A coffee-cup calorimeter is a polystyrene cup, a lid and a thermometer, and it measures enthalpy changes by refusing to let heat escape. The logic is a chain, and the chain has exactly three links. One: the solution's heat, $Q = mc\,\Delta T$, where $m$ is the mass of the whole SOLUTION in grams (not the solute — the water is what got warmed) and $c$ is water's $4.18\ \mathrm{J/(g\cdot^{\circ}C)}$. Two: the amount that reacted, from $C = \dfrac{n}{V}$ — C equals n over V — where $C$ is the molar concentration in moles per litre, $n$ the amount of solute in moles, and $V$ the volume of solution in litres. Millilitres must cross to litres before they touch that formula; the per-LITRE in mol/L is not negotiable. Three: $\Delta H = \dfrac{q}{n}$, reported in kJ/mol.

One flip lives between links one and three, and it is the whole reason calorimetry has a reputation. The heat you measured is what the WATER gained; the enthalpy you report belongs to the REACTION, and the reaction is the one that paid for it. So $q_{\text{rxn}} = -Q_{\text{solution}}$: warm cup, negative $\Delta H$. Assumptions worth saying out loud, because a good answer states them: the dilute solution has water's density and water's specific heat, and the cup itself absorbs nothing. Both are approximations, and both are the reason your answer lands near the textbook value rather than on it.