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Molarity is the concentration of a solution in terms of the number of moles of the solute in 1 dm3 (1 liter) of the solution.

The units of molarity are M or mol/L. A 1 M solution is said to be “one molar”.

A mole is the quantity of anything that has the same number of particles as 12 g of carbon-12. This equates to roughly 6.02x1023, also referred to as Avogadro's Number. So, 1 mole of hydrogen gas (H2) contains 6.02x1023 molecules, and 1 mole of glucose (C6H12O6.) also contains 6.02x1023 molecules, but as H2 is a much simpler molecule, 1 mole of H2 will have a much smaller mass (the molar mass) than 1 mole of C6H12O6.

The molar mass is the mass in grams of 1 mole of a particular molecule.

One mole of sodium (Na) is 22.99 g, and 1 mole of chlorine is 35.45 g. For sodium chloride (NaCl) they are in a ratio of 1:1 so the molar mass of NaCl is 22.99 + 35.45 = 58.44 g/mol.
For a compound like water (H2O), 1 mole of hydrogen (H) is 1.008 g/mol and 1 mole of oxygen (O) is 15.9994 g/mol. So, the molar mass of H2O is (2 x 1.008) + 15.9994 = 18.0154 g/mol.

Two solutions that have the same molarity will have the same number of molecules of the chemical per liter but are likely to contain differing masses of that chemical per liter to achieve this. Whereas two solutions at the same concentration will have the same mass of the chemical per liter of solution but are therefore likely to have differing numbers of molecules of that chemical per liter. Provided some additional information is known, one value can be deduced from the other using the equations below.

To calculate molarity or to calculate related values (including volume, mass, molar mass and concentration) from molarity, the following equations are utilized.

Number of moles (mol) = Mass (g) / Molar Mass (g/mol)
Concentration (g/L) = Mass (g) / Volume (L)
Molarity (M or mol/L) = Number of Moles (mol) / Volume (L)

The scenario in which most lab scientists will encounter this type of calculation is when making up solutions following a standard operating procedure (SOP) or a scientific paper. Here, the solution being used is typically defined by its molar concentration (M). For example;
You need to make a 0.5 M solution of NaCl, having decided you want 2 liters how much NaCl should you add?

  1. First you must calculate the number of moles in this solution, by rearranging the equation
    No. Moles (mol) = Molarity (M) x Volume (L)
    = 0.5 x 2
    = 1 mol
  2. For NaCl, the molar mass is 58.44 g/mol
    Now we can use the rearranged equation
    Mass (g) = No. Moles (mol) x Molar Mass (g/mol)
    = 1 x 58.44
    = 58.44 g
    So, to make 2 liters of a 0.5 M solution of NaCl, you would need to add 58.44 g of NaCl.

Rearranging the equation is not necessary as the calculator tool will do this for you.

As you are adding 58.44 g of NaCl to 2 L of water you could also express this value in terms of its concentration
Concentration (g/L) = Mass (g) / Volume (L)
= 58.44 / 2
= 29.22 g/L

Atomic mass (ma) – mass of one atom of that element which typically reflects the mass of the nucleus (protons plus neutrons). For example, hydrogen is 1. This used to be measured in atomic mass units (AMU) but is now typically expressed in Daltons (Da).

Molecular mass (also called molecular weight) - sum of the atomic weights of all atoms appearing in a given molecular formula. For example, glucose has a molecular formula of C6H12O6, the molecular weight of C is 12 Da, H is 1 Da and O is 16 Da. Therefore, the molecular mass of glucose is = (6 x 12) + (12 x 1) + (6 x 16)
= 72 + 12 + 96
= 180 Da
Numerically this is the same as molar mass differing only in the units in which they are expressed.

Relative atomic mass and relative molecular mass (MR) – values for atomic and therefore molecular masses are normally obtained relative to the mass of the isotope 12C (carbon-12), however “relative” is generally omitted from the title. Written correctly, the relative values have no units.

Formula mass (also called formula weight) - sum of the atomic weights of all atoms appearing in a given empirical formula. The empirical formula indicates the ratio of atoms of each element in a molecule rather than the actual number. For example, glucose (molecular formula C6H12O6) would therefore have the empirical formula CH2O, the molecular weight of C is 12 g, H is 1 g and O is 16 g. Therefore, the formula mass of glucose is
= (1 x 12) + (2 x 1) + (1 x 16)
= 12 + 2 +16
= 30 g
For molecules like H2O, where the formula is already in its simplest form, the formula mass and molecular mass are the same.

Normality – similar to molarity, however calculation of normality uses the number of mole equivalents rather than the number of moles. Units are N or eq/L. Normality is generally only used when a substance has more than one sub-species that can participate in a specified reaction such as a proton for acid/base reactions, an electron for oxidation/reduction reactions, or in precipitation reactions. For example, sulfuric acid (H2SO4) has two ionizable protons (H+) which can participate in the neutralization of a base such as sodium hydroxide (NaOH)
H2SO4(aq) + 2 NaOH → 2 Na+(aq) + 2 H2O + SO42+(aq)
If a sulfuric acid solution is 1 M, then it is 2 N expressed as normality. A 2 N solution is said to be “two normal”.

Molality – similar to molarity, however calculation of molality uses the mass rather than volume of the solvent used, making it temperature independent unlike molarity. Units are m or mol/kg.

NOTE: Fill in at least two values to obtain the result of another by clicking the 'calculate' button. The calculator will autopopulate other fields where possible. Click here to learn more about the molarity calculators.



(M or mol/L)