1. Chapter 4 - Electrons

2. Properties of Light

3. What is light? A form of electromagnetic radiation: energy that exhibits wavelike behavior as it travels through space

4. Electromagnetic radiation is classified into two types: Non-ionizing Radiation: Transfers energy causing vibrations, electron excitation, and heat Parts of the spectrum: radio, microwaves, infrared, and visible Ionizing Radiation: High energy that ejects electrons and transforms molecules into reactive unstable fragments Parts of the spectrum: UV, X-ray, Gamma Ray

5. Radiation is organized on the Electromagnetic Spectrum

6. Wave Properties Wavelength (λ) : distance between corresponding points on adjacent waves Unit: meters or nanometers Frequency (ν): number of waves that pass a given point per unit time (usually 1 sec) Unit: 1/s = s -1 = Hertz (Hz)

7. Properties of Light How are frequency and wavelength related? c = λν c : speed of light (m/s) c = 3.00 x 10 8 m/s λ : wavelength (m) ν : frequency of wave (s −1 = 1/s = Hz)

8. Max Planck proposed … Energy needed for electrons to move was quantized, or a quantum of energy is needed to move an electron. E = hν E : energy of light emitted (J) h : Planck’s constant (J·s) h = 6.626 × 10 −34 J·s ν : frequency of wave (s −1 = 1/s = Hz)

9. E = h Electrons can only move with a quantum of energy. Every color represents a different amount of energy released.

10. Louis de Broglie proposed … that electrons be considered waves confined to the space around an atomic nucleus. If electrons exist as waves, then they must have specific frequencies If they have specific frequencies, they also have specific energies So energy is quantized!

11. ELECTRON BEHAVIOR

12. Who made this model of the atom?

13. Electrons are always moving! The farther the electron is from the nucleus, the more energy it has. Electrons can change energy levels and emit a photon of light. Photon: packet of energy E photon = hν

14. Electron Movement Vocabulary Ground State: electrons in the lowest possible energy level Excited State: electrons absorb energy and move to a higher energy state Emission : when an electron falls to a lower energy level, a photon is released Absorption : energy must be added to an atom in order to move an electron from a lower energy level to a higher energy level How much energy must it absorb? Quantum: amount of energy needed to move between energy levels

15. Electron Movement

17. When electrons move from the excited state to the ground state they emit a photon of light The light emitted by an element is viewed as a bright line emission spectrum Each band of light represents the energy released by an electron when it moves from higher to lower energy Line Emission Spectrum of Hydrogen Atoms

19. Because electrons move so much… We cannot know an electrons exact speed and position at the same time Heisenberg Uncertainty Principle:

20. Schrodinger Wave Equation Equation that describes the wave & particle nature of an electron Don’t know the exact position… Predict where electron will be most of the time Where 90% of the e - density is found for the 1s orbital

21. Electron Location

22. DO NOW How would you describe the location of your home? Like giving your house an address, each electron in an atom is assigned an address or location inside the atom. Remember: Do electrons ever stop moving?

23. We do not know where the electrons are, but we can describe where they might be

24. The address of all the electrons in an atom is called an Electron Configuration Sodium #11 1s 2 2s 2 2p 6 3s 1

25. First Location: ENERGY LEVEL n = 1, 2, 3, 4 …

26. ENERGY LEVEL SUB- LEVEL

27. Second Location: SUBLEVEL These are the probability regions called: s, p, d, f

29. SUB- LEVEL ORBITAL

30. Energy 1s 2s 3s 4s 5s 2p 3p 4p 3d Third Location: ORBITAL s = 1p = 3d = 5 f = 7

31. Sublevels on the Periodic Table

32. Examples: What is the electron configuration of Mg? What is the electron configuration of Cl?

33. How to fill in an Orbital Diagram

34. Energy Step 1 – Aufbau Principle: Electrons are represented as arrows and drawn one at a time. Electrons start at the lowest possible energy. 1s 2s 3s 4s 5s 2p 3p 4p 3d Element: Hydrogen Number of Electrons: Electron Configuration:

35. Energy  Paired electrons Step 2 – Pauli Exclusion Principle: Electrons have opposite spins. Draw one up and one down. 1s 2s 3s 4s 5s 2p 3p 4p 3d Element: Helium Number of Electrons: Electron Configuration:

36. Energy Step 3 – Hund’s Rule: Electrons occupy equal energy orbitals so that a maximum number of unpaired electrons results Draw electrons in each orbital first then pair them.  Unpaired electrons 1s 2s 3s 4s 5s 2p 3p 4p 3d Element: Oxygen Number of Electrons: Electron Configuration:

37. What is the electron configuration of S? Energy 1s 2s 3s 4s 5s 2p 3p 4p 3d Element: Sulfur Number of Electrons: Electron Configuration:

38. What is the electron configuration of Sr? Energy 1s 2s 3s 4s 5s 2p 3p 4p 3d Element: Strontium Number of Electrons: Electron Configuration:

39. Valence Electrons

40. Why is location important? location influences behavior

41. Valence Electrons valence electrons: electrons in the outermost energy level can be lost, gained, or shared electrons that are responsible for reactions Elements in a group have similar properties because they have the same number of valence electrons

42. How many valence electrons in each group? Only the main block elements! 1 23 4 567 8

43. All electrons under the outer most level Inner Core Electrons