Solubility of a solid in a liquid
Make chashni for gulab jamun and you discover the effect of temperature on solubility without trying.
Sugar stirred into cold water soon stops dissolving and a heap sits at the bottom of the pan.
Heat the pan and the heap disappears; far more sugar dissolves in hot water.
Now try to wash a greasy kadhai with plain water: the grease does not budge, though a little kerosene lifts it at once.
Two questions are hiding in the kitchen.
Why does a substance dissolve in one liquid and not in another?
And why does temperature matter so much?
Solubility is the most that can dissolve: the maximum amount of a substance that can be dissolved in a stated amount of solvent at a stated temperature.
It depends on the nature of the solute and the solvent, on the temperature, and (for gases) on the pressure.
This section deals with solids dissolving in liquids; the next deals with gases.
Like dissolves like
Sodium chloride and sugar dissolve readily in water, but hardly at all in benzene.
Naphthalene (the white mothball) and anthracene do the opposite: they dissolve readily in benzene and hardly at all in water.
The pattern is summed up as like dissolves like: polar and ionic solutes dissolve in polar solvents, and non-polar solutes dissolve in non-polar solvents.
What the particles are doing. To dissolve a solid, three things must happen.
The solute particles must be pulled apart (this costs energy).
Some solvent particles must be pushed apart to make room (this also costs energy).
Then solute and solvent particles attract each other (this releases energy).
A solute dissolves well when the new solute-solvent attractions are similar in kind and strength to the attractions that were broken.
- In water, each Na⁺ ion is surrounded by the δ⁻ oxygen ends of water molecules and each Cl⁻ ion by the δ⁺ hydrogen ends. These strong ion-dipole attractions repay the energy needed to break up the crystal. Sugar molecules, full of OH groups, form hydrogen bonds with water just as water does with itself.
- In benzene, nothing can attract the ions strongly enough. The crystal's strong ionic bonds would be broken for almost no reward, so NaCl stays undissolved.
- Naphthalene molecules hold each other only by weak dispersion (London) forces, and so do benzene molecules. Mixing them swaps weak forces for similar weak forces, so naphthalene dissolves freely. In water it would have to break the strong hydrogen-bond network of water and could not replace it, so it stays out.
The same idea explains everyday choices.
Kerosene removes grease because both are non-polar hydrocarbons.
Iodine, made of non-polar I₂ molecules, gives a violet solution in carbon tetrachloride but dissolves only slightly in water.
Methanol and ethanol mix with water in all proportions because their OH groups hydrogen-bond with water.
You can see "like dissolves like" in a test tube.
Shake a crystal of iodine with water and carbon tetrachloride: the two liquids do not mix, the denser carbon tetrachloride settles at the bottom, and almost all the iodine moves into it, colouring it violet, while the water above stays nearly colourless.
Dissolving is a dynamic equilibrium
Add sugar a spoonful at a time to water at room temperature.
At first every spoonful dissolves; this forward process is dissolution.
But dissolved sugar molecules are moving about, and some of them strike the undissolved crystals and stick, coming out of solution again; this reverse process is crystallisation.
As the solution gets stronger, crystallisation speeds up.
Eventually the two rates become equal: sugar molecules leave the crystals exactly as fast as they return.
The amount dissolved stops changing although both processes continue.
This is a dynamic equilibrium:
Solute + Solvent ⇌ Solution
A solution in this state, in contact with undissolved solute, is saturated: it holds the maximum amount of solute that can dissolve at that temperature and pressure, so its concentration is the solubility.
A solution that could still take more solute at the same temperature is unsaturated.
A simple test tells you which you have.
Drop in one more crystal of the solute.
If it dissolves, the solution was unsaturated.
If it stays as it is, the solution was saturated.
The effect of temperature: Le Chatelier's principle
Because a saturated solution is an equilibrium, it obeys Le Chatelier's principle from Class XI: if a condition is changed, the equilibrium shifts in the direction that opposes the change.
The direction depends on whether dissolving takes in or gives out heat.
- Endothermic dissolution (ΔsolH > 0). Heat is taken in as the solid dissolves, so you can treat heat as a reactant: Solute + Solvent + heat ⇌ Solution. Raising the temperature supplies heat, the equilibrium shifts to the right, and solubility increases with temperature. Potassium nitrate, ammonium chloride, ammonium nitrate and sugar behave this way. Ammonium nitrate takes in so much heat as it dissolves that it is used in instant cold packs for sports injuries.
- Exothermic dissolution (ΔsolH < 0). Heat is given out: Solute + Solvent ⇌ Solution + heat. Raising the temperature shifts the equilibrium to the left, so solubility decreases with temperature. Cerium(III) sulphate and calcium hydroxide are examples.
- Small ΔsolH. Sodium chloride takes in only a little heat as it dissolves, so its solubility barely changes with temperature.
The solubility curves below show the contrast.
The steep curve of potassium nitrate is why it is purified by crystallisation: a hot saturated solution, cooled, throws out a large crop of crystals.
The flat curve of sodium chloride is why salt in the pans along the Gujarat coast is obtained by evaporating the water, not by cooling the brine.
The photograph shows salt pans of the kind found along the Gujarat coast.
Sea water or underground brine is let into shallow ponds and the sun evaporates the water over weeks; because the solubility of sodium chloride hardly changes with temperature, the salt crystallises only as the water is removed.
The effect of pressure
Pressure has no significant effect on the solubility of a solid in a liquid.
Solids and liquids are almost incompressible: squeezing them does not bring the particles noticeably closer or change how they attract each other, so the dissolution equilibrium is practically unaffected.
Gases are different, because a gas can be compressed, and that is the subject of Henry's law in the next section.
Interpreting a solubility value
A solubility of 31.6 g of KNO₃ per 100 g of water at 20 °C means that 131.6 g of saturated solution contains 31.6 g of the salt.
As a mass percentage that is 31.6/131.6 × 100 = 24.0%.
Be careful which "100 g" a solubility refers to: it is usually per 100 g of water, not per 100 g of solution.
- Solubility: Solubility of a substance is its maximum amount that can be dissolved in a specified amount of solvent at a specified temperature.
- Like dissolves like: In general, a solute dissolves in a solvent if the intermolecular interactions are similar in the two or we may say like dissolves like.
- Dissolution: solute going into solution as its concentration rises. This process is known as dissolution.
- Crystallisation: dissolved solute particles colliding with the solid and separating out of solution. This process is known as crystallisation.
- Saturated solution: Such a solution in which no more solute can be dissolved at the same temperature and pressure is called a saturated solution.
- Unsaturated solution: An unsaturated solution is one in which more solute can be dissolved at the same temperature.
- Effect of pressure on solids: Pressure does not have any significant effect on solubility of solids in liquids. It is so because solids and liquids are highly incompressible and practically remain unaffected by changes in pressure.
Section C, 3 marks (1 + 1 + 1).
Give reasons: (a) Iodine dissolves readily in carbon tetrachloride but only slightly in water.
(b) Glucose dissolves in water but not in hexane.
(c) A grease stain on a shirt is removed with kerosene, not with water.
1. (a) I₂ and CCl₄ are both non-polar and attract each other only by weak dispersion forces, so similar interactions are exchanged; I₂ cannot replace the hydrogen bonds it would have to break in water (1).
2. (b) Glucose has five OH groups and forms hydrogen bonds with water molecules; non-polar hexane cannot form such bonds, so the glucose-glucose hydrogen bonds are not repaid (1).
3. (c) Grease and kerosene are both non-polar hydrocarbons, so the grease dissolves in kerosene (like dissolves like); grease cannot mix with polar water (1).
Answer: each mark names the kind of interaction on both sides (polar or hydrogen bonding against non-polar or dispersion forces).
"Like dissolves like" alone, with no reason, is usually worth only ½.
Section B, 2 marks.
The dissolution of ammonium nitrate in water is endothermic.
Predict how its solubility changes when the temperature is raised, and explain using Le Chatelier's principle.
1. Write the equilibrium with heat on the left: NH₄NO₃(s) + water + heat ⇌ NH₄NO₃(aq) (½).
2. Raising the temperature supplies heat; the equilibrium shifts in the direction that absorbs it, the forward direction (1).
3. So the solubility increases with temperature (½).
Answer: solubility increases on heating, because the endothermic forward process is favoured.
If the dissolution were exothermic, the same reasoning would give a decrease.
Section C, 3 marks (2 + 1).
Use the solubility curves above (KNO₃: 110 g per 100 g water at 60 °C and 31.6 g per 100 g water at 20 °C).
(a) A saturated solution of KNO₃ made with 50 g of water at 60 °C is cooled to 20 °C. Calculate the mass of KNO₃ that crystallises.
(b) Calculate the mass percentage of KNO₃ in the saturated solution at 20 °C.
1. (a) At 60 °C, 50 g of water holds 110 × 50/100 = 55 g of KNO₃ (½).
2. At 20 °C, 50 g of water holds only 31.6 × 50/100 = 15.8 g (½).
3. Mass crystallising = 55 − 15.8 = 39.2 g (1; ½ deducted without the unit).
4. (b) Mass % = 31.6 / (31.6 + 100) × 100 = 24.0% (1).
Answer: 39.2 g; 24.0%.
Note in (b) that the whole solution is 131.6 g, not 100 g.
The same cooling with NaCl would give only about half a gram, which is why salt is not purified this way.
When asked how temperature affects the solubility of a solid, always link it to the sign of ΔsolH and to Le Chatelier's principle in one chain: "dissolution is endothermic, so raising the temperature shifts the equilibrium forward and solubility increases".
A bare "solubility increases with temperature" is not true for every solid and may not earn the mark.
- "All solids dissolve better in hot water." Only when dissolving is endothermic. Calcium hydroxide and cerium(III) sulphate dissolve less as the water warms, and NaCl hardly changes.
- "In a saturated solution, dissolving has stopped." Dissolution and crystallisation both continue, at equal rates. That is what dynamic equilibrium means.
- "Pressure helps a solid dissolve." Solids and liquids are practically incompressible, so pressure has no significant effect. Pressure matters for gases.
- Using 100 g of solution for a solubility quoted per 100 g of water. A solubility of 31.6 g per 100 g of water is 31.6 g in 131.6 g of solution.
Check your understanding
1. Which of the following dissolves most readily in benzene?
- (A) sodium chloride
- (B) glucose
- (C) naphthalene
- (D) urea [1]
2. Given below are two statements labelled Assertion (A) and Reason (R). Select the most appropriate answer from the options (A), (B), (C) and (D) given below.
Assertion (A): The solubility of sodium chloride in water hardly changes when the pressure above the solution is doubled.
Reason (R): Sodium chloride is an ionic compound made of Na⁺ and Cl⁻ ions.
- (A) Both Assertion (A) and Reason (R) are true and Reason (R) is the correct explanation of the Assertion (A).
- (B) Both Assertion (A) and Reason (R) are true, but Reason (R) is not the correct explanation of the Assertion (A).
- (C) Assertion (A) is true, but Reason (R) is false.
- (D) Assertion (A) is false, but Reason (R) is true. [1]
3. The dissolution of a salt in water is exothermic. When the temperature of its saturated solution is raised, the solubility of the salt
- (A) increases
- (B) decreases
- (C) does not change
- (D) first increases and then decreases [1]
4. (a) What is meant by a saturated solution? (b) How will you show in the laboratory that a given solution of potassium nitrate is unsaturated? (1 + 1) [2]
5. The solubility of a salt X is 80 g per 100 g of water at 70 °C and 30 g per 100 g of water at 20 °C. (a) Calculate the mass of X that crystallises when 90 g of its saturated solution at 70 °C is cooled to 20 °C. (b) Calculate the mass percentage of X in its saturated solution at 20 °C. (2 + 1) [3]
6. Give reasons: (a) Methanol mixes with water in all proportions. (b) Kerosene does not mix with water. (1 + 1) [2]
Answers
1. (C) / naphthalene (1). Naphthalene and benzene are both non-polar; the other three are ionic or form hydrogen bonds and need a polar solvent.
2. (B) / Both Assertion (A) and Reason (R) are true, but Reason (R) is not the correct explanation of the Assertion (A). (1). Both statements are true, but the reason pressure has no effect is that solids and liquids are practically incompressible, not that NaCl is ionic.
3. (B) / decreases (1). For an exothermic process, adding heat shifts the equilibrium Solute + Solvent ⇌ Solution + heat to the left.
4. (a) Such a solution in which no more solute can be dissolved at the same temperature and pressure is called a saturated solution. (1) (b) Add a few more crystals of potassium nitrate at the same temperature and stir: if they dissolve, the solution was unsaturated (1).
5. (a) At 70 °C, 180 g of saturated solution contains 80 g of X, so 90 g of solution contains 40 g of X and 50 g of water (1). At 20 °C, 50 g of water holds 15 g of X, so 40 − 15 = 25 g of X crystallises (1; ½ deducted without the unit). (b) Mass % = 30 / (30 + 100) × 100 = 23.1% (1).
6. (a) Methanol molecules have an OH group and form hydrogen bonds with water molecules, interactions similar to those in pure water and pure methanol (1). (b) Kerosene is a non-polar hydrocarbon; it cannot form hydrogen bonds with water or replace the hydrogen bonds between water molecules (1).