Solutions – NEET UG ChemistryPractice Questions & PYQs

Molarity is defined as the number of moles of solute per liter of solution, which can be determined using the mass percentage and density of the mixture. A 98% by weight solution contains 98 grams of sulfuric acid in 100 grams of solution, translating to exactly one mole of solute based on a molar mass of 98 grams per mole. Dividing the total mass of 100 grams by the density of 1.84 grams per cubic centimeter gives a volume of 54.35 milliliters, or 0.05435 liters. Dividing the one mole of sulfuric acid by this volume in liters results in a concentration of 18.4 molar.

Raoult's law establishes that the relative decrease in vapor pressure experienced by a solvent upon the addition of a non-volatile solute depends strictly on the relative number of solute particles present. Mathematically, this property is directly equated to the ratio of moles of the solute to the total moles in the mixture. Consequently, the extent of this vapor pressure depression acts as a direct measure of the concentration of the solute expressed in terms of its mole fraction.

Strong intermolecular attractions between unlike molecules in a binary liquid mixture restrict individual molecular escaping tendencies and cause negative deviations from Raoult's law. In a mixture of acetone and chloroform, the electron-rich oxygen atom of acetone acts as a hydrogen bond acceptor while the hydrogen atom of chloroform, polarized by three chlorine substituents, acts as a donor. The resulting hydrogen bonds formed between acetone and chloroform molecules are significantly stronger than the dipole-dipole interactions found in either pure liquid component. Consequently, the molecules are held together more tightly, leading to a reduced vapor pressure and negative deviation.

The relative lowering of vapour pressure for a solution containing a non-volatile solute is governed by Raoult's law, represented by the equation $\frac{P^\circ - P}{P^\circ} = \frac{w M}{m W}$. Substituting the given values into this expression, the mass of the solute is 0.5 grams, its molecular weight is 65, the mass of carbon tetrachloride is 100 grams, and its molecular weight is 154 alongside a pure vapour pressure of 143 mm Hg. Solving this proportional relationship for the vapour pressure of the solution yields 141.43 mm Hg.

Normality expresses concentration in gram equivalents per litre, meaning the equivalent weight of a dibasic acid is half its molar mass because it furnishes two replaceable hydrogen ions per molecule. Dividing the molecular weight of 200 by the basicity of 2 gives an equivalent weight of 100. Using the relation $\frac{w}{E} = \frac{V \cdot N}{1000}$, substitute the volume of 100 mL and normality of 0.1 N to find the required mass. This gives $\frac{w}{100} = \frac{100 \times 0.1}{1000}$, which simplifies to a mass of 1 g.

The depression in freezing point is governed by the relation $\Delta T_f = i K_f m$, where $i$ represents the van't Hoff factor, $K_f$ is the cryoscopic constant, and $m$ is the molality. A higher value of the product $im$ results in a greater freezing point depression, which in total lowers the freezing point. For 0.01 m NaCl, the van't Hoff factor is 2, yielding an effective particle concentration of 0.02. The other options yield effective particle concentrations of 0.005 for C2H5OH, 0.015 for MgI2, and 0.01 for MgSO4. Because the sodium chloride solution generates the maximum effective concentration, it undergoes the largest freezing point depression.

Ideal solutions are characterized by uniform intermolecular forces between unlike molecules that match those between like molecules, thereby satisfying Raoult's law across the entire concentration range. When polar substances like chloroform and acetone are mixed, they form intermolecular hydrogen bonds that restrict molecular escape and produce a negative deviation from ideal behavior. Conversely, mixtures of aniline and phenol also establish strong intermolecular hydrogen bonding, which similarly causes a negative deviation rather than a positive one.

Osmosis occurs when solutions separated by a semipermeable membrane have unequal solute concentrations, driving net solvent movement across the membrane. When red blood cells are suspended in a medium that exerts an identical osmotic pressure to the cell interior, no net flow of water occurs in or out of the cells. Consequently, the cellular membrane experiences no osmotic stress, allowing the cells to maintain their standard biconcave shape without swelling or crenating.

According to Raoult's law, the relative lowering of vapour pressure for a solution containing a non-volatile solute is equal to the mole fraction of the solute in the solution. For dilute solutions, the mole fraction of the solute can be approximated as the ratio of the number of solute moles or molecules to the number of solvent moles or molecules. This direct proportionality stems from the reduction in escaping tendency of solvent molecules at the surface caused by the presence of non-volatile solute particles.

Freezing point depression is a colligative property directly proportional to the total concentration of dissolved particles, governed by the relation $\Delta T_f = i K_f m$. Since all given solutions share the same molality, the extent of freezing point depression depends entirely on the number of ions produced per formula unit upon dissociation. Aluminium sulfate dissociates into the highest count of ions among the choices, maximizing its particle concentration. Consequently, this large particle count drives the freezing point down to the minimum value.

Osmotic pressure is a colligative property directly proportional to the number of moles of solute dissolved in a given volume, following the relation $\Pi = \frac{n}{V}RT$. Because all three substances have the same mass and are dissolved in the same volume, the osmotic pressure depends entirely on the inverse of their respective molar masses. Urea possesses the smallest molecular weight, yielding the highest number of moles and thus the highest osmotic pressure, followed by glucose, while sucrose features the largest molecular weight, resulting in the lowest osmotic pressure.

The relative lowering of vapour pressure is equal to the mole fraction of the non-volatile solute, linking the masses and molar masses of the components through Raoult's law. Substituting the given vapour pressures and mass values into the relationship $\frac{P^{\circ}-P_{s}}{P^{\circ}}= \frac{w \cdot M}{W \cdot m}$ alongside the standard molar mass of benzene at $78 \,\text{g/mol}$ establishes the working equation. Solving $\frac{640 - 600}{640} = \frac{2.175 \times 78}{39.08 \times m}$ directly isolates the unknown value. This calculation results in a molecular weight of 69.5 for the solid substance.

Mixing two liquids to form an ideal solution involves the complete energetic compensation of breaking existing intermolecular attractions and forming new ones, resulting in zero net enthalpy and volume changes represented as $\Delta_{\text{mix}} H = 0$ and $\Delta_{\text{mix}} V = 0$. Although these thermodynamic parameters equal zero, it is completely incorrect to assume that no interactions occur between the molecules. In reality, strong intermolecular forces actively operate between molecules P and Q, but they happen to be identical in magnitude to the forces present in the pure components. Therefore, the assertion accurately describes the thermodynamic properties of ideal mixing, whereas the reason falsely claims an absence of molecular interactions.

Ideal solutions strictly obey Raoult's law across all concentrations because the intermolecular forces between unlike molecules are identical to those between like molecules. Because these attractive forces remain unchanged upon combination, there is no heat absorbed or released, meaning $\Delta H_{\text{mixing}} = 0$. Furthermore, the uniform molecular spacing ensures that no expansion or contraction occurs, resulting in $\Delta V_{\text{mixing}} = 0$.

Concentration units that rely on solution volume are temperature-dependent because liquids expand or contract as thermal conditions change. Molality avoids this limitation by relating the amount of solute to the mass of the solvent rather than the volume of the solution. Since mass remains entirely invariant with temperature variations, molality stays constant regardless of heating or cooling.

According to Raoult's law, the fractional decrease in vapour pressure of an ideal solution equals the mole fraction of clones or non-volatile solute. Setting up the initial ratio $\frac{\Delta P}{P^{\circ}} = x_1$ with a pressure drop of 10 mm and a mole fraction of 0.2 gives $\frac{10}{P^{\circ}} = 0.2$, which determines the pure vapour pressure to be 50 mm. When the vapour pressure decrease increases to 20 mm, the new solute mole fraction becomes $\frac{20}{50} = 0.4$. Subtracting this solute mole fraction from one yields the required mole fraction of the solvent as 0.6.

The molar concentration of a solution is determined by the number of moles of solute present in one liter of solution. Calculating the molar mass of methanol yields $32$ grams per mole by combining the atomic masses of its constituent elements. Multiplying the molarity, the volume in milliliters, and the molar mass, then dividing by $1000$ to adjust for milliliters, gives the required mass. Substituting the values into the relation $\text{Mass} = \frac{M \times V \times \text{Molar Mass}}{1000}$ yields a total of $9.6$ grams of methanol.

Concentration metrics and gas solubility behaviors are determined by fundamental definitions where molality is calculated as moles of ethanoic acid divided by the mass of benzene in kilograms, yielding 0.04167 moles divided by 0.075 kilograms to give 0.556 molal. Similarly, molarity is found by dividing moles of sodium hydroxide by the solution volume in liters, resulting in 0.125 moles divided by 0.45 liters to give 0.278 M. Furthermore, lowering the temperature increases gas solubility according to Henry's law, which explains why aquatic species find cold water more hospitable due to higher dissolved oxygen levels. Conversely, gas solubility drops when pressure decreases rather than increases, and the mole fraction of component $B$ is properly expressed as $x_B = \frac{n_B}{n_A + n_B}$ with its own moles in the numerator.

Ideal solutions are formed when two components have similar structures, polarities, and intermolecular forces, allowing them to mix without any change in enthalpy or volume. Mixtures such as benzene with toluene, or chlorobenzene with bromobenzene, consist of structurally analogous molecules that form nearly ideal solutions. In contrast, pairing an alkyl halide like iodoethane with a strongly hydrogen-bonded polar liquid like ethanol introduces drastic differences in intermolecular interactions. This chemical dissimilarity causes strong deviations from ideal behavior, preventing the formation of an ideal solution.

Isotonic solutions possess identical osmotic pressures at the same temperature, which implies their molar concentrations must be equal. Equating the concentration terms for both solutions using the relation $\frac{W_1}{M_1 V_1} = \frac{W_2}{M_2 V_2}$ sets up the equality based on mass, molar mass, and volume. Substituting the given mass percentages and known molar mass into $\frac{5}{342 \times 0.1} = \frac{1}{M_2 \times 0.1}$ provides the necessary framework. Solving this equation for the unknown molar mass yields 68.4.