1. Electron Structure of the Atom Chapter 7

2. 7.1 Electromagnetic Radiation and Energy

3. Electromagnetic Radiation EM Radiation travels through space as an oscillating waveform. EM Radiation travels through a vacuum at a constant speed of 3.00×10 8 m/s

4. Properties of EM Radiation Wavelength ( λ, measured in nm) Frequency ( υ, measured in Hertz, Hz)

5. Relationship between λ and υ

6. Electromagnetic Spectrum

7. Mathematical Relationships υλ = c υ = Frequency of the light (1/s, or Hz) λ = Wavelength of light (nm or m) c = CONSTANT, Speed of light (3.00×10 8 m/s)

8. Mathematical Relationships E photon =hυ E photon =(hc)/ λ υ = Frequency of the light (1/s, or Hz) λ = Wavelength of light (nm or m) c = CONSTANT, Speed of light (3.00×10 8 m/s) h = Planck’s Constant (6.626×10 -34 J×s) E photon = Energy of a single photon (J)

9. Example

10. PROBLEM

12. Continuous vs. Line Spectra

13. 7.2 The Bohr Model of the Hydrogen Atom

14. Bohr Model of the Atom Propsed by Niels Bohr Explains the Emission Spectrum of Hydrogen Relies of quantitized energy levels. Does not work for atoms with more than one electron.

15. 7.3 The Modern Model of the Atom

16. Orbitals and Orbits Bohr’s model had electrons orbit in tight paths, but this only worked for Hydrogen. Schrödinger expanded the model by using 3 dimensional orbitals

17. Energy Levels and Orbital Shape Electrons are still in quantitized energy levels. Orbitals of roughly the same size are in the same overarching, or principal, energy level. There are four ground state orbital geometries: s, p, d and f.

18. Naming Orbitals Orbitals are named for their principal energy level and their orbital geometry. The n=1 principal energy level has only one geometry, s. The n=2 principal energy level has two geometries, s and p. n=3 is composed of s, p, and d n=4 is composed of s, p, d and f.

19. Orbital Geometries

20. Orbital Diagrams

21. Rules for Filling in Orbitals Ground State Atoms have the same number of electrons as protons. Aufbau Principle – Start with the lowest energy level. Pauli Exclusion Principle – Max of two electrons in each orbital with opposite spins Hund’s Rule – Electrons are distributed in orbitals of the same energy as to maximize the number of unpaired electrons.

22. Example Sodium p= 11 e= 11

23. PROBLEM Carbon

24. PROBLEM Titanium

25. Electron Configurations Orbital diagrams are informative but take a lot of space. Electron Configurations are a shorthand for these diagrams. Though they convey the same information, they do not show sublevel organization.

26. Example Sodium p= 11 e= 11 Na 1s 2 2s 2 2p 6 3s 1

27. PROBLEM Nitrogen

28. PROBLEM Iron

29. 7.4 Periodicity of Electron Configuration

30. Periodic Table

31. Another Way to Look at It

32. 7.5 Valance Electrons in the Main Group Elements

33. Main Group Elements

34. Valance Electrons Valance Electrons are those electrons in the last filled principal energy level. Core Electrons are those below the valance level. Valance Electrons for Main Group Elements are those in the highest s and p orbitals. Main Elements in the same group have the same number of valance electrons.

35. 7.6 Electron Configurations for Ions

36. Example Sodium ion p= 11 e= 10 Na 1s 2 2s 2 2p 6

37. Ion Electron Configurations Ion charges are as they are due to the role of orbitals. Ions are stable at 1+, 2+, or such because that gets the electron configuration to a completed principal energy shell (for main group elements). Na (1+) is isoelectronic with Neon (a completed n=2)

38. 7.7 Periodic Properties of Atoms

39. Valance Electrons and Chemistry Valance electrons are the ones participating in chemical reactions. Compounds are stabilized by reaching a filled principal energy level. We will return to this next chapter.

40. Ionization Energy Ionization Energy, the amount of energy required to remove en electron from an gaseous atom (kJ/mol) The lower the ionization energy the more reactive a compound is.