History of the Atom

440 BC Leucippus and his pupil Democritus originate the atom concept, stating:

1. There exist indivisible particle called atoms

“a” meaning not; “tomos” meaning cut

2. There exists empty space between atoms

3. Atoms are completely solid

4. Atoms are homogeneous, with no internal structure

5. Atoms differ in size shape and weight (the last criteria, weight was added by democritus)

1789 Antoine Lavoisier discovers the law of conservation of mass

1800’s Joseph Proust observed that specific substances always contain elements in the same ratio by mass – the law of definite proportions

1808John Dalton proposes his theory of atoms, stating:

1. Each element is made up of tiny particles called atoms

2. The atoms of a given element are identical

3. Chemical compounds form when atoms combine with each other AND a given compound always has the same relative numbers and types of atoms.

4. Chemical reactions involve reorganization of the atoms – changes in the way they are bound together. The atoms themselves are not changed in a chemical reaction.

Dalton’s theory thus explains the laws of conservation of mass and energy

1810 Dalton states a second law – the Law of Multiple Proportions – the ratio of masses of one element that combine with a constant mass of another element can be expressed in small whole numbers. I.e. atoms react as whole units

1897 JJ Thomson works with cathode (negatively charged) rays and discovers the electron

1900 Max Planck states his quantum hypothesis

1904 Thomson develops the “plum pudding” model of the atom, described as a sphere of positive electricity (the pudding) with electrons (the bits of plum) scattered in it.

1906 Ernest Rutherford announced that alpha particles can be scattered by air and goes on to discover the presence of the nucleus

1913 Robert Millikan determines the charge on the electron

1913 Rutherford, along with Niels Bohr, proposes the planetary model of the atom

Late 1910’s Thomson, Millikan and their associates continue their work with the atom and show that there are also positive particles in the atom – with the same amount of electrical charge as an electron. These became known as protons

1920 Rutherford predicts the presence of a third particle, but evidence is not found until 1930, by Walter Bothe.

1924 Wolfgang Pauli states the quantum exclusion principle

1926 Werner Heisenberg discovers that it is impossible to accurately predict both the position and the momentum of any object (including an electron) at the same time – the Heisenberg uncertainty principle

1926 Erwin Schrodinger, building on the work of Loius de Broglie, tries to overcome the uncertainty principle by considering the electron’s behavior as like that of a wave rather than a particle

1926 Max Born, using the work of Shrodinger and deBroglie applies the Pauli exclusion principle to identify the position of an electron in a cloud surrounding the nucleus

1932 James Chadwick repeats Bothe’s work and finds high energy particles with essentially the same mass as a proton and no charge – the neutron

Planck’s quantum theory – energy is given off in packets rather than continuously. He called these packets quanta. Quanta of radiant energy (ex. Solar) are often called photons

In developing the planetary model of the atom, Bohr theorized that electrons release energy as they move from one energy level to another. The energy levels further from the nucleus have the highest energy. The energy level closest to the nucleus is the ground state.

PAULI EXCLUSION PRINCIPLE

Identifies four quantum numbers to describe the position occupied by an electronin the atom.

n – the principle quantum number, the energy level occupied by an electron the total number of electrons in an energy level can be represented by 2n²

l – defines the shape of the orbital, the subshells s, p, d, and f correspond to the values 0, 1, 2 and 3

m- describes the number and orintation of the orbitals within a subshell, can have a value from –1 to +1

s – identifies the direction of the spin of each of the two electrons in the orbital, +1/2 or –1/2

Dalton’s theory of the solid atom was disproved by the discovery of subatomic particles. These particles help to identify and differentiate the elements. The number of protons determines the identity of the element. Dalton’s idea that all atoms of a given element was also disproven as examples of the same element with different mass – isotopes – were discovered. The number of neutrons determines the particular isotope of that element

A nuclide is described as a particular kind of atom with a definite number of protons and neutrons. The particles in the nucleus of any nuclide are the nucleons. The nucleons are represented by the mass number, while the number of protons which identifies the element is represented by the atomic number.