Molar concentration

"Molarity" redirects here. It is not to be confused with Molality or Morality.

For the comic strip, see Molarity (comic strip).

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).

Definition

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:^[1]

c={\frac {n}{V}}={\frac {N}{N_{{{\rm {A}}}}\,V}}={\frac {C}{N_{{{\rm {A}}}}}}.

Here, n is the amount of the solute in moles,^[2] N is the number of molecules present in the volume V (in litres), the ratio N/V is the number concentration C, and N_A 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.^[2]

The reciprocal quantity represents the dilution (volume) which can appear in Ostwald's law of dilution.

Units

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:

mol/m³ = 10⁻³ mol/dm³ = 10⁻³ mol/L = 10⁻³ M = 1 mmol/L = 1 mM.

The words "millimolar" and "micromolar" refer to mM and μM (10⁻³ mol/L and 10⁻⁶ mol/L), respectively.

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
yoctomolar	yM^[3]	10⁻²⁴ mol/dm³ (1 particle per 1.6 L)	10⁻²¹ mol/m³

Related quantities

Number concentration

The conversion to number concentration $C_{i}$ is given by:

C_i = c_i \cdot N_{\rm A}

where $N_{\rm A}$ is the Avogadro constant, approximately 6.022×10²³ mol⁻¹.

Mass concentration

The conversion to mass concentration $\rho _{i}$ is given by:

\rho_i = c_i \cdot M_i

where $M_{i}$ is the molar mass of constituent $i$ .

Mole fraction

The conversion to mole fraction $x_{i}$ is given by:

x_i = c_i \cdot \frac{M}{\rho} = c_i \cdot \frac{\sum_i x_i M_i}{\rho}

x_i= c_i \cdot \frac{\sum x_j M_j}{\rho - c_i M_i}

where $M$ is the average molar mass of the solution, $\rho$ 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.

x_i = \frac{c_i}{c} = \frac{c_i}{\sum c_i}

Mass fraction

The conversion to mass fraction $w_{i}$ is given by:

w_i = c_i \cdot \frac{M_i}{\rho}

Molality

The conversion to molality (for binary mixtures) is:

b_2 = \frac{{c_2}}{{\rho - c_2 \cdot M_2}} \,

where the solute is assigned the subscript 2.

For solutions with more than one solute, the conversion is:

b_i = \frac{{c_i}}{{\rho - \sum c_i \cdot M_i}} \,

Properties

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.

\sum_i c_i \cdot \bar{V_i} = 1

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 :

c_i = \frac {{c_{i,T_0}}}{{(1 + \alpha \cdot \Delta T)}}

where $c_{i,T_0}$ is the molar concentration at a reference temperature, $\alpha$ is the thermal expansion coefficient of the mixture.

Spatial variation and diffusion

Molar and mass concentration have different values in space where diffusion happens.

Examples

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.

Example 4: A typical protein in bacteria, such as E. coli, may have about 60 copies, and the volume of a bacterium is about $10^{-15}$ L. Thus, the number concentration C is:

C = 60 / (10⁻¹⁵ L)= 6×10¹⁶ L⁻¹

The molar concentration is:

c={\frac {C}{N_{A}}}

= 6×10¹⁶ L⁻¹ / (6×10²³ mol⁻¹) = 10⁻⁷ mol/L = 100 nmol/L

Reference ranges for blood tests, sorted by molar concentration.

Formal concentration

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.

References

↑ IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version: (2006–) "amount concentration, c".
1 2 Kaufman, Myron (2002). Principles of thermodynamics. CRC Press. p. 213. ISBN 0-8247-0692-7.
↑ David Bradley. "How low can you go? The Y to Y".

External links

Mole concepts

Physical quantities	Mass concentration Molar concentration Molality Mass Volume Density Mole fraction Mass fraction Amount of substance Molar mass Atomic mass Particle number Avogadro constant Pressure Kelvin temperature Gas constant Molar volume Specific volume

Laws	Charles's law Boyle's law Gay-Lussac's law Combined gas law Ideal gas law

Articles related to solutions

Solution	Ideal solution Aqueous solution Solid solution Buffer solution Flory–Huggins Mixture Suspension Colloid Phase diagram Eutectic point Alloy Saturation Supersaturation Serial dilution Dilution (equation) Apparent molar property Miscibility gap

Concentration and related quantities	Molar concentration Mass concentration Number concentration Volume concentration Normality Percentage solution Molality Mole fraction Mass fraction Isotopic abundance Mixing ratio

Solubility	Solubility equilibrium Total dissolved solids Solvation Solvation shell Enthalpy of solution Lattice energy Raoult's law Henry's law Solubility table (data) Solubility chart

Solvent	(Category) Acid dissociation constant Protic solvent Inorganic nonaqueous solvent Solvation List of boiling and freezing information of solvents Partition coefficient Polarity Hydrophobe Hydrophile Lipophilic Amphiphile Lyonium ion

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