Solutions: common doubts

Question 1

Which concentration terms do not change with temperature?

Accepted Answer

Molality, mole fraction, mass percentage and parts per million stay the same at any temperature, because they are built only from masses or amounts in moles. Molarity and volume percentage change, since they involve the volume of the solution, which grows on warming and shrinks on cooling. That is why molality is used in boiling and freezing point calculations.

Question 2

What is ppm and when is it used?

Accepted Answer

Parts per million gives how many parts of a component are present in a million parts of the solution, usually by mass. It is used when a solute is present in tiny amounts, such as pollutants or fluoride in water, where a percentage would need many decimal places. For example, about 1 ppm of fluoride helps protect teeth from decay, while levels above about 1.5 ppm cause mottling.

Question 3

Why does gas solubility decrease as temperature rises?

Accepted Answer

Dissolving a gas in a liquid usually releases heat, since gas molecules lose energy as they settle into the liquid. By Le Chatelier's principle, heating pushes this equilibrium back towards undissolved gas, so less of it stays in solution; in Henry's law terms, K_H rises with temperature. That is why a warm soft drink goes flat faster and aquatic life needs cool water.

Question 4

Does a larger Henry's law constant mean a gas is more soluble?

Accepted Answer

No, a larger K_H means the gas is less soluble. Henry's law is p = K_H x, so the mole fraction dissolved is x = p/K_H. At the same partial pressure, the gas with the bigger constant ends up with the smaller mole fraction in the liquid. Many students assume the opposite when they first see a table of K_H values.

Question 5

Why do scuba divers breathe air diluted with helium?

Accepted Answer

Deep under water a diver breathes air at high pressure, so by Henry's law more nitrogen dissolves in the blood. If the diver surfaces quickly, the pressure drops and the nitrogen comes out of solution as bubbles that block capillaries, causing the painful condition called the bends. Diluting the air with helium, which is much less soluble in blood, reduces the risk.

Question 6

What does Raoult's law say for a mixture of two volatile liquids?

Accepted Answer

It says the partial vapour pressure of each liquid equals its pure vapour pressure times its mole fraction in the solution: p_i = p_i° x_i. The total vapour pressure is the sum of the two partial pressures. The vapour above the mixture is richer than the liquid in the more volatile component, because its vapour mole fraction is y_i = p_i/p_total.

Question 7

Is Raoult's law a special case of Henry's law?

Accepted Answer

Yes. Henry's law makes the partial pressure of a volatile component proportional to its mole fraction, p = K_H x, and Raoult's law has exactly the same form, p = p° x. Raoult's law is simply the case in which the proportionality constant equals the vapour pressure of the pure component. Both say vapour pressure grows in step with mole fraction.

Question 8

What is the difference between ideal and non-ideal solutions?

Accepted Answer

An ideal solution follows Raoult's law at every composition, with no heat absorbed or released and no change in total volume on mixing; this happens when attractions between unlike molecules match those between like ones, as with benzene and toluene. A non-ideal solution departs from Raoult's law, and mixing it changes enthalpy or volume, because unlike molecules attract each other more or less strongly.

Question 9

Why does ethanol with acetone show positive deviation while chloroform with acetone shows negative deviation?

Accepted Answer

Acetone molecules push between ethanol molecules and break some of their hydrogen bonds, so attractions in the mixture are weaker and molecules escape more easily, giving a higher vapour pressure than Raoult's law predicts. Chloroform's hydrogen forms a new hydrogen bond with acetone's oxygen, so molecules are held more firmly and the vapour pressure is lower than predicted, a negative deviation.

Question 10

Why can't pure ethanol be obtained by distilling an ethanol-water mixture?

Accepted Answer

Because ethanol and water form an azeotrope, a mixture that boils at a constant temperature and gives vapour of the same composition as the liquid. For this pair it occurs at about 95% ethanol by volume and boils lower than either liquid, so fractional distillation stalls there. Mixtures with negative deviation, such as 68% nitric acid in water, form maximum boiling azeotropes instead.

Question 11

Why is the relative lowering of vapour pressure equal to the mole fraction of the solute?

Accepted Answer

With a non-volatile solute present, only the solvent contributes vapour, and Raoult's law gives p₁ = p₁° x₁. Rearranging, (p₁° − p₁)/p₁° = 1 − x₁ = x₂, which is the mole fraction of the solute. The result depends only on how many solute particles are present, not what they are, which makes it a colligative property.

Question 12

Why does a non-volatile solute raise the boiling point of a solvent?

Accepted Answer

A liquid boils when its vapour pressure equals the pressure above it. A dissolved non-volatile solute lowers the vapour pressure, so the solution has to be heated further before its vapour pressure reaches atmospheric pressure. For dilute solutions the rise is ΔTb = Kb m, proportional to molality; a 1 mol kg⁻¹ sucrose solution in water boils 0.52 K above pure water, at 373.52 K.

Question 13

What is the ebullioscopic constant and what are its units?

Accepted Answer

The ebullioscopic constant Kb is the rise in boiling point produced by a solution whose molality is 1 mol kg⁻¹, so its unit is K kg mol⁻¹. Its value depends only on the solvent, not on the solute; for water it is 0.52 K kg mol⁻¹. The cryoscopic constant Kf does the same job for freezing point depression and is 1.86 K kg mol⁻¹ for water.

Question 14

Why is ethylene glycol added to the water in car radiators?

Accepted Answer

It lowers the freezing point of the coolant so it does not freeze in cold weather. A dissolved solute reduces the liquid's vapour pressure, so solid and liquid reach equal vapour pressure only at a lower temperature. The more glycol dissolved, the lower the freezing point, so the coolant stays liquid on cold nights instead of turning to ice that expands and cracks the engine.

Question 15

How do you find the molar mass of a solute from freezing point depression?

Accepted Answer

Combine ΔTf = Kf m with m = (w₂ × 1000)/(M₂ × w₁), where w₂ is the solute's mass and w₁ the solvent's mass in grams, to get M₂ = (Kf × w₂ × 1000)/(ΔTf × w₁). If 1.0 g of a non-electrolyte in 50 g of benzene lowers the freezing point by 0.40 K, with Kf = 5.12 K kg mol⁻¹, then M₂ = 256 g mol⁻¹.

Question 16

Why is osmotic pressure preferred for finding the molar mass of proteins?

Accepted Answer

Osmotic pressure gives a sizeable, easily measured reading even for very dilute solutions, and it is measured at room temperature. A protein's high molar mass means a sample contains very few moles, so the boiling or freezing point changes would be too small to read reliably. Many biomolecules also break down on heating, so a room-temperature method is safer.

Question 17

Why is normal saline used for intravenous injections?

Accepted Answer

Normal saline, 0.9% (m/V) sodium chloride, is isotonic with the fluid inside blood cells, meaning both have the same osmotic pressure, so no net water moves in or out of the cells. In a more concentrated, hypertonic solution, water leaves the cells and they shrink; in a dilute, hypotonic one, such as pure water, water enters and the cells swell and may burst.

Question 18

What is reverse osmosis and how does it purify sea water?

Accepted Answer

Reverse osmosis is the flow of solvent out of a solution through a semipermeable membrane, forced by applying a pressure greater than the osmotic pressure to the solution side. It is used to desalinate sea water through membranes such as cellulose acetate, which let water molecules pass but hold back dissolved salts, so fresh water collects on the other side.

Question 19

What is the van't Hoff factor and why is it needed?

Accepted Answer

The van't Hoff factor i compares the observed colligative effect with the value expected if the solute neither split up nor clumped together, and it enters the formulas as, for example, ΔTf = i Kf m. In dilute solution NaCl gives i close to 2, since each unit releases two ions, and K₂SO₄ gives i close to 3. Leaving out i badly underestimates electrolyte effects.

Question 20

Why does acetic acid show an abnormally high molar mass in benzene?

Accepted Answer

In benzene, acetic acid molecules pair up into dimers held by two hydrogen bonds, which nearly halves the number of particles in solution. Colligative properties depend on particle count, so the freezing point drops only about half as much as expected. The van't Hoff factor is close to 0.5 and the molar mass calculated comes out about double the true 60 g mol⁻¹.

Question 21

How do you find the degree of dissociation or association from the van't Hoff factor?

Accepted Answer

For a weak electrolyte AB that splits into two ions, i = 1 + α, so α = i − 1; for one that gives n ions, i = 1 + (n − 1)α. For dimerisation the particle count falls and i = 1 − x/2, where x is the fraction of molecules that associate. Work out i from the measured colligative property first, then solve for the degree.

Question 22

How is the Solutions chapter usually tested in NEET?

Accepted Answer

Expect numericals on Raoult's law and vapour composition, Henry's law, boiling point elevation, freezing point depression and osmotic pressure, often requiring the van't Hoff factor for electrolytes. Conceptual items cover positive and negative deviations, azeotropes and isotonic solutions. Typical slips are using grams instead of kilograms of solvent, molality instead of molarity in π = CRT, and forgetting i.

Solutions: common doubts, answered

About the chapter