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GECH119 Atomic Properties

GECH119 Atomic Properties. Dr. Ralph C. Gatrone Virginia State University Department of Chemistry and Physics. Chapter Objectives. Where are the electrons? Light Quantum Mechanics Models of the atom Valence Electrons Electron Configuration. Assignment.

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GECH119 Atomic Properties

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  1. GECH119Atomic Properties Dr. Ralph C. Gatrone Virginia State University Department of Chemistry and Physics

  2. Chapter Objectives • Where are the electrons? • Light • Quantum Mechanics • Models of the atom • Valence Electrons • Electron Configuration

  3. Assignment • Read Chapters 3 in Investigating Chemistry: A Forensic Science Perspective • For future tests and quizzes you should be able to do problems 1 – 16; 20 – 24; 30 - 37 in Chapter 3.

  4. Developing Model of the Atom • Nucleus • Small • Positively charged • Protons and neutrons • Atomic number = protons • Mass number = protons + neutrons • Where are the electrons? • Need to consider experiments done on light

  5. Light • Radiant energy • Electromagnetic energy • Moves through time and space • Constant speed (c = 3.00 X 108 m/s) • A wave

  6. Light

  7. Light • Wavelength – distance between peaks • Unit is the meter (m) • Frequency – number of peaks that pass a given point in one second • Unit is cycles/s = s-1 = Hz (Hertz)

  8. Light • Speed of light is a constant in a vacuum • Designated c • Equal to the wavelength X frequency • c = l X n • Energy is related to these terms by • E = hc/ l

  9. Light • Therefore, • Long wavelengths have lower frequencies • Short wavelengths have higher frequencies • High frequencies have high energy • Low frequencies have low energy

  10. Observations • Heat a metal bar • As the metal bar gets hot • Emits light • Tungsten: light is white • Electric stove: glows red • What is the relationship between light intensity and heat? • Newtonian physics can’t explain this observation • Need something new

  11. Quantum Mechanics • Classical mechanics • Proposed by Newton • Energy is continuous • In 1900 Max Planck broke with tradition • Proposed a new physics • Quantum Mechanics • Energy is no longer continuous • Energy is emitted in chunks of minimum size • A quantum • hn= E (v = frequency, h = Planck’s constant)

  12. Quantum Mechanics • Energy is released or absorbed in integer (multiples) of hv • You are familiar with these concepts • Violin – continuous sounds between notes • Piano – quantized notes

  13. Color and Heat • Color of light emitted • Related to energy of light • Energy released is quantized • Energy is released in chunks • Multiples of hv

  14. Observation • Shine light onto a metal • Electricity is observed. • Electrons are observed being released • Photo – light and electric – electron • Photoelectric Effect • Each metal has a distinct frequency at which electrons are emitted

  15. The photoelectric effect.

  16. Photoelectric Effect Explanation • 1905, A. Einstein proposed an explanation • Light strikes metal surface as a packet of energy • Given by hn • Consider this packet of energy as a particle • Einstein called it the photon • The energy of the photon is given by • E = hn • Light is quantized

  17. Summary • Light is quantized • Energy is quantized • E = hn • Energy expression used is the same.

  18. Summary • Light is quantized • E = hn • Energy is quantized • E = hn • Energy expression is the same.

  19. Result • hn (photon) strikes metal • Energy is transferred to an electron • Energy must be greater than force holding the electron for electron to be ejected • Excess Energy observed as movement of the electron • High v (low l) (e.g. x-rays and gamma rays) have high energy and cause tissue damage • Light can be treated as a wave and particle

  20. Atoms Emit Light • Each element gives off a distinctive colored glow when heated. • This can be done by putting a platinum wire dipped in a solution of the salt of a particular element. • Examples:

  21. Flames of lithium. Photo courtesy of James Scherer.

  22. Flames of sodium. Photo courtesy of James Scherer.

  23. Flames of strontium. Photo courtesy of James Scherer.

  24. Atoms Emit Light • Visualized by a spectroscope • Atomic spectrum is observed • Hydrogen has a very simple pattern • Simplicity was noted by Balmer and Rydberg

  25. Emission (line) spectra of some elements.

  26. Transitions of the electron in the hydrogen atom.

  27. Quantum Mechanics and Atomic Spectra • Niels Bohr • Energy increases as the electron gets further away from nucleus • Atom absorbs a photon the electron absorbs energy • Electron now has higher energy and moves further from the nucleus • Electron must release energy (as light) in order to return to low energy (preferred) state • Electron’s energy must be quantized • Electron can only be at certain energy levels • Each energy level is given a principal quantum number • Corresponds to the periods in the Periodic Table • Bohr’s model is the familiar planetary model

  28. The Bohr Model

  29. Bohr Model • Bohr model explains • The energy of emitted photons as the difference between the two orbits • The observed spectral lines for H

  30. Bohr Model • Bohr Model doesn’t explain • Why are energy levels quantized? • Electrons have negative charge and an orbiting charge will spiral downward into the nucleus.

  31. Dual Properties of Matter • If light can be a wave or particle • What about matter? • Louis de Broglie • Matter has wave properties • Mass is so large we don’t see these • Electron is very small • Characteristic wavelength • l = h/mv where m = mass

  32. Consider de Broglie Equation • l = h/mv where m = mass • Mass is in the denominator • As mass gets bigger • l (the wavelength) gets smaller • As mass gets smaller • l (the wavelength) gets bigger

  33. The Electron • The electron is very small • Mass is very nearly zero (1/1800 amu) • Using the de Broglie equation • Wavelength of the electron can be calculated • Approximately the same as an X-ray • Basis for the electron microscope • Electron in an atom is moving very fast • Electrons have wavelike properties • Electrons have particle like properties

  34. Scanning electron microscope..

  35. The Electron as a Wave • Consider a rolling ball • We can use classical physics to calculate • Position, direction, speed (momentum) • However, if matter behaves like a wave • Cannot know exact position and its momentum • Heisenberg stated this as the Uncertainty Principle

  36. The Bohr Model

  37. Bohr Model • Electrons are in orbits at some distance from the nucleus specified by the principal quantum number, n. • Electrons are moving around the nucleus • We know their position and their momentum. • Heisenberg states we cannot know both pieces of information if the electron is a wave. • This is another weakness of the Bohr Model • A new approach is needed.

  38. A New Approach • Must consider new concepts • Erwin Schrodinger • Quantum Mechanics • HY= EY • H is an operator • Y is the function describing the wave • Solutions generate the wave function Y • Y2 is the probable area of finding an electron • For hydrogen the solutions using this approach generate the Bohr Model data

  39. The Quantum Model • Generates 4 quantum numbers • Principal Quantum number, n, is same as Bohr proposed and corresponds to the periods in the Periodic Table • The other 3 numbers provide the probable location of finding any one electron, and • The shapes of orbitals (not orbits) for the most probable region around the nucleus that an electron could be found

  40. The Atomic Orbitals • 1s, 2s, and 2p orbitals

  41. The Atomic Orbitals • 3d orbitals (5 orbitals are modeled)

  42. Electrons • Quantum Mechanics • Electrons have wave and particle properties • Electrons are located in orbitals near the nucleus • Energies are quantized • Electrons are in lowest energy orbital • Ground State

  43. Excited Electrons • Electron in ground state absorbs energy • Electron moves to higher energy orbital • The Excited State • Electron releases energy • Returns to Ground State

  44. Electron Configuration • Electrons are arranged in orbitals • Increasing energy with distance from nucleus • All electrons are the same • Each electron has same charge • Each electron has same mass • Electrons differ only in location with respect to the nucleus • Inner electrons are held tightly • Outermost electrons are held the least • Outmost electrons are the valence electrons

  45. Electron Configuration • A model • Valence electrons feel less pull from the nucleus • H: 1s1 (there is 1 electron in the 1s orbital) • He: 1s2 (2 electrons in the 2s orbital) • Li: 1s2 2s1 • Be: 1s2 2s2 • B: 1s2 2s2 2p1

  46. Electron Configuration • Calcium • Ca has atomic number 20 • 20 protons, 20 electrons • Ground state electron configuration • 1s2 2s22p6 3s23p6 4s2 • 4s2 – 2 valence electrons

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