# How does the enthalpy of neutralisation vary for different sodium salts?
+## Introduction
+
+Many chemical reactions involve a transfer of heat energy, exxpressed by the change in enthalpy /\H. This value is positive for endothermic reactions (absorption of energy) and negative for exothermic reactions (release of energy). A characteristic /\H value for a particular chemical reaction can be determined through thermal calorimetry, which measures the temperature change caused by a reaction in an (approximately) thermally-isolated system.
+
+Enthalpy of neutralisation refers to the enthalpy change caused by the reaction with an acid and base undergoing a neutralisation reaction. These reactions convert an acid and a base to a salt and water. The reactants are fully dissociated, so only the cation of the acid and the anion of the base are involved in the reaction, resulting in the following process:
+
+H+ + OH- --> H2O
+
+The aim of this experiment is to determine the enthalpy of neutralisation for three sodium salts (NaOH, Na2CO3 and Na2SO4), thereby showing the effect that different anions and cations can have on the /\H value for a thermochemical equation.
+
## Method
-The calorimeter was first calibrated with 80.0 mL of deionised water. The water was stirred as it was heated, and the average voltage and current draw were recorded for calculation of the calibration constant.
+A copper calorimeter insulated with polystyrene was first calibrated with 80.0 mL of deionised water. The water was stirred as it was heated, and the average voltage and current draw were recorded for calculation of the calibration constant. Power was applied for exactly 300 seconds, and the maximum temperature was recorded with a mercury thermometer. Care was taken when handling this thermometer to minimise risk of breakage. This calibration was repeated three times and an average calibration constant was calculated.
+
+Two of the test solutions, Na2CO3 and Na2SO4, were then made up by diluting appropriate masses of each compound with deionised water to make 1.0 M solutions. The solid sodium hydroxide posed moderate risk since it is highly corrosive and releases heat when dissolving in water. Standard PPE mitigated this risk.
-After emptying and rinsing the calorimeter, 40.0 mL of 1.0 M NaOH was added to the calorimeter and its temperature recorded. 80.0 mL of 1.0 M HCl was then addedd and its temperature also recorded. Power was applied to the calorimeter for exactly 180 seconds, and the highest temperature of the solution was recorded. The solution was constantly mixed to minimise the error due to slow convection.
-The previous procedure was repeated for the other two compounds, Na2CO3 and Na2SO4.
+After emptying and rinsing the calorimeter, 40.0 mL of 1.0 M NaOH was added to the calorimeter and its initial temperature recorded. 40.0 mL of 1.0 M HCl was then addedd, and the temperature of the resulting reaction environment was closely observed. Temperature readings were taken every 30 s until 60 s after a peak temperature was reached. The solution was constantly mixed to minimise the error due to slow convection.
+The previous procedure was repeated for the other two compounds, substituting NaOH with Na2CO3 and then Na2SO4.
## Results
Assumed density of all solutions is 1 g/mL (=> mass = volume)
Assumed heat capacity of solution is the same as water (4.18 JK/g)
+
+As expected, the neutralisation of Na2CO3 and NaOH were significantly exothermic, due to the full ionisation of the reactants. For example, NaOH consists of sodium and hydroxide ions dissolved in water. Since both ions are dissolved, the strong covalent bonds between the hydrogen in the water and the hydroxide ion are broken which releases energy, raising the temperature of the environment.
+
+Na2SO4 was tested to confirm the theory that there should be no heat released in this reaction. This is because sodium sulfate is a neutral salt, i.e. its pH is 7 under SLC. In this case, the Na+ ion only weakly bonds to the water molecules, so cannot be neutralised. Interestingly, an energy change of 8.6 kJ/mol was obserbed with this reaction, which is expected to be caused by distortion of the temperature readings when the solution is added to the calorimeter.