1. Laws of Definite Proportions & Multiple Proportions
PreAP Chemistry

2. Law of Definite Proportion
Put forth by Joesph Louis Proust ( ) in 1800
States that different samples of the same compound always contain its constituent elements in the same proportion by mass regardless of the source.
It is an intensive property.
Example CO and CO2
CO = O/C = 1/1
CO2 = O/C =2/1
CO to CO2 = 1:2

3. Calculation Mass Fraction = Mass of element Mass of Compound
Mass Percent = Mass of element x 100
Mass of element in a sample =
Mass of compound x Mass of element in a sample Mass of Compound

4. Law of Multiple Proportions
Derived by John Dalton ( ) in between
States if two elements A and B combine to form more than one compound, the masses of B that can combine with a given mass of A are in a ratio of small whole numbers.
It is an intensive property.
Example – H2O and H2O2

5. Ratio of O mass with a constant mass of Sn
Example
Compound
Mass of Tin (Sn) in g
Mass of Oxygen (O) in g
Ratio of O mass with a constant mass of Sn
Tin (II) Oxide SnO
119 g
16 g 1
Tin (IV) Oxide SnO2
32 g 2

6. Difference between Laws of multiple proportions and definite proportions
Both laws have to do with relating to Dalton's Atomic Theory. The only difference is that the Law of Definite Proportions deals with elements combining to form ONE compound in a simple whole number ratio. The
Law of Multiple Proportions is comparing the same 2 elements that make up 2 different compounds, the division of these 2 ratios should equal a simple whole number ratio.
For example: Carbon and oxygen can combine to form carbon monoxide and carbon dioxide. If you calculated each compounds ration of oxygen to carbon you would get the following ratios: compound A would equal a combining ratio of 1.34:1 (O:C). Compound B would equal a combining ratio of 2.67:1 (O:C).
If you divided the bigger ratio by the smaller ratio you would have that oxygen combines with a ratio of 2.67/1.34 which would equal 1.99:1, which is close enough to 2:1.