A mole (symbol: mol) is one of the seven SI base units. It is defined as the amount of substance of a system which contains as many elementary entities as there are atoms in 0.012 kilogram of carbon-12. When the mole is used, the elementary entities must be specified. Entities may be:

See also list of particles, chemistry and physics

Put more colloquially, the mole is a convenient way of counting large numbers of particles. The number defined above ("as many elementary entities . . . ") is known as Avogadro's number, and is approximately 6.02 x 1023. If you are dealing with this many atoms or eggs or artichoke hearts, then you have a mole of atoms or eggs or artichoke hearts. If you have half this number of such entities, then you have half a mole of such entities.

A mole of atoms or molecules is also called a 'gram atom' or 'gram molecule', respectively.

Moles and calculations

Moles are very useful in chemical calculations, as they enable the calculation of yields and other values when dealing with particles of different mass. In this example, moles are used to calculate the mass of CO2 given off when 1g of ethane is burnt. The formula involved is:

3.5O2 + C2H6 → 2CO2 + 3H2O

Here 3.5 moles of oxygen reacts with 1 mole of ethane, to give 2 moles of CO2 and 3 moles of H2O. Notice that the amount of moles does not need to balance on either side of the equation. This is because a mole does not count mass or the number of atoms involved, simply the number of individual particles. In our calculation it is first necessary to work out the number of moles of ethane that has been burnt. The mass of one mole of a substance is defined as being equal to its atomic or molecular mass. The atomic mass of hydrogen is 1g, and the atomic mass of carbon is 12g, so the molecular mass of C2H6 is: 2×12 + 6×1 = 30g. One mole of ethane is 30g. The amount burnt was 1g, or 1/30th of a mole. The molecular mass of CO2 (when the atomic mass of carbon is 12g, and oxygen is 16g) is: 2×16 + 12 = 44g, so one mole of carbon dioxide is 44g. From the formula we know that:

We also know the masses of a mole of both ethane and carbon dioxide, so:

  • 30g of ethane gives off 2×44g carbon dioxide.

It is necessary to multiply the mass of carbon dioxide by two because two moles are produced. However, we also know that just 1/30th of a mole of ethane was burnt. Again:

  • 1/30th of a mole of ethane gives off 2×1/30th of a mole of carbon dioxide.

So finally

  • 30×1/30 g ethane gives off 44×2/30 g of carbon dioxide = 2.93g

From "Mol":

A measure of quantity used in chemistry (pronounced "mole"). Weight is not a very useful unit in chemistry because reactions take place between atoms (two hydrogen atoms and one oxygen atom makes one molecule of water) which have very different weights (one oxygen atom weighs almost 16 times as much as a hydrogen atom). The numbers of atoms in a reaction is also quite useless, because they are simply too large. Just one milliliter of water contains over 30,000,000,000,000 billion molecules.

1 Mol = 6.02214×1023 parts. (This number is called Avogadro's number)

Example: The relative atomic mass of nitrogen is 14u. The rule is "mol times atomic mass equals grams": 1 mol times 14u equals 14 grams

The number of parts in a mol was originally chosen so that 14 grams of nitrogen make up one mol; however, the definition of the mol and the atomic mass unit are currently set such that one mol of carbon, which has a relative atomic mass of 12u, is exactly 12 grams. (This does lead to a small difference, because of Strong interaction)

See also: Mole Day