1. The Modern Atomic Model After Thomson: Bohr, Placnk, Einstein, Heisenberg, and Schrödinger

2. Electromagnetic Radiation All forms move at a constant speed, 3.0  10 8 m/s. 3.0  10 8 m/s is considered the speed of light through air. It is represented by the variable c.

3. Waves Waves are repetitive in nature Wavelength variable:

4. Waves The number of waves that pass through a specific point over a period of time is called the frequency (variable: ) of the wave, measured in Hertz (Hz) The relationship between and  c =

5. Electromagnetic Spectrum The electromagnetic spectrum is a chart showing the frequencies and wavelengths of all forms of electromagnetic radiation. electromagnetic spectrum A closeup of the visible light portion:

6. Photoelectric Effect The photoelectric effect refers to emission of e - from a metal when light shines on the metal

7. Photoelectric Effect There is a minimum energy, called the work function (  ), below which no e - are emitted, regardless of how long light is shined on the metal.work function (  ), Classical theory suggested that a metal that is continuously bombarded by waves should eventually eject an electron. Max Planck’s suggestion: Objects emit energy in small, specific amounts called quanta. A quantum is the minimum amount of energy that can be gained/lost in an atom.

8. Light as a Particle Relationship between quantum and the frequency of radiation: E = h Where h = 6.626  10 34 J  s (Planck’s constant)

9. Light as a Particle In 1905, Einstein described the dual wave/particle nature for electromagnetic radiation. A massless particle carrying a quantum of energy – photon Therefore, E photon = h If the energy of a photon is less than the work function for a metal, electrons will not be emitted.

10. Electron Energy States Ground State – lowest energy state of an atom Excited State – state to which electron moves when it has greater energy than the ground state An electron moving from the excited state to the ground state gives off energy in the form of a photon

11. Photons and Electrons E photon = E excited - E ground The diagram to the right shows different energy values for various energy levels Note: the electronvolt (eV) is the unit used here  1 eV = 1.602  10 -19 J

12. Hydrogen Atom Line Emission Spectra When a beam of colored light is passed through a prism, the beam is separated into several frequencies of visible light. Collectively, these frequencies are called the emission spectrum.

13. Electron Energy States E excited and E ground are fixed values for every identical atom of a given element. Silicon

14. Line Emission Spectra for Selected Elements

15. Line Emission Spectra for Hydrogen

16. Bohr Model of the Atom The Bohr Model explained the spectral lines for hydrogenThe Bohr Model It did NOT explain: –other atoms’ spectral lines –the chemical behavior of atoms

21. Electrons as Waves 1924 – Louis de Broglie suggests electrons are waves confined to space around the nucleus Therefore, electrons can only exist at certain frequencies, which correspond to specific energies, those of Bohr’s orbits.

22. Heisenberg Uncertainty Principle How does one find the electron “waves” in an atom – strike it with a photon! However, once the photon strikes the electron, its energy is changed, and therefore its path is changed. You can never know both the velocity (speed and direction) and position of an electron simultaneously. In fact, the more you know about position, the less is known about velocity (and vice versa).

23. Schrödinger Equation Solutions to the equation, called wavefunctions (  ), can only be specific values. Treats quantization as a natural consequence, not an assumption Electrons (in all probability) exist in 3-D regions around the nucleus called orbitals