1. Chapter 4

2. Everything you ever wanted to know about where the electrons hang out!

3. Building on the Atomic Theory What did Thompson determine? (Think cathode ray tubes!) What did Rutherford’s gold foil experiment prove?

4. Section 1: Early 1900’s Scientists started doing a lot of experiments looking at the absorption and emission of light by matter. Found that there is a relationship between light and an atom’s electrons.

5. Light behaves as a wave Transfer of energy

6. Draw the Wave! Amplitude: height of the wave from the origin to the crest Wavelength ( ) : the distance between the crests (m, cm, nm) Frequency (v): number of waves to pass a given point per unit of time (waves/second = Hz)

7. Slinky Demo

8. An Important Relationship The frequency and wavelength of all waves, including light, are inversely related. As the wavelength of light increases, the frequency decreases.

9. C = v Where: C= speed of light 3.00 x 10 8 m/sec = wavelength (m, cm, nm…) v = frequency (1/sec or sec -1 )

10. What is the frequency of a wave that has a wavelength of 670nm? C = v

11. Electromagnetic Radiation Includes radio waves, radar, microwaves, visible light, infrared light, ultraviolet light, X-rays, and gamma rays

12. Wave Particle Duality http://www.youtube.com/watch?v=DfPeprQ7oGc

13. Sometimes Light Acts Like Particles! What would happen if the frequency of the wave increased so much that you could hardly tell where one wave ended and another began? Light would start acting more like a particle than a wave.

14. Photoelectric Effect Looks at the emission of electrons from a metal when light shines on the metal. Light causes electrons to be ejected from the metal.

15. The Photon Photon- a particle of electromagnetic radiation having no mass, carrying a quantum of energy.

16. Max Plank Objects emit small packets of energy- Quanta Quantum- the minimum quantity of energy that can be lost or gained by an atom. E = hv E = Energy h = 6.626 x 10 -34 Js (Joule x sec) V = frequency (1/sec)

17. What is the energy of a wave that has a frequency of 4.5 x 10 14 Hz? E = hv

18. You should be ready to do the WS…. Let the units be your guide!!!!!

19. So, what happens when photons hit an atom and eject an electron? The electron goes from the ground state to an excited state. As the electron returns to the ground state, it gives off the energy that it gained- LIGHT

20. Energy Levels Energy levels are not evenly spaced Energy levels become more closely spaced the greater the distance from the nucleus

21. Work on you Electrons, Energy and Light Pogil.

22. Warm Up You have 20 minutes to finish up the POGIL.

23. Another Cool Illustration

24. Flame Test Lab The flame you see is orange in color and you determine th wavelength is about 590nm. a. Calculate the frequency. b. Calculate the energy. You will looking at the excitement of electrons of the metals in several ionic solutions

25. What did you really see? The light you saw, was really a combination of all the colors that were produced when the electrons on the metal were excited.

26. Warm Up—Pass it up! Many of you didn’t turn in the Unit 2 Work— Vegium Mole Activity (with work) Mole Problems (the hard sheet) Unit 3 Review Check IC and turn it in by Monday!!

27. Spectral Analysis of Emitted Light from Excited Atoms When the emitted light from excited atoms is passed through a prism, a spectrum of discrete lines of different colors (separate energies) is observed rather than a continuous spectrum of ROY G BIV. Different elements show different line spectra. Line spectra are used to identify the presence of different elements

28. Test Question Spectra

29. Each element has a unique line-emission spectra

30. Emission Spectra Atomic Line Spectrum

32. Interpretation of Atomic Spectra The line spectrum is related to energy transitions in the atom. Absorption = atom gaining energy Emission = atom releasing energy All samples of an element give the exact same pattern of lines. Every atom of that element must have certain, identical energy states

33. Atomic Spectrum Activity

34. Using Atomic Spectral Data Bohr Model Electrons orbit around a nucleus Each orbit has a fixed energy and because of this cannot lose energy and fall into the nucleus Energy Level of an electron is the region around the nucleus where the electron is likely to be moving

35. This helped explain the spectral lines Absorption- the electron gains energy and moves to a higher energy level. Emission- when the electron falls to a lower energy level.

36. The Quantum Model Finally– the truth (as we know it!) Electrons can behave as both waves and particles. Electrons can be considered waves with specific frequencies confined to the space around the nucleus. Electrons can also be considered negatively charged particles.

37. Schrodinger Wave Equation Developed an equation that treated electrons as waves and described the location of electrons. Helped lay the foundation for modern quantum theory (atomic model).

38. Quantum Theory Estimates the probability of finding an electron in a certain position We denote the position of the electron as a “fuzzy” cloud This volume of space where an electron is most likely to be found is called an orbital. The atomic orbitals have distinct shapes