Short Answer
The law of multiple proportions states that when two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other element are in simple whole-number ratios. This law was proposed by John Dalton and strongly supports the atomic theory.
For example, carbon and oxygen form carbon monoxide (CO) and carbon dioxide (CO₂). For a fixed mass of carbon, the mass of oxygen in CO₂ is exactly twice the mass of oxygen in CO. This simple ratio (1:2) shows how elements combine in definite, whole-number proportions.
Detailed Explanation
Law of Multiple Proportions
The law of multiple proportions is a key principle in chemistry that explains how elements combine to form different compounds. Proposed by John Dalton in 1803, this law states that if two elements form more than one compound, then the masses of one element that combine with a fixed mass of the other element will always be in simple whole-number ratios.
This law is one of the strongest pieces of evidence supporting Dalton’s atomic theory, which says that atoms combine in whole numbers and cannot be divided into smaller pieces during chemical reactions.
Meaning of the law
To understand the law clearly:
- Take two elements, say element A and element B.
- They may combine to form more than one compound.
- If the mass of element A is kept constant, the masses of element B that combine with A will form simple ratios like 1:2, 2:3, 3:4, etc.
This shows that atoms combine in fixed, whole-number ratios and cannot form fractions of atoms.
Example 1: Carbon and oxygen
Carbon (C) and oxygen (O) form two common compounds:
- Carbon monoxide (CO)
- 12 g carbon combines with 16 g oxygen
- Carbon dioxide (CO₂)
- 12 g carbon combines with 32 g oxygen
If we keep carbon fixed at 12 g, the oxygen masses are:
- CO → 16 g
- CO₂ → 32 g
The ratio of oxygen masses = 16 : 32 = 1 : 2
This simple whole-number ratio shows the law of multiple proportions.
Example 2: Nitrogen and oxygen
Nitrogen (N) and oxygen (O) form several compounds:
- NO
- NO₂
- N₂O
- N₂O₃
- N₂O₅
For simplicity, compare NO and NO₂:
- In NO: 14 g nitrogen combines with 16 g oxygen
- In NO₂: 14 g nitrogen combines with 32 g oxygen
Oxygen ratio = 16 : 32 = 1 : 2
Again, a simple whole-number ratio.
Example 3: Hydrogen and oxygen
Hydrogen and oxygen form:
- Water (H₂O)
- Hydrogen peroxide (H₂O₂)
If we fix hydrogen mass:
- In H₂O → hydrogen:oxygen = 2 g : 16 g
- In H₂O₂ → hydrogen:oxygen = 2 g : 32 g
Oxygen ratio = 16 : 32 = 1 : 2
This fits the law perfectly.
Why the law supports atomic theory
Dalton’s atomic theory states:
- Matter is made of small, indivisible atoms.
- Atoms combine in simple whole-number ratios.
- Atoms of each element are identical.
Because elements combine through whole atoms, compounds can only form in simple integer proportions. This explains why oxygen in CO and CO₂ has a 1:2 mass ratio. Fractional atoms cannot exist in chemical reactions, so ratios must be whole numbers.
Importance of the law of multiple proportions
The law is important because:
- It distinguishes compounds from mixtures
Compounds follow fixed whole-number ratios; mixtures do not. - It supports the idea of atoms
Simple ratios are possible only if atoms combine in whole numbers. - It helps establish chemical formulas
For example, the difference between CO and CO₂ is explained by this law. - It helps determine empirical and molecular formulas
Mass ratios lead to exact formula determination. - It shows consistency in chemical behaviour
This consistency supports fundamental chemical principles.
Connection with other chemical laws
The law of multiple proportions is connected to:
- Law of definite proportions
A compound always has the same composition. - Law of conservation of mass
Atoms rearrange but are not created or destroyed.
Together, these laws form the foundation of modern chemistry.
Applications in chemistry
This law is used in:
- Chemical formula determination
- Identifying unknown compounds
- Understanding chemical bonding
- Analysing reaction patterns
- Calculating percent composition
The concept is also essential in quantitative chemistry.
Conclusion
The law of multiple proportions states that when two elements form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios. This law supports atomic theory, helps distinguish compounds, and is essential for determining chemical formulas. It explains why compounds like CO and CO₂ differ and shows that atoms combine in definite, whole-number proportions.