Exploring Mixtures and their Separation

Why This Matters

Look at a packet of sugar. It is white, sweet and made of neat little crystals. Now picture a tall green sugarcane plant in a field. How do we get those clean white crystals out of a juicy green plant? Somewhere between the field and your kitchen, a lot of separating happened.

Here is another one. A doctor takes just a few drops of your blood and can tell whether you have malaria or anaemia. How? They spin the blood very fast in a machine until it splits into layers. Each layer can then be tested.

Both of these are the same science — the science of taking a mixture apart. Salt from seawater, petrol and diesel from crude oil, clean water from dirty water, perfume from flower petals — all of it is separation.

You already learnt in earlier classes what a mixture is and a few simple ways to split one up. In this chapter we go deeper. We will sort mixtures into clear types, learn the exact tools used to separate each type, and — most importantly — understand why each method works. By the end you should be able to look at any mixture and say, “I know how to take that apart, and I know why that method is the right one.”

The Big Idea

The Big Idea: A mixture is two or more substances jumbled together, and they keep their own properties. Mixtures come in two big families. A homogeneous mixture (a solution) is the same all the way through — you cannot see the parts. A heterogeneous mixture is not the same throughout — you can spot the different parts. To separate a mixture, you find a property where the two parts differ — size, boiling point, density, solubility, or whether one part sticks to paper — and you use a method that takes advantage of that one difference.

Before we go further, let us refresh the words you will use on every page of this chapter.

How Do We Classify Mixtures?

Stir a spoon of sugar into water and taste it. The first sip and the last sip are equally sweet. The sugar has spread out evenly and you cannot see it any more. A mixture that is the same throughout like this is called a homogeneous mixture, or simply a solution. Vinegar (acetic acid in water) and soda (carbon dioxide in water) are other examples.

Now stir a spoon of sand into water. This is different. The sand does not vanish. You can see the grains floating about, and if you wait, they sink to the bottom. A mixture that is not the same throughout — where you can see the separate parts — is called a heterogeneous mixture.

Figure 5.1 puts the two side by side.

This relies on telling “uniform” from “not uniform”, so the picture below makes the contrast clear:

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Here is a quick test you can do in class — it is Activity 5.1 from your book. Take three glasses. In one, stir common salt into water. In the second, stir chalk powder into water. In the third, add a few drops of milk to water. Now ask three questions of each: Can I see the particles? If I shine a laser through it, does the beam show up? If I leave it alone, does anything settle? You will find all three glasses behave differently. That is the clue that these are three different kinds of mixture — solution, suspension and colloid. We will meet all three properly in this chapter.

The full family tree of matter is shown in Figure 5.2.

To see where each type fits, here is the whole classification at a glance:

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Solutions and Concentration

You know that a solution is made of a solute dissolved in a solvent. But how much solute? That amount matters a great deal.

Think about ORS — the salt-and-sugar drink doctors give when someone is sick and dehydrated. You must add the exact amounts of salt and sugar to a fixed amount of water. Too much or too little, and it is no longer proper ORS. It may even be harmful. The same is true when a farmer mixes pesticide with water. Too little pesticide and the crop is not protected. Too much and it damages the crop, the soil and the environment.

An Indian paediatrician, Dilip Mahalanabis, developed and spread the simple ORS treatment for dehydration from diseases like diarrhoea and cholera. The WHO took it worldwide. It has saved millions of lives. Getting the proportion right is exactly why ORS works.

So the amount of solute matters. The amount of solute dissolved in a given amount of solvent or solution is called the concentration of the solution. A concentrated solution has a lot of solute; a dilute one has little.

How do we express concentration?

A common and simple way to state concentration is as a percentage. There are three percentage methods. They differ only in whether we measure each part by mass (grams) or by volume (millilitres).

Here are all three side by side so you can see when each is used:

A small but important point: the mass of solution = mass of solute + mass of solvent. The solution includes both parts, not just the water.

Let us work through one of each so the formulas feel real.

We start with mass by mass — this is the one used on food packets:

Next, mass by volume — used for medicines, where we measure the liquid’s volume:

And finally volume by volume — used when two liquids are mixed:

A saline drip in hospitals is usually 0.9% m/v sodium chloride (common salt) in water. That means 0.9 g of salt in every 100 mL of solution. This exact concentration is safe to mix with blood and helps replace lost body fluids.

Solubility — how much can dissolve

There is a limit to how much solute a solvent can hold. Let us refresh that idea, because crystallization (coming up soon) depends on it.

A graph of solubility against temperature is called a solubility curve. If one substance’s curve climbs steeply, its solubility rises a lot as the water gets hotter. This will matter in a moment.

Separating Solutions (Homogeneous Mixtures)

The parts of a solution are mixed at the particle level, so they are the hardest to separate. We cannot just filter them. Instead we use clever tricks based on solubility and boiling point.

Crystallization

Imagine you make a saturated solution of a solid in hot water — packed to the limit. Now you cool it slowly. As the water cools, it can hold less solute (remember the recap). So the extra solute has nowhere to go but out of the liquid, and it comes out as a pure solid — often as beautiful, regular crystals. A crystal is a solid whose particles are arranged in a neat, repeating pattern.

This is crystallization: forming pure crystals from a saturated solution by cooling it. It is used to get a pure solid out of a solution, and to purify a solid by leaving its impurities behind in the liquid.

Figure 5.3 shows the three steps using copper sulfate (the blue salt).

Let us walk through how it is actually done in the lab, step by step:

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One more thing the graph teaches us. If a substance’s solubility drops a lot on cooling, a lot of solid will crystallize out. If it barely changes (like common salt, which dissolves almost the same hot or cold), then cooling gives you very little. That is why common salt is got from seawater by evaporating the water away, not by cooling.

Crystallization separates solids that came out of the same solvent. But what if we want to recover the liquid solvent too? Evaporation throws the water away as vapour. To save it, we need a different method.

Distillation

Suppose a student accidentally mixes acetone and water. These two are miscible — they mix completely, so there are no layers to separate. A separating funnel is useless here. How do we split them?

The trick is boiling point. Acetone boils at about 56 °C; water boils at 100 °C. That is a big gap. If we heat the mixture gently, the acetone boils off first, long before the water turns to vapour. We catch that acetone vapour, cool it back into liquid, and collect it — pure. This is distillation.

Distillation works when two miscible liquids differ in boiling point by at least about 25 °C. It can also pull a pure liquid out of a solution of a solid (the solid stays behind in the flask).

Figure 5.4 shows the apparatus. Follow the path of the liquid through it.

Here is the full set-up, labelled part by part:

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Why does the condenser have water flowing the “wrong” way — in at the bottom, out at the top? So the cold water always meets the hot vapour at the far end and stays coldest where it is most needed. This cools the vapour fully and turns all of it back to liquid.

What if two liquids have boiling points that are close together — less than 25 °C apart? Then simple distillation cannot cleanly separate them. We use fractional distillation, a taller version with extra surfaces inside that re-boil and re-cool the vapour many times. This is how a petroleum refinery splits crude oil into petrol, kerosene, diesel and LPG — each “fraction” boils off at a slightly different temperature.

Distillation is old and useful. In Kannauj, Uttar Pradesh — the “perfume capital of India” — the lovely earthy smell of the first rain is captured into a perfume called Mitti ka Ittar, using a traditional distillation method called Deg-Bhapka. The fragrance is distilled out of the soil and flowers.

Paper Chromatography

Drop water on something written with a black sketch pen. The black smudges out — and often you see hidden colours appear: blue, pink, yellow. That black ink was actually a mixture of dyes all along. Paper chromatography separates such a mixture of coloured substances.

Here is the idea. You put a small spot of the ink near the bottom of a paper strip and dip the very bottom into water (or another solvent). The water climbs up the paper and carries the dyes along with it. But each dye moves at its own speed — some cling more to the paper, some travel more easily with the water. So they spread out into separate bands. The dye that travels fastest ends up highest.

Figure 5.5 shows the three steps.

The figure below walks through marking the spot, dipping the strip, and the colours separating:

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The name says it all: chroma is Greek for “colour” and graphein means “to write” — so chromatography means “writing with colour”. You can use it to find how many dyes are in a food colour, or to separate the green and yellow pigments in spinach leaves. Sometimes plain water will not carry a dye, so chemists use alcohol or another solvent instead.

Separating Heterogeneous Mixtures

These are easier, because the parts are not mixed at the particle level — you can often see them. We pick whichever property the two parts differ in.

Two immiscible liquids — the separating funnel

Pour oil into water and they refuse to mix. They form two layers. Liquids that do not mix are called immiscible. To separate them we use a separating funnel — a pear-shaped glass funnel with a tap (stopcock) at the bottom.

We let the mixture settle into its two layers, then open the tap to drain off the lower layer carefully. When the lower liquid is almost gone, we close the tap, throw away the tiny in-between portion, and then drain the upper liquid separately.

But why do they form two layers, with oil always on top? Because of density. Let us refresh that.

Figure 5.6 shows the separating funnel in action.

This is the apparatus used to separate mustard oil from water:

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What about gases? Gas particles move freely in all directions, so most gas-gas mixtures are homogeneous — like the hydrogen-oxygen mix used as rocket fuel. But smoke (solid bits in air), fog (tiny water drops in air) and dust in the air are heterogeneous mixtures with a gas as one part.

Sublimation — for solids that turn straight to vapour

Take a mixture of camphor and sand. Camphor has a special trick: when you heat it gently, it turns directly from solid into vapour, skipping the liquid stage altogether. This is called sublimation. Sand does not do this. So if we heat the mixture and catch the vapour somewhere cool, the camphor leaves the sand behind and re-forms as solid on the cool surface. The vapour turning back into solid (without becoming liquid) is called deposition.

Figure 5.7 shows how this is set up.

Here is the labelled apparatus for separating camphor from sand:

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Other substances that sublime include naphthalene (mothballs) and dry ice (solid carbon dioxide), which is used to keep ice cream cold and turns straight into gas without making a wet puddle.

A note on alloys: metals do not dissolve into each other at room temperature. But melt them together at high heat and they mix into a single liquid that solidifies into a new material — an alloy. Brass is about 80% copper and 20% zinc; bronze is about 80% copper and 20% tin; stainless steel is iron with chromium, nickel and a little carbon. An alloy is a homogeneous mixture, but ordinary physical methods cannot separate it back into its metals.

Suspensions — and getting the solid out

Look again at sand in water (Figure 5.1b). No matter how hard you stir, the sand stays visible and eventually settles. A heterogeneous mixture where solid particles stay spread through a liquid without dissolving is called a suspension. The particles are large — bigger than in a solution — and visible to the naked eye. Tea leaves in water and sawdust in water are suspensions too.

Now, how do we clean muddy water? You can let it stand so the heavy mud settles, then filter it. But very fine particles often slip through the filter and leave the water cloudy. Filtration alone is not always enough. For the stubborn fine particles we use centrifugation or coagulation.

Centrifugation

Hold a friend’s hands and spin around fast. You feel pulled outwards. Centrifugation uses that same outward pull. A mixture is spun in a tube at very high speed. The spinning throws the heavier particles outward and down to the bottom of the tube, while the lighter liquid stays on top.

This is exactly how blood is separated, as Figure 5.8 shows.

The figure compares the blood sample before and after spinning:

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Centrifugation is used in labs to separate blood components and in many industries. Cleverly, a low-cost hand-spun device called a paperfuge, made from paper and string like a toy, can do the same job without electricity — helping detect malaria and anaemia in remote villages.

Coagulation

Some particles are too fine and too light to settle even with spinning. For these we add a substance that makes them clump together. Add powdered alum (fitkari) to muddy water. The alum makes the tiny scattered particles stick together into bigger, heavier clumps. This clumping is called coagulation, and the alum is the coagulant. The big clumps then sink by gravity (sedimentation) and can be removed by decantation (pouring off the clear water) or filtration.

You have seen coagulation at home: when paneer (cheese) is made, lemon juice or vinegar is added to milk. The acid coagulates the milk proteins into solid paneer.

Colloids

Here is a puzzle. Blood can be split by centrifuging, like a suspension. Yet you cannot see blood cells with the naked eye, like a solution. So what is blood? It is neither — it is a colloid. Milk, tomato sauce and ice cream are colloids too.

The difference between a solution, a colloid and a suspension is really about particle size:

• Solution: particles smaller than 1 nm. The smallest.
• Colloid: particles 1 to 1000 nm. Medium.
• Suspension: particles bigger than 1000 nm. The largest.

(A nanometre, nm, is one-millionth of a millimetre — incredibly tiny.) Like a solution, a colloid’s particles stay evenly spread and do not settle over time. But like a suspension, the particles are big enough to do something special with light, as we will see next.

Figure 5.9 lines up the three by particle size.

The size scale below shows exactly where each type sits:

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The Tyndall Effect

Go back to that three-glass experiment. When you shone a laser through the salt solution, the beam path was invisible inside it. But through the chalk suspension and the milk colloid, the beam lit up — you could see its path glowing from the side. The scattering of light by particles, which makes the beam visible, is called the Tyndall effect, named after the scientist John Tyndall.

Figure 5.10 shows the three glasses with the laser.

Compare the three beakers below — watch which beam path shows up:

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Now for the question NCERT does not fully answer: why does a colloid scatter light but a solution does not? It comes down to a contest between the size of the particle and the size of the light wave.

Light travels as a wave. In a solution, the particles are far, far smaller than the light wave. The wave simply sweeps past them, the way an ocean swell rolls straight over a tiny pebble without being disturbed. No scattering — so no visible beam.

In a colloid, the particles are about the same size as the light wave. Now the light cannot ignore them. It hits each particle and bounces off in all directions. Some of that bounced light reaches your eye from the side — and that glow is the beam you see. Figure 5.11 shows this contest.

The picture below shows exactly why the light behaves differently in each:

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You have seen the Tyndall effect outside the lab. The bright rays of sunlight slanting through gaps in tree leaves, the visible beam of light entering a dusty dark room, the cone of light above a stadium floodlight at night — all are light scattering off dust, smoke or water droplets in the air.

The two parts of a colloid have special names. The scattered particles (the solute-like part) are the dispersed phase, and the medium they float in is the dispersion medium. When both parts are liquids — like oil droplets spread through water — the colloid is called an emulsion. Milk and vanishing cream are oil-in-water emulsions; butter and cold cream are water-in-oil emulsions.

Here is the full comparison of all three mixture types — keep this table for revision:

Common Mistakes

Quick Check

Practice Problems

Easy

Medium

Challenge

Summary

You can now explain:

• What separates a homogeneous mixture (a solution, uniform throughout) from a heterogeneous one (not uniform), and where solutions, suspensions and colloids fit in the family tree of matter.
• How to express the concentration of a solution three ways — % m/m, % m/v and % v/v — and how to calculate each, remembering that mass of solution = solute + solvent.
• How solubility depends on temperature, and why that makes crystallization work — cooling a saturated solution forces pure crystals out.
• How distillation separates miscible liquids using their different boiling points, and how fractional distillation handles close boiling points (as in a petroleum refinery).
• How paper chromatography separates coloured substances by their different speeds up the paper, and how a separating funnel separates immiscible liquids by density.
• How sublimation separates a solid that turns straight to vapour, and how centrifugation and coagulation pull fine particles out of a suspension.
• What the Tyndall effect is, and why a colloid scatters light (its particles match the light wave’s size) while a true solution does not (its particles are far too small).

What’s Next

You have spent this chapter taking mixtures apart using physical differences like density, boiling point and particle size. Next you will switch from matter to motion. In Chapter 6 — How Forces Affect Motion you will learn what a force really does — how it makes things speed up, slow down, change direction or stop. The idea of density you met here (mass packed into volume) connects to mass, which is the very thing a force has to push around.