmetic, are what went wrong.