1. The Atom

2. Greek Model
One group of Greek philosophers who lived in the forth and fifth centuries B.C. believed that matter is composed of tiny indestructible particles which they called atoms. However, this theory was never verified by experiments and was discarded. More than 2000 years later, in 1804, John Dalton, an English school teacher, reintroduced the model or theory and developed it to such an extent that it was able to explain the laws of chemical change.

3. Dalton Model
Dalton proposed that elements are composed of identical, indivisible atoms in his “billiard ball model” of the atom.
He devised a system of atomic symbols.
He correctly predicted the Law of Multiple Proportions (if the same elements combine to form different compounds, the ratio of the elements in the compounds are in simple multiples.

4. The discovery of subatomic particles showed that Dalton’s proposal that atoms were indivisible was wrong.

5. J.J. Thomson 1900
Thomson revised Dalton’s theory with his raisin bun or plum pudding model of the atom.
Thomson proposed that an atom could be considered a sphere of positive electricity in which negative electrons are embedded like raisins in a bun.
Most of the mass of the atom in this model is associated with the positive electricity.

6. Rutherford Model
In 1911, Rutherford’s gold foil experiment lead to another revision of the atomic model.
Rutherford proposed that
An atom has a nucleus in which its positive charge (protons) and mass are concentrated.
The vast majority of the atom’s volume would be empty space occupied only by the moving negatively charged electrons.

7. Rutherford also postulated the existence of uncharged particles (neutrons)

8. Problems with Rutherford’s Model
(1) If electrons are not moving, then the attraction of the negative electrons for the positive nucleus should collapse the atom.
(2) If the electrons are moving (to counteract the pull of the nucleus), then the electrons should radiate energy and in time spiral down to the nucleus.

9. Bohr Model 1913
Bohr studied the atomic spectra of the elements and discovered the each spectrum showed a series of lines of definite energies.
Bohr proposed that electrons of specific energy moved in circular orbits around the central atomic nucleus and that electrons could not exist between the orbits.

10. Bohr’s model worked well for hydrogen, but it did not work well for multi-electron atoms.

11. Quantum Mechanical Model 1920’s
This is the most recent model of the atom.
This model supports Bohr for the most part, but suggests that electrons do not exist in fixed orbits or a fixed definite path. Instead, they electrons exist anywhere within an electron cloud.
Determining where an electron will be at any given moment is very difficult and can only be theorized using mathematical equations.

12. The Atom
The model of the atom has evolved from Dalton’s “billiard ball” model to a highly complicated quantum mechanical model. It was first believe that an atom could not be broken down into smaller parts. We now know that it can, mostly by nuclear reactions. The three main sub-atomic particles are the proton, electron, and neutron.

13. Charge and Mass of Subatomic Particles
Mass (kg)
Mass (u)
Proton (p) +1
1.672 x 10-27
Neutron (n)
1.675 x 10-27
Electron -1
9.110 x 10-31

14. The proton and neutron are about equal in mass
The proton and neutron are about equal in mass. These two types of particles are found together in the nucleus (center) of the atom. This explains Rutherford’s gold foil experiment.
The electron is much smaller then the other two particles. It is found in orbits or shells surrounding the nucleus and does not contribute significantly to the mass of the atom.

15. Atomic Number
Atomic number refers to the number of protons in the nucleus of the atom. It is represented by the symbol Z.
Atomic number determines the identity of the element. Every atom of a given element has the same unique number of protons.
In a neutral atom, atomic number is also equal to the number of electrons surrounding the nucleus.

16. Atomic Mass
An atom’s mass is expressed in Atomic Mass Units (u or a.m.u.)
It was impossible for scientists to determine the mass of individual atoms (they are too small) so they assigned relative atomic masses that agreed with the known composition of compounds.

17. Atomic Mass continued
A new unit was developed to mass atoms. Carbon-12 (C-12) was chosen as the reference standard. An atom of C-12 was arbitrarily assigned a mass of 12 atomic mass units.
The masses of all other atoms are compared with the mass of this type of carbon atom. According to this definition, an atomic mass unit is defined as 1/12 the mass of a carbon-12 atom.

18. Mass Number
Mass number is the total number of protons and neutrons in an atom.
The symbol for mass number is A.
For now, we will round the atomic mass to the nearest whole number to get the mass number of an element.

19. Elements and the Periodic Table
All elements can be represented as symbols that are organized in the periodic table.
In addition to the symbols, the periodic table provides us with the atomic number and the mass number of an element.

20. Summary Atomic Number (Z) = # of protons
Mass Number (A) = # of protons + # of neutrons
Neutral Atom: # of protons = # of electrons