So I have been working on my chemistry, and I'm stuck on understanding one of the enthalpy problems. I've tried to get ahold of my teacher but he has'nt gotten back to me.
Anyway.
This is the question.
Determine the molar heat of a chemical reaction,and compare the obtained value with the widely accepted value for the neutralization of sodium hydroxide.
30.0mL of 1.0 mol/L h(2)SO(4) (aq) is combined with 50.0 mL of 1.0 mol/L NaOH(aq). Assume the secific heat capacity and density is identical to water.The initial temp of H(2)SO(4) is 28.2*C, and NaOH is 27.9*C. The final temp of the combined solution is 35.9*C.
Q=mc^T
mass: 30mL + 50mL= 80mL(or grams)
capacity: 4.19 g*C (was told that was the specific heat capacity of water that we would be using)
change in temperature:28.2+27.9=56.1/2=28.05 average initial temp. temp. temp final minus temp initial =7.85*C
30.0mL of 1 mol/L solution is 0.03 moles
50.0mL of 1mol/L solution is 0.05 moles
Together this is 0.08 mols.
So: (80 x 4.19 x 7.85) / 0.08=32891.5 J convert that to kJ, you have 32.9 kJ
The widely accepted value is -57 kJ, so : 32.9/57= 0.5771929825, meaning the percent error is approx. 58% for the calorimeter used?
If I'm right please tell me if I've made a mistake or are just completely wrong please help. I'm getting so frustrated trying to learn this on my own.