1. Development of Atomic Theory 400 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 (1 st Atomic Theory proposed) i- all matter is made up of small particles (atoms) ii- all atoms in an element are identical iii- atoms in different elements are different iv- atoms can’t be created/destroyed v- atoms combine in simple whole number ratios to make compounds

2. Dalton’s Model was known as the Billiard Ball Model since atoms were nondescript spheres. Points 4 & 5 were proposed from experimental work by: Lavoisier – Law of Conservation of Mass Proust – 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.

3. 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

4. 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.

5. Observations: i) most particles (99.99%) went straight through foil ii) some particles (0.01%) were slightly deflected iii)a few particles deflected straight back (0.0001%) Conclusions: i) most of the atom is empty space ii) something positively charged to deflect “+” alpha particles iii)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?

6. 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)

7. - experimental evidence from line spectra of elements supported this. Energy is released (emission) or absorbed (absorption) by electrons at certain wavelengths.

8. - 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!

9. 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 = mc 2 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.

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

12. 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.

13. 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.