1. HISTORY OF THE ATOM JJ Thomson Neils Bohr Earnest Albert Rutherford
Einstein
HISTORY OF THE ATOM

2. JJ Thomson- Plum Pudding Model
Electrons are scattered throughout a positively charged body.

3. DISCOVERY OF RADIATION
Radioactive Material- unstable material that emits high energy electromagnetic waves or particles when broken down.

4. Henri Becquerel-1896- discovered radiation is emitted from uranium.

5. Marie Curie- isolated polonium and radium
“Neither of us could foresee that in beginning this work we were to enter the path of a new science which we should follow for all our future.”

6. Ernest Rutherford-1911- determined that there are 3 types of radiation
Alpha Particles (α)- Helium nucleus. ( 24He ) Beta Particle (β)- High speed electron (0-1e ) Gamma Rays (γ)-High energy electromagnetic wave moving at the speed of light.

8. RUTHERFORD GOLD FOIL EXPERIMENT
-Alpha particles

9. GOLD FOIL EXPERIMENT
Most particles passed through foil without deflection. Indicated that the atom is mostly empty space.
Some particles were deflected backwards. Indicated that the atom has a central (+) mass.
Few deflections occurred at large angles. Indicated that the nucleus is small and dense.
Electrons don’t have enough mass OR charge to change the path of the alpha particle

10. RUTHERFORD MODEL OF THE ATOM
Small, dense, positively charged nucleus.
Electrons are located outside the nucleus.

11. GOLD FOIL LIMITATIONS
Didn’t identify entire mass. (1932 Chadwick found the neutron)
Didn’t explain why the electrons aren’t attracted to the nucleus, which would cause the atom to collapse.
Didn’t account for light emission as electrons orbit.
Didn’t account for different spectrums of light produced by different atoms.

12. Maxwell Planck Proposed that light is an electromagnetic wave that forms bundles of energy called photons. This brought about the idea of the dual nature of light and led to further investigation of the atom

13. Neils Bohr 1927
Developed the planetary model of the atom where the electrons orbit the nucleus in fixed positions without radiating light.

14. Bohr Model The (+) nucleus and (-) electrons give the atom its energy.
As distance from the nucleus increases, more energy is needed to hold electrons in orbit.
Atoms emit a photon when an electron drops to a lower energy level.
Energy of the emitted photon is equal to the change in energy levels.

15. Ionization energy- the energy needed to move an electron to ground state. The (-) value indicates the electron energy is controlled by the nucleus.
On reference table

16. If electrons move to a higher energy level, they absorb a photon.
If they drop to a lower energy level they give off energy in the form of a light wave.

17. Photon energy Ei = initial level Ef= final level Ephoton = Ei – Ef
unit: eV (electronvolts) or J (joules)
1 eV= 1.6 x 10-19J

18. Ex- An electron falls from the 2nd energy level to ground state of a hydrogen atom. If it gives off a photon, what is the energy of the photon?
Ephoton = -3.4eV − eV = 10.2ev
Ephoton = (10.2eV)(1.6 x J/eV)
Ephoton = 1.63 x J

19. Spectrum- the light pattern given off when a photon is emitted.
Lyman Series- electron moves to ground state. (ultraviolet range)
Paschen Series- Electron falls to the 3rd energy level
( Infrared range)
Balmer Series- electron falls to 2nd energy level.(visible light)

20. Bohr Model Limitations
Limited electrons to specific levels.
Didn’t explain how the electrons could have centripetal acceleration without extra energy.
Goes against Newton’s Laws

21. DeBroglie Model Showed that the electrons orbit the nucleus in a wave pattern. Doesn’t account for electrons between orbitals

22. Cloud Model- Erwin Schrodinger- instead of levels, electrons are located in areas of high or low probability called energy states
Low
Probability
High
probability

23. Albert Einstein-1905
Proposed that the mass of a body is the measure of its energy content. This could help explain model limitations and showed a mass discrimination in the atom.

24. Einstein’s energy equation
E = m c2 E- energy (Joules) c- speed of light m- mass (kg) 3 x 108 m/s celeritas (c) - extreme swiftness Ex) What is the energy content of a 70kg student? E = 70kg(3x108)2 E = 6.3 x 1018 J

25. Ex) What is the mass of a subatomic particle having 0.66eV of energy?
a) First convert eV to joules: E= 0.66eV(1.6 x J/eV) E = x J b) Then Use E=mc2 : x J = m (3x108m/s)2 m= 1.17 x kg

26. Ex: What is the energy content of a carbon atom in eV and Joules?
Atoms are measured in AMU (atomic mass units)
1AMU= 931MeV
Carbon has a mass of 12 AMU
E= 12 AMU(931MeV/AMU)
E = 11,172 MeV

27. b) 1 MeV = 1x106eV
E= 11172MeV(1x106eV/MeV)
E = x 1010 eV
1 eV = 1.6 x J
E = x 1010 eV( 1.6x10-19 J/eV)
E = 1.78 x 10-9J