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I conducted an experiment in which I added Sodium Nitrate, Potassium Nitrate, Potassium Chloride, and Sodium Hydroxide - all separately into water - then measured the temperature change of each. The results were endothermic reactions for the first three reactions, and it was exothermic for Sodium Hydroxide. (sulfuric acid was used as a catalyst for the exothermic reaction).

Information that I have: Mass/volume for each substance, Initial Temperature, and Final Temperature.

Now I want to find the q (energy) for each solution, the enthalpy change for each solution, and the enthalpy change for the reaction.

I'm thinking I can use the heat capacity of water (4.18 J/gK) to calculate, however, I'm not sure.

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  • $\begingroup$ What relationships (mathematical equations) are you familiar with that relate heat (q) and the heat capacity of water or a solute in water? $\endgroup$ – bobthechemist Nov 6 '13 at 20:25
  • $\begingroup$ I am familiar with q= mct , -q reactants = q surroundings , Hess's Law ... Grade 12 chemistry level. $\endgroup$ – Ds.109 Nov 7 '13 at 3:14
  • $\begingroup$ Yes. The water is absorbing the heat from the reaction, so 4.18 J/g is the appropriate value for c to use. I presume you also know the mass of water used so you can use $q=mc \Delta T$ $\endgroup$ – bobthechemist Nov 7 '13 at 3:39
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Assuming that the heat capacity of the system is constant, you could use the temperature differences you measured and calculate the enthalpy change: $\Delta H = C_P\Delta T$. At constant pressure (which you presumably have), the enthalpy change is equal to the heat entering or leaving the system, i.e. $q=\Delta H$.

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  • $\begingroup$ I don't have any pressure information - however I assume that the pressure was constant.. it was in a coffee cup calorimeter. ... I would use q=mct - however I don't know if I can use 4.18 J/g C as the heat capacity.. can I because its dissolved in water ? $\endgroup$ – Ds.109 Nov 7 '13 at 3:16
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    $\begingroup$ @Ds.109 The coffee cup calorimeter is open to the atmosphere (lab). Do you think that the pressure of the lab changed appreciably because of the reaction you performed? If not, then your assumption is correct and the reaction was performed at a constant pressure. A similar experiment, bomb calorimetry, the pressure is not constant so q would not equal the enthalpy in this case. $\endgroup$ – bobthechemist Nov 7 '13 at 3:43

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