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Chapter 5

Chapter 5. Electrons In Atoms. Topics to Be Covered. 5.1 Light and Quantized Energy 136-145 5.2 Quantum Theory and the Atom 146-155 5.3 Electron Configuration 156-162. Section 5.1. Light and Quantized Energy. The Atom & Unanswered Questions. Early 1900s

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Chapter 5

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  1. Chapter 5 Electrons In Atoms

  2. Topics to Be Covered • 5.1 • Light and Quantized Energy • 136-145 • 5.2 • Quantum Theory and the Atom • 146-155 • 5.3 • Electron Configuration • 156-162

  3. Section 5.1 Light and Quantized Energy

  4. The Atom & Unanswered Questions • Early 1900s • Discovered 3 subatomic particles • Continued quest to understand atomic structure • Rutherford’s model • Positive charge in nucleus • Fast moving electrons around that • No accounting for differences and similarities in chemical behavior

  5. The Atom and Unanswered Questions • Example: • Lithium, sodium, and potassium have similar chemical behaviors (explained more in next chapter) • Early 1900s • Scientists began to unravel mystery • Certain elements emitted visible light when heated in a flame • Analysis revealed chemical behavior depends on arrangement of electrons

  6. The Wave Nature of Light • Electromagnetic radiation • A form of energy that exhibits wavelike behavior as it travels through space • Visible light is a type of ER

  7. Characteristics of Waves • All waves can be described by several characteristics • Wavelength • Frequency • Amplitude

  8. Wavelength • Represented by lambda λ • Shortest distance between equivalent points on a continuous waves • Measure crest to crest or trough to trough • Usually expressed in m, cm, or nm

  9. Frequency • Represented by nu ν • The number of waves that pass a given point per second • Given in the unit of hertz (Hz) • 1 Hz = 1 wave per second

  10. Amplitude • The wave’s height from the origin to a crest or from the origin to a trough • Wavelength and frequency do not affect amplitude

  11. Speed • All electromagnetic waves in a vacuum travel at a speed of 3.00 x 108 m/s • This includes visible light • The speed of light has its own symbol • C • C= λν

  12. Electromagnetic Spectrum • Also called the EM spectrum • Includes all forms of electromagnetic radiation • With the only differences in the types of radiation being their frequencies and wavelengths

  13. Electromagnetic Spectrum • Figure 5.5

  14. Problems • Page 140 • Calculating Wavelength of an EM Wave

  15. Particle Nature of Light • Needed to explain other properties of light • Heated objects emit only certain frequencies of light at a given temperature • Some metals emit electrons when light of a specific frequency shines on them

  16. Quantum Concept • When objects are heated they emit glowing light • 1900 • Max Planck began searching for an explanation • Studied the light emitted by heated objects • Startling conclusion

  17. Quantum Concept • Planck discovered: • Matter can gain or lose energy only in small specific amounts • These amounts are called quanta • Quantum—is the minimum amount of energy that can be gained or lost by an atom

  18. Example • Heating a cup of water • Most people thought that you can add any amount of thermal energy to the water by regulating the power and duration of the microwaves • In actuality, the temperature increases in infinitesimal steps as its molecules absorb quanta of energy, which appear to be a continuous manner

  19. Quantum Concept • Planck proposed that energy emitted by hot objects was quantized • Planck further demonstrated mathematically that a relationship exists between energy of a quantum and a frequency

  20. Energy of a Quantum • Equantum=hv • Equantum represents energy • h is Planck’s constant • v represents frequency

  21. Planck’s Constant • Symbol = h • 6.626 x 10-34 J*s • J is the symbol for joule • The SI unit of energy • The equation shows that the energy of radiation increases as the radiation’s frequency, v, increases.

  22. Planck’s Theory • For given frequencies • Matter can emit/absorb energy only in whole number multiples of hv • 1hv, 2hv, 3hv, 4hv etc. • Matter can have only certain amounts of energy • Quantities of energy between these values do not exist

  23. The Photoelectric Effect • Photoelectric effect • electrons, called photoelectrons • are emitted from a metal’s surface • when light of a certain frequency, or higher than a certain frequency shines on the surface

  24. Light’s Dual Nature • Einstein proposed in 1905 that light has a dual nature • photon—a massless particle that carries a quantum of energy

  25. Energy of a Photon • Ephoton=hv • Ephoton represents energy • h is Planck’s constant • v represents frequency

  26. Light’s Dual Nature • Einstein proposed • Energy of a photon must have a certain threshold value to cause the ejection of a photoelectron from the surface of a metal • Even small #s of photons with energy above the threshold value will cause the photoelectric effect • Einstein won Nobel Prize in Physics in 1921

  27. Sample Problems • Page 143 • Sample Problem 5.2 • Calculating Energy of a Photon

  28. Atomic Emission Spectra • See page 145

  29. Section 5.2 Quantum Theory and The Atom

  30. Bohr’s Model of the Atom • Dual-nature explains more • Atomic Emission Spectra • Not continuous • Only certain frequencies of light • Explained the Atomic Emission Spectra

  31. Energy States of Hydrogen • Bohr proposed certain allowable energy states • Bohr proposed electrons could travel in certain orbitals

  32. Energy states of Hydrogen • Ground State • Lowest allowable energy state of an atom • Orbit size • Smaller the orbit, the lower the energy state/level • Larger the orbit, the higher the energy state/level

  33. Energy states of Hydrogen • Hydrogen can have many excited states • It only has one electron • Quantum Number • Number assigned to each orbital • n • Look at Table 5.1

  34. The Hydrogen Line Spectrum • Hydrogen Ground State • Electron is in n=1 orbit • Does not radiate energy • Hydrogen Excited State • Energy is added to the atom from outside source • Electron moves to a higher energy orbit

  35. The Hydrogen Line Spectrum • Only Certain Atomic Energy Levels Possible • Example Our Classroom • Balmer Series • Electron transitions from higher-energy orbits to the second orbit • Account for visible lines

  36. The Hydrogen Line Spectrum • Lyman Series • Ultraviolet • Electrons drop into n=1 orbit • Paschen Series • Infrared • Electrons drop into n = 3 orbit

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