Mount Everest rises to a height of 8,850 m above sea level. At a base camp on the mountain the atmospheric pressure is measured to be 314.0 mm Hg. At what temperature (in °C) will water boil at base camp ? The vapor pressure of water at 373 K is 760.0 mm Hg. (ΔH°vap for H2O = 40.7 kJ/mol and R = 8.314 J/mol K).
a. 344°C.
b. 70.8°C .
c. 2.91E-3°C .
d. 79.8°C.
e. 57.8°C.

Water will boil at .

Explanation:

According to clausius-clapeyron equation for liquid-vapor equilibrium:

where, and are vapor pressures of liquid at (in kelvin) and (in kelvin) temperatures respectively.

Here, = 760.0 mm Hg, = 373 K, = 314.0 mm Hg

Plug-in all the given values in the above equation:

                       or,

So,

Hence, at base camp, water will boil at

the ideal gas law equation is as follows

pv = nrt

where p - pressure

v - volume

n - number of moles

r - universal gas constant

t - temperature

so if the pressure, volume and temperature are already known

we are left with n and r

since r is the universal gas constant that has a known fixed value then r too is known

so we are left with 'n'

once we know temperature volume and pressure

we can find the number of moles of gas present using the ideal gas law equation

this is not my own answer, this is an answer for the same question by the user blahblahmali. i do not claim ownership of this account.