Development of Atomic Theory400 B.C. - Democritus was first to use the word : atom

atomos meaning “indivisible”

Aristotle (famous philosopher) disputed atoms;

all matter is made up of: earth, air, water, and fire

1808 - Dalton (1st Atomic Theory proposed)i- all matter is made up of small particles

(atoms)ii- all atoms in an element are identicaliii- atoms in different elements are

differentiv- atoms can’t be created/destroyedv- atoms combine in simple whole number

ratios to make compounds

Dalton’s Model was known as the Billiard Ball Modelsince atoms were nondescript spheres.

Points 4 & 5 were proposed from experimental work by:Lavoisier – Law of Conservation of MassProust – Law of Constant Composition

However, points 2 & 4 were determined to be not entirely correct because of the existence of isotopes and subatomic particles, respectively.

1897 – Thomson (Raisin Bun Model)

-negative particles were embedded in sphere of positive charge

-discovered the electron through the use of cathode ray tubes or Crookes tube

-only able to measure charge/mass ratio of electron

Millikan subsequently measured charge and mass of electron from famous Oil Drop Experiment

1909 – Rutherford (Nuclear Model)

Gold foil experiment where a stream of positively charged alpha particles shot at a micro thin sheet of gold foil and pathways detected on coated screen.

Observations:i) most particles (99.99%) went straight through foilii) some particles (0.01%) were slightly deflectediii)a few particles deflected straight back (0.0001%)

Conclusions:i) most of the atom is empty spaceii) something positively charged to deflect “+” alpha particlesiii)dense, positive core in atom to cause massive deflection

nucleus describes this region of atom; also contains neutrons discovered by Chadwick in 1932.

Problem: Where are the

electrons?

1914 – Bohr Model (Planetary Model)

- electrons travel around the nucleus in specific pathways called orbits

- concept of energy levels came from Planck & Einstein who proposed that energy is quantized (specific values) in packets called quanta (photons for light)

- experimental evidence from line spectra of elements supported this. Energy is released (emission) or

absorbed (absorption) by electrons at certain wavelengths.

- electrons travel in these orbits without losing energy (stationary state) but could gain energy and jump into a higher orbit (excited or transition state) or could lose energy and fall to a lower orbit, or the lowest orbit (ground state)

- there is a specific number of electrons that can fit into each energy level or orbit: 2,8,18,32

Problem: Only explained Hydrogen!

Quantum Atomic Theory / Wave Mechanical Model (1924)

Following the groundwork of Bohr, DeBroglie (along with Planck & Einstein) noticed the dual behaviour of electrons as both a form of energy and as a particle of matter.

E = h E = mc2

Heisenberg added to this concept that the position and velocity of an electron could never be simultaneously determined (Uncertainty Principle). All of these concepts/findings led to the development of our current model of the atom proposed by Schrodinger.

Schrodinger’s model of the atom is just a more in depth approach to Bohr’s model, but involves mathematically derived differential equations.

These calculations simply suggest the maximum probability of finding an electron in a given region of space with a

particular quantity of energy. This region is known as an orbital.

Orbitals can have different sizes, shapes, orientations and properties. There are 4 parameters that define the characteristics of these orbitals and the electrons within. These parameters can be defined as quantum numbers and provide the basis of our understanding of chemical bonding.