1. Atom Model History

2. Democritus Fifth century B.C. Greek philosopher
All matter composed of indivisible and indestructible particles called atoms (Greek for uncuttable).
Though later challenged, many ideas agreed with later theory

3. Dalton’s Atomic Theory
Early 1800’s - Billiard Ball Model 
Viewed the atom as a small, solid, indivisible sphere.
Atoms of each element identical in mass and properties.
Atoms of one element differ from another atom. 
Got the "ball" rolling for modern chemistry!

5. J.J. Thomson Late 1800’s-Plum Pudding Model
Atom was a sphere of positive electricity which was diffuse with negative particles imbedded throughout
Discovered the electron
Experiments that passed electric currents through gases at low pressure
Attracted to + charge so particles of current must be –
Nobel Prize in physics in 1906.

6. What he did…..

7. Rutherford Solar System Model
Discovered the nucleus- Gold Foil Experiment
Atom mostly empty space with dense positively charged nucleus surrounded by negative electrons. 
Nobel Prize in chemistry in 1908

8. What this shows… flying dots = Alpha particles( +) emitted
green wall = detecting screen (where particles hit)
yellow window in the middle = gold foil
as you see, 1/8000 alpha particles are repelled backwards, that's because the nucleus is VERY small, only few alpha particles hit it completely and repel back.
Thus Rutherford knew that the nucleus was super duper small!

9. Bohr
Modified Rutherford model- why electrons don’t crash into the nucleus if + attract –
Energy levels - electrons traveled in circular orbits and that only certain orbits were allowed. 
Close to nucleus- lower energy level e- occupies
Far from nucleus- higher energy level e- occupies
Electron in one energy level or another not in between – like a ladder
Electrons gain or lose energy by changing energy level
Difference in energy between levels (to move e-) = quantum of energy
Nobel Prize in physics in 1922
Like rungs on a ladder and potential energy as you rise, you increase potential energy can’t stand between a rung
Can’t have continuous range of values but certain definite values

11. Electrons and Light Light- moving waves and particles
Frequency- occurrences per time (1/s or s-1)
(v measured in Hz or cycles/second)
Speed- of light ( c ) x 10 8 m/s
Wavelength- distance wave repeats – 2 consecutive peaks (λ measured in m)
Stream of particles – energy Is determined by light’s frequency To remove e-, particle of light needs min. energy and a min frequency

12. Electromagnetic Spectrum – broad range of wavelengths
Visible Spectrum: 400 nm (violet) to 700 nm (red)
Red: low frequency/ long wavelength
Violet: high frequency/ short wavelength
The product of frequency and wavelength always equals a constant, c, the speed of light
C = λv
Frequency and wavelength are inversely proportional

13. Light provides info… Ground State – lowest possible energy
Excited State - electron absorbs energy & moves levels
A quantum of energy in the form of light is emitted when the electron drops back to a lower energy level – an abrupt step.
The light emitted by an electron from a higher to lower energy level has a frequency directly proportional to the energy change of the electron.
Or….

14. E = h x f Where E = energy measured in Joules (J) h = Planck’s constant in J/Hz (J-s) = 6.63 x J-s f = frequency of light measured in Hz If, c = λv Then f = c λ And, E = h c

15. Quantum Models
How likely to find an electron in various locations – probability
Uncertainty Principle:
-exact position and momentum of e- is unknown; instead look at probability … think of blades on an airplane propeller or ceiling fan
Quantum Number – specifies the properties of electrons
Electron Configuration- specific dispersal of electrons among subshells (or sublevels)

16. Follow 3 principles:
Aufbrau- electrons occupy lowest energy orbital first
Pauli Exclusion- orbital may describe at most two electrons
Hund’s Rule – orbital of a sublevel fill up by a single electron before pairing

17. Electron Capacity
2n2 where n = quantum number designation and indicates energy level
Quantum Number (n)
Shell Capacity (2n2 ) 1 2 8 3 18 4 32

18. The periodic table is structured so that elements with the same type of valence electron configuration are arranged in columns.

19. The left-most columns include the alkali metals and the alkaline earth metals. In these elements the valence s orbitals are being filled
On the right hand side, the right-most block of six elements are those in which the valence p orbitals are being filled
These two groups comprise the main-group elements
In the middle is a block of ten columns that contain transition metals. These are elements in which d orbitals are being filled
Below this group are two rows with 14 columns. These are commonly referred to the f-block metals. In these columns the f orbitals are being filled

20. Quantum Number Quantum # Principle Angular Momentum (Azimuthal)
Magnetic
Spin
Symbol n l ml ms
Possible Values
1,2,3,4,…
s, p, d, f
… -1,0,+1…
+,- 1/2
Characteristics
Size and energy level *
*bigger number, higher energy level
n=1, ground state
Shape
Orientation in space
Magnetic spin

21. Electron Configuration Details
Address of Atom
Periodic Table is the Map
Ions-Isoelectronic
Cations  Previous noble gas as core
Anions  next noble gas as core
Paramagnetic-one or more unpaired electrons
Diamagnetic – all electrons paired