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Chapter 4 Electron Configurations

Chapter 4 Electron Configurations. Early thoughts. Much understanding of electron behavior comes from studies of how light interacts with matter. Early belief was that light traveled as a wave, but some experiments showed behavior to be as a stream of tiny fast moving particles. Waves.

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Chapter 4 Electron Configurations

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  1. Chapter 4 Electron Configurations

  2. Early thoughts • Much understanding of electron behavior comes from studies of how light interacts with matter. • Early belief was that light traveled as a wave, but some experiments showed behavior to be as a stream of tiny fast moving particles

  3. Waves • Today scientists recognize light has properties of waves and particles • Waves: light is electromagnetic radiation and travels in electromagnetic waves.

  4. 4 Characteristics of a wave: • 1) amplitude - height of the wave. For light it is the brightness • 2) Wavelength ()– distance from crest to crest. • For light – defines the type of light • Visible light range – 400-750nm

  5. Properties continued • 3) Frequency ()– measures how fast the wave oscillates up and down. • It is measured in number per second. • Hertz = 1 cycle per second • FM radio = 93.1 MHz or 93.1 x 106 cycles per second • Visible light = 4 x 1014 cycles per second to 7 x 1014 cycles per second • 4) speed – 3.00 x 108 m/s

  6. Short wavelength, high frequency • Long wavelength, low frequency • Visible Spectrum • ROY G BIV • Longer wavelength shorter wavelength

  7. Electromagnetic spectrum (meters) • 10-11 gamma • 10-9 x-rays • 10-8 UV • 10-7 visible light • 10-6 infrared • 10-2 microwave • 1 TV

  8. Wavelength and frequency • Wavelength and frequency are inversely related!! •  = c/ • Where  is the wavelength, c is the speed of light and  is the frequency • Speed of light = Constant = • 3.00 x 108m/sec

  9. Example • Example: An infrared light has a wavelength of 2.42 x 10-6m. Calculate the frequency of this light. •  = c/  • = 3.0 x 108m/sec = • 2.42 x 10-6m • = 1.2 x 1014 waves/sec

  10. Wavelength and frequency • ****Remember  and  are inverse. Therefore short wavelength = high frequency!!

  11. Quantum Theory • Needed to explain why certain elements when heated give off a characteristic light (certain color) • 1900- Max Planck – idea of quantum

  12. Quantum Theory • - The amount of energy (electromagnetic radiation) an object absorbs/emits occurs only in fixed amounts called quanta (quantum) • - Quanta – finite amount of energy that can be gained or lost by an atom.

  13. 1905 Einstein’s theory • Einstein explained photoelectric effect by proposing that light consists of quanta of energy called PHOTONS • Photon = discrete bit of energy • Consider light traveling as photons

  14. Energy equation • Amount of energy of a photon described as • E = h • Where • E = energy •  = frequency • h = Planck’s constant = 6.6262 x 10-34 J s • Joule = SI unit for energy

  15. Photoelectric effect – • electrons are ejected from the surface of a metal when light shines on the metal. The frequency determines the amount of energy. The higher the frequency, the more energy per photon.

  16. Photoelectric effect • Ex. X-rays have a high frequency; therefore can damage organisms while radio waves have a low frequency.

  17. Dual nature of radiant energy • Photons act both like particles and waves. • Line Spectra: A spectrum that contains only certain colors, or wavelengths

  18. Question: How are electrons arranged in atoms? • Note: All elements emit light when they are vaporized in an intense flame or when electricity is passed through their gaseous state.

  19. How are electrons arranged in atoms? • Explanation: Bohr atom: 1911 • postulated that atoms have energy levels in which the electrons orbit • energy levels and/or orbits are labeled by a quantum number, n. • lowest energy level = ground state n=1

  20. How are electrons arranged in atoms? • when an electron absorbs energy, it jumps to a higher level (known as the excited state) n = 2,3 or 4 • Bohr model of an atom • Hydrogen • Red- e falls from 3 to 2 • Blue e falls from 4 to 2 • Violet e falls from 6 to 2

  21. 1924 – Louis DeBroglie • – If waves of light can act as a particle, then particles of matter should act like a wave. Found to be true.

  22. DeBroglie • Matter waves = wavelike behavior of particles. • From equation – wave nature is inversely related to mass therefore we don’t notice wave nature of large objects. However, electrons have a small mass so they have a larger wave characteristic

  23. Schroedinger’s wave equation • predicted probability of finding an electron in the electron cloud around nucleus. Gave us four quantum numbers to describe the position.

  24. Heisenberg’s Uncertainty Principle The position and momentum of a moving object cannot simultaneously be measured and known exactly. • Cannot know where it is and where its going at the same time.

  25. Quantum mechanical model of an atom – • Treats the electrons as a wave that has quantized its energy • Describes the probability that electrons will be found in certain locations around the nucleus.

  26. Orbitals • Orbitals – atomic orbital is a region around the nucleus of an atom where an electron with a given energy is likely to be found (high probability) • Have characteristic shapes, sizes and energies. • Four different kinds of orbitals s, p, d, f.

  27. S and p orbitals

  28. Orbitals and energy • Principal Energy Level = quantum number n. • Principal Quantum number • Value of n = 1,2,3,4,5,6,7 • Tells you the distance from the nucleus

  29. sublevels • Each level is divided into sublevels • # of sublevels = value of n • N=1 1 sublevel • n=2 2 sublevels • s,p,d or f shows the shape

  30. Orbital shape • Each sublevel has a certain number of orbitals which are directed three dimensionally • S = 1 sphere • P = 3 figure 8 (along the x,y or z axis) • D = 5 figure 4-27 p. 145 • F = 7

  31. Each electron in an orbital will have a spin – 2 options clockwise, vs. counter clockwise. • Electron configuration • Pauli Exclusion Principle – each orbital in an atom can hold at most 2 electrons and their electrons must have opposite spin.

  32. Aufbau Principle- electrons are added one at a time to the lowest energy orbitals available • Hund’s Rule – electrons occupy equal energy orbitals so that the maximum number or unpaired electrons result. • Occupy singly before pairing

  33. Diagonal Rule: for order of sublevels: • must remember 1s, 2s, 2p, 3s, 3p, 4s • Exceptions to Aufbau Principle • Cr Z=24 • Cu Z = 29

  34. Energy levels • Number of electrons per energy level • 1st = 2 • 2nd = 8 • 3rd = 18 • 4th = 32

  35. Number of electrons per orbital • 2 • Difference between paired and unpaired electrons • Paired = 2 electrons in the same orbital • Unpaired = 1 electron in the orbital

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