Answering your question as it stands is impossible, as we readers do not have enough information about your actual experiment. What do you mean with 'enthalpy of 1-butanol'? You can only measure changes in enthalpy, so you need to specifiy what 'change' you're looking at.
You could mean the enthalpy of evaporation or the reaction enthalpy of some chemical reaction or hundreds of other possibilities. It's essential to clarify what process you're referring to, so one can associate the corresponding starting (e.g. liquid) and end state (-> gas) to this energy change.
However, you stated that 200 ml of solution were heated up by whatever your process was. At constant pressure, the enthalpy required to change the temperature by $\Delta T$ is
$$\Delta H = m c_p \Delta T$$
where $c_p$ is the isobaric heat capacity. Knowing the value of $c_p$, you can calculate this energy change. Now where did that energy come from? Presumably, from your process. At this point, assume that there are no energy losses (f.e. neglect the heat lost to the surrounding air). Then by conservation of energy, the enthalpy increase measured by the temperature increase must be equal in size to the energy released by the process.
Finally, note that by convention we say if a process releases energy, it's enthalpy change is negative (we could've just done it the other way around; think of it as the same idea as defining a sea level).