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

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. ( 2 4 He ) Beta Particle (β)- High speed electron ( 0 -1 e ) Gamma Rays (γ)-High energy electromagnetic wave moving at the speed of light.

8. RUTHERFORD GOLD FOIL EXPERIMENT Alpha particles http://www.youtube.com/watch?v=Q8RuO2ek NGw

9. GOLD FOIL EXPERIMENT  Most particles passed through foil without deflection. Indicated that the atom is mostly empty space.  Some particle 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-1900- 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 light.

17. Photon energy E photon = E i – E f E i = initial level E f = final level unit: eV (electronvolts) or J (joules) 1 eV= 1.6 x 10 -19 J

18. Ex- An electron falls from the 2 nd energy level to ground state of a hydrogen atom. If it gives off a photon, what is the energy of the photon? E photon = -3.4eV − -13.6eV = 10.2ev E photon = (10.2eV)(1.6 x 10 -19 J/eV) E photon = 1.63 x 10 -18 J

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

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-1924- 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 High probability Low 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 c 2 E- energy (Joules) c- speed of light m- mass (kg) 3 x 10 8 m/s Ex) What is the energy content of a 70kg student? E = 70kg(3x10 8 ) 2 E = 6.3 x 10 18 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 10 -19 J/eV) E = 1.056 x 10 -19 J b) Then Use E=mc 2 : 1.056 x 10 -19 J = m (3x10 8 m/s) 2 m= 1.17 x 10 -36 kg

26. Ex: What is the energy content of a carbon atom in eV and Joules? a)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 = 1x10 6 eV E= 11172MeV(1x10 6 eV/MeV) E = 1.1172 x 10 10 eV c)1 eV = 1.6 x 10 -19 J E = 1.1172 x 10 10 eV( 1.6x10 -19 J/eV) E = 1.78 x 10 -9 J