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Determining the enthalpy of a reaction
Determine the enthalpy of this reaction:
Mg(OH)2(s) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)
Method
Measure out 50 mL of 1 mol L-1 sulfuric acid (an excess) into a styrofoam cup.
Weigh accurately about 1.5 g of Mg(OH)2 powder.
Measure the temperature of the sulfuric acid.
Quickly add the Mg(OH)2 powder, stir and measure the maximum temperature of the reaction mix.
Measure 50 mL of about 1 mol L-1 sulfuric acid.
Pour the acid into a styrofoam cup (two stacked together will give even better insulation).
Weigh accurately about 1.5 g of Mg(OH)2 powder.
Measure the initial temperature of the sulfuric acid solution.
Tinitial = 16.5 °C
Add the Mg(OH)2 to the acid.
Stir well and watch the temperature. All the powder should dissolve since we have excess acid.
Measure the maximum temperature change.
Tfinal = 28.0 °C
Calculate:
• The amount, in moles, of Mg(OH)2 reacting.
• The total energy released during the reaction.
• The energy change per mole of Mg(OH)2 reacting.
The energy released by the reaction has made the temperature of the solution rise by
28.0 °C – 16.5 °C = 11.5 °C
All dilute aqueous solutions — and 1 mol L-1 is still considered to be dilute — contain very much more water than they do any other reagent. We therefore assume that the specific heat of the solution is equal to that of water, which is 4.18 J °C-1 g-1. That is, it takes 4.18 J of energy to change the temperature of 1 g of the solution by 1 °C.
The mass of 1 mL of water at room temperature is 1 g. We assume that 50 mL of acid has a specific heat equivalent to 50 g of water.
Energy =
mass of
water heated
×temperature change
× 4.18 J °C-1 g-1
Mg(OH)2(s) + H2SO4(aq) → MgSO4(aq) + 2H2O(l)
∆H = -92.9 kJ mol-1
The temperature rose, so the reaction is exothermic and the ∆H is negative: