A 109.2 mL sample of 1.00 M NaOH is mixed with 54.6 mL of 1.00 M H2SO4 in a large Styrofoam coffee cup; the cup is fitted with a lid through which passes a calibrated thermometer.

Question

The temperature of each solution before mixing is 22.65 °C. After adding the NaOH solution to the coffee cup and stirring the mixed solutions with the thermometer, the maximum temperature measured is 30.30 °C. Assume that the density of the mixed solutions is 1.00 g/mL, that the specific heat of the mixed solutions is 4.18 J/(g·°C), and that no heat is lost to the surroundings.

Write a balanced chemical equation for the reaction that takes place in the Styrofoam cup. Remember to include phases in the balanced chemical equation.

Calculate the enthalpy change per mole of H₂SO₄ in the reaction.

Answer 1

Balanced equation for the reaction:

2NaOH(aq) + H₂SO₄(aq) ==> Na₂SO₄(aq)+ 2H₂O(l)

q = mC∆T

q = heat = ?

m = mass = 109.2 ml + 54.6 ml = 163.8 ml x 1 g/ml = 163.8 g (this does not include the mass of the NaOH and H2SO4. In more advanced chem. classes, they will include these. If your professor teaches to include these, then you will have to add 9.8 g to the 163.8 g to include these masses.)

C = specific heat = 4.18 J/g/deg

∆T = change in temperature = 30.30 - 22.65 = 7.65 degrees

q = (163.8 g)(4.18 J/g/deg)(7.65 deg)

q = 5238 J for the reaction

To find ∆H per mole H2SO4 we will divide this calculated heat by the moles of H2SO4 used in the reaction.

moles H2O4 = 54.6 ml x 1 L/1000 ml x 1.00 mol/L = 0.0546 moles

∆H = 5238 J / 0.0546 moles = 95,934 J/mole = 95.9 kJ/mol (to 3 significant figures)