mass
'no magic to the mole'
amount
molar mass
concentration
solution volume
gas volume
molar gas volume
Avogadro
constant, L
number of
entities, N
Now try the following question.
CALCULATING ON THE SIDE OF SUCCESS – Module F(2022): SOLUTION CONCENTRATION & AMOUNT OF SUBSTANCE
F. CONCENTRATION OF SOLUTIONS
F1. INTRODUCTION TO SOLUTION CONCENTRATION & SOLUBILITY
For pure solids the amount of substance is calculated from the mass of substance and its molar mass;
for gases, the amount of substance is calculated from the volume of gas and the molar gas volume.
In practice, however, the majority of inorganic chemical reactions met at this stage are likely to occur in aqueous solution.
Chemistry and other Sciences must therefore have a method for keeping tabs on the amount of substance present in a solution. One very convenient way of doing so is by considering the amount of substance present per unit volume of solution. This takes us to the idea of CONCENTRATION. It is worth pausing, however, to recap on some important terminology encountered when dealing with solutions.
When a solid, liquid or gas is placed in a suitable solvent - at a sufficient temperature - it may dissolve
as it disperses into nano-scopic particles (ions, molecules or atoms) and passes into solution. The new homogeneous mixture of the combination is termed a solution. Should one of the substances have changed state then that qualifies it as the solute. If neither component has changed state then that present in the greater quantity is referred to as the solvent; the minor component is referred to as the solute.
Take two relatively simple cases: one a covalent, molecular substance like sucrose, C12H22O11(s); the other ionic, like sodium chloride, NaCℓ(s).
Examples of other less obvious solute / solvent phase mixtures that might be encountered include:
Where solutions are colourless, there is no obvious visual indication of how much of a solute is present. With coloured solutions, however, some visual estimation may be possible as depicted opposite in the various plastic bottle containing aqueous CuSO4 stock solutions.
The solubility of solids is often defined as the mass of a particular
solid substance that will dissolve in a known mass of solvent
at a specified temperature, often expressed in units g / 100 g solvent.
SATURATED SOLUTION - one which contains as much dissolved solute as possible at a specified temperature.
Solubility data and their temperature dependence for KCℓO3, NaCℓ, & PbCℓ2 are as follows.
These figures represent the conditions at saturation. Values below these figures, at a given temperature, indicate an unsaturated (or under-saturated) solution.
Solubility data and their temperature dependence for a given substance in a specified solvent are conveniently presented graphically in the shape of a so-called solubility curve. Those for three inorganic salts - lead(II) chloride, sodium chloride, and potassium chlorate are shown below.
Any point on a plot represents a saturated solution.
Data falling above the curve points towards either over-saturated or supersaturated solutions.
In the former, crystallization would be expected to occur.
Data falling in the area under the curve point towards under-saturated solutions.
Lead(II) chloride is generally considered
to be insoluble. Values below 3 g/dL are typical of 'insoluble' salts.
The solubility of NaCℓ does not vary greatly with temperature.
The following descriptive categories are in use to class varying degrees of solubility, depending upon the mass of solute that dissolves in unit mass of solvent. While these find particular use in pharmacy and drug solubility, the terms are convenient elsewhere also. A few illustrative examples of salts of varying solubility that are encountered in inorganic chemistry are tabulated below.
In almost all scientific disciplines, whether pure or applied, a knowledge of how the solubility of substances depend on solvent, temperature, and pressure are of considerable importance. Solubility in aqueous solution of sparingly soluble metal salts such as carbonates plays an important role in chemical processes. Solubility phenomena like dissolution and precipitation reactions frequently control procedures for preparing, separating, and purifying chemicals. Moreover, since so many chemical processes occur in important aqueous systems, e.g., the human body, it is vital that the chemist, biochemist, chemical engineer, pharmacist, or physician knows, and can account for, the solubility of the materials with which they work.
Like KCℓO3, hydrated copper(II) sulphate crystals, CuSO4.5H2O(c), possess a solubility in water that varies markedly with temperature.
At 20 °C, a maximum mass of 32 g of CuSO4.5H2O(c), will dissolve in 100 g of water. Any mass in excess of 32 g will, in this case, remain un-dissolved.
3
3
Since the concentration of a solution depends upon its volume, one needs to be comfortable dealing with the units for this physical quantity, and particularly the inter-conversion of cm and dm . These have the capacity to confuse. For the student who applies quantity algebra systematically, the procedures involved are entirely routine. In any event, however, it is likely to be beneficial for commonly used units of volume to be considered at the outset.