Calculate the quantity of heat gained or lost in the following change:
0.44 mol of water evaporates at 100 degrees Celsius
- So what I did was:
1 mol -> 6.01 kJ
0.44 mol -> x
x = 2.64 kJ of energy gained
But apparently this is incorrect? Could somebody please explain to me why and with the steps corrected? I might have a slight idea as to why it could be incorrect, like maybe the molar enthalpy of vaporization is not 6.01 kJ/mol. But then how would I know that value without researching this information? Thank you.
The heat of vaporization is 40.66 kJ/mol for water. I believe your 6.01 must be heat fusion for H2O.
To calculate the quantity of heat gained or lost during a change, you need to consider the enthalpy of vaporization of water. The molar enthalpy of vaporization represents the amount of heat required to convert one mole of a substance from its liquid phase to its gaseous phase at a specific temperature.
In the case of water at 100 degrees Celsius, the molar enthalpy of vaporization is not 6.01 kJ/mol, as you mentioned. Its actual value is approximately 40.7 kJ/mol.
To correct your calculation, you could use the following steps:
1. Multiply the number of moles of water (0.44 mol) by the molar enthalpy of vaporization (40.7 kJ/mol):
0.44 mol × 40.7 kJ/mol = 17.9 kJ
So, the correct answer would be that approximately 17.9 kJ of heat is gained or lost when 0.44 mol of water evaporates at 100 degrees Celsius.
To find accurate values for molar enthalpies, it's usually necessary to consult reliable sources or references, such as textbooks or reputable online databases.