IB Chemistry home > Syllabus 2016 > Stoichiometry > The 'mole' concept

 Syllabus ref: 1.2 Atoms and molecules are pretty small (understatement!) and, as scientists are interested in being able to describe quantities of matter in terms of the mass and number of particles contained per unit mass, this poses a problem. Measuring the mass of individual atoms, we find that one hydrogen atom has a mass of about 1.66 x 10-27g This number is far too small to be useful and so it makes sense to deal with quantities of atoms which can be measured in the laboratory. Amedeo Avogadro

The hydrogen standard

Hydrogen is the smallest atom and it was originally used as the standard by which all the other atoms were compared. It was assigned a value of 1 unit and other atoms masses calculated compared to hydrogen atoms.

The 1H isotope has a mass assigned a value of exactly 1 atomic mass unit. This was the original reference.

 The 1H isotope

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The carbon 12 standard

Nowadays the 12C isotope is used as a reference for comparison of relative atomic masses. This isotope has the assigned mass of 12.00000, all other atoms are measured relative to 12C.

 The 12C isotope

Measured on this scale hydrogen atoms (on average) have a relative mass of 1.00797

Carbon atoms have a relative mass of 12.01111 on average. Although this seems strange at first sight, it is because carbon has several isotopes 12C, 13C, and 14C and the relative mass of carbon is given as the weighted average of all of the isotopes in a naturally occuring sample. Clearly the average must be greater than 12.0000.

Most tables use the relative atomic masses rounded up to one or two decimal places.

Carbon atoms have a mass approximately 12 times that of a hydrogen atom, therefore they have a RELATIVE mass of 12 (there are no units as it is a comparison - see relative measures)

number of hydrogen atoms
number of carbon atoms
mass of hydrogen atoms
mass of carbon atoms
mass of carbon:mass of hydrogen ratio
1
1
1
12
1:12
2
2
2
24
1:12
10
10
10
120
1:12
20
20
20
240
1:12
50
50
50
600
1:12

Provided the number of carbon atoms is equal to the number of hydrogen atoms the mass of carbon is always 12 times the mass of hydrogen.

Clearly there will be a specific number of hydrogen atoms that when weighed have a mass of 1g and that the same number of carbon atoms MUST have a mass of 12g. This number, named after its discoverer is called:

 Avogadro's constant = 6.02 x 1023

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Avogadro's number or constant is the number to which the mass of an atom must be multiplied to give a mass in grams numerically equal to its relative atomic mass.

 Example Hydrogen has a relative atomic mass of 1 therefore 6.02 x 1023 hydrogen atoms have a mass of 1g Carbon has a relative mass of 12 therefore 6.02 x 1023 carbon atoms have a mass of 12g Magnesium has a relative atomic mass of 24 therefore 6.02 x 1023 magnesium atoms have a mass of 24g

This gives rise to two important definitions

• The amount of any substance containing an Avogadro number of particles of that substance is called a mole.
• 1 mole of any substance has a mass equal to its relative mass expressed in grams
 Example 1 mole of magnesium contains 6.02 x 1023 magnesium atoms 1 mole of magnesium has a mass of 24g 12g of magnesium is equivalent to 1/2 moles = 0.5 moles of magnesium 12g of magnesium contains 1/2 moles of magnesium atoms = 0.5 x 6.02 x 1023 = 3.01 x 1023 magnesium atoms

The relationship between moles, mass and number of particles can be expressed by simple formulae:

These formulae can be used to find any quantity when the other two quantities are known.

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