1. How did we discover electron arrangement in an atom? ELECTROMAGNETIC RADIATION ! ! !

2. Waves  Repeated disturbance through a medium (air, liquid) from origin to distant points.  Medium does not move  Ex. Ocean waves, sound waves

3. Characteristics of WavesCharacteristics of Waves  Wavelength  Distance between 2 points within a wave cycle  2 peaks  Frequency  # of wave cycles passing a point for a particular time unit  Usually seconds.

5. Wavelength and frequency are inversely proportional.

6. c = νλ c = speed of light, 3.0 x 10 8 m/s  Constant ν = frequency (s -1 or Hz) λ = wavelength (m)

7. Example 1:Example 1:  Find the frequency of a green light that has a wavelength of 545 nm.

8. Electromagnetic WavesElectromagnetic Waves  Produced from electric charge movement  Changes within electric and magnetic fields carried over a distance  No medium needed

9. Electromagnetic SpectrumElectromagnetic Spectrum  Contains full range of wavelengths and frequencies found with electromagnetic radiation  Wavelength/frequency changes cause color changes  Mostly invisible, visible range (390 nm -760 nm)  Different materials absorb/transmit the spectrum differently.

11. Types of SpectraTypes of Spectra  What is a spectra?  Spectrum– white light/radiation split into different wavelengths and frequencies by a prism  Continuous spectrum  No breaks in spectrum  Colors together  Line spectrum  Line pattern emitted by light from excited atoms of a particular element  Aided in determining atomic structure

12. Line SpectrumLine Spectrum  Pattern emitted by light from excited atoms of an element  Specific for each element  Only certain wavelengths of visible spectrum present  Used for element identification

15. Flame TestsFlame Tests  Some atoms of elements produce visible light if heated  Each element has a specific flame color  Examples: Li, Na, Cs, Ca

16. A Bit of Quantum Theory……

17. Max PlanckMax Planck  1900  Related energy and radiation  E = h ν  h= 6.626 x 10 -34 J  s (Planck’s constant)  E = energy per photon (J)  Quantum ---smallest amount of energy  Atoms can only absorb/emit specific quanta

18. Albert EinsteinAlbert Einstein  1905  Added to Planck’s concept  Photons—  Bundles of light energy  Same energy as quantum  E = h ν (energy of photon)  Photons release energy and electrons gain energy  Threshold frequency– minimum amount of energy needed by photon to extract electron

19. THEREFORE ………THEREFORE ………  Light is in the form of electromagnetic waves  Photons can resemble particles  Gave raise to the possibility of thinking about wave AND particle qualities of subatomic particles (electron)

20. Example 1Example 1 Calculate the energy found in a photon of red light with a wavelength of 700.0 nm

21. Example 2Example 2 How much energy (in joules) is found in the radiation of the hydrogen atom emission spectrum with a 656.3 nm wavelength?

22. Example 3:Example 3:  A sodium atom emits yellow light with a wavelength of 589 nm when it is excited. Find the energy per photon of this light.

24. Coulomb’s LawCoulomb’s Law  Describes the attractive force between negative electrons and positive nucleus.  Force is directly related to the charge of electron and nucleus  Force is inversely related to distance between particles F = q e x q p r 2 (IE…an electron’s energy is dependent on distance from nucleus)

25. Early Models of the Atom Bohr  1913—hydrogen atom structure  Physics + quantum theory  Electrons move in definite orbits around the positively charged nucleus— planetary model  Does not apply as atoms increase in electron number

26.  Electrons orbit nucleus in different energy levels  Lower energy levels, closest to nucleus (n = 1)  Higher energy levels increase electron’s distance from nucleus  Electrons can “transition” or jump between energy levels through photons  Gain/absorb photon—higher energy level  Lose/emit photon—lower energy level Bohr ModelBohr Model

28. Energy States in an AtomEnergy States in an Atom  Atoms can gain or loss energy.  Specific energy states within an atom.  Can be counted  Ground State = lowest energy state  Excited State = higher energy level than ground, gained energy

30. So, where does the Bohr Model fit in?  Electrons orbit around the nucleus at different energy levels/orbits.  Electron’s energy level = orbit level where electron is located.  Light absorption = electron moves from a state of low energy to high energy. “becomes excited”  Light Emitted = electron falls from an “excited” state of energy to a lower energy level.

31. Ex. LiEx. Li

32. Erwin SchrödingerErwin Schrödinger  Quantum mechanics  1926---wave equation  Electrons behave more like waves than particles

33. Heisenberg’s Uncertainty Principle  Electron’s location and direction cannot be known simultaneously  Electron as cloud of negative charge

34. Modern Model of the Atom The electron cloud  Sometimes called the wave model  Electron as cloud of negative charge  Spherical cloud of varying density  Varying density shows where an electron is more or less likely to be

35.  Treats electron’s location as wave property  Defined by quantum numbers  Quantum numbers  Provide information about size, shape, and orientation of atomic orbitals  Define atomic orbitals from general to specific Quantum TheoryQuantum Theory

36.  Determines orbital size and electron energy  Same as “n” value/orbital in Bohr model  Positive whole number, NOT 0  Shells – orbitals with same value  n = 1, 2, 3, 4, etc. Principal Quantum Number (n)

37.  Defines orbital shape for a particular region of atom  Think as “subshell”  l = n-1  # of orbitals/subshells = principal quantum # Orbital Angular Momentum Quantum Number (l)

38. l Orbital/Subshell 0s 1p 2d 3f

39.  Describes orbital orientation within an atom  Range from –l to +l, 0 is possible  m l = 0, ± 1, ± 2, etc.  m l = 2l + 1 (number of orientations) Magnetic Quantum Number (m l )

40.  2p  4f How do you specify orbitals?

41.  s orbital  1 possible orbital orientation, spherical shape  n value determines size  Charge cloud found near center, likely electron location  p orbital  3 possible orbital orientations, dumbbell shape  p X, p y, p z Orbital ShapesOrbital Shapes

45. PP What does atomic structure REALLY look like?

46.  Describes the motion of an electron, spinning  As electron moves, magnetic field induced  Electrons with opposite spins, cancel magnetic field of other  Values: +1/2, -1/2 Electron SpinElectron Spin

47.  Read lab procedure  Read pp. 267-289 (for Friday) Homework