Struggling with solubility and solubility product in chemistry? Learn the easy-to-understand guide with examples, formulas, and tips that make solving Ksp problems simple and exam-ready. Perfect for Class XI & XII students!
Welcome to Learn Chemistry by Inam Jazbi! Today, we’re diving into one of the most important topics in chemistry – Solubility and Solubility Product (Ksp). Don’t worry if it sounds tricky – I’ll break it down step by step, with simple explanations, examples, and tips that make it super easy to understand. By the end of this post, you’ll be able to solve any solubility problem like a pro, boost your exam confidence, and actually enjoy chemistry!
Solubility and Factors Affecting Solubility
Definition of
Saturated Solution
A solution that contains the maximum equilibrium amount of solute at a specific temperature is called saturated solution. A saturated solution cannot dissolve any more amount of solute at a given temperature.
Definition and Unit of Solubility
The solubility (dissolution) of a substance in a given solvent is a physical property. Solubility is a measure of the amount of solute that will dissolve in a solvent at a specific temperature. It is the extent (tendency) of a solvent to dissolve solute.
Solubility is the maximum equilibrium amount of solute (expressed either in grams or mole) needed to prepare a
saturated solution in a given quantity of solvent (usually 100 g or 100 ml or
in one litre or dm3 or 1000 mL of a solvent) at specified
temperature.
[In fact, the solubility of a solute in a solvent is the concentration of the saturated solution at a specified temperature. Thus it denotes maximum amount of solute which be held in a solution at a particular temperature.]
Miscible and Immiscible Liquids
Pairs of liquids that are mutually soluble in all proportions are
said to be miscible e.g. ethyl alcohol and water are miscible.
Pairs of liquids that do not mix are said to be immiscible e.g. water and methane are immiscible.
Unit of Solubility
Solubility is usually expressed as:
g/litre or g/dm3→
grams of solute per litre of solvent
mol/L (Molarity) → moles of solute per litre of solution i.e. Molar solubility (s) (practical and S.I unit)
g/100 g water or g/100 ml water → commonly used in data tables and solubility curves
Formula to calculate Solubility
Solubility of Some Solutes at 20°C
1. KNO₃ = 37 g (31.6 g)
2. NaCl = 36.0 g
3. KCl = 34.0 g
4. NaBr = 94.32 g
5. AgNO₃ = 222.0 g
6. NaNO₃ = 91.2 g
Types of Compounds according to Solubility
1. Soluble compounds; With solubility 10 or more gram per litre or
Solubility > 0.1 M
2. Slightly soluble compounds; With solubility 0.1 to 10 gram per litre or
Solubility = 0.01-0.1
M
3. Insoluble compounds; With solubility less than 0.1 gram per litre or Solubility < 0.01 M
Factors Affecting Solubility
Solubility depends upon following factors:
1. Temperature
2. Pressure
3. Nature of solute and solvent
1. Effect of Nature of Solute and Solvent
1. Substances with similar types of polarity or intermolecular
attractive forces tend to be soluble in one another
(i.e. substances may mix and dissolve in each other if they have approximately
the same type of polarity).
This generalization is a guiding rule of solubility and is often simply stated
as “Like dissolves like” which means
non-polar substances are soluble in non-polar solvent (e.g. benzene is soluble in CCl4 as both are
non-polar. Similarly, naphthalene or methane being a non-polar solute is more soluble in non-polar solvents like
benzene or carbon tetrachloride than in water) while polar or ionic solutes are soluble in polar
solvents. (e.g common salt (NaCl) being an ionic compound dissolves more readily in polar
solvent like water but it is insoluble in non-polar (organic) solvents like benzene or petrol).
The polar (or ionic) nature of polar solutes and polar
solvent like water produces particularly strong solute-solvent interactions in
the form of ion-dipole forces or dipole-dipole forces thereby favouring
solubility. On the other hand, non-polar solutes and polar solvents like water
have weak solute-solvent interactions in the form of dispersion forces which is
not sufficient enough to overcome the stronger hydrogen bonding initially
present in water.]
2. Hydrogen bonding interactions between solute and solvent greatly increase the solubility of non-polar compounds in water. e.g. ethanol is completely miscible with water due to hydrogen bonding(and also both have same type of structure having polarized –OH group). Similarly sugar, glucose, glycerine although non-polar in nature are very soluble in water due to hydrogen bonding as both have similar type of structure containing –OH groups].
3. The
solubility of the compound increases with the increase of –OH group along the
carbon chains of organic compounds.
e.g.
(i) glucose with
five –OH groups on a six-carbon framework is very soluble in water (83 g/100 mL of
water at 17.5°C).
(ii) Similarly, sugar with 10 –OH groups on a twelve-carbon framework is highly soluble in water (179 g/100 mL of water at 0°C).
4. The solubility of gases increases with increasing molecular mass
e.g.
i). Solubility of N2 gas (molecular mass 28 amu) is 0.63 × 10‒3 M
ii) Solubility of CO gas (molecular mass 28 amu) is 1.04 × 10‒3 M
iii) Solubility of O2 gas (molecular mass 32 amu) is 1.38 × 10‒3 M
iv) Solubility of Ar gas (molecular mass 40 amu) is 1.50 × 10‒3 M
v) Solubility of Kr gas (molecular mass 83.7amu) is 2.79 × 10‒3 M
5. The solubility of gas increases if it reacts with solvent. e.g. the solubility of Cl2 in water is 0.102M (much higher than predicted) as it reacts with water on dissolution. Similarly, solubility of CO2 and NH3 are also much higher than expected because of their reaction with solvent water.
Cl2 + H2O → HOCl + HCl
NH3 + H2O ⇌ NH4+ + OH‒
CO2 + H2O ⇌ H2CO3
6.The
solubility of the alcohols decreases with increasing molecular masses (as the –OH group is more tightly bound and molecules become more like a hydrocarbon).
e.g.
(i) solubility of butanol is 0.11 mol/100 g
H2O at 20°C
(ii) solubility
of pentanol is 0.03 mol/100 g H2O at 20°C.
Summary of Effect of Nature of Solute and
Solvent on Solubility
The rule “Like dissolves like” applies:
Polar solutes dissolve in polar solvents (e.g., NaCl in water).
Non-polar solutes dissolve in non-polar solvents (e.g., wax in
benzene).
Polar substances dissolve well due to ion–dipole or hydrogen
bonding interactions.
2. Effect of
Temperature
Solubilities are temperature-dependent. The solubilities of most ionic and molecular solids in liquids (or solubility of partially miscible liquids) usually increase with increasing temperature though the solubilities of some (NaCl, NaBr) are almost unchanged and the solubilities of others [CaSO4, Na2SO4, Ce2(SO4)3] decrease i.e. solubility of solids in liquids is directly proportional to temperature. The dissolution for ionic solids is endothermic (heat is absorbed), so by Le-Chatelier’s principle the solubility increases with increasing temperature.
When a solid substance is dissolved in water, either heat is
evolved (exothermic) or heat is absorbed (endothermic).
(i) For endothermic
solubility process, solubility increases with increase in temperature.
(ii) For exothermic solubility process, solubility decreases with
increase in temperature.
For example;
i) Solubility
of sugar in water at 0°C is 179 g/100 ml of water whereas 100°C it is 487 g/100
ml of water.
ii) Similarly the solubility of KNO3 at 0°C is 13.5 g/100 g of water but at 100°C it is 247 g/100 g of water.
However, solubility of some solutes in liquids decreases with the
increase in temperature. e.g. CaSO4, Na2SO4,
Ce2(SO4)3. The dissolution for some solids is
exothermic (heat is released), so by Le-Chatelier’s principle the solubility
decreases with rising temperature.
e.g. solubility of sodium sulphate decahydrate (Na2SO4.10H2O) increases with the rise of temperature till 32.4°C reached and then its solubility decreases. This is due to the decomposition of hydrated salt to the anhydrous Na2SO4.
The solubility of gases in a liquid (unlike solids) decreases with
increasing temperature i.e. gases are more soluble in cold water than in hot
water. Thus when a solution comprising of gas in a liquid is heated, gases are
evolved
e.g.
(i) the solubility of oxygen in water at
0°C is 4.8 cm3/100 cm3 water while at 100°C it is only
1.72 cm3.
(ii) at 0°C, the
solubility of CO2 in water is 171.3 cm3 but at 60°C its
solubility is reduced to 35.9 cm3.
consequences of rising
temperatures on gas solubility
Some consequences of rising temperatures on gas solubility are
discussed below:
i). When a glass of cold tape water is warmed, the bubbles of dissolved air are seen on the inside of the glass. [The boiled water has a characteristic flat taste because dissolved air (and also the salts) has been expelled during boiling. That is why it is advised to fill boiled water in bottles after cooling].
ii). The effect of rise in temperature on gas solubility is obvious when carbonated drinks bubble continuously as they warm up to room temperature after being refrigerated. Soon carbonated beverages go ‘flat’ due to escape of dissolved CO2 gas. At 0°C, the solubility of CO2 in water is 171.3 cm3 but at 60°C its solubility is reduced to 35.9 cm3. Thus it is advised by soft drink companies that carbonated beverages give their best taste when chilled.]
iii). A much more important consequence is the damage to aquatic life that can result from the decrease in dissolved oxygen when hot water is discharged from power stations into lakes and rivers, an effect known as Thermal Pollution. [The effect is particularly serious in deep lakes because warm water is less dense than cold water. It therefore, tends to remain on top of cold water, at the surface. This situation impedes the dissolving of oxygen into the deeper layers, thus stifling the respiration of all aquatic life needing oxygen. Fish may suffocate and die in these circumstances].
iv). On a hot summer day, an experienced fisherman knowing rules of gas solubility usually picks a deep spot in the river or lake to cast the bait because the oxygen content is greater in the deeper, cooler region, most fish will be found there.
Summary of Effect of Temperature on Solubility
Effect on
Solids in Liquids:
Solubility usually
increases with temperature because heat provides more kinetic energy to
particles, helping them separate and mix.
Example:
Solubility of KNO₃ increases as temperature rises.
Effect on Gases
in Liquids:
Solubility decreases
with temperature because gas molecules gain energy and escape from the liquid.
Example: Warm water holds less CO₂ (why cold drinks go flat when warm).
3. Effect of Pressure
The solubilities of solids and liquids in liquid solvent are not affected by pressure (i.e. pressure has practically no effect on the solubility of solids and liquids). The solubility of gases in liquids considerably increases with increasing pressure.
The quantitative relationship between gas solubility and pressure
is given by Henry’s Law which states that the solubility of gas (i.e. the
amount of gas dissolved in a given amount of liquid solvent) at constant
temperature is directly proportional to the partial pressure of the gas over
the solution.
m ∝ Pg And m = k Pg Where;
m = amount of gas
dissolved.
P = Partial pressure of the gas over the solution.
k = Henry’s Law Constant (with unit mol/liter-atm),
which is temperature dependent and characteristic
of a specific gas.
e.g.
at S.T.P. (standard temperature, 0°C and standard pressure, 1 atmosphere), 0.335 g of CO2 dissolves per 100 cm3 of water but if pressure is doubled 0.67g of CO2 will dissolve.]
This effect is used in the manufacture of soft drinks bottled such as Coca-Cola, 7-up etc. These are bottled in which CO2 is filled under high pressure slightly greater than 1 atm. [CO2 gas is slightly soluble in water at S.T.P., so CO2 gas is filled in soda water bottles under high pressure].
The most common example of Henry’s law behaviour occurs when a can or bottle of soda or other carbonated drinks (in which CO2 is filled under pressure slightly greater than 1 atm) are opened to air, bubbles of CO2 gas comes fizzing out of the solution with effervescence because the pressure of CO2 above the solution drops and CO2 suddenly becomes less soluble
A more serious example of Henry’s law behaviour occurs when a deep-sea diver surfaces too quickly and develops a painful and life-threatening condition called the “bend”. Deep-sea divers rely on compressed air for their oxygen supply. If a diver is suddenly exposed to atmospheric pressure (where the solubility of gases is less), large amounts of nitrogen dissolved in the blood at high underwater pressure, form bubbles in the blood-stream blocking capillaries and inhibiting blood flow, and affect nerve impulses giving rise to the condition known as the bends or decompression sickness. [The bends can be prevented by using an oxygen/helium mixture (98% He and 2% O2) for breathing rather than air (O/N2), because helium has a much lower solubility in blood than nitrogen].
Summary of Effect of Pressure on Solubility
Pressure affects only gases dissolved in liquids.
According to Henry’s Law:
“The solubility of a gas in a liquid is directly proportional to
the pressure of that gas above the liquid.”
Example:
⇒ Soft drinks are bottled under high pressure to keep CO₂
dissolved.
⇒ When opened, pressure drops → gas escapes → fizz appears.
4. Common Ion
Effect
Solubility of a
sparingly soluble salt decreases when a common ion is added from another
source.
This is
explained by Le-Chatelier’s Principle — the system shifts toward the solid
phase to reduce ion concentration.
Example:
AgCl(s)⇌Ag⁺ + Cl⁻
If NaCl is
added → [Cl⁻] increases → equilibrium shifts left → less AgCl dissolves.
5. pH of the
Solution
The solubility
of salts of weak acids (like carbonates, sulfides, and hydroxides) increases in
acidic medium because H⁺ reacts with the anion.
Examples:
CaCO₃ (s)+2H⁺ →Ca2⁺ + CO₂ + H₂O
FeS(s)+2H⁺→Fe2⁺ + H₂S↑
6. Formation of
Complex Ions
Solubility of
some sparingly soluble salts increases in the presence of a complexing agent,
because new soluble complexes form.
Example:
AgCl(s)+2NH₃→[Ag(NH₃)₂]⁺ + Cl−
Here, AgCl
dissolves due to the formation of the soluble diammine silver complex.
7. Lattice Energy and Hydration Energy
Lattice Energy
(U): Energy needed to separate ions in a crystal lattice.
Hydration
Energy (H): Energy released when ions are surrounded by water molecules.
A salt
dissolves if hydration energy > lattice energy.
Example:
NaCl dissolves easily (hydration > lattice), but AgCl does not (lattice >
hydration).
Saturated Solution and Solubility Product (Ksp)
Solubility
Solubility is the extent (tendency) of a solvent to dissolve
solute. It is defined as the amount of solute in gram, which is required to
prepare a saturated solution of 1 liter (1 dm3). It is expressed in
g/liter or g/dm3. The
practical unit is mol/liter.
Types of Compounds according to Solubility
Salts are
classified into three types on the basis of their solubility in the following
table.
Saturated Solution
A solution that contains the maximum equilibrium amount of solute and cannot dissolve any more amount of solute at a specific temperature is called saturated solution. A saturated solution is in a state of dynamic equilibrium between the dissolved, dissociated, ionic compound and the undissolved solid.
Definition of
Solubility Product
The value of Ksp is a measure of the solubility of an ionic salt (ionic compound). The larger the value of Ksp, the greater is the concentration of ions in the solution and hence greater is the solubility of the salt and vice versa. The smaller the value of Ksp, the more insoluble is the salt.
The product of
molar ionic concentration of its positive and negative ions of dissolved
sparingly soluble salt, each raised to an appropriate power equal to its
co-efficient in balanced ionized equilibrium equation in a saturated solution
at a specific temperature is called Solubility Product represented by Ksp. It is
expressed in chemical unit of concentration i.e. molar concentration (mol/dm3).
OR
It is the product of the concentration of the ions of a
dissolved salt produced in a saturated solution; each concentration is raised
to power equal to its co-efficient given in the balanced equation at a given
temperature.
Ksp is the equilibrium constant for the equilibrium between a solid ionic solute and its ions in a saturated solution.
It represents the equilibrium between solid phase and ions in a saturated solution of ionic compounds of relatively low solubility. Solubility products of a salt in a saturated solution remains constant. Like other equilibrium constants it changes with temperature.
General Representation of Solubility Product
For any sparingly soluble salt like AmBn, the Ksp expression is given as:
Explanation
Ksp represents the
maximum value of the ionic product (i.e. Ksp represents the
saturated state of the solution). Thus a salt cannot be dissolved in a solvent
after the achievement of its Ksp. The precipitation of a salt can
only occur when its ionic product is greater than its Ksp.
The solubility product characterizes the solubility of a salt at a given temperature. If there are two similar salts (e.g. CaSO4, BaSO4), the solubility is greater for that salt whose solubility product is greater.
Unit of Solubility Product
Ksp is represented in chemical unit of concentration i.e. molar
concentration (moles/dm3). It is expressed in mol2/dm6
(or M2) or mol3/dm9 (or M3) etc.
Factor affecting
Solubility Product
Ksp varies directly
with temperature
(as increase in temperature increases solubility, which in turn raises
ionic concentrations).
Derivation for a General Ksp Expression
Various ionic compounds such as AgCl, BaSO4, Ca3(PO4)2
etc. are practically very slightly soluble in water and commonly known as
sparingly soluble salts. When a sparingly or slightly soluble salt like AmBn
is dissolved in water, its very small part becomes ionized and it forms a
saturated after certain time. At this stage an equilibrium is established
between dissolved ions of salt and its solid phase (undissolved salt). By
applying law of mass action, the equilibrium constant (Kc) can be
given as:
Since for a saturated solution, the concentration of undissolved
solid salt (AmBn) remains constants, therefore; [AmBn(s)]
= K, it is not included in the equilibrium expression and thus Kc is
replaced by Ksp which is known as solubility product constant or
simply solubility product. By applying law of mass action, the equilibrium
constant (Kc) and the equilibrium expression of this ionic dissolution
equilibrium can be given as:
Derivation for Ksp Expression
Various ionic compounds such as AgCl, BaSO4, Ca3(PO4)2
etc. are practically very slightly soluble in water and commonly known as
sparingly soluble salts. CaSO4 is slightly soluble salt, when it is
dissolved in water, its very small part becomes ionized and it forms a
saturated after certain time. At this stage an equilibrium is established
between dissolved ions of salt and its solid phase (undissolved salt).
Since for a saturated solution, the concentration of undissolved
solid salt (CaSO4) remains constants, therefore; [CaSO4(s)]
= K, it is not included in the equilibrium expression and thus Kc is
replaced by Ksp which is known as solubility product constant or
simply solubility product. By applying law of mass action, the equilibrium
constant (Kc) and the equilibrium expression of this ionic
dissolution equilibrium can be given as:
When a sparingly or slightly soluble salt like AgCl is dissolved in
water, then it forms saturated solution after certain time. The amount of salt
in a saturated solution represents the maximum limit of dissolved salt at given
temperature. At this stage an equilibrium is established between dissolved ions
of salt and undissolved salt. By applying law of mass action, the equilibrium
constant (Kc) can be given as:
Applications of Solubility Product
Precipitation is a kind of double displacement reaction in which
two solutions of different salts mixed together to form two products, one of
these product is insoluble in solution and called precipitate (ppt).
In order to predict precipitation, we have to calculate Ionic
Product or reaction quotient (Qsp or QI)
denoted by Qsp
or QI which is the product of
initial ionic molar concentrations of a dissolved salt (not necessarily correspond to
those at equilibrium) each raised to the power by
its coefficient mentioned in net ionic equation of ionization. Ionic product is relevant to
both saturated and unsaturated solution while Ksp is only applicable
to saturated solution (which involves a dynamic equilibrium between an
insoluble salt and its aqueous ions).
The solubility product (Ksp) of salts helps to predict
whether the precipitation will occur or not from the solution of known ionic
concentrations. Ksp value represents the saturated state of the
solution. The following three
possible relations between Qsp (ionic product) and Ksp
are used to predict the nature of solution even if the salt forms precipitate
by mixing two different salt solutions.
