The normal boiling point for acetone is 56.5°C. At an elevation of 5300 ft the atmospheric pressure is 630. torr. What would be the boiling point of acetone (ÄHvap = 32.0 kJ/mol) at this elevation?
I need to know how to do this!!
what formula do u use for this? is it
the clausius-claperyron? or no? what is the cl-cl equation for anyway?
Yes, use the Clausis-Clapeyron equation.
What pressure would have to be applied to steam at 325°C to condense the steam to liquid water (ÄH vap = 40.7 kJ/mol)?
do u use the Clausius claperyron for this also???
how do u do this?
Please see the correction I made below to the 6.02 kJ/mol for heat vaporization for water. It's under the screen name of Rachelle. Sorry about the error. Check me out on all of this.
To calculate the boiling point of acetone at a specific elevation given the atmospheric pressure, you can use the Clausius-Clapeyron equation. The Clausius-Clapeyron equation relates the boiling point of a substance to its enthalpy of vaporization and the pressure.
The Clausius-Clapeyron equation is given by:
ln(P2/P1) = (∆Hvap / R) * (1/T1 - 1/T2)
Where:
P1 and P2 are the initial and final pressures, respectively
∆Hvap is the enthalpy of vaporization of the substance in joules or kilojoules per mole
R is the ideal gas constant (8.314 J/(mol·K) or 0.008314 kJ/(mol·K))
T1 and T2 are the initial and final temperatures, respectively (in Kelvin)
To use this equation, you need to convert the given temperature to Kelvin because the equation requires temperature in Kelvin.
First, convert the normal boiling point of acetone to Kelvin:
T1 = 56.5°C + 273.15 = 329.65 K
Next, convert the given atmospheric pressure to atm:
P1 = 630 torr / 760 torr/atm = 0.828 atm
Now, you need to solve for the final temperature (T2) using the given information. Rearrange the Clausius-Clapeyron equation to solve for T2:
T2 = ∆Hvap / (R * (1/T1 - ln(P2/P1)))
Plug in the values:
∆Hvap = 32.0 kJ/mol
R = 0.008314 kJ/(mol·K)
P2 = atmospheric pressure at elevation = 0.828 atm
P1 = normal atmospheric pressure = 1 atm
T1 = 329.65 K
T2 = (32.0 kJ/mol) / (0.008314 kJ/(mol·K) * (1/329.65 K - ln(0.828/1)))
Now, you can calculate T2 using a scientific calculator or any software that can handle logarithmic functions.
It is important to note that there might be some error associated with using the equation at high altitudes due to variations in the atmosphere. However, for most practical purposes, this equation provides a reasonable estimate.