1. Development of Atomic Models

2. Democritus
Greek philosopher
400 BC
“Atomos” concept

3. Can matter can be divided forever?
Eventually, a piece would be “indivisible”
“Atomos,” meaning “not to be cut,” is smallest piece of matter

4. John Dalton (early 1800’s)
Coined the term “atom”.

5. Dalton’s Atomic Theory
Matter made of tiny indivisible particles called “atoms”.
Atoms of one element are alike, and different from atoms of other elements.

6. Page from Dalton’s Journal

7. Compounds form when different atoms combine in fixed proportions.
Chemical reactions involve rearrangement of atoms.
Atoms can’t be created or destroyed, but are conserved in chemical reactions.

8. Dalton’s Atomic Theory called
“Hard Spheres Model”

9. JJ Thomson (1897)

10. Thomson’ Experiments
Studied “cathode rays” (electric current) in a “Crooke’s Tube”.
Fluorescent screen, shows how cathode ray behaved in a magnetic field.
Lets draw a typical Crooke’s Tube in our notes.

11. Cathode Rays were negatively charged
Cathode Ray Tube and Magnet
Cathode Rays were negatively charged
They bent toward (+) plate

12. Cathode Rays were particles
They couldn’t pass through matter.

13. JJ is Awesome
Concluded the negative “cathode ray” particles came from within atoms.
Discovered first subatomic particle (electron).

14. What about the Positive?
But…matter is neutral.
Therefore:
A positive charge must exist to balance the negative.

15. Plum Pudding Model
Atoms are positively charged spheres with negatively charged particles scattered throughout.

16. Yummy… Brian Cox: Thompson and Discovery of Electron

17. Ernest Rutherford (1908)
Physicist who worked in new field of radioactivity.

18. Found 3 Different Types of Radiation
Used magnetic field to isolate three types of radiation.
Alpha (α)
Beta (β)
Gamma (γ)

19. Identify the charge each type of radiation has.
Charges of Radiation
The radiation had different charges.
Identify the charge each type of radiation has.

20. Gold Foil Experiment
Shot alpha particles, at very thin piece of gold foil.
Alpha particles have a positive charge, and a mass of 4 amu
Fluorescent screen shows where the particles went.
Rutherford Gold Foil

21. Observation:
Most alpha particles passed straight
through gold foil.
Conclusion:
Atom’s volume is mostly empty space.

22. small, dense positively charged nucleus.
Observation:
A few alpha particles
deflected at an angle
or bounced back.
Conclusion:
Atoms have a very
small, dense positively
charged nucleus.

23. Deflections happened rarely (1/8000).
Nucleus is extremely small compared to the size of the atom as a whole.
Deflections happened rarely (1/8000).
Modern Example of Gold Foil Experiment in Action

24. The Nuclear Model Rutherford’s Model is called the “Nuclear Model”
Brian Cox: Rutherford and the Nucleus

25. Comparison to Thomson Positively charge only contained in nucleus.
Negatively particles scattered outside nucleus.
Charge is not disbursed evenly.

26. Niels Bohr (1913)
Came up with the “Planetary Model”

27. Bohr’s Theory
Electrons circle nucleus in specific energy levels or “shells”.
The higher the “energy level,” the higher the electron’s energy.

28. Energy Levels
Different energy levels can contain different numbers of electrons.

29. 2n2 = maximum number of electrons an energy level can hold.
How many per level?
n = the number of the energy level
2n2 = maximum number of electrons an energy level can hold.
Ex: Level 3 can hold 2(3)2 = 18 electrons

30. Draw a Bohr Atom Ex: The Fluorine Atom (F) Protons = 9 Neutrons = 10
Electrons = 9
How many energy levels do you draw?
How many electrons in each level?
Human Bohr Model

31. Draw a Bohr Ion
They only difference is that one or more electrons gets added or taken out of the outer energy level.
Ex: The Magnesium Ion (Mg+2)
Protons = 12
Neutrons = 12
Electrons = 10

32. (+) Ions (cations)
(+) ions are smaller
Lost electron(s)

33. (-) Ions (anions)
(-) ions are larger
Gained electron(s)

34. How Did Bohr Come Up With His Model?
Studied the spectral lines emitted by various elements (especially Hydrogen)

35. What are Spectral Lines?
Energy gets absorbed by an atom causing it to emit a unique set of colored lines.
Used to identify what elements are present in a sample. (elemental “Fingerprint”)

36. Spectral Lines are Different for Each Element

37. Answer: 1

38. What Causes Spectral Lines?
Jumping Electrons!!
Video of Line Spectra of Hydrogen

39. Jumping Electrons
Electrons normally exist in the lowest energy level possible called the “ground state”. (stable)
“Ground state” e- configurations are written on the periodic table for each element.
Ex: Aluminum is 2-8-3
Calcium is

40. An Electron Gets “Excited”
Electrons can absorb a photon (or “quanta”) of energy and “jump up” to a higher energy level farther from the nucleus.
This is called the “excited state”. (unstable)

41. Jumping Electrons
They quickly “fall back down” to the ground state. (stable)
They emit a photon (or “quanta”) of energy that corresponds to how far they jumped.
Spectral Lines

42. This photon of energy is seen as a spectral line!
Each spectral line corresponds to a specific photon of energy that is released.
Model Of Hydrogen Atom and Electrons Jumping

43. REMEMBER
Absorb Energy
Jump Up
Emit Energy
Fall Down

44. Excited vs Ground State
Periodic table lists ground state electron configurations for neutral atoms.
To recognize an “excited state” configuration, count the electrons and see if the configuration matches the one on the table.
Ex: = 20 electrons
Calcium (atomic # 20) is
So this must be showing one of the ways calcium could be in the excited state.

45. Valence Electrons Electrons in highest occupied energy level.
Involved in forming bonds with other atoms.
Atoms are most stable when they obtain a “stable octet” of 8 valence electrons
Noble Gases: (Group 18)
Have stable octet already and are “inert” and unreactive
Ex: Argon 2-8-8, Neon 2-8

46. Valence Electrons
Look at the last number in the atom’s electron configuration to determine the number of valence electrons.
Ex:
Al valence
Ca valence
F valence

47. Lewis Dot Diagrams
Shows the number of valence electrons an atom has as “dots” around the atom’s symbol.
Phosphorus is 2-8-5

48. Kernel Nucleus and non-valence electrons
Inner part of atom not involved directly in reactions
Ex:
Al has 10 kernel electrons
and 3 valence electrons

49. The Nature of Light
Study of light has provided important information about the structure of atoms.
Dual Nature of Light:
behaves as both waves and as particles (depending on what type of experiment is being performed.)
Speed of Light: all light waves travel at the same velocity
C = 3.0 x 108 meters/sec

50. Electromagnetic Spectrum
Spectral lines can come from all areas of the EM Spectrum.
Lines of visible colors make up only a small part of the spectrum.

51. Which wave has higher energy?
EM waves carry different amounts of energy based upon their wavelength and frequency.
Wavelength (λ): distance between two peaks of a wave
Frequency (γ): number of peaks that pass per second. (Hertz (Hz) or cycles/sec)
Which wave has higher energy?

52. Relationship of Frequency, Wavelength and Energy of colored line

53. Good Overview Videos Crash Course: History of Atomic Theory
Quantum Mechanics and the Bohr Model

54. Calculating the Energy of a Spectral Line (HONORS)
STEP 1:
If you know the wavelength of the spectral line you can find it’s frequency.
c = λ x ү
c = the speed of light = 3 x 108 meters/sec
λ = wavelength (in meters)
ү = frequency of the wave

55. Calculating the Energy of a Spectral Line (HONORS)
STEP 2:
Using the frequency find the energy of the line (in Joules)
E = h x ү
E = energy in Joules
h = Planck's constant = 6.63 × kg x m2 / sec
ү = frequency of the wave