Shine a Torch Through Milk and Through Salt Water and Only One Glows
Learn to separate true solutions, colloids and suspensions by particle size and behaviour, tell saturated from unsaturated and supersaturated with a test, calculate solubility per hundred grams, and read a solubility curve.
Why does a torch beam show up in milk but not in salt water?
Fill one glass with salt dissolved in water and another with a little milk in water. Both look uniform. Now shine a torch through each in a dark room and look from the side.
The salt water shows nothing — the beam passes straight through and is invisible from the side.
The milk shows the beam as a visible shaft of light running through the liquid.
That difference is the Tyndall effect, and it is a direct report on particle size. Salt dissolves into particles far too small to interfere with light. The fat and protein particles in milk are large enough to scatter it sideways into your eye, and small enough to stay suspended indefinitely.
So three quite different things can look alike in a glass:
- a true solution, whose particles are too small to see or to scatter light
- a colloid, whose particles scatter light but never settle
- a suspension, whose particles are big enough to see and heavy enough to settle out
One torch separates the first from the other two, and standing the glass overnight separates the last.
This page covers the first part of the ICSE Class 9 Chemistry chapter on water — the three kinds of mixture, saturation, solubility and solubility curves.
The salt water shows nothing — the beam passes straight through and is invisible from the side.
The milk shows the beam as a visible shaft of light running through the liquid.
That difference is the Tyndall effect, and it is a direct report on particle size. Salt dissolves into particles far too small to interfere with light. The fat and protein particles in milk are large enough to scatter it sideways into your eye, and small enough to stay suspended indefinitely.
So three quite different things can look alike in a glass:
- a true solution, whose particles are too small to see or to scatter light
- a colloid, whose particles scatter light but never settle
- a suspension, whose particles are big enough to see and heavy enough to settle out
One torch separates the first from the other two, and standing the glass overnight separates the last.
This page covers the first part of the ICSE Class 9 Chemistry chapter on water — the three kinds of mixture, saturation, solubility and solubility curves.
How do you tell a true solution, a colloid and a suspension apart?
By particle size, and every other difference follows from it.
A true solution.
- Particle size **below about nanometre
- Homogeneous — the same throughout
- Transparent, and the particles cannot be seen even under a microscope
- Passes through filter paper unchanged
- Does not scatter light, so shows no Tyndall effect
- Stable — never settles on standing
- Examples: salt in water, sugar in water, copper sulphate solution, air, brass
A colloid.
- Particle size roughly between and nanometres
- Appears homogeneous to the eye but is really heterogeneous
- Translucent rather than clear
- Passes through filter paper but is held back by a parchment membrane or an ultrafilter
- Scatters light — shows the Tyndall effect
- Fairly stable — does not settle on standing
- Examples: milk, starch solution, gum, ink, soap solution, blood, jelly, fog, smoke, clouds
A suspension.
- Particle size above about nanometres
- Heterogeneous, and visibly so
- Opaque, and the particles can be seen with the naked eye or a hand lens
- Does not pass through filter paper — the particles are retained
- Scatters light
- Unstable — settles on standing
- Examples: chalk in water, muddy water, sand in water, flour in water, paint before stirring
The two practical tests, and what each separates.
- Filtration separates a suspension from the other two, because only a suspension is held back by filter paper
- The Tyndall effect separates a true solution from the other two, because only a true solution fails to scatter light
So neither test alone identifies all three** — you need both, and together they give a clean answer:
Worked classification. Muddy water is held back by filter paper, so it is a suspension, and standing it overnight leaves clear water with silt at the bottom. Milk passes the filter and scatters light, so it is a colloid, and it never separates however long it stands. Copper sulphate solution passes the filter and scatters nothing, so it is a true solution.
A colloid is not a weak suspension, and this is the point most often missed. A suspension settles because its particles are heavy enough for gravity to win; a colloid's particles are constantly jostled by the molecules of the liquid and never settle at all. So the difference is permanent rather than a matter of waiting longer — leaving milk to stand for a year does not produce a layer of fat at the bottom and clear water above.
A true solution.
- Particle size **below about nanometre
- Homogeneous — the same throughout
- Transparent, and the particles cannot be seen even under a microscope
- Passes through filter paper unchanged
- Does not scatter light, so shows no Tyndall effect
- Stable — never settles on standing
- Examples: salt in water, sugar in water, copper sulphate solution, air, brass
A colloid.
- Particle size roughly between and nanometres
- Appears homogeneous to the eye but is really heterogeneous
- Translucent rather than clear
- Passes through filter paper but is held back by a parchment membrane or an ultrafilter
- Scatters light — shows the Tyndall effect
- Fairly stable — does not settle on standing
- Examples: milk, starch solution, gum, ink, soap solution, blood, jelly, fog, smoke, clouds
A suspension.
- Particle size above about nanometres
- Heterogeneous, and visibly so
- Opaque, and the particles can be seen with the naked eye or a hand lens
- Does not pass through filter paper — the particles are retained
- Scatters light
- Unstable — settles on standing
- Examples: chalk in water, muddy water, sand in water, flour in water, paint before stirring
The two practical tests, and what each separates.
- Filtration separates a suspension from the other two, because only a suspension is held back by filter paper
- The Tyndall effect separates a true solution from the other two, because only a true solution fails to scatter light
So neither test alone identifies all three** — you need both, and together they give a clean answer:
Worked classification. Muddy water is held back by filter paper, so it is a suspension, and standing it overnight leaves clear water with silt at the bottom. Milk passes the filter and scatters light, so it is a colloid, and it never separates however long it stands. Copper sulphate solution passes the filter and scatters nothing, so it is a true solution.
A colloid is not a weak suspension, and this is the point most often missed. A suspension settles because its particles are heavy enough for gravity to win; a colloid's particles are constantly jostled by the molecules of the liquid and never settle at all. So the difference is permanent rather than a matter of waiting longer — leaving milk to stand for a year does not produce a layer of fat at the bottom and clear water above.
What is the difference between saturated, unsaturated and supersaturated?
An unsaturated solution can dissolve more; a saturated one cannot; a supersaturated one is holding more than it should.
An unsaturated solution contains less solute than it is capable of dissolving at that temperature. Add more and it dissolves.
A saturated solution contains the maximum amount of solute it can dissolve at that temperature. Any further solute added stays undissolved at the bottom, in equilibrium with the dissolved solute.
A supersaturated solution contains more dissolved solute than a saturated solution would hold at that temperature. It is made by saturating the solution at a higher temperature and then cooling it slowly and without disturbance. It is unstable — drop in a single crystal or give it a jolt, and the excess solute crystallises out at once.
The test for saturation. Add a little more of the same solute at the same temperature and stir.
- If it dissolves, the solution was unsaturated
- If it settles at the bottom undissolved, the solution was saturated
An equivalent test: drop a crystal of the solute in. In a saturated solution it neither grows nor dissolves; in an unsaturated one it dissolves; in a supersaturated one it grows rapidly as the excess comes out.
Worked observation. Stir sugar into a glass of water at room temperature, spoon by spoon. For a while every spoonful disappears — the solution is unsaturated. Eventually a spoonful sits on the bottom and will not go, however long you stir: the solution is now saturated. Warm the glass and that residue dissolves too, because the solubility has risen.
Now let the warm solution cool very slowly and undisturbed, and it becomes supersaturated. Drop in one sugar crystal and a shower of crystals forms through the liquid.
Saturation is defined at a stated temperature and nowhere else. A solution saturated at is very far from saturated at — it is supersaturated there. So "this solution is saturated" is an incomplete statement, and every saturation question must name the temperature.
And a saturated solution is not a concentrated one. Some substances are so sparingly soluble that a saturated solution of them contains very little solute indeed, while a sugar solution can be concentrated and still unsaturated. "Saturated" says the solution has reached its limit; "concentrated" says how much is in it — and those are two different claims, which is why the next section needs a number.
An unsaturated solution contains less solute than it is capable of dissolving at that temperature. Add more and it dissolves.
A saturated solution contains the maximum amount of solute it can dissolve at that temperature. Any further solute added stays undissolved at the bottom, in equilibrium with the dissolved solute.
A supersaturated solution contains more dissolved solute than a saturated solution would hold at that temperature. It is made by saturating the solution at a higher temperature and then cooling it slowly and without disturbance. It is unstable — drop in a single crystal or give it a jolt, and the excess solute crystallises out at once.
The test for saturation. Add a little more of the same solute at the same temperature and stir.
- If it dissolves, the solution was unsaturated
- If it settles at the bottom undissolved, the solution was saturated
An equivalent test: drop a crystal of the solute in. In a saturated solution it neither grows nor dissolves; in an unsaturated one it dissolves; in a supersaturated one it grows rapidly as the excess comes out.
Worked observation. Stir sugar into a glass of water at room temperature, spoon by spoon. For a while every spoonful disappears — the solution is unsaturated. Eventually a spoonful sits on the bottom and will not go, however long you stir: the solution is now saturated. Warm the glass and that residue dissolves too, because the solubility has risen.
Now let the warm solution cool very slowly and undisturbed, and it becomes supersaturated. Drop in one sugar crystal and a shower of crystals forms through the liquid.
Saturation is defined at a stated temperature and nowhere else. A solution saturated at is very far from saturated at — it is supersaturated there. So "this solution is saturated" is an incomplete statement, and every saturation question must name the temperature.
And a saturated solution is not a concentrated one. Some substances are so sparingly soluble that a saturated solution of them contains very little solute indeed, while a sugar solution can be concentrated and still unsaturated. "Saturated" says the solution has reached its limit; "concentrated" says how much is in it — and those are two different claims, which is why the next section needs a number.
Formula
How do you calculate the solubility of a solute?
**Solubility is the mass of solute in grams that saturates grams of solvent at a stated temperature.**
The temperature must always be quoted with the answer, and the unit is **grams per grams of solvent.
Worked example 1.** g of a salt just saturates g of water at :
Worked example 2. g saturates g of water:
Worked example 3. g saturates g of water:
Worked example 4 — working forwards from a known solubility. A salt has a solubility of g per g of water at . How much dissolves in g of water?
Worked example 5 — when the SOLUTION mass is given. A saturated solution weighing g is found to contain g of solute. The mass of solvent is therefore
and the solubility is
The denominator is the SOLVENT, not the solution. Dividing by would give g, which is the percentage by mass of the solution and not the solubility. That substitution is the commonest error in this topic, and the fix is to subtract the solute from the solution mass first.
Worked example 6 — how much crystallises out on cooling. A saturated solution is made with g of water at , where the solubility is g per g. It is cooled to , where the solubility is g per g. How much solute crystallises?
The water can now hold only g, so the excess is
That calculation is the whole principle of purification by crystallisation — dissolve hot, cool, and collect what the cooler solvent can no longer hold.
Worked example 7 — a larger quantity of solvent. The same pair of temperatures with g of water:
Solubility is a property of a pair, not of a substance. Salt is highly soluble in water and hardly soluble in petrol; iodine is the other way round. So a solubility figure must name the solvent as well as the temperature, and "the solubility of sugar" is an incomplete quantity until both are stated.
The temperature must always be quoted with the answer, and the unit is **grams per grams of solvent.
Worked example 1.** g of a salt just saturates g of water at :
Worked example 2. g saturates g of water:
Worked example 3. g saturates g of water:
Worked example 4 — working forwards from a known solubility. A salt has a solubility of g per g of water at . How much dissolves in g of water?
Worked example 5 — when the SOLUTION mass is given. A saturated solution weighing g is found to contain g of solute. The mass of solvent is therefore
and the solubility is
The denominator is the SOLVENT, not the solution. Dividing by would give g, which is the percentage by mass of the solution and not the solubility. That substitution is the commonest error in this topic, and the fix is to subtract the solute from the solution mass first.
Worked example 6 — how much crystallises out on cooling. A saturated solution is made with g of water at , where the solubility is g per g. It is cooled to , where the solubility is g per g. How much solute crystallises?
The water can now hold only g, so the excess is
That calculation is the whole principle of purification by crystallisation — dissolve hot, cool, and collect what the cooler solvent can no longer hold.
Worked example 7 — a larger quantity of solvent. The same pair of temperatures with g of water:
Solubility is a property of a pair, not of a substance. Salt is highly soluble in water and hardly soluble in petrol; iodine is the other way round. So a solubility figure must name the solvent as well as the temperature, and "the solubility of sugar" is an incomplete quantity until both are stated.
What do the solubility curves of three salts tell you?
A solubility curve plots solubility against temperature, and its slope says how strongly temperature matters for that substance.
**Potassium nitrate, — a steeply rising curve. Its solubility increases sharply with temperature, so a great deal more dissolves hot than cold.
Consequence: potassium nitrate is excellent for purification by crystallisation. Dissolve the impure solid in hot water, filter off any insoluble dirt, and cool — a large mass of pure crystals comes out, because the cold water can hold so much less. The soluble impurities stay behind in the liquid.
Sodium chloride, — a nearly flat curve. Its solubility rises only slightly with temperature.
Consequence: cooling a saturated salt solution yields almost nothing, so common salt is obtained by evaporating sea water rather than by cooling it. The flatness of the curve is why salt pans are shallow and open to the sun, and not cooling tanks.
Calcium sulphate, — a falling curve. Its solubility decreases as the temperature rises, which is the exception among common salts.
Consequence: calcium sulphate comes out of solution when water is heated, and it deposits as hard scale on the inside of kettles, boilers and hot-water pipes. The furring of a kettle is a solubility curve sloping the wrong way, and it is the reason boiler feed water has to be softened.
The general rule and its exceptions.
- For most solids the solubility increases with temperature, because dissolving them absorbs heat
- For a few, including calcium sulphate and calcium hydroxide, it decreases
Worked reading of a curve.** Suppose a curve shows a solubility of g per g at and g per g at . Then cooling a solution saturated at down to deposits
So the vertical gap between two points on a curve is the yield of a crystallisation — which is how a curve is used rather than merely read.
Gases behave in the opposite way, and on two counts.
- The solubility of a gas decreases as the temperature rises
- The solubility of a gas increases as the pressure rises
That is why a fizzy drink loses its gas when it warms up, why bubbles appear in a pan of water long before it boils, and why aquatic life suffers in water warmed by industry — warm water holds less dissolved oxygen. And it is why a sealed bottle stays fizzy and hisses when opened: releasing the pressure reduces the solubility at once.
A rising curve is not a law of nature. The direction depends on whether dissolving that particular substance takes heat in or gives it out, so the curves of two salts in the same beaker can slope opposite ways. Assuming "hot water dissolves more" is right for most salts and wrong for calcium sulphate, and the exception is precisely the one that fills a kettle with scale.
**Potassium nitrate, — a steeply rising curve. Its solubility increases sharply with temperature, so a great deal more dissolves hot than cold.
Consequence: potassium nitrate is excellent for purification by crystallisation. Dissolve the impure solid in hot water, filter off any insoluble dirt, and cool — a large mass of pure crystals comes out, because the cold water can hold so much less. The soluble impurities stay behind in the liquid.
Sodium chloride, — a nearly flat curve. Its solubility rises only slightly with temperature.
Consequence: cooling a saturated salt solution yields almost nothing, so common salt is obtained by evaporating sea water rather than by cooling it. The flatness of the curve is why salt pans are shallow and open to the sun, and not cooling tanks.
Calcium sulphate, — a falling curve. Its solubility decreases as the temperature rises, which is the exception among common salts.
Consequence: calcium sulphate comes out of solution when water is heated, and it deposits as hard scale on the inside of kettles, boilers and hot-water pipes. The furring of a kettle is a solubility curve sloping the wrong way, and it is the reason boiler feed water has to be softened.
The general rule and its exceptions.
- For most solids the solubility increases with temperature, because dissolving them absorbs heat
- For a few, including calcium sulphate and calcium hydroxide, it decreases
Worked reading of a curve.** Suppose a curve shows a solubility of g per g at and g per g at . Then cooling a solution saturated at down to deposits
So the vertical gap between two points on a curve is the yield of a crystallisation — which is how a curve is used rather than merely read.
Gases behave in the opposite way, and on two counts.
- The solubility of a gas decreases as the temperature rises
- The solubility of a gas increases as the pressure rises
That is why a fizzy drink loses its gas when it warms up, why bubbles appear in a pan of water long before it boils, and why aquatic life suffers in water warmed by industry — warm water holds less dissolved oxygen. And it is why a sealed bottle stays fizzy and hisses when opened: releasing the pressure reduces the solubility at once.
A rising curve is not a law of nature. The direction depends on whether dissolving that particular substance takes heat in or gives it out, so the curves of two salts in the same beaker can slope opposite ways. Assuming "hot water dissolves more" is right for most salts and wrong for calcium sulphate, and the exception is precisely the one that fills a kettle with scale.
Exam tip
Exam tip: divide by the solvent, and always state the temperature
Solubility divides by the SOLVENT, not the solution. Given a solution of g containing g of solute, first find the solvent as g.
Write the unit in full — *grams per grams of water* — and state the temperature with every answer.
Name the solvent too. Solubility is a property of a solute-and-solvent pair.
For a crystallisation yield, subtract the two solubilities: g per g of water.
**Scale the answer if the water is not g**: with g of water the yield becomes g.
Two tests, two jobs. Filtration picks out the suspension; the Tyndall effect picks out the true solution. Neither alone identifies all three.
Quote the size bands — below nm, between and nm, above nm — and let every other property follow from them.
A colloid never settles, and that is permanent rather than slow. A suspension settles.
For the saturation test, add more of the same solute at the same temperature and say what each outcome means.
Saturated is not concentrated — one states a limit, the other an amount.
And learn the three curves: potassium nitrate rises steeply, sodium chloride is nearly flat, calcium sulphate falls — with a consequence for each.
Write the unit in full — *grams per grams of water* — and state the temperature with every answer.
Name the solvent too. Solubility is a property of a solute-and-solvent pair.
For a crystallisation yield, subtract the two solubilities: g per g of water.
**Scale the answer if the water is not g**: with g of water the yield becomes g.
Two tests, two jobs. Filtration picks out the suspension; the Tyndall effect picks out the true solution. Neither alone identifies all three.
Quote the size bands — below nm, between and nm, above nm — and let every other property follow from them.
A colloid never settles, and that is permanent rather than slow. A suspension settles.
For the saturation test, add more of the same solute at the same temperature and say what each outcome means.
Saturated is not concentrated — one states a limit, the other an amount.
And learn the three curves: potassium nitrate rises steeply, sodium chloride is nearly flat, calcium sulphate falls — with a consequence for each.
Did you know
Why a kettle furs up and a salt pan does not
Boil water in a kettle day after day and a hard white crust builds up inside it. Evaporate sea water in a shallow pan and you collect salt. Both involve heating water and both leave a solid behind — and the two processes are chemically opposite.
The kettle's crust is largely calcium sulphate and calcium carbonate, and it forms because calcium sulphate becomes less soluble as the water gets hotter. So heating the water forces the calcium sulphate out of solution, and it deposits precisely where the water is hottest — on the heating element and the base.
That is a falling solubility curve doing damage. Nothing has evaporated; the water has simply become a worse solvent as it warmed.
The salt pan is the ordinary case. Sodium chloride's curve is nearly flat, so heating the sea water dissolves scarcely any more salt and cooling it would deposit scarcely any. The only way to get the salt out is to remove the water, which is what the sun does. Evaporation, not cooling.
And potassium nitrate is the third possibility. Its curve climbs so steeply that cooling alone is enough — dissolve it hot, let it cool, and crystals pour out with no evaporation needed at all.
So three salts and three methods, each dictated by the shape of one curve:
- steeply rising — purify by cooling a hot saturated solution
- nearly flat — recover only by evaporating the solvent
- falling — the substance comes out when heated, which is a nuisance rather than a method
Nobody chose those methods by preference. Each is the only one the curve allows, and that is why a solubility curve is a working document for anyone purifying a solid rather than a graph to be memorised.
The kettle's crust is largely calcium sulphate and calcium carbonate, and it forms because calcium sulphate becomes less soluble as the water gets hotter. So heating the water forces the calcium sulphate out of solution, and it deposits precisely where the water is hottest — on the heating element and the base.
That is a falling solubility curve doing damage. Nothing has evaporated; the water has simply become a worse solvent as it warmed.
The salt pan is the ordinary case. Sodium chloride's curve is nearly flat, so heating the sea water dissolves scarcely any more salt and cooling it would deposit scarcely any. The only way to get the salt out is to remove the water, which is what the sun does. Evaporation, not cooling.
And potassium nitrate is the third possibility. Its curve climbs so steeply that cooling alone is enough — dissolve it hot, let it cool, and crystals pour out with no evaporation needed at all.
So three salts and three methods, each dictated by the shape of one curve:
- steeply rising — purify by cooling a hot saturated solution
- nearly flat — recover only by evaporating the solvent
- falling — the substance comes out when heated, which is a nuisance rather than a method
Nobody chose those methods by preference. Each is the only one the curve allows, and that is why a solubility curve is a working document for anyone purifying a solid rather than a graph to be memorised.
Exam relevance
How do solutions and solubility feed into JEE Main and NEET?
Because solubility becomes a quantitative equilibrium, and the colloid material is a chapter in its own right.
This is the foundation for Class 12 Chemistry Solutions and Surface Chemistry, examined in JEE Main and NEET. The solubility calculated per g of water here is replaced by molarity, molality and mole fraction, and questions convert between them — so the habit of dividing by the solvent rather than the solution is exactly what molality requires, since molality is moles of solute per kilogram of solvent while molarity uses the volume of solution. Confusing those two is the Class 12 version of the trap on this page.
The temperature effect becomes a law. Class 12 states that the solubility of a gas falls with temperature and rises with pressure, and gives Henry's law with a constant for each gas. The fizzy-drink and dissolved-oxygen observations here are that law applied, and numericals on it are recurring JEE Main material. The exception of calcium sulphate is explained by the sign of the enthalpy of solution, using Le Chatelier's principle from Class 11 Equilibrium — an endothermic dissolving is helped by heating and an exothermic one is hindered.
Saturation becomes the solubility product. Class 11 Equilibrium treats a saturated solution of a sparingly soluble salt as an equilibrium between the solid and its ions, with a solubility product . The undissolved solute sitting in equilibrium with the dissolved solute, described on this page, is that equilibrium, and the precipitation reactions of the previous chapter are predicted from it.
Colloids get a whole chapter. Class 12 Surface Chemistry classifies colloids, explains the Tyndall effect by scattering, treats Brownian motion — the jostling that keeps colloidal particles from settling, which is why milk never separates — and covers coagulation, emulsions and micelles. The three-way classification learnt here is the entry point, and assertion-reason questions on the Tyndall effect appear in both papers.
Crystallisation as purification reappears in Class 11 Organic Chemistry as a standard purification technique, chosen exactly when the solubility curve is steep.
For NEET Biology, colloids matter directly: cytoplasm is a colloidal system, blood is a colloid, and the behaviour of colloidal particles underlies osmosis, dialysis and the working of the kidney. Dialysis in Class 12 Chemistry and the nephron in NEET Biology use the same membrane principle flagged here — a parchment membrane holds back colloidal particles while letting a true solution through.
What the questions look like. For board work, expect distinguish the three mixtures in a table, classify given examples, define saturated, unsaturated and supersaturated, describe the saturation test, calculate solubility including from a solution mass, and interpret the three named curves with a consequence. For JEE Main and NEET, expect concentration-unit conversions, Henry's law, solubility product, and colloid classification.
How board and competitive emphasis differ. A board paper rewards the table of differences with particle sizes and the stated temperature on every solubility answer. A competitive paper assumes all of it and tests the arithmetic of molality against molarity, or a calculation.
The single trap that costs the most marks. Dividing by the mass of the solution. A g saturated solution containing g of solute has a solubility of g per g of water, and dividing by gives — which is a correct percentage by mass and a wrong solubility. The defence is to write the subtraction as its own line, g of solvent, before the formula is touched.
This is the foundation for Class 12 Chemistry Solutions and Surface Chemistry, examined in JEE Main and NEET. The solubility calculated per g of water here is replaced by molarity, molality and mole fraction, and questions convert between them — so the habit of dividing by the solvent rather than the solution is exactly what molality requires, since molality is moles of solute per kilogram of solvent while molarity uses the volume of solution. Confusing those two is the Class 12 version of the trap on this page.
The temperature effect becomes a law. Class 12 states that the solubility of a gas falls with temperature and rises with pressure, and gives Henry's law with a constant for each gas. The fizzy-drink and dissolved-oxygen observations here are that law applied, and numericals on it are recurring JEE Main material. The exception of calcium sulphate is explained by the sign of the enthalpy of solution, using Le Chatelier's principle from Class 11 Equilibrium — an endothermic dissolving is helped by heating and an exothermic one is hindered.
Saturation becomes the solubility product. Class 11 Equilibrium treats a saturated solution of a sparingly soluble salt as an equilibrium between the solid and its ions, with a solubility product . The undissolved solute sitting in equilibrium with the dissolved solute, described on this page, is that equilibrium, and the precipitation reactions of the previous chapter are predicted from it.
Colloids get a whole chapter. Class 12 Surface Chemistry classifies colloids, explains the Tyndall effect by scattering, treats Brownian motion — the jostling that keeps colloidal particles from settling, which is why milk never separates — and covers coagulation, emulsions and micelles. The three-way classification learnt here is the entry point, and assertion-reason questions on the Tyndall effect appear in both papers.
Crystallisation as purification reappears in Class 11 Organic Chemistry as a standard purification technique, chosen exactly when the solubility curve is steep.
For NEET Biology, colloids matter directly: cytoplasm is a colloidal system, blood is a colloid, and the behaviour of colloidal particles underlies osmosis, dialysis and the working of the kidney. Dialysis in Class 12 Chemistry and the nephron in NEET Biology use the same membrane principle flagged here — a parchment membrane holds back colloidal particles while letting a true solution through.
What the questions look like. For board work, expect distinguish the three mixtures in a table, classify given examples, define saturated, unsaturated and supersaturated, describe the saturation test, calculate solubility including from a solution mass, and interpret the three named curves with a consequence. For JEE Main and NEET, expect concentration-unit conversions, Henry's law, solubility product, and colloid classification.
How board and competitive emphasis differ. A board paper rewards the table of differences with particle sizes and the stated temperature on every solubility answer. A competitive paper assumes all of it and tests the arithmetic of molality against molarity, or a calculation.
The single trap that costs the most marks. Dividing by the mass of the solution. A g saturated solution containing g of solute has a solubility of g per g of water, and dividing by gives — which is a correct percentage by mass and a wrong solubility. The defence is to write the subtraction as its own line, g of solvent, before the formula is touched.
Key takeaways
Solutions, saturation, solubility and its curves: quick revision
- True solution: particles below about ** nm, homogeneous, transparent, passes filter paper, no Tyndall effect, never settles. Salt or sugar in water.
- Colloid: particles about to nm, looks homogeneous but is not, translucent, passes filter paper but held by parchment, shows the Tyndall effect, never settles. Milk, ink, starch, blood, fog, smoke.
- Suspension: particles above about nm, heterogeneous, opaque, held back by filter paper, scatters light, settles on standing. Chalk or mud in water.
- Filtration picks out the suspension; the Tyndall effect picks out the true solution — neither test alone does all three.
- A colloid never settles at all — that is Brownian jostling, not slow settling.
- Unsaturated can dissolve more; saturated holds the maximum at that temperature, with any extra staying undissolved in equilibrium; supersaturated holds more than it should, made by cooling a hot saturated solution slowly and undisturbed, and it crystallises at a jolt or a seed crystal.
- Test for saturation: add more of the same solute at the same temperature — dissolving means unsaturated, settling means saturated. A crystal added neither grows nor dissolves in a saturated solution.
- Saturation is always at a stated temperature, and saturated is not the same as concentrated.
- Solubility **, in grams per g of solvent at a stated temperature.
- g in g gives ** g per g**; g in g gives ****; g in g gives ****.
- A solubility of g per g means g dissolves in g of water.
- From a solution mass, subtract first: a g solution with g of solute has g of solvent, so the solubility is g per g — not .
- Crystallisation yield is the difference of two solubilities: g per g of water, or g with g of water.
- Solubility names a PAIR — a solute and a solvent — as well as a temperature.
- Potassium nitrate: curve rises steeply, so it purifies beautifully by cooling a hot saturated solution.
- Sodium chloride: curve nearly flat, so salt is recovered by evaporating sea water, not by cooling it.
- Calcium sulphate: curve falls, so it comes out when water is heated and deposits as scale in kettles and boilers.
- Most solids dissolve better hot because dissolving absorbs heat; calcium sulphate and calcium hydroxide are the exceptions.
- The vertical gap between two points on a curve is a crystallisation yield — g per g of water.
- Gases are opposite on both counts: solubility falls with temperature and rises with pressure — hence flat warm drinks, bubbles before boiling, and less oxygen in warm water for fish.
Shine a phone torch sideways through a glass of milk and then through a glass of sugar water, and see which one shows the beam — you have just done the Tyndall test.
- Colloid: particles about to nm, looks homogeneous but is not, translucent, passes filter paper but held by parchment, shows the Tyndall effect, never settles. Milk, ink, starch, blood, fog, smoke.
- Suspension: particles above about nm, heterogeneous, opaque, held back by filter paper, scatters light, settles on standing. Chalk or mud in water.
- Filtration picks out the suspension; the Tyndall effect picks out the true solution — neither test alone does all three.
- A colloid never settles at all — that is Brownian jostling, not slow settling.
- Unsaturated can dissolve more; saturated holds the maximum at that temperature, with any extra staying undissolved in equilibrium; supersaturated holds more than it should, made by cooling a hot saturated solution slowly and undisturbed, and it crystallises at a jolt or a seed crystal.
- Test for saturation: add more of the same solute at the same temperature — dissolving means unsaturated, settling means saturated. A crystal added neither grows nor dissolves in a saturated solution.
- Saturation is always at a stated temperature, and saturated is not the same as concentrated.
- Solubility **, in grams per g of solvent at a stated temperature.
- g in g gives ** g per g**; g in g gives ****; g in g gives ****.
- A solubility of g per g means g dissolves in g of water.
- From a solution mass, subtract first: a g solution with g of solute has g of solvent, so the solubility is g per g — not .
- Crystallisation yield is the difference of two solubilities: g per g of water, or g with g of water.
- Solubility names a PAIR — a solute and a solvent — as well as a temperature.
- Potassium nitrate: curve rises steeply, so it purifies beautifully by cooling a hot saturated solution.
- Sodium chloride: curve nearly flat, so salt is recovered by evaporating sea water, not by cooling it.
- Calcium sulphate: curve falls, so it comes out when water is heated and deposits as scale in kettles and boilers.
- Most solids dissolve better hot because dissolving absorbs heat; calcium sulphate and calcium hydroxide are the exceptions.
- The vertical gap between two points on a curve is a crystallisation yield — g per g of water.
- Gases are opposite on both counts: solubility falls with temperature and rises with pressure — hence flat warm drinks, bubbles before boiling, and less oxygen in warm water for fish.
Shine a phone torch sideways through a glass of milk and then through a glass of sugar water, and see which one shows the beam — you have just done the Tyndall test.