3.5.1 Thermodynamics - Enthalpy change (ΔH) of solution
Students should:
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Theory
Energy is required to break the bonds in the solid form of the solute and to separate the particles from one another. This is an endothermic process. The water (or solvent) particles then bond to the solute particles - generally called solvation, or hydration in the case of water. It is an exothermic process.
The overall energy change is the sum of the endothermic and exothermic processes.
MgSO4(s) 
Mg2+(g) + SO42-(g)
Mg2+(g) + SO42-(g)
+ (aq) Mg2+(aq)
+ SO42-(aq)
- 1 - the lattice enthalpy of magnesium sulfate
- 2 - the sum of the hydration enthalpies of the magnesium and sulfate ions
| Enthalpy change of solution = lattice enthalpy + hydration enthalpies |
Experimental
A known mass of a solute is stirred into a sample of solvent and the temperature change in the solvent recorded. The energy change in the solution may be calculated by:
| Enthalpy change = mass x specific heat capacity x temperature change |
| The moles of the solute may then be calculated from the relative molecular mass and the energy change per mole calculated by: |
| Energy per mole = energy change / number of moles |
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Example: 10g of ammonium nitrate (NH4NO3) is stirred into 100cm3 of water and the temperature change recorded. If the temperature decreased by 7ºC what is the enthalpy of solution of the ammonium nitrate. (specific heat capacity of the solution = 4.2 kJ kg-1 ºC-1) Energy change in the water = mass x specific heat capacity x temperature change Energy change = 0.1 x 4.2 x 7 Energy change = +2.94 kJ (the sign is positive because the water temperature decreases, i.e. heat is being transferred from the water to the ammonium nitrate - it is an endothermic change) Moles of ammonium nitrate = 10/80 = 0.125 moles ∴ enthalpy of solution of ammonium nitrate = 2.94/0.125 = +23.52 kJ mol-1 |