Vapor pressure is the pressure exerted by the vapor that is in equilibrium with its own liquid in a closed space. It arises because molecules at the liquid surface are constantly escaping (evaporating) while vapor molecules are returning (condensing); equilibrium is reached when these two rates are equal, and the corresponding vapor pressure is the saturation pressure at that temperature.

Raising the temperature increases vapor pressure because of what happens at the molecular level:

• Temperature is a measure of the average kinetic energy of the molecules. Heating the liquid widens the distribution of molecular speeds and raises that average.
• To leave the liquid, a surface molecule must have enough energy to overcome the intermolecular attractive forces holding it in the liquid. At higher temperature a larger fraction of molecules exceed this escape energy.
• More molecules therefore enter the vapor phase before equilibrium is re-established, and a greater vapor density means a higher equilibrium vapor pressure. This is the trend described quantitatively by the Clausius–Clapeyron relation, in which vapor pressure rises roughly exponentially with temperature.

So the correct explanation is that molecules gain kinetic energy and escape into the vapor. The statement that surface tension remains constant is both inaccurate (surface tension actually falls slightly as temperature rises) and irrelevant to why vapor pressure grows. Viscosity reducing exponentially is a separate consequence of heating and does not cause evaporation. Density reducing in a straight-line manner is not the mechanism either, and the liquid's density does not fall linearly; these describe side effects, not the driving cause of increased vapor pressure.