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  1. Why is the sky blue?

  2. Chemistry Chapter 11 The 1998 Nobel Prize in Physics was awarded "for the discovery of a new form of quantum fluid with fractionally charged excitations." At the left is a computer graphic of this kind of state. Arrangement of Electrons in Atoms

  3. The Puzzle of the Atom • Protons and electrons are attracted to each other because of opposite charges • Electrically charged particles moving in a curved path give off energy • Despite these facts, atoms don’t collapse (Rutherford’s Model did not explain why! We need a new model!)

  4. What is light? Light is made up of fluctuating electric and magnetic fields that do not require the existence of matter. Visible light is a kind of electromagnetic radiation (form of E with wavelike behavior as travels through space.) All form of light travel at a speed of 3.0 x 108 m/s

  5. Wave-Particle Duality JJ Thomson won the Nobel prize for describing the electron as a particle. His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron. The electron is a particle! The electron is an energy wave!

  6. Confused??? You’ve Got Company! “No familiar conceptions can be woven around the electron; something unknown is doing we don’t know what.” Physicist Sir Arthur Eddington The Nature of the Physical World 1934

  7. Electromagnetic radiation propagates through space as a wave moving at the speed of light. (ALL FORMS!) c =  C = speed of light, a constant (3.00 x 108 m/s) = frequency, in units of hertz (hz, sec-1)  = wavelength, in meters

  8. Spectroscopic analysis of the visible spectrum… …produces all of the colors in a continuous spectrum

  9. Types of electromagnetic radiation:

  10. Light as Particles • 1900’s German physisist Max Planck studied emission of light by hot objects. • Suggested objects emitted in small specific amounts called quanta. • Quantum: the minimum quantity of E that can be lost or gained by an atom. • Proposed E = hv • h is “Planck’s Constant” 6.626 x 10-34 Js

  11. The energy (E ) of electromagnetic radiation is directly proportional to the frequency () of the radiation. E = h E= Energy, in units of Joules (kg·m2/s2) h= Planck’s constant (6.626 x 10-34 J·s) = frequency, in units of hertz (hz, sec-1)

  12. Photoelectric Effect (early 1900’s) • Could NOT be explained by the wave theory of light. • Refers to the emission of e’ from a metal when light shines on it. • For a given metal, no e’ were emitted if the light was below a certain v • Wave theory said any v could supply enough E to eject any e’

  13. What are wave-like and particle like properties???

  14. Albert Einstein (1905) • Suggested dual wave-particle nature • Each particle carried a quantum of E. • Einstein called these particles “photons” • E photon = hv • For each e’ to be removed, must be struck by single photon with minimum E required to knock e’ loose. • e’ of different metals require different minimum v to exhibit photoelectric effect.

  15. Hydrogen-Atom Line Emission Spectrum • When current is passed through a gas at low P, the potential E of some of the gas atoms increases. • Ground state: lowest E state of an atom • Excited state: atom has higher potential E than the ground state. • When at atom returns to ground state, it gives off E. (neon lights) • Results in Hydrogen’s Line Emission Spectrum.

  16. Spectroscopic analysis of the hydrogen spectrum… …produces a “bright line” spectrum

  17. Electron transitionsinvolve jumps of definite amounts ofenergy. This produces bands of light with definite wavelengths.

  18. The Bohr Model of the Atom I pictured electrons orbiting the nucleus much like planets orbiting the sun. But I was wrong! They’re more like bees around a hive. WRONG!!! Neils Bohr

  19. What does a line emission spectrum tell you? • Proves that e’ are constantly moving from one E level to another. • Proves that the distance between the two levels is always the same. • Use • Identification of unknown samples

  20. Bohr’s Model of H Atom • Linked an atom’s e’ to photon emission • Said that e’ can circle the nucleus only is specific paths or orbits. • Orbits are separated from the nucleus by large empty spaces (nodes) where the e- cannot exist. (Rungs on a ladder) • Remember, E increases as e’ are furtherfrom the nucleus.

  21. Bohr’s Model of H Atom • While an e’ is in orbit, it cannot gain nor lose E • It can, move to a higher E level IF it gains the amount of E = to the difference between the higher E orbit and the lower energy orbit. • When falls, it emits a PHOTON = to the difference. • Absorption gains E/Emission gives off E

  22. Take the good with the bad (Bohr’s Model) • Good: • Led scientists to believe a similar model could be applied to all atoms. • Bad: • Yikes!!! Did not explain the spectra of atoms with more than one e’ or chemical behavior of atoms!

  23. The Wave-like Electron The electron propagates through space as an energy wave. To understand the atom, one must understand the behavior of electromagnetic waves. Louis deBroglie

  24. Schrodinger Wave Equation Equation for probability of a single electron being found along a single axis (x-axis) Erwin Schrodinger

  25. The electron as a standing wave: • Standing waves do not propagate through space • Standing waves are fixed at both ends

  26. An orbital is a region within an atom where thereis a probability of finding an electron. This is a probability diagram for the s orbital in the first energy level… Orbital shapes are defined as the surface that contains 90% of the total electron probability.

  27. Heisenberg Uncertainty Principle “One cannot simultaneously determine both the position and momentum of an electron.” . … Werner Heisenberg

  28. Sizes of s orbitals Orbitals of the same shape (s, for instance) grow larger as n increases… Nodes are regions of low probability within an orbital.

  29. The s orbital has a spherical shape centered around the origin of the three axes in space. s orbital shape

  30. P orbital shape There are three dumbbell-shaped p orbitals in each energy level above n = 1, each assigned to its own axis (x, y and z) in space.

  31. Things get a bit more complicated with the five d orbitals that are found in the d sublevels beginning with n = 3. To remember the shapes, think of “double dumbells” d orbital shapes …and a “dumbell with a donut”!

  32. Shape of f orbitals

  33. Orbital filling table

  34. Electron configuration of the elements of the first three series

  35. Irregular confirmations of Cr and Cu Chromium steals a 4s electron to half fill its 3d sublevel Copper steals a 4s electron to FILL its 3d sublevel

  36. Pauli Exclusion Principle No two electrons in an atom can have the same four quantum numbers. Wolfgang Pauli

  37. Quantum Numbers Each electron in an atom has a unique set of 4 quantum numbers which describe it. • Principal quantum number • Angular momentum quantum number • Magnetic quantum number • Spin quantum number

  38. Principal Quantum Number Generally symbolized by n, it denotes the shell (energy level) in which the electron is located. Number of electrons that can fit in a shell: 2n2

  39. Angular Momentum Quantum Number The angular momentum quantum number, generally symbolized by l, denotes the orbital (subshell) in which the electron is located.

  40. Magnetic Quantum Number The magnetic quantum number, generally symbolized by m, denotes the orientation of the electron’s orbital with respect to the three axes in space.

  41. Spin Quantum Number Spin quantum number denotes the behavior (direction of spin) of an electron within a magnetic field. Possibilities for electron spin:

  42. Assigning the Numbers Principle: n = shell (period) Angular Momentum: l = shape (s=0, p=1, d=2) Magnetic: ml= +l to -l Spin: +1/2 or – 1/2

  43. Chlorine • Write the orbital notation • Write the electron configuration • So, can we now write the quantum #’s for each e’?

  44. Electron Conf. and Quant #’s

  45. Key Terms • Ground State Electron Configurations (1s22s2…/Orbital Notations __ __ __ __ __ • Aufbau Principle (lowest first/periodic guide) • Hund’s Rule: Single e’ before pairing begins • Pauli Exclusion (up/down—no 2 e’ same set of quant. #’s) • Highest Occupied Level • Inner-Shell e’ • Noble Gas Notation

  46. Orbitals in outer energy levels DO penetrate into lower energy levels. Penetration #1 This is a probability Distribution for a 3s orbital. What parts of the diagram correspond to “nodes” – regions of zero probability?

  47. Which of the orbital types in the 3rd energy level Does not seem to have a “node”? Penetration #2 WHY NOT?

  48. Principle, angular momentum, and magnetic quantum numbers: n, l, and ml

  49. Chemistry Chapter 5 The Periodic Law

  50. Law of Mendeleev: • Properties of the elements recur in regular cycles (periodically) when the elements are arranged in order of increasing atomic mass.