Why is it not possible to obtain pure ethanol by fractional distillation? What general name is given to binary

NCERT Exemplar · Q58

Q.Why is it not possible to obtain pure ethanol by fractional distillation? What general name is given to binary mixtures which show deviation from Raoult's law and whose components cannot be separated by fractional distillation. How many types of such mixtures are there?

Andhra Pradesh BieapLong· 3mImportance★★★★★

95% · 124/131 Questions

🔒 Locked · start free trial →

You're viewing a preview — the full solution, concept, methods & PYQ mapping are locked.

Start your 14-day free trial to unlock the full solution →

Concept understanding — Molality Calculation

Molality: The Concentration That Ignores Temperature

Imagine you're making a cup of sweet tea. You add sugar to hot water, stir, and taste. If you let the tea cool to room temperature, the amount of sugar hasn't changed — but the volume of the liquid has shrunk slightly. If you measured concentration as "grams of sugar per litre of solution," that number would change just because the temperature changed. That's annoying if you're a chemist who needs a reliable, temperature-independent way to describe how much solute is present.

Molality was invented to solve exactly this problem.

The Intuition

Instead of measuring the volume of the solution (which expands and contracts with temperature), molality measures the mass of the solvent. Mass doesn't change with temperature. So molality gives you a concentration that stays the same whether your solution is hot or cold.

Think of it this way:

Molarity = moles of solute per litre of solution (temperature-sensitive)
Molality = moles of solute per kilogram of solvent (temperature-independent)

The solvent is the substance doing the dissolving — usually water. The solute is what gets dissolved — sugar, salt, etc.

The Precise Definition

\text{Molality } (m) = \frac{\text{moles of solute}}{\text{kilograms of solvent}}

The symbol for molality is a lowercase m (not to be confused with M for molarity).

Key points to remember:

The denominator is solvent mass, not solution mass
The unit is mol/kg (often written as simply "m")
It is independent of temperature because mass doesn't change with temperature

Worked Example

Problem: 36 g of glucose ( $C_6H_{12}O_6$ , molar mass = 180 g/mol) is dissolved in 500 g of water. Calculate the molality of the solution.

Step 1: Find moles of solute

\text{Moles of glucose} = \frac{36 \text{ g}}{180 \text{ g/mol}} = 0.2 \text{ mol}

Step 2: Convert solvent mass to kilograms

500 \text{ g} = 0.5 \text{ kg}

Step 3: Apply the formula

m = \frac{0.2 \text{ mol}}{0.5 \text{ kg}} = 0.4 \text{ m}

Note

The answer is 0.4 m (or 0.4 mol/kg). Notice we used the mass of water (500 g), not the mass of the solution (which would be 536 g).

Common Mistake to Avoid

Watch out

Do not use the mass of the solution in the denominator. The formula specifically asks for the mass of the solvent alone. If the problem gives you the total mass of the solution, subtract the mass of the solute to find the solvent mass.

When Do You Use Molality?

Molality is the star in two important situations: …

Why this formula?

Molality Calculation: Why the Formula Works

Molality is a measure of concentration that is temperature-independent — this is its key advantage over molarity. Let's understand why the formula takes the form it does.

The Definition First

Molality ( $m$ ) is defined as:

m = \frac{\text{moles of solute}}{\text{mass of solvent in kg}}

The unit is mol/kg, often written as m (e.g., 0.5 m glucose solution).

Why Mass of Solvent, Not Solution?

This is the critical conceptual point.

The Reasoning

Molarity uses volume of solution → volume changes with temperature (expansion/contraction). So molarity changes with temperature.
Molality uses mass of solvent → mass is invariant with temperature. So molality remains constant regardless of temperature changes.

Key insight: By using the solvent's mass (not the solution's volume), we eliminate temperature dependence. This is why molality is preferred for colligative properties (boiling point elevation, freezing point depression) — these properties depend on the number of solute particles, not on temperature.

Deriving the Formula Step-by-Step

Step 1: Moles of Solute

If you have $w_{\text{solute}}$ grams of solute with molar mass $M_{\text{solute}}$ (g/mol):

\text{moles of solute} = \frac{w_{\text{solute}}}{M_{\text{solute}}}

Step 2: Mass of Solvent in kg

If the solvent mass is $W_{\text{solvent}}$ grams:

\text{mass of solvent in kg} = \frac{W_{\text{solvent}}}{1000}

Step 3: Putting It Together

m = \frac{\frac{w_{\text{solute}}}{M_{\text{solute}}}}{\frac{W_{\text{solvent}}}{1000}}

Simplifying:

m = \frac{w_{\text{solute}} \times 1000}{M_{\text{solute}} \times W_{\text{solvent}}}

The Final Formula (Exam-Ready)

\boxed{m = \frac{w_{\text{solute}} \times 1000}{M_{\text{solute}} \times W_{\text{solvent}}}}

Where:

$w_{\text{solute}}$ = mass of solute in grams
$M_{\text{solute}}$ = molar mass of solute in g/mol
$W_{\text{solvent}}$ = mass of solvent in grams

Why the ×1000 Factor? …

Start your 14-day free trial to unlock the full solution →

Pure ethanol cannot be obtained by fractional distillation because ethanol and water form a minimum-boiling azeotrope at 95.6% ethanol by mass. Such mixtures are called azeotropes, and they come in two types: minimum-boiling and maximum-boiling azeotropes.

Why fractional distillation fails for ethanol

Fractional distillation works because different liquids have different boiling points. When you heat a mixture, the vapour is richer in the more volatile component. By repeatedly condensing and re-vaporising, you can separate the components almost completely — provided the vapour composition keeps shifting away from the liquid composition.

Ethanol (boiling point 78.4 °C) is more volatile than water (100 °C). So you’d expect that distilling a dilute ethanol solution would give you progressively stronger ethanol. And it does — up to a point.

But at about 95.6% ethanol by mass (roughly 96% by volume), something strange happens: the vapour has exactly the same composition as the liquid. The mixture boils at a constant temperature of 78.2 °C — lower than either pure ethanol or pure water. This is a minimum-boiling azeotrope.

Once you reach this composition, further distillation does nothing. The liquid and vapour are identical in composition, so no enrichment occurs. You cannot cross this barrier by simple or fractional distillation. The mixture behaves as if it were a single pure substance.

Watch out

A common mistake is to think that because ethanol boils at 78.4 °C and water at 100 °C, you can keep concentrating ethanol indefinitely. The azeotrope at 78.2 °C is lower than ethanol’s own boiling point — that’s the trap. The mixture boils more easily than either pure component.

The general name: azeotropes

Binary mixtures that deviate from Raoult’s law so strongly that their vapour and liquid compositions become equal at some point are called azeotropes (from Greek: “to boil unchanged”). They cannot be separated into pure components by fractional distillation alone.

x_i = y_i \quad \text{for all components } i

Two types of azeotropes

There are exactly two types of azeotropic mixtures:

Minimum-boiling azeotropes — The mixture boils at a temperature lower than either pure component. Ethanol–water is the classic example. These arise from positive deviations from Raoult’s law (the A–B interactions are weaker than A–A and B–B interactions, so the mixture is “easier” to boil). …

Start your 14-day free trial to unlock the full solution →

Here is a breakdown of the common mistakes students make on this specific concept, along with the correct reasoning and exam-ready answers.

The Core Concept: Azeotropes

First, let's establish the correct answer to the question, as understanding this is the only way to avoid mistakes.

Why can't we get pure ethanol by fractional distillation?

Ethanol and water form a minimum-boiling azeotrope. At a specific composition (approximately 95.6% ethanol, 4.4% water by volume), the mixture boils at a constant temperature (78.15°C), which is lower than the boiling point of pure ethanol (78.37°C) and pure water (100°C).

When you distill a dilute ethanol solution, the vapor is richer in ethanol, but once the liquid in the flask reaches the azeotropic composition, the vapor has the same composition as the liquid. Further distillation yields only the azeotrope, not pure ethanol.
General name for such mixtures: Azeotropes (or azeotropic mixtures).
How many types? Two types:
1. Minimum-boiling azeotropes (positive deviation from Raoult's law) — e.g., ethanol-water.
2. Maximum-boiling azeotropes (negative deviation from Raoult's law) — e.g., nitric acid-water.

Common Mistakes & How to Avoid Them

Mistake 1: Saying "Pure ethanol cannot be obtained because it forms a constant boiling mixture" without explaining why it boils constantly.

Why it's wrong: This is a memorized phrase without understanding. Examiners want to see that you know the mechanism: at the azeotropic composition, the vapor and liquid phases have identical composition, so separation by simple or fractional distillation stops.
How to avoid: Always link the constant boiling point to equal composition of vapor and liquid. Write:

"At the azeotropic composition, the mole fraction of ethanol in the vapor equals that in the liquid. Hence, no further enrichment occurs."

Mistake 2: Confusing "fractional distillation" with "simple distillation" in the context of azeotropes.

Why it's wrong: Some students think that using a fractionating column (fractional distillation) can separate azeotropes. It cannot. Fractional distillation is more efficient for separating close-boiling liquids, but it fails when the vapor and liquid compositions become identical.
How to avoid: Remember: No distillation method (simple or fractional) can separate an azeotrope into its pure components. The only way is to break the azeotrope using a different technique (e.g., adding benzene to form a ternary azeotrope, or using molecular sieves).

Mistake 3: Forgetting to mention that ethanol-water is a minimum-boiling azeotrope.

Why it's wrong: The question asks "Why is it not possible?" — the answer must specify the type of deviation. If you just say "it shows deviation from Raoult's law," you lose marks.
How to avoid: Always classify the azeotrope:

"Ethanol-water shows positive deviation from Raoult's law, forming a minimum-boiling azeotrope."

Mistake 4: Stating that there are "three types" of azeotropes (e.g., positive, negative, and ideal).

Why it's wrong: There is no "ideal azeotrope." An ideal solution obeys Raoult's law exactly and can be separated by fractional distillation. Azeotropes are defined by their deviation from Raoult's law, which gives exactly two types: minimum-boiling (positive deviation) and maximum-boiling (negative deviation).
How to avoid: Memorize the two types with one example each:
- Minimum-boiling: Ethanol-water, isopropanol-water.
- Maximum-boiling: Nitric acid-water, hydrochloric acid-water (at constant pressure). …

Start your 14-day free trial to unlock the full solution →

Same / Similar Concept — real previous-year questions on the same or a closely similar concept, not this exact question.

Showing the 12 most recent of 22 on this concept.

AP EAPCET 2026Set eng-2026-05-12-FN1 markMCQ
$\mathrm{CH_3OH}$
›Reveal solutionSolution
Convert mole fraction of solute directly to molality using $m=\dfrac{x_2}{x_1}\cdot\dfrac{1000}{M_1}$ , taking $M_1=18\ \mathrm{g/mol}$ for water. Answer: 1.13 m.

Concept and Intuition

Molality ( $m$ = moles solute per kg solvent) and mole fraction ( $x_2$ = moles solute per total moles) are both intensive composition measures, so one converts to the other purely through mole-count and molar-mass arithmetic — no need to assume any solution volume or density, which is exactly why molality/mole-fraction problems are solvable without extra data (unlike molarity, which needs density).

Step-by-Step Solution

Given $x_2$ (mole fraction of $\mathrm{CH_3OH}$ ) $=0.02$ , so $x_1$ (mole fraction of water) $=1-0.02=0.98$ .

Take a basis of 1 mole total solution: moles of $\mathrm{CH_3OH}$ , $n_2=0.02$ ; moles of water, $n_1=0.98$ .

Mass of water (solvent) $= n_1 \times M_1 = 0.98 \times 18\ \mathrm{g/mol} = 17.64\ \mathrm{g} = 0.01764\ \mathrm{kg}$ . …
Start your 14-day free trial to see the full solution →
AP EAPCET 2026Set eng-2026-05-13-AN1 markMCQ
$g\ mL^{-1}$
›Reveal solutionSolution
Find the NaOH needed to neutralize 0.5 mol HCl (= 0.5 mol NaOH), back out the water present at $x_{NaOH}=0.02$ , and convert that water's mass to volume using its density — giving ≈441 mL.

Concept and Intuition

Mole fraction directly links the moles of solute to the moles of solvent present. Once we know exactly how many moles of NaOH must be present (fixed by the stoichiometric neutralization requirement), the mole-fraction relation pins down exactly how much water accompanies it — and hence, via water's known density, the volume of solution.

Step-by-Step Solution

Moles of HCl = $1\ L \times 0.5\ mol/L = 0.5$ mol.

$NaOH + HCl \rightarrow NaCl + H_2O$ is 1:1, so moles of NaOH required = 0.5 mol.

Mole fraction: $x_{NaOH} = \dfrac{n_{NaOH}}{n_{NaOH}+n_{H_2O}} = 0.02$ .

$\Rightarrow n_{NaOH}+n_{H_2O} = \dfrac{0.5}{0.02} = 25 \Rightarrow n_{H_2O} = 24.5$ mol.

Mass of water = $24.5 \times 18\ g/mol \approx 441$ g. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2026Set ap-2026-05-20-FN1 markMCQ
$Q.A solid solute is dissolved in water. The mole fraction of solute is 0.02. What is the molality of the solution? (A) 2.133 m (B) 2.5 m (C) 1.5 m (D) 1.133 m$
›Reveal solutionSolution
Convert mole fraction to molality by picking a convenient basis (1 mol total) and computing the solvent's mass. Answer: 1.133 m.

Concept and Intuition

Mole fraction and molality are related but different concentration scales: mole fraction is a ratio of moles, while molality is moles of solute per kilogram of solvent. To convert, assume a convenient total amount (1 mole of solution), find the moles of each component from the given mole fraction, convert the solvent's moles to a mass, and then compute molality directly.

Step-by-Step Solution

Given: mole fraction of solute $x_{solute}=0.02$ , so mole fraction of water $x_{water}=1-0.02=0.98$ .

Assume 1 mol of total solution: $n_{solute}=0.02$ mol, $n_{water}=0.98$ mol.

Mass of water $= n_{water}\times M_{water} = 0.98\times18 = 17.64$ g $= 0.01764$ kg. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2025Set eng-2025-05-21-FN1 markMCQ
$\mathrm{C_2H_6O_2}$
›Reveal solutionSolution
This is a direct molality calculation: moles of solute per kilogram of solvent.

Concept and Intuition

Molality is defined as moles of solute dissolved per kilogram of solvent (not solution), making it temperature-independent and ideal for colligative-property calculations like antifreeze formulations. Ethylene glycol's molar mass must first be computed from its formula to find the moles present.

Step-by-Step Solution

Molar mass of $C_2H_6O_2$ : $2(12)+6(1)+2(16) = 24+6+32 = 62\ \mathrm{g\,mol^{-1}}$ .

Moles of ethylene glycol $= \dfrac{248}{62} = 4$ mol.

Mass of water (solvent) $= 200\ \mathrm{g} = 0.200\ \mathrm{kg}$ . …
Start your 14-day free trial to see the full solution →
AP EAPCET 2025Set eng-2025-05-26-AN1 markMCQ
$^{-1}$
›Reveal solutionSolution
Molality uses moles of solute per kg of solvent (not solution). For 10% w/w glucose, 100 g solution = 10 g glucose + 90 g water, giving molality $\approx 0.62$ m.

Concept and Intuition

Molality is defined relative to the mass of solvent only, unlike mass percent (w/w) which is defined relative to total solution mass. So the first step is always to extract the solvent mass from the given composition before applying $m = \dfrac{n_{solute}}{\text{mass of solvent (kg)}}$ .

Step-by-Step Solution

Basis: 100 g of solution. 10% (w/w) glucose means 10 g glucose and $100-10=90$ g water.

Moles of glucose $= \dfrac{10\text{ g}}{180\text{ g mol}^{-1}} = 0.05\overline{5}$ mol.

Mass of solvent (water) in kg $= 90\text{ g} = 0.090$ kg. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2025Set eng-2025-05-27-FN1 markMCQ
$= 62\ \mathrm{g\ mol^{-1}}$
›Reveal solutionSolution
This is a direct molality calculation. 124 g of ethylene glycol is 2 mol; requiring a 10 m solution fixes the water mass at 200 g.

Concept and Intuition

Molality (m) is defined per kilogram of solvent, not solution — it is temperature-independent and depends only on moles of solute and mass of solvent.
$m = \frac{n_{solute}}{w_{solvent}(\text{kg})}$
Step-by-Step Solution

Moles of ethylene glycol $= \dfrac{124\ \text{g}}{62\ \text{g mol}^{-1}} = 2\ \text{mol}$ .

Let mass of water $= x$ g $= \dfrac{x}{1000}$ kg.

Given molality $= 10$ m: $10 = \frac{2}{x/1000}$ …
Start your 14-day free trial to see the full solution →
AP EAPCET 2024Set ap-2024-05-17-AN1 markMCQ
$\mathrm{dm^3}$
›Reveal solutionSolution
Converting molarity to molality requires the mass of solvent (not solution), obtained by subtracting the solute's mass from the total solution mass computed via density.

Concept and Intuition

Molarity (mol/L of solution) and molality (mol/kg of solvent) are only related once you know the density of the solution, because you need to convert the solution's volume to solution mass and then subtract the solute mass to get solvent mass.

Step-by-Step Solution

Molarity: $x = \dfrac{82\ \mathrm{g}}{60\ \mathrm{g/mol}} = 1.3\overline{6}$ mol per litre of solution (so moles of acetic acid per litre solution $= x$ ).

Mass of 1 L (1 dm³) of solution $= 1.01\ \mathrm{kg/dm^3} \times 1\ \mathrm{dm^3} = 1.01\ \mathrm{kg} = 1010$ g.

Mass of solvent (water) $= 1010 - 82 = 928$ g $= 0.928$ kg. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2023Set ap-2023-05-22-FN1 markMCQ
$\mathrm{Na_2CO_3}$
›Reveal solutionSolution
Direct molality calculation: moles of solute divided by kilograms of solvent gives 0.02 m.

Concept and Intuition

Molality is defined as moles of solute per kilogram of solvent (not solution), making it temperature-independent and convenient for colligative property calculations.

Step-by-Step Solution

Moles of $\mathrm{Na_2CO_3} = \dfrac{1.06\ \mathrm{g}}{106\ \mathrm{g\,mol^{-1}}} = 0.01$ mol.

Mass of solvent (water) = 500 g = 0.500 kg.

Molality $= \dfrac{0.01\ \mathrm{mol}}{0.500\ \mathrm{kg}} = 0.02$ mol/kg = 0.02 m. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2023Set eng-2023-05-15-FN1 markMCQ
$Q. The molarity of one molal glucose solution having density of 1.2 g/mL is (A) 0.101 M (B) 1.01 M (C) 2.01 M (D) 0.001 M$
›Reveal solutionSolution
Convert 1 molal glucose solution to molarity using the solution's density; molality and molarity differ because molarity is based on total solution volume, not solvent mass. Answer: 1.01 M.

Concept and Intuition

Molality (mol/kg solvent) and molarity (mol/L solution) are numerically close for dilute aqueous solutions but not identical, because molality ignores the solute's contribution to the total solution mass/volume. Density lets us convert between the two using $\text{Molarity} = \dfrac{1000 \times m \times d}{1000 + m\times M_w}$ , where $m$ is molality, $d$ is density (g/mL), and $M_w$ is the solute's molar mass.

Step-by-Step Solution

Take 1 kg (1000 g) of water as solvent, containing 1 mol glucose ( $M_w = 180\ \mathrm{g/mol}$ , so mass = 180 g).

Total solution mass $= 1000 + 180 = 1180$ g. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2023Set eng-2023-05-17-FN1 markMCQ
$\mathrm{H_2O}$
›Reveal solutionSolution
Molality is moles of solute per kilogram of solvent; with 0.1 mol glucose in 0.018 kg water, molality $=5.55\,\text{m}$ .

Concept and Intuition

Molality is defined using the mass of solvent (not volume, and not total solution mass), which is exactly why it is temperature-independent and preferred in colligative-property calculations. Here both the solute (glucose) and the solvent (water) are given in grams, so both must first be converted appropriately — glucose to moles, water to kilograms.

Step-by-Step Solution

Moles of glucose ( $M=180\,\text{g/mol}$ ): $n = \dfrac{18\,\text{g}}{180\,\text{g/mol}} = 0.1\,\text{mol}$ .

Mass of water in kg: $18\,\text{g} = 0.018\,\text{kg}$ . …
Start your 14-day free trial to see the full solution →
AP EAPCET 2022Set ap-2022-07-11-AN1 markMCQ
$^{-1}$
›Reveal solutionSolution
Converting 15 ppm (by mass) of a solute of molar mass 120 g/mol into molality gives $1.25\times10^{-4}$ mol kg $^{-1}$ .

Concept and Intuition

ppm (parts per million) by mass means grams of solute per million grams of solution; for a dilute aqueous solution, the solution mass is essentially the water mass. Molality is moles of solute per kilogram of solvent, so we convert the mass basis to moles and then to per-kg terms.

Step-by-Step Solution

15 ppm by mass $\Rightarrow$ 15 g of carcinogen per $10^6$ g ( $=1000$ kg) of water.

Moles of carcinogen $= \dfrac{15\text{ g}}{120\text{ g/mol}} = 0.125$ mol. …
Start your 14-day free trial to see the full solution →
AP EAPCET 2022Set ap-2022-07-11-FN1 markMCQ
$\mathrm{mL^{-1}}$
›Reveal solutionSolution
Converts a 10% w/w NaOH solution into molarity using its density and molar mass.

Concept and Intuition

Working with exactly 1 L (1000 mL) of solution makes the arithmetic simplest: the total mass of that litre comes from its density, 10% of that mass is NaOH, and dividing by NaOH's molar mass gives the number of moles present in that one litre — which is the molarity.

Step-by-Step Solution

Mass of 1 L (1000 mL) of solution $=1000\times1.11=1110$ g.

Mass of NaOH in it (10% w/w) $=0.10\times1110=111$ g.

Moles of NaOH $=\dfrac{111}{40}=2.775$ mol. …
Start your 14-day free trial to see the full solution →

←Previous questionQ57 — Explain the terms ideal and non-ideal solutions in the light of forces…→Next questionQ59 — When kept in water, raisin swells in size. Name and explain the phenom…

Q.Why is it not possible to obtain pure ethanol by fractional distillation? What general name is given to binary mixtures which show deviation from Raoult's law and whose components cannot be separated by fractional distillation. How many types of such mixtures are there?

Why fractional distillation fails for ethanol

The general name: azeotropes

Two types of azeotropes

🎓Unlock everything free for 14 days

Unlock everything free for 14 days