Confusion between vapor pressure vs. vapor-liquid equilibrium

Question

I have confused myself when thinking about vapor-liquid equilibrium. The typical way that I saw PV diagrams explained is that the left edge of the two-phase region is the boiling point of the liquid (bubble point is the same as boiling point for pure liquids?). But as we know, liquids will always generate some vapor pressure when sitting in a container even well below its boiling point. So what is different about existing inside the two-phase part of the diagram?

Is the description of bubble point = boiling point incorrect, and the only way to land in the region from A to B in the picture is to compress the liquid so that there is simply no room for gas to exist?

Answer 1

Let's see how this can be quantified.

For an equilibrium (saturated) mixture of liquid water and water vapor in the tank (2 phase region), let

$m=\mathrm{mass\ of\ water\ in\ tank}$
$x=\mathrm{mass\ fraction\ vapor\ in\ tank}$
$V=\mathrm{tank\ volume}$
$T=\mathrm{temperature}$
$P(T)=\mathrm{equilibrium\ vapor\ pressure\ at temperature}\ T$
$v_L(T)=\mathrm{specific\ volume\ of\ saturated\ liquid\ water\ at\ temperature}\ T$
and
$v_V(T)=\mathrm{specific\ volume\ of\ saturated\ water\ vapor\ at\ temperature}\ T$

then you can calculate the mass fraction of vapor of in the tank by writing: $$m(1-x)v_L+m(x)v_V=V$$ That is, the volume of saturated water vapor in the tank plus the volume of saturated liquid water in the tank must equal the (constant) tank volume. If you change the temperature, the fraction of vapor, the fraction of liquid, and the masses of vapor and liquid will change in order to preserve to total volume of the rigid container.

Answer 2

It would help to label the axes appropriately: the x-axis is volume at a given temperature for each isothermal line [A-B-D and N]; the y-axis is pressure of the gas, of the gas on the liquid or the mechanical force on the liquid from right to left. There is no air or inert gas present. This diagram is explained nicely in Maron and Prutton, "Principles of Physical Chemistry" pp 48 f.

At the right low gas pressure means no liquid; increasing pressure below the critical T eventually induces liquid formation in equilibrium; attempts to increase pressure simply increase liquid while decreasing vapor at constant pressure. Eventually there is no more vapor and pressure on the liquid measures the compressibility of the liquid.

Above the critical T the curves show the behavior of the gas to the various gas laws [whether empirical or actual laws depends on the author] eventually Boyles or the Ideal Gas Law. Again this is explained understandably in M&P.

Answer 3

Following the orange arrow as the volume is reduced along the isotherm in the diagram below:

We start with vapor which increases in pressure at constant T as it is compressed. At point D the substance begins to condense into liquid (therefore dew point). The amount of liquid increases between D and B as described in Chet Miller's answer until at point B only liquid is present. Further reduction in the volume results in a rapid increase in p as liquids are relatively incompressible.

If the direction of the process is reversed the liquid expands and pressure drops at constant T until point B is encountered and vapor begins to form (fittingly named the bubble point).