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Atoms and molecules faithfully obey certain sets of laws that enable the study of chemistry (or indeed other science subjects) in a systematic way. Below listed some of the basic laws that lay down the foundation for the study of chemistry.

Dalton's Atomic Theory - This is a set of fundamental rules about the nature of atoms of which all other rules are based on. It was put forth by John Dalton (1766-1844), an English scientist.
(1) Matter consists of definite particles called atoms.
(2) Atoms are indestructable. They can rearrange in chemical reactions but they do not themselves break apart.
(3) Atoms of a particular element are indistinguishable from one another. They are all identical in mass, as well as other properties.
(4) Atoms of different elements (or types) differ in mass (and other properties).
(5) When atoms of different elements combine to form compounds, new and more complex particles (molecules) are formed. Their constituent atoms are always present in a definite numerical ratio.

The theory breaks down when going beyond chemical reactions. For example, when large amount of energy is introduced (much larger than a typical chemical energy), atoms will break apart. For example, this occurs in particle colliders or 'atom-smashing' machines that crack atomic nuclei, releasing even more fundamental constituents.

Law of Conservation of Mass - No detectable gain or loss in mass occurs in chemical reactions. However, the state of a substance may change in a chemical reaction. For example, substances involving in a chemical reaction can change from solid states to gaseous states but the total mass will not change. Note that the energy released (exothermic) or adsorbed (endothermic) in a chemical reaction is a result of energy transfer between atoms and their environment.

Law of Definite Proportion - The elements in a given compound are always combined in the same proportion by mass. This law forms the basis for the definition of a chemical compound. For example, a water molecule (H2O) consists of two hydrogen atoms each of relative mass of 1 and one oxygen atom of relative mass of 16 (rounded to nearest integer number). By putting a sensible unit measurement this means that there are 2 g of hydrogen and 16 g of oxygen in a sample of 18 g of water. The ratio is 1 to 8. Thus, a sample of, say, 51.435 g of water always contain (51.435 x 1/9) 5.715 g of hydrogen and (51.435 x 8/9) 45.720 g of oxygen. The ratio, again, is 1 hydrogen to 8 oxygen. This rule applies for water found anywhere in the universe and the mass proportion is always the same for any given unit measurement (kilogram, pound etc.).

Law of Multiple Proportion - Whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other element are in the ratio of small whole numbers. Take an example of two mineral samples iron pyrite (FeS2) and iron troilite (FeS). Both contain iron and sulfur atoms. However, for a given fixed amount of iron it requires exactly twice the mass of sulfur needed to make pyrite than that of troilite with the same amount of iron.

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