Laws of Chemical Combinations

Full Notes:

1.5.1 Law of Conservation of Mass

This was one of the first laws established in chemistry, based on the work of Antoine Lavoisier in 1789.

Statement of the Law:
“Matter can neither be created nor destroyed.”

In other words, during any chemical reaction, the total mass of the reactants equals the total mass of the products.

Example:

BaCl₂ + H₂SO₄ → BaSO₄ + 2HCl
If we weigh the total mass before and after the reaction (in a closed system), it remains exactly the same.

This law applies strictly only to ordinary chemical changes, not to nuclear reactions where mass-energy conversions occur.

1.5.2 Law of Definite Proportions

Also known as Proust’s Law, after Joseph Proust (1799).

Statement of the Law:
“A given compound always contains the same elements in the same proportion by mass.”

This law tells us that no matter how a compound is prepared, or from where it is obtained, the ratio of the masses of its elements remains constant.

Example:

Two samples of pure water – collected from different sources – both contain hydrogen and oxygen in the mass ratio of 1:8.

18 g of water → 2 g hydrogen, 16 g oxygen
This 1:8 ratio is constant, no matter the source.

Proust demonstrated this using compounds like copper carbonate, showing consistent mass proportions regardless of source.

Note: This law is not valid for non-stoichiometric compounds like FeO, which may vary in composition due to defects in structure.

1.5.3 Law of Multiple Proportions

Formulated by John Dalton in 1803, this law extends the idea of definite proportions to situations where two elements form more than one compound.

Statement of the Law:
“If two elements combine to form more than one compound, the different masses of one element that combine with a fixed mass of the other element are in the ratio of small whole numbers.”

Example:

Carbon forms two oxides:

• CO (carbon monoxide): 12 g carbon combines with 16 g oxygen
• CO₂ (carbon dioxide): 12 g carbon combines with 32 g oxygen

The mass of oxygen combining with a fixed mass of carbon (12 g) is in a simple ratio: 16:32 = 1:2

This law strongly supports the atomic theory, since atoms combine in simple whole number ratios.

1.5.4 Gay Lussac’s Law of Gaseous Volumes

Proposed by Joseph Louis Gay Lussac in 1808, this law deals specifically with gaseous reactions.

Statement of the Law:
“When gases react together, they do so in volumes which bear a simple whole number ratio to each other and to the volumes of the products (if gaseous), all measured under the same conditions of temperature and pressure.”

Example:

2 volumes of hydrogen + 1 volume of oxygen → 2 volumes of water vapour
(Measured at same T and P)

The volume ratio is 2:1:2, which is a simple whole number ratio.

This observation was crucial in supporting the idea of discrete, countable molecules.

However, this law could not be explained by Dalton’s atomic theory at the time, since it implied division of atoms, which Dalton’s model didn’t allow. The explanation came later through Avogadro’s hypothesis.

1.5.5 Avogadro’s Law

Proposed by Amedeo Avogadro in 1811, this law resolved the issues that Dalton’s theory couldn’t explain – especially Gay Lussac’s volume observations.

Statement of the Law:
“Equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules.”

Molecules could be diatomic (like H₂, O₂), which wasn’t considered in Dalton’s atomic theory.

NCERT Implication:
Using this law, Avogadro concluded that:

• 1 molecule of oxygen must split and combine with 2 molecules of hydrogen

So water is formed as: 2H₂ + O₂ → 2H₂O
This supports both the volume ratio and the molecular model

This idea became foundational in determining molecular formulas and in the development of Avogadro’s number (6.022 × 10²³) later on.

Matt’s exam tip

Most students memorise the five laws but fail to apply them. Make sure you understand which laws relate to mass vs volume and recognise that Gay Lussac’s law applies only to gases

Summary

• Conservation of mass states total mass remains constant in chemical reactions.
• Definite proportions fix element mass ratios in a compound regardless of source.
• Multiple proportions give simple whole number mass ratios across related compounds.
• Gay Lussac’s law gives simple whole number volume ratios for reacting gases.
• Avogadro’s law links equal gas volumes to equal molecule numbers at the same conditions.