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Understanding Electrons in Atoms: Light, Quantum Theory, and Electron Configuration

Chapter 5 delves into the intricate relationship between light and the structure of atoms. It introduces the concept of quantized energy, highlighting the early discoveries of subatomic particles and the limitations of Rutherford's atom model. Electromagnetic radiation's wave characteristics—including wavelength, frequency, and amplitude—are covered, leading to the quantum theory which explains energy exchange in specific amounts called quanta. The dual nature of light is explored through Planck’s findings and Einstein’s photon concept, culminating in an examination of atomic emission spectra and electron configuration.

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Understanding Electrons in Atoms: Light, Quantum Theory, and Electron Configuration

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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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