Denver, Colorado is about 1600 m above sea level, where atmospheric pressure is lower than $1\mathrm{atm}$. As a result, the boiling point of water in Denver is closest to which value?

Use the Clausius-Clapeyron equation to relate the change in boiling point to the change in pressure: \[\ln\left(\frac{P_2}{P_1}\right) = -\frac{\Delta H_{vap}}{R} \left(\frac{1}{T_2} - \frac{1}{T_1}\right)\] where \(P_1\) and \(T_1\) are the vapor pressure and boiling point at standard conditions (1 atm and 373 K), \(P_2\) is the atmospheric pressure at Denver, and \(T_2\) is the boiling point at Denver.

Gather or estimate the atmospheric pressure at Denver (less than 1 atm, approximately 0.83 atm) and use the known enthalpy of vaporization of water (\(\Delta H_{vap}\)) and the gas constant \(R\) to solve for \(T_2\).