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 A principle stated in 1811 by the Italian chemist Amadeo Avogadro (1776-1856) that equal volumes of gases at the same temperature and pressure contain the same number of molecules regardless of their chemical nature and physical properties. 

This number (Avogadro’s number) is 6.023 X 1023.

 It is the number of molecules of any gas present in a volume of 22.41 L and is the same for the lightest gas (hydrogen) as for a heavy gas such as carbon dioxide or bromine.

Avogadro’s number is one of the fundamental constants of chemistry. It permits one to compare the different atoms or molecules of given substances where the same number of atoms or molecules are being compared.

It also makes possible determination of how much heavier a simple molecule of one gas is than that of another, as a result the relative molecular weights of gases can be ascertained by comparing the weights of equal volumes.

In a chemical reaction, atoms or molecules react in a specific ratio. Let us consider the following examples

Reaction 1     : C + O2 → CO2

Reaction 2     : CH4 + 2 O2  → CO2 + 2 H2O

In the first reaction, one carbon atom reacts with one oxygen molecule to give one carbon dioxide molecule. In the second reaction, one molecule of methane burns with two molecules of oxygen to give one molecule of carbon dioxide and two molecules of water. It is clear that the ratio of reactants is based on the number of molecules. Even though the ratio is based on the number of molecules it is practically difficult to count the number of molecules. Because of this reason it is beneficial to use ‘mole’ concept rather than the actual number of molecules to quantify the reactants and the products. We can explain the first reaction as one mole of carbon reacts with one mole of oxygen to give one mole of carbon dioxide and the second reaction as one mole of methane burns with two moles of oxygen to give one mole of carbon dioxide and two moles of water. When only atoms are involved, scientists also use the term one gram atom instead of one mole.



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