What is the primary difference between a solution and a suspension or colloid?

The primary difference lies in the particle size of the dispersed phase and the homogeneity of the mixture. In a true solution, the solute particles are extremely small (less than 1 nm), invisible to the naked eye, and uniformly distributed, forming a homogeneous mixture that does not scatter light and cannot be separated by filtration. Suspensions have much larger particles (greater than 100 nm) that settle over time, are visible, and form heterogeneous mixtures. Colloids have intermediate particle sizes (1-100 nm), appear homogeneous but are microscopically heterogeneous, scatter light (Tyndall effect), and do not settle.

Why is molality preferred over molarity for expressing concentration in colligative property calculations?

Molality (moles of solute per kg of solvent) is preferred because it is temperature-independent. The mass of the solvent does not change with temperature. In contrast, molarity (moles of solute per litre of solution) is temperature-dependent because the volume of the solution changes with temperature. Since colligative properties are often studied at varying temperatures, using a temperature-independent concentration unit like molality ensures consistency and accuracy in calculations.

Explain why the solubility of gases in liquids decreases with increasing temperature.

The dissolution of a gas in a liquid is generally an exothermic process, meaning it releases heat. According to Le Chatelier's principle, if an equilibrium system is subjected to a change in temperature, it will shift in a direction that counteracts the change. If the temperature is increased, the equilibrium will shift to favor the reverse process (gas coming out of solution) to absorb the added heat. This results in a decrease in the solubility of the gas. This phenomenon is evident in aquatic environments, where warmer water holds less dissolved oxygen, impacting marine life.

What are azeotropes, and how do they relate to ideal and non-ideal solutions?

Azeotropes are constant boiling mixtures that distill without any change in their composition. They behave like pure compounds during distillation, even though they are mixtures. Azeotropes are formed by non-ideal solutions that show significant deviations from Raoult's Law. Solutions showing positive deviation form minimum boiling azeotropes (e.g., ethanol-water), where the boiling point is lower than either pure component. Solutions showing negative deviation form maximum boiling azeotropes (e.g., nitric acid-water), where the boiling point is higher than either pure component. Ideal solutions, by definition, do not form azeotropes.

How does the Van't Hoff factor account for abnormal molar masses?

The Van't Hoff factor (i) is introduced to correct colligative properties for solutions where the solute undergoes dissociation (like electrolytes) or association (like carboxylic acids in non-polar solvents). When dissociation occurs, the number of particles increases, leading to a lower observed molar mass (abnormal molar mass) and higher colligative property values than expected. When association occurs, the number of particles decreases, leading to a higher observed molar mass and lower colligative property values. The Van't Hoff factor is the ratio of the observed colligative property to the calculated colligative property (or normal molar mass to abnormal molar mass), effectively quantifying the extent of dissociation or association and allowing accurate prediction of colligative properties.

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Solutions

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Version 1Updated 22 Mar 2026

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A solution is defined as a homogeneous mixture of two or more non-reacting chemical substances whose composition can be varied within certain limits. The components of a solution are typically classified as the solute and the solvent. The solvent is the component present in the largest quantity, or the one that determines the physical state of the solution. The solute is the component present in a…

Quick Summary

Solutions are homogeneous mixtures where a solute is uniformly dispersed in a solvent. They can be solid, liquid, or gaseous. Key ways to express solution concentration include mass percentage, volume percentage, parts per million (ppm), mole fraction (ratio of moles of a component to total moles), molarity (moles of solute per liter of solution, temperature-dependent), and molality (moles of solute per kg of solvent, temperature-independent).

Solubility, the maximum amount of solute that dissolves, is affected by the nature of solute/solvent, temperature, and pressure (for gases, Henry's Law). Raoult's Law describes the vapor pressure of solutions, leading to concepts of ideal and non-ideal solutions (positive/negative deviations) and azeotropes.

Colligative properties (relative lowering of vapor pressure, elevation in boiling point, depression in freezing point, osmotic pressure) depend only on the number of solute particles, not their identity.

For electrolytes, the Van't Hoff factor (i) corrects these properties for dissociation or association, accounting for abnormal molar masses.

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Key Concepts

Molarity (M) and its Calculation

Molarity is a fundamental concentration unit, representing the number of moles of solute present in exactly…

Molality (m) and its Calculation

Molality is another concentration unit, defined as the number of moles of solute dissolved per kilogram of…

Osmotic Pressure (

\Pi

) and its Application

Osmotic pressure is one of the colligative properties, defined as the pressure that must be applied to the…

Solution: — Homogeneous mixture of solute and solvent.

Concentration Units:

- Mass %: $(\text{mass of solute} / \text{mass of solution}) \times 100$ - Volume %: $(\text{volume of solute} / \text{volume of solution}) \times 100$ - ppm: $(\text{mass of solute} / \text{mass of solution}) \times 10^6$ - **Mole Fraction ( $\chi$ ):** $n_A / (n_A + n_B)$ - Molarity (M): $\text{moles of solute} / \text{volume of solution (L)}$ (Temperature-dependent) - Molality (m): $\text{moles of solute} / \text{mass of solvent (kg)}$ (Temperature-independent)

Henry's Law: —

P_{\text{gas}} = K_H \chi_{\text{gas}}

Raoult's Law: —

P_A = P_A^0 \chi_A

; for non-volatile solute:

\frac{P^0 - P_s}{P^0} = \chi_{\text{solute}}

Colligative Properties (for non-volatile solute):

- RLVP: $\frac{P^0 - P_s}{P^0} = i \chi_{\text{solute}}$ - ** $\Delta T_b$ (Elevation in BP):** $\Delta T_b = i K_b m$ - ** $\Delta T_f$ (Depression in FP):** $\Delta T_f = i K_f m$ - ** $\Pi$ (Osmotic Pressure):** $\Pi = i CRT$

Van't Hoff Factor (i): —

i = \frac{\text{observed CP}}{\text{calculated CP}}

i = 1 + (n-1)\alpha

(dissociation) or

i = 1 + (1/n-1)\alpha

(association).

To remember the four Colligative Properties and their dependence on 'i':

Really Low Vapor Pressure Elevates Boiling Points Depresses Freezing Points Osmotic Pressure Increases

Each of these is directly proportional to the I (Van't Hoff factor) and the concentration (molality for $\Delta T_b$ , $\Delta T_f$ ; molarity for $\Pi$ ; mole fraction for RLVP). So, think of 'RLVP, EBP, DFP, OPI' as a sequence, all linked by 'i'.

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