1. The Wave Nature of Light All forms of NRG/Light have characteristic wavelengths (λ) and frequency (υ). Inversely related λ υ = c (the speed of light) Light visible to the naked eye exists as a tiny portion of the electromagnetic spectrum

2. Electromagnetic Spectrum

3. Max Planck Transfer of energy was not continuous Only came in certain values (quantized) ΔE = hν h = Planck’s constant = 6.626 x 10-34 Js Packets of energy (quantum

4. Albert Einstein Proposed that electromagnetic radiation was quantized and made up of a stream of particles Photons The dual nature of light λ = h/mv (deBroglie equation)

5. Electrons as Waves Louis de Broglie ~1924 Louis de Broglie (1924) Applied wave-particle theory to electrons electrons exhibit wave properties QUANTIZED WAVELENGTHS Standing Wave Second Harmonic or First Overtone Fundamental mode 200 150 100 50 0 - 50 -100 -150 -200 200 150 100 50 0 - 50 -100 -150 -200 200 150 100 50 0 - 50 -100 -150 -200 0 50 100 150 200 0 50 100 150 200 0 50 100 150 200 Adapted from work by Christy Johannesson www.nisd.net/communicationsarts/pages/chem

6. Electrons as Waves QUANTIZED WAVELENGTHS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

7. Electrons as Waves ELECTRONS VISIBLE LIGHT Evidence: DIFFRACTION PATTERNS Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem Davis, Frey, Sarquis, Sarquis, Modern Chemistry2006, page 105

8. Spectrum Light is diffracted (bent) through different objects Hydrogen Spectra gave interesting results

9. Niels Bohr In the Bohr Model (1913) the neutrons and protons occupy a dense central region called the nucleus, and the electrons orbit the nucleus much like planets orbiting the Sun. They are not confined to a planar orbit like the planets are.

10. Bohr Model Planetary model After Rutherford’s discovery, Bohr proposed that electrons travel in definite orbits around the nucleus.

11. Bohr Cont. When the electrons are in their lowest possible NRG level, they are in their ground state. Electrons absorb NRG and go to a higher NRG level(Excited State) NRG (Light) is released when the electron jumps from a higher NRG level to a lower NRG level(Fluorescence) Constantly happening

12. An unsatisfactory model for the hydrogen atom According to classical physics, light should be emitted as the electron circles the nucleus. A loss of energy would cause the electron to be drawn closer to the nucleus and eventually spiral into it. Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 294

13. Quantum Mechanical Model Niels Bohr & Albert Einstein Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).

14. Modern View The atom is mostly empty space Two regions Nucleus protons and neutrons Electron cloud region where you might find an electron

15. Quantum Mechanics Werner Heisenberg ~1926 Heisenberg Uncertainty Principle Impossible to know both the velocity and position of an electron at the same time g Microscope Electron

16. Quantum Mechanics SchrödingerWave Equation(1926) finite # of solutions quantized energy levels defines probability of finding an electron Erwin Schrodinger ~1926 Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

17. Quantum Mechanics Orbital Orbital (“electron cloud”) Region in space where there is 90% probability of finding an electron 90% probability of finding the electron Electron Probability vs. Distance 40 30 20 Electron Probability (%) 10 0 100 150 200 250 50 0 Distance from the Nucleus (pm) Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

18. Quantum Numbers UPPER LEVEL Four Quantum Numbers: Specify the “address” of each electron in an atom Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

19. Quantum Numbers Principal Quantum Number( n) Angular Momentum Quantum #( l) Magnetic Quantum Number( ml) Spin Quantum Number( ms)

20. Quantum Numbers 1. Principal Quantum Number( n) Energy level Size of the orbital Integral values 1s 2s 3s Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

21. Quantum Numbers s p d f 2. Angular Momentum Quantum #( l) Energy sublevel Shape of the orbital Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

22. Angular Momentum Quantum #( l ) Has integral values from 0 to n-1 Related to the shape of the orbital l = 0 is called s l = 1 is called p l = 2 is called d l = 3 is called f

23. Shapes of s, p, and d-Orbitals

24. Shape of f-orbitals

25. Quantum Numbers 3. Magnetic Quantum Number( ml) Orientation of orbital Specifies the exact orbital within each sublevel Has values between l and -l Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

26. Quantum Numbers 4. Spin Quantum Number( ms) Electron spin  +½ or -½ An orbital can hold 2 electrons that spin in opposite directions. Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem