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Electron Structure of the Atom

This chapter delves into the principles of electromagnetic (EM) radiation, exploring its properties such as wavelength, frequency, and speed. It emphasizes the critical mathematical relationships governing EM radiation, including the equation that relates energy, frequency, and wavelength. The text also discusses the Bohr model and modern atomic theories, including orbitals and electron configurations. Practical examples illustrate how to calculate frequency and energy for different wavelengths, highlighting the significance of valence electrons in chemical properties of elements.

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Electron Structure of the Atom

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  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 3.00×108 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 (3.00×108 m/s)

  8. Mathematical Relationships Ephoton=hυEphoton=(hc)/λ υ = Frequency of the light (1/s, or Hz) λ = Wavelength of light (nm or m) c = CONSTANT, Speed of light (3.00×108 m/s) h = Planck’s Constant (6.626×10-34 J×s) Ephoton = Energy of a single photon (J)

  9. Example • Assume we want to determine the frequency of orange light and the energy of a single photon of this light. • Orange light = 600 nm = 6.00×10-7 m • υλ = c, therefore υ= c/λ • = 5.00×1014 Hz • Ephoton=hυ=(6.626×10-34J×s)(5.00×1014Hz) • Ephoton=3.31×10-19 J

  10. PROBLEM • Calculate the frequency and photon energy for an X-ray of wavelength 1.00 nm. • X-Ray= 1.00 nm = 1.00×10-9m • υλ = c, therefore υ= c/λ • = 3.00×1017Hz • Ephoton=hυ=(6.626×10-34J×s)(3.00×1017Hz) • Ephoton=1.99×10-16J

  11. PROBLEM • What color is laser with a frequency of 6.0×1014 Hz? • therefore • = 5.00×10-7 m = 500 nm • 500 nm = Green Light

  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 quantitizedenergy 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 1s2 2s2 2p6 3s1

  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 1s2 2s2 2p6

  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.

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