1. Atomic Theory

2. The Atom
Democritus
Believed the smallest indivisible and indestructible particle was the atom

3. The Atom
Atom: the smallest particle of an element that retains the properties of that element.
John Dalton
Studied ratios of element
combinations

4. Dalton’s Atomic Theory
All elements are composed of indivisible particles called atoms.
Atoms of the same elements are identical, and atoms of different elements are different.
Atoms of different elements can physically mix or chemically combine with one another in simple whole number ratios to make compounds.
Chemical reactions occur when atoms are separated, joined, or rearranged. Atoms of one element never change into another element.

5. The Atom
Dalton’s Atomic Theory was found to be incorrect in several aspects;
1. Sub atomic particles – atoms divisible
2. Isotopes – same elements,
different masses
3. Radio active decay –
atoms changing

6. The Atom
J.J. Thomson
Applied an electric current through a gas
The current had a positive end (anode) and a negative end (cathode)
When the gas was charged, a beam of light formed.
A cathode ray= a stream of negatively charged particles
The negatively charged particles were named electrons

7. One of the first cathode ray tubes

8. The Atom E. Goldstein Used Thomson’s cathode ray
Found rays traveling in the opposite direction (called canal rays) to be positively charged
These particles were called protons

9. The Atom James Chadwick
Found an atomic particle that had mass, but no charge
These particles were called neutrons

10. Review of Subatomic Particles
Symbol
Charge
Mass
Actual Mass
Electron e- 1-
1/1840
9.11x10-28
Proton p+ 1+ 1
1.67x10-24
Neutron no

11. Atomic Model

12. Dalton’s model of the atom
Dalton – called the atomic model because the atom was the smallest particle

13. Thomson’s Model
proposed electrons were stuck in a lump of positively charged material.
This is called the plum-pudding model because that is what it is thought to look like, (personally, I like the chocolate chip cookie dough model, but that’s me

14. Positive Charge Negative Charge
Plum Pudding…
Chocolate chip cookie dough!!!

15. Rutherford’s Model of the Atom
Ernest Rutherford
Used alpha particles α (He without the 2 e-’s)
Shot toward a piece of gold foil which it passed through with only slight deflection.
Found the Atom is mostly empty space!
It is the mass and + charge in the nucleus that deflects α particles

16. Rutherford’s Gold Foil Experiment

17. Rutherford’s model of the atom
Rutherford identified the mass and positive charge as a nucleus
The electron’s surround the dense nucleus
Most of the atom is
empty space

18. Draw a model of the atom…
If your model looks like a target,the model you have drawn came from Niels Bohr…

19. Bohr’s Model of the Atom
Niels Bohr ( )
Suggested in 1913 that electrons orbit around the nucleus
This was called the PLANETARY MODEL
Electrons cannot fall into the nucleus due to fixed energy levels (they do not lose energy)

20. The Nature of Light

21. Properties of Light
Wave characteristics
Particle characteristics

22. Light as a wave
Wave properties

23. Amplitude - height of the wave.
Wavelength - distance between any two corresponding points on successive waves (m)
Amplitude - height of the wave.
frequency - number of waves that pass a point in space during any time interval (1/s or s-1)
Wave equation
c = f
c = speed of light (3.0x108 m/s)
 = wavelength (m)
f = frequency (1/s or s-1)

24. White light is refracted (bent) or passed through a prism
Separates the light by color – continuous spectrum

25. Light as a particle
Light is a form of energy that is packaged into photons
Energy from light is both absorbed and given off
Equation for energy of a photon
E = hf E = energy of a photon (J)
h = planck’s constant (6.626x10-34J*s)
f = frequency (1/s or s-1)

26. Electromagnetic Spectrum

27. Spectrums Two types of light spectrums Continuous Bright line
Bright line spectra
States there are mandatory energy levels for atoms
Different atoms have different energy levels, therefore give off different light combinations
Fingerprint of all atoms, elements, compounds

28. Bigger jump (n=1 to 4), more energy, higher frequency, shorter wavelength
Smaller jump (n=2 to 3), less energy, lower frequency, longer wavelength

29. Quantum Mechanical Model
Erwin Schrödinger
Created a mathematical equation to tell the location and energy of an electron in an atom
This model does not define a specific path the electron travels. Rather, it estimates the probability of finding an electron in a given position.

30. Quantum Mechanical Model
Electrons are represented by a fuzzy cloud, (the darker the cloud, the higher the probability of finding an electron in that position)

31. The Atom

32. Atomic Symbols C  Element symbol Atomic #  6 12 13 39 238
 Mass #  12
C  Element symbol
Atomic #  6
- Calculate the number of subatomic particles:
- p+ = atomic #
- e- = atomic #
- n = atomic mass – atomic #
E.G.: 6 C , C , 9 K , U

33. Isotopes Isotopes - Alternate form of an element Same Different
# of p # of n
Atomic # Mass #
Element
- Often isotopes will be identified by their atomic mass

34. Atomic Mass Calculating Atomic Mass
- Weighted average of the masses of all isotopes of an element
E.G.: Carbon
C12 + C13 + C14 =
C C12 has most abundance: the amu will be closest to 12
Calculating Atomic Mass
Step 1: Multiply the % abundance/100 of each isotope by its atomic mass
Step 2: Add the values from step 1 together
12.011 6

35. Electron Configuration

36. Energy Levels-Review Like a ladder, the lowest ring=lowest energy
Electron can jump from one level to another
Quantum= amount of energy needed to change levels
To change energy levels, electrons must gain or lose the right amount of energy

38. Electron Orbitals
Orbitals are where electrons reside (think of them as the electron’s homes)
These are considered principal energy levels
sublevel # of orbitals Hold # of e-‘s shape
s sphere
p dumbbell
d lobes
f lobes

39. Rules Orbital diagram follows an elements electron configuration
Aufbau Principle: electrons enter orbitals of low energy first.
Pauli Exclusion Principle: Only two electrons per orbital (represented as a line), and when paired electrons (represented as arrows) have an opposite spin due to repulsion (up and down arrows)
Hund’s Rule: If multiple orbitals are present (sublevels p, d, and f) each orbital in that sublevel needs one electron before electrons start pairing.

40. Electron Configuration/Orbital Examples
Electron configuration: 1s1
Orbital diagram: ___
Electron configuration: 1s2 2s22p2
Orbital diagram: __ __ __ __ __ 
Electron configuration: 1s2 2s22p63s23p64s23d7
  Orbital diagram: __ __ __ __ __ __ __ __ __ __ __
__ __ __ __