Vapour Pressure of Liquid Solutions

Q: What is the primary factor affecting the vapour pressure of a pure liquid?

The primary factor affecting the vapour pressure of a pure liquid is temperature. As the temperature increases, the kinetic energy of the liquid molecules also increases. More molecules gain sufficient energy to overcome the intermolecular forces holding them in the liquid phase and escape into the vapour phase. This leads to a higher concentration of vapour molecules above the liquid, resulting in a higher vapour pressure at equilibrium. The nature of the liquid (intermolecular forces) also plays a crucial role, but for a given liquid, temperature is the variable that changes its vapour pressure.

Q: What defines an ideal solution in terms of vapour pressure?

An ideal solution is one that strictly obeys Raoult's Law over the entire range of concentrations and temperatures. In terms of vapour pressure, this means that the partial vapour pressure of each volatile component in the solution is directly proportional to its mole fraction and its vapour pressure in the pure state ($P_A = X_A P_A^0$). Crucially, for an ideal solution, the intermolecular forces between solute-solvent molecules are identical to the average of the intermolecular forces between solute-solute and solvent-solvent molecules. This results in no heat change and no volume change upon mixing.

Q: How do positive and negative deviations from Raoult's Law differ at the molecular level?

Positive deviations occur when the intermolecular forces between solute and solvent molecules (A-B) are weaker than the average of the forces between pure components (A-A and B-B). This makes it easier for molecules to escape into the vapour phase, leading to a higher vapour pressure than predicted. Conversely, negative deviations occur when A-B interactions are stronger than A-A and B-B interactions. These stronger attractions make it harder for molecules to escape, resulting in a lower vapour pressure than predicted by Raoult's Law.

Q: Can azeotropes be separated by simple fractional distillation?

No, azeotropes cannot be separated into their pure components by simple fractional distillation. An azeotrope is a special type of liquid mixture that boils at a constant temperature and distills without changing its composition. This happens because the composition of the vapour phase is exactly the same as the composition of the liquid phase at its boiling point. Since distillation relies on differences in volatility (and thus vapour phase composition), azeotropes behave like pure compounds during distillation, making their separation by this method impossible.

Q: What is the significance of the relative lowering of vapour pressure?

The relative lowering of vapour pressure ($\frac{P^0 - P_s}{P^0}$) is a colligative property, meaning it depends only on the number of solute particles, not their identity. Its significance lies in its direct proportionality to the mole fraction of the non-volatile solute ($X_B$). This relationship allows us to determine the molar mass of an unknown non-volatile solute by experimentally measuring the lowering of vapour pressure. It's a fundamental concept linking macroscopic properties to microscopic particle count.

Updated 24 Mar 2026

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Vapour pressure of a liquid solution is defined as the pressure exerted by the vapours of the volatile components of the solution in equilibrium with the liquid phase at a given temperature. This equilibrium is dynamic, meaning that the rate of evaporation equals the rate of condensation. For solutions containing a non-volatile solute, the vapour pressure is primarily due to the solvent. For solut…

Quick Summary

Vapour pressure is the pressure exerted by the vapour in equilibrium with its liquid phase at a given temperature. For pure liquids, it increases with temperature. When a non-volatile solute is added to a volatile solvent, the vapour pressure of the solution decreases because fewer solvent molecules are available at the surface to vaporize.

Raoult's Law quantifies this, stating that the partial vapour pressure of a volatile component in a solution is proportional to its mole fraction ( $P_A = X_A P_A^0$ ). For solutions with multiple volatile components, the total vapour pressure is the sum of their partial pressures ( $P_{total} = X_A P_A^0 + X_B P_B^0$ ).

Ideal solutions strictly obey Raoult's Law, exhibiting no heat or volume change on mixing, with similar intermolecular forces. Non-ideal solutions deviate: positive deviations show higher vapour pressure (weaker A-B interactions, \Delta H_{mixing} > 0, \Delta V_{mixing} > 0), while negative deviations show lower vapour pressure (stronger A-B interactions, \Delta H_{mixing} < 0, \Delta V_{mixing} < 0).

Significant deviations can lead to azeotropes, constant boiling mixtures that cannot be separated by fractional distillation.

Full explanation

The concept of vapour pressure is fundamental to understanding the physical properties of liquid solutions. It describes the tendency of molecules to escape from the liquid phase into the gaseous phase. For a pure liquid, this pressure is a characteristic property at a given temperature. When we introduce a solute to form a solution, the vapour pressure of the system changes, and this change is governed by the nature of the solute and solvent.

1. Vapour Pressure of Pure Liquids:

Before delving into solutions, let's reiterate that for a pure liquid in a closed container, molecules are constantly transitioning between liquid and vapour phases. At equilibrium, the pressure exerted by the vapour is its vapour pressure. This value increases with temperature because a higher temperature provides more kinetic energy to molecules, enabling more of them to overcome intermolecular forces and escape into the gas phase.

2. Raoult's Law for Solutions:

Raoult's Law is a cornerstone principle that quantitatively describes the vapour pressure of ideal solutions. It can be applied in two primary scenarios:

a) Solutions Containing a Non-Volatile Solute:

When a non-volatile solute (one that does not vaporize significantly at the given temperature, e.g., glucose, urea, common salts) is dissolved in a volatile solvent, the vapour pressure of the solution is solely due to the solvent.

Raoult's Law states that for such a solution, the partial vapour pressure of each volatile component (in this case, only the solvent) in the solution is directly proportional to its mole fraction in the solution.

Mathematically, for a solvent A and a non-volatile solute B:

P_A = X_A P_A^0

Where: *

P_A

is the partial vapour pressure of the solvent in the solution. *

X_A

is the mole fraction of the solvent in the solution.

* $P_A^0$ is the vapour pressure of the pure solvent at the same temperature.

Since $X_A < 1$ (as $X_A + X_B = 1$ and $X_B > 0$ ), it implies that $P_A < P_A^0$ . This means the vapour pressure of the solution is always lower than that of the pure solvent. This phenomenon is known as the lowering of vapour pressure.

The relative lowering of vapour pressure is a colligative property, meaning it depends only on the number of solute particles, not their identity. Relative lowering of vapour pressure is given by:

\frac{P_A^0 - P_A}{P_A^0} = \frac{P_A^0 - X_A P_A^0}{P_A^0} = \frac{P_A^0(1 - X_A)}{P_A^0} = 1 - X_A

Since

1 - X_A = X_B

(mole fraction of solute), we get:

\frac{P_A^0 - P_A}{P_A^0} = X_B

This equation is extremely useful for determining the molar mass of a non-volatile solute.

b) Solutions Containing Two or More Volatile Components:

When a solution consists of two or more volatile liquids (e.g., benzene and toluene), each component contributes to the total vapour pressure. Raoult's Law extends to this scenario, stating that for each volatile component, its partial vapour pressure in the solution is proportional to its mole fraction in the solution. For a binary solution of volatile components A and B: * Partial vapour pressure of A: $P_A = X_A P_A^0$ * Partial vapour pressure of B: $P_B = X_B P_B^0$

According to Dalton's Law of Partial Pressures, the total vapour pressure of the solution ( $P_{total}$ ) is the sum of the partial vapour pressures of its components:

P_{total} = P_A + P_B

Substituting Raoult's Law expressions:

P_{total} = X_A P_A^0 + X_B P_B^0

Since

X_B = 1 - X_A

, we can also write:

P_{total} = X_A P_A^0 + (1 - X_A) P_B^0

P_{total} = (P_A^0 - P_B^0) X_A + P_B^0

This equation shows that the total vapour pressure of an ideal solution varies linearly with the mole fraction of one of its components.

3. Ideal Solutions:

An ideal solution is one that obeys Raoult's Law over the entire range of concentrations and temperatures. For an ideal solution, the intermolecular forces between solute-solvent molecules (A-B) are identical to the intermolecular forces between solute-solute (A-A) and solvent-solvent (B-B) molecules.

This means: * \Delta H_{mixing} = 0: No heat is absorbed or released when components are mixed. * \Delta V_{mixing} = 0: There is no change in volume upon mixing. * Interactions: A-B interactions are similar in magnitude to A-A and B-B interactions.

Examples: Benzene and toluene, n-hexane and n-heptane, chloroethane and bromoethane.

4. Non-Ideal Solutions (Deviations from Raoult's Law):

Most real solutions do not behave ideally and show deviations from Raoult's Law. These deviations arise when the intermolecular forces between solute and solvent molecules are significantly different from those between pure components.

a) Positive Deviation from Raoult's Law:

* Observation: The total vapour pressure of the solution is higher than that predicted by Raoult's Law. * Molecular Explanation: In these solutions, the A-B intermolecular forces are weaker than the average of A-A and B-B interactions.

This means molecules of A and B find it easier to escape from the solution surface into the vapour phase compared to their pure states. Consequently, the partial vapour pressures of A and B, and thus the total vapour pressure, are higher than ideal.

* Thermodynamic Consequences: \Delta H_{mixing} > 0 (endothermic, heat is absorbed) and \Delta V_{mixing} > 0 (volume increases upon mixing). * Examples: Ethanol and water, acetone and ethanol, carbon disulphide and acetone.

b) Negative Deviation from Raoult's Law:

* Observation: The total vapour pressure of the solution is lower than that predicted by Raoult's Law. * Molecular Explanation: Here, the A-B intermolecular forces are stronger than the average of A-A and B-B interactions.

This enhanced attraction makes it more difficult for molecules of A and B to escape from the solution surface into the vapour phase. As a result, the partial vapour pressures of A and B, and the total vapour pressure, are lower than ideal.

* Thermodynamic Consequences: \Delta H_{mixing} < 0 (exothermic, heat is released) and \Delta V_{mixing} < 0 (volume decreases upon mixing). * Examples: Acetone and chloroform (due to hydrogen bonding), nitric acid and water, acetic acid and pyridine.

5. Azeotropes:

Non-ideal solutions that show significant deviations from Raoult's Law can form azeotropes. An azeotrope (or constant boiling mixture) is a liquid mixture that boils at a constant temperature and distills without change in composition. This means the composition of the vapour phase is the same as that of the liquid phase. Azeotropes cannot be separated into their pure components by fractional distillation.

a) Minimum Boiling Azeotropes: Formed by solutions showing large positive deviations from Raoult's Law. At a specific composition, the vapour pressure is maximum, leading to a minimum boiling point. Example: Ethanol (95.6%) and water (4.4%) mixture boils at 351.3 K, lower than pure ethanol (351.5 K) or pure water (373 K).

b) Maximum Boiling Azeotropes: Formed by solutions showing large negative deviations from Raoult's Law. At a specific composition, the vapour pressure is minimum, leading to a maximum boiling point. Example: Nitric acid (68%) and water (32%) mixture boils at 393.5 K, higher than pure nitric acid (359 K) or pure water (373 K).

Real-World Applications:

Distillation — Understanding vapour pressure is crucial for separating liquid mixtures by distillation. Components with higher vapour pressure (lower boiling point) vaporize more readily.

Humidity — The partial pressure of water vapour in the air contributes to atmospheric pressure and determines humidity levels.

Boiling Point Elevation/Freezing Point Depression — These colligative properties are direct consequences of the lowering of vapour pressure by a non-volatile solute.

Industrial Processes — Many chemical processes, from petroleum refining to pharmaceutical manufacturing, involve controlling and predicting vapour pressures of mixtures.

Common Misconceptions:

Vapour pressure is always lowered by adding any solute — This is true for non-volatile solutes. For volatile solutes, the vapour pressure can be higher or lower than the pure components depending on their relative volatility and mole fractions.

Ideal solutions are common — Most real solutions are non-ideal. Ideal solutions are theoretical constructs that provide a baseline for understanding deviations.

Azeotropes are compounds — Azeotropes are mixtures, not pure compounds, even though they boil at a constant temperature and have a fixed composition during distillation. They can be separated by other means, like azeotropic distillation or extractive distillation.

NEET-Specific Angle:

For NEET, questions frequently test the direct application of Raoult's Law for both non-volatile and volatile solutes. Identifying ideal vs. non-ideal solutions based on given properties (\Delta H_{mixing}, \Delta V_{mixing}, or intermolecular forces) is common.

Understanding the characteristics and examples of positive and negative deviations, and the formation of azeotropes, are high-yield areas. Numerical problems often involve calculating the vapour pressure of a solution, the mole fraction of a component, or the molar mass of a non-volatile solute using the relative lowering of vapour pressure formula.

Key Concepts

Raoult's Law for Non-Volatile Solutes

When a non-volatile solute is dissolved in a volatile solvent, only the solvent contributes to the vapour…

Raoult's Law for Volatile Solutes and Total Vapour Pressure

For a solution containing two or more volatile components, each component contributes to the total vapour…

Deviations from Raoult's Law (Molecular Basis)

Real solutions often deviate from ideal behavior. Positive deviations occur when the attractive forces…

Often confused with

Side-by-side differences the NEET paper likes to test.

Vapour Pressure of Liquid Solutions vs Non-Ideal Solutions

Aspect	Vapour Pressure of Liquid Solutions	Non-Ideal Solutions
Obedience to Raoult's Law	Obeys Raoult's Law over the entire range of concentration.	Does not obey Raoult's Law over the entire range of concentration.
Intermolecular Forces	A-B interactions are similar to A-A and B-B interactions.	A-B interactions are different from A-A and B-B interactions (either weaker or stronger).
Enthalpy of Mixing (\Delta H_{mixing})	Zero (no heat absorbed or released).	Non-zero (either positive for endothermic or negative for exothermic).
Volume of Mixing (\Delta V_{mixing})	Zero (no change in volume upon mixing).	Non-zero (either positive for expansion or negative for contraction).
Vapour Pressure	Total vapour pressure is intermediate between pure components, varying linearly with mole fraction.	Total vapour pressure is either higher (positive deviation) or lower (negative deviation) than predicted by Raoult's Law.
Examples	Benzene + Toluene, n-Hexane + n-Heptane, Chloroethane + Bromoethane.	Positive deviation: Ethanol + Water, Acetone + Ethanol. Negative deviation: Acetone + Chloroform, Nitric acid + Water.

Ideal solutions represent a theoretical benchmark where components mix without any change in energy or volume, and their vapour pressures strictly follow Raoult's Law due to identical intermolecular forces.

In contrast, non-ideal solutions are the norm in reality, exhibiting deviations from Raoult's Law because the intermolecular forces between solute and solvent differ significantly from those within the pure components.

These differences lead to measurable heat and volume changes upon mixing, and consequently, the observed vapour pressure is either higher or lower than what Raoult's Law would predict for an ideal mixture.

Why it is tested: For NEET, understanding the distinction between ideal and non-ideal solutions is crucial. Questions frequently test the conditions for ideality (\Delta H_{mixing}, \Delta V_{mixing}, intermolecular forces) and the characteristics of positive and negative deviations, including their examples. This forms the basis for predicting solution behavior and solving related numerical problems.

Questions students ask

6 answered on this topic.

What is the primary factor affecting the vapour pressure of a pure liquid?

The primary factor affecting the vapour pressure of a pure liquid is temperature. As the temperature increases, the kinetic energy of the liquid molecules also increases. More molecules gain sufficient energy to overcome the intermolecular forces holding them in the liquid phase and escape into the vapour phase.

This leads to a higher concentration of vapour molecules above the liquid, resulting in a higher vapour pressure at equilibrium. The nature of the liquid (intermolecular forces) also plays a crucial role, but for a given liquid, temperature is the variable that changes its vapour pressure.

Why does adding a non-volatile solute lower the vapour pressure of a solvent?

When a non-volatile solute is added to a volatile solvent, the solute particles occupy some of the surface area of the solvent. This reduces the number of solvent molecules exposed at the surface that can escape into the vapour phase.

Consequently, the rate of evaporation of the solvent decreases. While the rate of condensation of solvent molecules from the vapour phase remains largely unaffected initially, a new equilibrium is established at a lower concentration of vapour molecules, leading to a lower vapour pressure compared to the pure solvent.

What defines an ideal solution in terms of vapour pressure?

An ideal solution is one that strictly obeys Raoult's Law over the entire range of concentrations and temperatures. In terms of vapour pressure, this means that the partial vapour pressure of each volatile component in the solution is directly proportional to its mole fraction and its vapour pressure in the pure state ( $P_A = X_A P_A^0$ ).

Crucially, for an ideal solution, the intermolecular forces between solute-solvent molecules are identical to the average of the intermolecular forces between solute-solute and solvent-solvent molecules.

This results in no heat change and no volume change upon mixing.

How do positive and negative deviations from Raoult's Law differ at the molecular level?

Positive deviations occur when the intermolecular forces between solute and solvent molecules (A-B) are weaker than the average of the forces between pure components (A-A and B-B). This makes it easier for molecules to escape into the vapour phase, leading to a higher vapour pressure than predicted.

Conversely, negative deviations occur when A-B interactions are stronger than A-A and B-B interactions. These stronger attractions make it harder for molecules to escape, resulting in a lower vapour pressure than predicted by Raoult's Law.

Can azeotropes be separated by simple fractional distillation?

No, azeotropes cannot be separated into their pure components by simple fractional distillation. An azeotrope is a special type of liquid mixture that boils at a constant temperature and distills without changing its composition.

This happens because the composition of the vapour phase is exactly the same as the composition of the liquid phase at its boiling point. Since distillation relies on differences in volatility (and thus vapour phase composition), azeotropes behave like pure compounds during distillation, making their separation by this method impossible.

What is the significance of the relative lowering of vapour pressure?

The relative lowering of vapour pressure ( $\frac{P^0 - P_s}{P^0}$ ) is a colligative property, meaning it depends only on the number of solute particles, not their identity. Its significance lies in its direct proportionality to the mole fraction of the non-volatile solute ( $X_B$ ).

This relationship allows us to determine the molar mass of an unknown non-volatile solute by experimentally measuring the lowering of vapour pressure. It's a fundamental concept linking macroscopic properties to microscopic particle count.

Revise in 30 seconds

Vapour Pressure (VP) — Pressure by vapour in equilibrium with liquid.

Raoult's Law (Non-volatile solute) —

P_s = X_{solvent} P_{solvent}^0

Relative Lowering of VP —

\frac{P_{solvent}^0 - P_s}{P_{solvent}^0} = X_{solute}

Raoult's Law (Volatile solutes) —

P_A = X_A P_A^0

P_B = X_B P_B^0

Total VP —

P_{total} = P_A + P_B = X_A P_A^0 + X_B P_B^0

Ideal Solution — Obeys Raoult's Law,

\Delta H_{mixing} = 0

\Delta V_{mixing} = 0

, A-B forces similar to A-A, B-B.

Positive Deviation —

P_{total} > P_{ideal}

, A-B forces < A-A/B-B,

\Delta H_{mixing} > 0

\Delta V_{mixing} > 0

. (e.g., Ethanol + Water)

Negative Deviation —

P_{total} < P_{ideal}

, A-B forces > A-A/B-B,

\Delta H_{mixing} < 0

\Delta V_{mixing} < 0

. (e.g., Acetone + Chloroform)

Azeotropes — Constant boiling mixtures, cannot be separated by fractional distillation.

- Minimum Boiling Azeotrope: From large positive deviation. - Maximum Boiling Azeotrope: From large negative deviation.

For deviations from Raoult's Law:

Positive Deviation: People Drink Ethanol & Water (Ethanol + Water). They feel Weaker (A-B forces weaker), get Hot ( $\Delta H_{mixing} > 0$ ), and Volume Increases ( $\Delta V_{mixing} > 0$ ). Their Vapour Pressure is High.

Negative Deviation: No Drinking Acetone & Chloroform (Acetone + Chloroform). They feel Stronger (A-B forces stronger), get Cold ( $\Delta H_{mixing} < 0$ ), and Volume Decreases ( $\Delta V_{mixing} < 0$ ). Their Vapour Pressure is Low.