1 / 37

Chapter 5

Chapter 5. Electrons In Atoms. Electromagnetic Radiation. Any energy that moves in waves. Wavelength ( λ ). Frequency ( ν ) – the number of times a wavelength will pass a given point in a given time. The Speed of Light in a Vacuum. c = λν c is the speed of light It is a constant

levia
Download Presentation

Chapter 5

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 5 Electrons In Atoms

  2. Electromagnetic Radiation • Any energy that moves in waves. Wavelength (λ) • Frequency (ν) – the number of times a wavelength will pass a given point in a given time.

  3. The Speed of Light in a Vacuum • c = λν • c is the speed of light • It is a constant • 3.00 x 108 m/s • This is an inverse relationship • Meaning if λ goes up ν goes down • Therefore, if you have a large wavelength you will have small frequency. If you have a large frequency you will have a small wavelength.

  4. The Sun’s Light • How long does sunlight take to get here? • 8 minutes • How can I show calculations for this? • Distance = (rate)(time) • t = d/r • t = d/r (distance = 1.45 x 1011 m • t = 1.45 x 1011 m / 3.00 x 108 m/s • t = 483 seconds or 8 minutes

  5. The Electromagnetic Spectrum • Pg. 120 • Very little is visible • What makes white light? • A combination of all visible light • ROYGBIV • What gives each color its color? • Its wavelength • What seperates white light? • Prism • How does it work? • Refraction – bending of light • What is different about each color of light that allows it to be bent differently?

  6. Things to Know • λ is calculated in meters. • ν is calculated in s-1, 1/s, or Hertz (Hz) • They all mean the same thing • Math skills • x/y ÷ x = 1/y or y-1 • λ is called lambda • v is called gnew

  7. Pg. 121 Example Problems • λ = 4.90 x 10-7 • ν = ? • c = λν • v = c/λ • v = 3.00 x 108 m/s / 4.90 x 10-7 m • v = 6.12 x 1014 s-1 • You do #4 • 3.17 m • This illustrates how your radio works. You tune to the frequency of the wave.

  8. Planck– Study of Energy • Plank • He saw that light was emitted from heated objects • He thought that energy was emitted in small amounts • He called these energy packets quantums • A quantum is the minimum amount of energy lost or gained by an object • Equantum = hv • h is Planck’s constant • 6.626 x 10-34 J · s

  9. Einstein – Study of Energy • He explained the photoelectric effect • Light reflects when shined on a metal surface and photoelectrons are emitted • He is credited for the discovery of photons and how they behave. • Photon • Packet of light energy • Have both wave and particle characteristics • Ephoton = hv • Einstein’s work explaining the photoelectric effect also explained why electrons behave the way they do. • They have wavelike characteristics and are tiny bundles of energy which gives them particle characteristics

  10. Color of Gases • Different gases emit different colors • It is because of their electrons • What about their electrons makes them emit different colors? (Apply what you know about different colors of light) • Each gas emits a certain color because it’s electrons jump up and down at specific distances. • Electron behaving like a wave • Electrons have a mass • Electrons have both wave and particle characteristics

  11. Pg. 124 • 5b • v = 9.50 x 1013 Hz • What is the energy of a photon if wavelength is 1.10 x 10-12 m?

  12. Homework • Pg 147 65-76

  13. Atom Models and Scientists • Bohr • Showed atom at ground state (lowest energy level) • As the electron gained energy it moved to a higher state • His model only worked for H • His model did not account for the chemical behavior of atoms.

  14. Bohr Models

  15. de Broglie • Predicted that all moving particles have wave characteristics • de Broglie equation • λ = h / mv

  16. Heisenberg Uncertainty Principal • States that it is impossible to know the position and velocity of a moving object at the same time. • Example • You are driving down the road and telling someone your exact location. • When you tell them exactly where you are you are no longer there.

  17. Schrodinger Wave Equation • Allows you to calculate probability density function • Allows you to see where an electron might be but not where it is

  18. Electron Probability Density Function • The greater the density of dots, the greater the likelihood of finding an electron.

  19. Electron Energy Levels and Primary Shapes • Principle Energy Levels • Principle Quantum Number • 1st definitional location of electrons • 1, 2, 3, … • Energy Sub Levels • s, p, d, and f • Each sublevel holds a specific number of electrons • s – 2, p – 6, d – 10, f – 14 • Each sublevel has a primary shape • S – spherical, p – dumbbell, d – planer, f – don’t need to know • Electrons build from the inside to outside

  20. Matryoshka dolls are put together similar to how electrons fill energy levels.

  21. Pg. 873 Section 5.1 # 1-4

  22. What gives each proton its uniqueness? • Protons • Why are neutrons important? • Stability • Why are electrons important? • Bonding

  23. How can the shapes of energy sublevels be determined by the electron density function?

  24. Atoms build out • Levels nest inside of other levels • As you move out, energy increases

  25. Pauli Exclusion Principle • States that a maximum of two electrons may occupy any sublevel; paired electrons must have opposite spin • Degenerate levels • S – 1, p – 3, d – 5, f – 7 • Each degenerate level can hold a maximum of two electrons. • s has one degenerate level and can hold 2 electrons. p has 3 degenerate levels and can hold 6 electrons.

  26. Hund’s Rule • Each degenerate level will receive one electron before any level receives the 2nd electron of opposite spin. • Example: p orbital electrons

  27. Two Ways to Denote Electrons in an Atom • Orbital diagram • Electron configuration

  28. Each principal energy level contains the sublevels but it has to be a big enough atom in order to contain certain sublevels. • 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, …

  29. Orbital Diagrams • Oxygen contains 8 e- • Example: C • You do: F

  30. Electron Configuration • Oxygen has 8 e- • 1s2, 2s2, 2p4 • Example: C • You do: F

  31. Aufbaw Diagram • Pg 138 (in yellow) • Device to show how electrons fill

  32. Examples • Orbital diagram and electron configuration for Cl. • You do Selenium.

  33. Noble Gas Shorthand • You will use the noble gas that precedes the element. • Ex. Selenium • [Ar] 4s2 3d10 4p4 • What are valence electrons? • Electrons in the outermost energy level • Determined by looking at the highest principle energy level • How many valence electrons does Selenium have? • 6 • 4s2 and 4p4

  34. Lewis Structures • aka Electron dot structures • Rules for illustrating Lewis Structures • Determine the number of valence electrons • Add one valence e- to each position before adding the second. • Example: Iodine • You do: Sulfur

  35. Homework • Pg. 147 78-81

More Related