The concept of atoms began almost 2000 years ago in ancient Greece. Greek philosophers based their conclusions not on evidence but from philosophical reasoning. The concept of atoms remained a philosophical belief until the discovery of two quantitative laws of chemical combination - law of conservation of mass and law of constant proportions.
A molecule is a combination of atoms bonded in a certain way to form a small collection or cluster of atoms that are all hooked together. The atoms in the molecule might all be the same (element) or they might be different (compound).
The Law of Conservation of Mass was established in 1789 by French Chemist Antoine Lavoisier. It states that mass is neither created nor destroyed in any ordinary chemical reaction. Or more simply, the mass of substances produced (products) by a chemical reaction is always equal to the mass of the reacting substances (reactants).
The law of definite proportions, sometimes called Proust's Law, states that a chemical compound always contains exactly the same proportion of elements by mass. An equivalent statement is the law of constant composition, which states that all samples of a given chemical compound have the same elemental composition.
A early nineteenth century English scientist, John Dalton (1766-1844), reasoned that if atoms really exist, they must have certain properties to account for the two laws of chemical combination. He developed the first useful atomic theory of matter around 1803. Some of the details of Dalton's original atomic theory are now known to be incorrect. But the core concepts of the theory (that chemical reactions can be explained by the union and separation of atoms, and that these atoms have characteristic properties) are foundations of modern physical science.
Dalton's atomic theory makes the following assumptions:
All matter consists of tiny particles.
The existence of atoms was first suggested more that 2000 years before Dalton's birth. Atoms remained pure speculation through most of this time, although Newton used arguments based on atoms to explain the gas laws in 1687. (Newton's speculations about atoms in the Principia were carefully copied by hand into Dalton's notebooks.)
Atoms are indestructible and unchangeable.
Atoms of an element cannot be created, destroyed, broken into smaller parts or transformed into atoms of another element. Dalton based this hypothesis on the law of conservation of mass and on centuries of experimental evidence.
With the discovery of subatomic particles after Dalton's time, it became apparent that atoms could be broken into smaller parts. The discovery of nuclear processes showed that it was even possible to transform atoms from one element into atoms of another. But we don't consider processes that affect the nucleus to be chemical processes. The postulate is still useful in explaining the law of conservation of mass in chemistry. A slightly more restrictive wording is "Atoms cannot be created, destroyed, or transformed into other atoms in a chemical change".
Elements are characterized by the weight of their atoms.
All atoms of the same element have identical weights, Dalton asserted. Atoms of different elements have different weights. (Dalton used the word "weight" rather than mass, and chemists have called atomic masses "atomic weights" ever since).
We now know that atoms of the same element sometimes have slightly different masses, but always have identical nuclear charge. In modern atomic theory, the postulate has been amended to read: "Elements are characterized by the nuclear charge of their atoms".
When elements react, their atoms combine in simple, whole-number ratios.
This postulate suggested a practical strategy for determining relative atomic weights from elemental percentages in compounds. Experimental atomic weights could then be used to explain the fixed mass percentages of elements in all compounds of those elements. By suggesting that compounds contained characteristic atom-to-atom ratios, Dalton effectively explained the law of definite proportions.
When elements react, their atoms sometimes combine in more than one simple, whole-number ratio.
Dalton used this postulate to explain why the weight ratios of nitrogen to oxygen in various nitrogen oxides were themselves simple multiples of each other. Even Dalton's critics were impressed by the power and simplicity of his explanation, and it persuaded many of them that his atomic theory was worthy of further investigation. This law is also known as the law of multiple proportions.
Unfortunately, Dalton included an additional postulate that prevented his theory from being accepted for many years. When atoms combine in only one ratio, Dalton said, "..it must be presumed to be a binary one, unless some cause appear to the contrary". He had no experimental evidence to support this postulate, and it lead him to mistakenly assume that the formula of water was OH and the formula of ammonia was NH. As a result, Dalton's atomic weights for oxygen and nitrogen were incorrect and his experimental data did not support many of the conclusions he drew from it.
As chemists tried to deduce formulas for more and more compounds, the flaws in Dalton's atomic weights and in his rule of simplicity became more and more obvious. No one could come up with dependable method of deciding on chemical formulas. Of the three pieces of molecular information - combining weights of the elements, atomic weights of the elements, and molecular formulas - any one could be calculated if the other two were known. Yet only one, the combining weight, was directly measurable.
In 1808, Joseph Gay-Lussac (1778-1850) began a series of experiments with the volumes of reacting gases. He found that two volumes of hydrogen react with one of oxygen to form two volumes of steam; three volumes of hydrogen react with one of nitrogen to yield two volumes of ammonia; and one volume of hydrogen reacts with one of chlorine to produce two volumes of HCl gas. An initial excess of either gas is left over at the end of the reaction.
(This is now known as the law of combining volumes - The ratio between the volumes of the reactant gases and the products can be expressed in simple whole numbers.)
Dalton used Gay-Lussac's data to "prove" that equal volumes of gas do not have equal numbers of molecules, another wrong turn, like his rule of simplicity. The Italian physicist Amedeo Avogadro (1776-1856) saw another path. He began by assuming that equal volumes of gas (at the same temperature and pressure) contain equal numbers of molecules. He also proposed that a molecule was a combination of atoms.
The man who settled the entire issue was the Italian Stanislao Cannizzaro (1826-1910) and a mole was defined as the number of grams of a compound equal to its molecular weight (the weight of a molecule is the sum of the weights of the atoms of which it is made) on Cannizzaro's scale. It was realized that a mole of any compound would have the same number of molecules. Although the value of that number was not then known, it was named Avogadro's number, N, in belated recognition of his contribution.
Accurate determinations of Avogadro's number became possible for the first time when American physicist Robert Millikan measured the charge on an electron in 1910. The charge of a mole of electrons is the constant called the Faraday and had been known since 1834 when Michael Faraday published his works on electrolysis. Using the two the value of the Avogadro constant was determined to be 6.022×1023 mol−1.
1 mole of any substance contains 6.022×1023 particles of that substance and its mass is equal to the relative weight of that substance in grams.
Dalton used symbols to represent his atoms. However, he did not use the same symbols that we use. He used circles with markings to represent the various individual atoms. He used circles with dots, lines, crosses and shading in them. When he ran out of marks he put letters in the circles to represent the elements. Each different symbol represented a different kind of atom--the atom of a different element.
About ten years later, in Sweden, Berzelius suggested just using letters to represent atoms of each element and also to represent the elements in general. These are the symbols that we use today.