Solutions – NEET UG ChemistryPractice Questions & PYQs

Osmotic pressure behaves as a colligative property that scales directly with the total number of dissolved particles or ions in a given volume. While covalent compounds like glucose and urea remain intact as single molecules in solution, ionic electrolytes dissociate into multiple distinct ions. Since calcium chloride breaks apart into three individual ions per formula unit, whereas potassium chloride yields only two, the calcium chloride solution generates the greatest overall particle concentration and therefore the highest osmotic pressure.

Colligative properties are those thermodynamic properties of solutions that rely strictly on the total number of solute particles rather than their specific chemical nature. While boiling point, vapour pressure, and freezing point require measuring a change or difference relative to the pure solvent, osmotic pressure is itself measured directly as a colligative property. This makes osmotic pressure uniquely categorized as a full colligative property among the available alternatives.

The elevation in boiling point is a colligative property directly proportional to the molality of the solution and the molal elevation constant of the solvent. Applying the boiling point elevation formula $\Delta T_b = \frac{1000 \cdot K_b \cdot w}{m \cdot W}$ isolates the unknown molar mass from the given experimental parameters. Substituting the known values into $0.216 = \frac{1000 \times 2.16 \times 0.15}{m \times 15}$ sets up the final calculation. Rearranging this expression yields a molecular weight of 100 for the solute.

The van't Hoff factor measures the degree of particle multiplication that occurs when an ionic compound dissolves and dissociates completely in water. Potassium ferricyanide breaks apart into four total ions, consisting of three potassium cations and one complex hexacyanoferrate anion. Aluminum nitrate similarly undergoes complete dissociation to release one aluminum cation and three nitrate anions, yielding an identical total count of four independent particles.

Solutions possessing identical molar concentrations and identical particle counts exert equal osmotic pressures when separated by a semipermeable membrane. Because both glucose and urea are non-ionizing solutes that produce the same number of dissolved particles at equal molarity, the solutions are perfectly isotonic. Consequently, solvent molecules pass through the membrane at equal rates in both directions, resulting in zero net transfer.

Boiling point elevation is a colligative property governed by the equation $\Delta T_b = i K_b m$, making the temperature increase directly proportional to the product of the van't Hoff factor and molality. Sodium sulfate dissociates into three ions, giving a van't Hoff factor of 3 and a particle product of 0.03. Both glucose and urea remain undissociated with a factor of 1, resulting in lower products of 0.015 and 0.01 respectively. Potassium nitrate dissociates into two ions, producing a product of 0.02. Consequently, sodium sulfate establishes the highest particle concentration and exhibits the highest boiling point.

Solutions are categorized by the physical states of their constituent solute and solvent components. Examining each pair reveals that humidity represents liquid water dispersed in air, alloys are solid in solid mixtures, amalgams form when liquid mercury dissolves in a solid metal, and smoke consists of solid particulate matter suspended in a gas. Connecting these specific phases directly yields the pairing where humidity corresponds to liquid in gas, alloys correspond to solid in solid, amalgams correspond to liquid in solid, and smoke corresponds to solid in gas.

Raoult's law dictates that the total vapor pressure of an ideal binary mixture can be determined using the formula $P_{\text{total}} = X_x P_x^{\circ} + X_y P_y^{\circ}$ Substituting the given mole fractions and pure component pressures yields $(\frac{5}{15} × 63) + (\frac{10}{15} × 78)$, which simplifies to $21 + 52$, giving a calculated ideal pressure of $73$ torr. Because the experimentally observed vapor pressure is only $70$ torr, the actual value falls short of the ideal prediction. This decrease happens when solute-solvent interactions exceed individual component attractions, causing the system to exhibit a negative deviation from Raoult's law.

Solutions that exhibit identical osmotic pressures at a given temperature are termed isotonic, meaning their molar concentrations are directly equivalent. This relationship is represented by the formula $C_1 = C_2$ which compares the mass and molar mass ratios over volume for both solutes. Equating the concentration of glucose to the given concentration of urea sets up the proportional relation $\frac{m}{180 × 1} = \frac{15}{60 × 1}$ using their respective molar masses. Solving this expression isolates the required mass of glucose as $45$ grams.

Chloroform acts as a weakly polar organic solvent, meaning its dissolving capacity is governed by polarity compatibility and intermolecular forces. Highly ionic or strongly polar substances exhibit poor compatibility with chloroform because their internal electrostatic forces strongly resist solvation by a weakly polar medium. As the dipole moment or ionic character of a solute increases, its affinity for chloroform drops significantly. Consequently, ionic sodium chloride is the least soluble, followed by polar acetonitrile and methanol, while entirely non-polar cyclohexane dissolves most readily.

Ferrous ions dissolved in water have a strong tendency to undergo oxidation and hydrolysis, which leads to the formation of unwanted precipitates unless an appropriate acidic environment is maintained. Introducing a specific mineral acid supplies hydronium ions that suppress this forward hydrolysis pathway without interfering with the redox properties of the salt. Dilute sulphuric acid is chosen for this purpose because it provides the necessary acidic medium while avoiding undesired side reactions or oxidation that stronger or oxidizing acids might cause.

Osmotic pressure follows the relation $\Pi = CRT$. Comparing this equation to a linear equation shows that the slope corresponds to the product of the gas constant and absolute temperature, expressed as $\text{Slope} = RT$. Substituting the given slope and gas constant values yields an absolute temperature of $\frac{25.73}{0.083}$, resulting in approximately $310 \, \text{K}$. Converting this Kelvin temperature to the Celsius scale by subtracting $273$ gives $37^{\circ}\text{C}$.

Henry's law establishes an inverse relationship between gas solubility in a liquid and its corresponding Henry's law constant at a constant temperature. A higher constant value indicates that a greater partial pressure is required to dissolve the gas, which fundamentally translates to a lower overall solubility in water. Evaluating the given constants shows that gas A has the highest value, followed by gas C, and gas B possesses the smallest value. Reversing this sequence to reflect solubility places gas B at the top, followed by gas C, and gas A as the least soluble.

Elevation in boiling point is a colligative property directly proportional to the effective particle concentration of the solution, expressed as $\Delta T_{b} = iK_{b}m$. For $0.05 \, \text{M} \, \text{NaCl}$, the product of the van 't Hoff factor and molarity is $0.1$, which is lower than the values produced by the other solutions due to smaller concentrations or lesser ionic counts. Consequently, the solution with this lowest effective particle concentration exhibits the smallest boiling point elevation.

The solubility of a gas in water is inversely related to its Henry's law constant, expressed as $\text{Solubility} \propto \frac{1}{K_H}$. Examining the provided constants, argon has the highest value at 40.3, followed by carbon dioxide at 1.67, methane at 0.413, and formaldehyde with the smallest value of $1.83 \times 10^{-5}$. Because a lower constant indicates a higher tendency to dissolve, arranging these values from largest to smallest directly yields the reverse order of their dissolution capacity in water. Consequently, the solubility increases progressively from argon through carbon dioxide and methane, culminating with formaldehyde as the most readily dissolved substance.

Molality is calculated using the formula $m = \frac{\text{Moles of solute}}{\text{Mass of solvent in kg}}$. Substituting the given values of $1$ molal for molality and $0.5$ mole for the solute yields a solvent mass of $\frac{0.5}{1}$, which equals $0.5 \, \text{kg}$. Converting this mass from kilograms to grams by multiplying by $1000$ gives $500 \, \text{g}$ of solvent.

The calculation of mass from density and volume requires applying the rule that the final result of a multiplication must be rounded to match the number with the least significant figures. Mass is determined using the relation $\text{Mass} = \text{Density} \times \text{Volume}$. Multiplying $2.15 \, \mathrm{g} \, \mathrm{mL}^{-1}$ by $2.5 \, \mathrm{mL}$ produces an unrounded value of $5.375$, which rounds to $5.4 \, \mathrm{g}$ because the volume term contains only two significant figures.

The total vapour pressure of an ideal binary solution is given by Raoult's law, combining the partial pressures of each volatile component according to their mole fractions. From the given molar ratio of 3 : 2, the mole fractions for benzene and octane are calculated as 3/5 and 2/5 respectively. Substituting these mole fractions and the respective pure vapour pressures of 280 mm Hg and 420 mm Hg into $p_{\text{total}} = p^\circ_{\text{benzene}}\chi_{\text{benzene}} + p^\circ_{\text{octane}}\chi_{\text{octane}},$ the calculation yields 168 plus 168, resulting in a total vapour pressure of 336 mm of Hg.

Osmotic pressure adheres to the relationship $\Pi = iCRT$, meaning it increases directly with the molar concentration of the solution. Because equal masses of the solutes are dissolved in identical volumes, a smaller molar mass yields a higher molar concentration and a greater osmotic pressure. Arranging the solutes by increasing molar mass directly gives the reverse sequence for their respective osmotic pressures.

Positive deviations from Raoult's law occur when the intermolecular forces between solute and solvent molecules are weaker than those in the pure components. Pure ethanol features strong intermolecular hydrogen bonding that holds its molecules tightly together. When acetone is introduced, it intercalates between ethanol molecules and disrupts these hydrogen bonds, leading to weaker overall attractive interactions and higher escaping tendency.