Molar concentration, also called molarity, amount concentration or substance concentration, is a measure of the concentration of a solute in a solution, or of any chemical species, in terms of amount of substance in a given volume. A commonly used unit for molar concentration used in chemistry is mol/L. A solution of concentration 1 mol/L is also denoted as 1 molar (1 M).
- 1 Definition
- 2 Units
- 3 Related quantities
- 4 Properties
- 5 Spatial variation and diffusion
- 6 Examples
- 7 Formal concentration
- 8 References
- 9 External links
Molar concentration or molarity is most commonly expressed in units of moles of solute per litre of solution. For use in broader applications, it is defined as amount of solute per unit volume of solution, or per unit volume available to the species, represented by lowercase c:
Here, n is the amount of the solute in moles, N is the number of molecules present in the volume V (in litres), the ratio N/V is the number concentration C, and NA is the Avogadro constant, approximately 6.022×1023 mol−1.
Or more simply: 1 molar = 1 M = 1 mole/litre.
In thermodynamics the use of molar concentration is often not convenient, because the volume of most solutions slightly depends on temperature due to thermal expansion. This problem is usually resolved by introducing temperature correction factors, or by using a temperature-independent measure of concentration such as molality.
In the International System of Units (SI) the base unit for molar concentration is mol/m3. However, this is impractical for most laboratory purposes and most chemical literature traditionally uses mol/dm3, or mol dm−3, which is the same as mol/L. These traditional units are often denoted by a capital letter M (pronounced molar), sometimes preceded by an SI prefix to denote sub-multiples, for example:
|Name||Abbreviation||Concentration||Concentration (SI unit)|
|millimolar||mM||10−3 mol/dm3||100 mol/m3|
|micromolar||μM||10−6 mol/dm3||10−3 mol/m3|
|nanomolar||nM||10−9 mol/dm3||10−6 mol/m3|
|picomolar||pM||10−12 mol/dm3||10−9 mol/m3|
|femtomolar||fM||10−15 mol/dm3||10−12 mol/m3|
|attomolar||aM||10−18 mol/dm3||10−15 mol/m3|
|zeptomolar||zM||10−21 mol/dm3||10−18 mol/m3|
(1 particle per 1.6 L)
The conversion to number concentration is given by:
The conversion to mass concentration is given by:
where is the molar mass of constituent .
The conversion to mole fraction is given by:
where is the average molar mass of the solution, is the density of the solution and j is the index of other solutes.
A simpler relation can be obtained by considering the total molar concentration namely the sum of molar concentrations of all the components of the mixture.
The conversion to mass fraction is given by:
The conversion to molality (for binary mixtures) is:
where the solute is assigned the subscript 2.
For solutions with more than one solute, the conversion is:
Sum of molar concentrations – normalizing relations
The sum of molar concentrations gives the total molar concentration, namely the density of the mixture divided by the molar mass of the mixture or by another name the reciprocal of the molar volume of the mixture. In an ionic solution, ionic strength is proportional to the sum of molar concentration of salts.
Sum of products molar concentrations-partial molar volumes
The sum of products between these quantities equals one.
Dependence on volume
Molar concentration depends on the variation of the volume of the solution due mainly to thermal expansion. On small intervals of temperature the dependence is :
where is the molar concentration at a reference temperature, is the thermal expansion coefficient of the mixture.
Spatial variation and diffusion
Molar and mass concentration have different values in space where diffusion happens.
Example 1: Consider 11.6 g of NaCl dissolved in 100 g of water. The final mass concentration ρ(NaCl) will be:
- ρ(NaCl) = 11.6 g / (11.6 g + 100 g) = 0.104 g/g = 10.4 %
The density of such a solution is 1.07 g/mL, thus its volume will be:
- V = (11.6 g + 100 g) / (1.07 g/mL) = 104.3 mL
The molar concentration of NaCl in the solution is therefore:
- c(NaCl) = (11.6 g / 58 g/mol) / 104.3 mL = 0.00192 mol/mL = 1.92 mol/L
Here, 58 g/mol is the molar mass of NaCl.
Example 2: Another typical task in chemistry is the preparation of 100 mL (= 0.1 L) of a 2 mol/L solution of NaCl in water. The mass of salt needed is:
- m(NaCl) = 2 mol/L x 0.1 L x 58 g/mol = 11.6 g
To create the solution, 11.6 g NaCl are placed in a volumetric flask, dissolved in some water, then followed by the addition of more water until the total volume reaches 100 mL.
Example 3: The density of water is approximately 1000 g/L and its molar mass is 18.02 g/mol (or 1/18.02=0.055 mol/g). Therefore, the molar concentration of water is:
- c(H2O) = 1000 g/L / (18.02 g/mol) = 55.5 mol/L
Likewise, the concentration of solid hydrogen (molar mass = 2.02 g/mol) is:
- c(H2) = 88 g/L / (2.02 g/mol) = 43.7 mol/L
The concentration of pure osmium tetroxide (molar mass = 254.23 g/mol) is:
- c(OsO4) = 5.1 kg/L / (254.23 g/mol) = 20.1 mol/L.
- C = 60 / (10−15 L)= 6×1016 L−1
The molar concentration is:
- = 6×1016 L−1 / (6×1023 mol−1) = 10−7 mol/L = 100 nmol/L
If the concentration refers to original chemical formula in solution, the molar concentration is sometimes called formal concentration. For example, if a sodium carbonate solution has a formal concentration of c(Na2CO3) = 1 mol/L, the molar concentrations are c(Na+) = 2 mol/L and c(CO2−
3) = 1 mol/L because the salt dissociates into these ions.