# What is the differences between partial pressure and vapour pressure?

Was looking at Henry's law and Raoult's law constants and there seemto be lots of equations involved. Henry's law involves partial pressure and the latter involves the vapor pressure. Wondering what the difference is?

In a mixture of gases, each gas has a partial pressure which is the hypothetical pressure of that gas if it alone occupied the volume of the mixture at the same temperature.

What does this mean? For example, if we have a mixture of gases $A$, $B$ and $C$ in an isolated room, then, according to Dalton's law, the pressure exerted by the gases will be the sum of their partial pressures : $$P = p_A + p_B + p_C$$ where $p_A$, $p_B$ and $p_C$ are the partial pressures of each gas. Also, if we have a moles of $A$, b moles of $B$ and c moles of $C$, we can express the partial pressure of each gas as below: $$p_A = \frac{a}{a+b+c} P$$ $$p_B = \frac{b}{a+b+c} P$$ $$p_B = \frac{c}{a+b+c} P$$

Vapor pressure or equilibrium vapor pressure is the pressure exerted by a vapor in thermodynamic equilibrium with its condensed phases (solid or liquid) at a given temperature in a closed system.

As the definition states, if we took an equilibrium as $\ce{A(l) <=> A(g)}$, where $A(l)$ and $A(g)$ are the liquid and the vapor state of the compound $A$ . The vapor pressure of $A$ is the pressure exerted by the gaseous state ($A(g)$) over the liquid state ($A(l)$). Now, if we have a mixture of $A$, $B$ and $C$, we will have the equilibrium between the liquid and gaseous state : $\ce{(A, B, C)(l) <=> (A, B, C)(g)}$ and the vapor pressure $P_v$ (I noted $P_v$ so there won't be any confusion between it and the total pressure $P$ I mentioned before). According to Raoult's law, the vapor pressure $P_v$ is the sum of the products between the vapor pressures of each compound taken separately and the molar fraction of the compound in the mixture (i.e. in the solution) : $$P_v = x_A p_{v^A} + x_B p_{v^B} + x_C p_{v^C}$$

This may look just like Dalton's law, but there are different pressures involved. The partial pressure of a gas is the pressure exerted by a gas in the volume occupied by a mixture of gases, while the vapor pressure of a gas is the pressure exerted by a gas over its condensed phase. Although :

The vapor pressure that a single component in a mixture contributes to the total pressure in the system is called partial pressure.

The vapor pressure still refers to the pressure of a gas over its liquid (or solid) phase in a phase transition (condensed phase$\ce{<=>}$non-condensed phase), while the notion of partial pressure can be extended to a mixture of gases only.

All the definitions are linked to Wikipedia, and they can really cause confusion throughout chemistry learners..

Partial pressure is the pressure exerted by the individual gas in the mixture of different gases at the same temperature.

Vapour pressure is the pressure exerted by the condensed(cooling) gas by lowering the temperature.

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• I don't understand your definition of vapor pressure. It can either mean the total pressure of a mixed gas above a mixture, or more often, the equilibrium pressure of the vapor of a pure compound in equilibrium in with a pure liquid of the same compound. – Curt F. Mar 25 '15 at 18:20