1 / 25

Chapter 5

Chapter 5. Electrons in Atoms. The Bohr Model. An electron is found only in specific circular paths, or orbits, around the nucleus. Each orbit has a fixed energy. The orbits are called ‘energy levels.’. Energy Levels. Energy levels are like the rungs of a ladder:

nigel
Télécharger la présentation

Chapter 5

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. Chapter 5 Electrons in Atoms

  2. The Bohr Model • An electron is found only in specific circular paths, or orbits, around the nucleus. • Each orbit has a fixed energy. The orbits are called ‘energy levels.’

  3. Energy Levels • Energy levels are like • the rungs of a ladder: • You can move up or down by going from rung to rung. • You can’t stand in-between rungs. • For an electron to change energy levels it must gain or lose exactly the right amount of energy.

  4. A Quantum • A quantum of energy is the amount needed to move an electron from one energy level to another. • The energy of an electron is said to be “quantized.” • Energy levels in an atom are not all equally spaced.

  5. An Airplane Propeller • The blurry picture of an airplane propeller represents the area where the actual propeller blade can be found. • Similarly, the electron cloud of an atom represents the locations where an electron is likely to be found.

  6. The Model Quantum Mechanical • Comes from the mathematical solution to the Schrodinger equation. • Determines allowed energies an electron can have & how likely it is to find the electron in various locations around the nucleus. • Uses probability

  7. Atomic Orbitals • A region in space in which there is a high probability of finding an electron. • Energy levels of electrons are labeled by principal quantum numbers (n) n = 1, 2, 3, 4 …

  8. s Orbitals are spherical

  9. p Orbitals are dumbbell- shaped

  10. d Orbitals 4 out of the 5 d orbitals have clover leaf shapes

  11. f Orbitals are more complicated

  12. Atomic Orbitals The number and kinds of atomic orbitals depend on the energy sub level. • N=1 has 1 sublevel called 1s • N=2 has 2 sublevels called 2s and 2p • N=3 has 3 sublevels called 3s, 3p, and 3d • N=4 has 4 sublevels 4s, 4p, 4d, and 4f The maximum number of electrons that can occupy a principle energy level is 2n2. (n=principle quantum #)

  13. Electron Configurations • Electrons in an atom try to make the most stable arrangement possible (lowest energy) • The Aufbau Principle, the Pauli Exclusion Principle, and Hund’s Rule are guidelines that govern electron configurations in atoms

  14. Aufbau Principle • Electrons occupy the orbitals of lowest energy first

  15. Pauli Exclusion Principle • An orbital can hold at most 2 electrons • Does it make sense that two negatively charged particles will ‘want’ to share the same space? • This phenomenon is made possible because electrons possess a quantum mechanical property called spin

  16. Electron Spin • Spin may be thought of as clockwise or counter-clockwise • An arrow indicates an electron and its direction of spin • An orbital containing paired electrons is written

  17. Hund’s Rule • When filling orbitals of equal energy, one electron enters each orbital until all the orbitals contain one electron with similar spin

  18. Hund’s Rule • How would you put 2 electrons into a p sublevel? • How would you put 7 electrons into a d sublevel?

  19. Light • Now that we understand how electrons are arranged in atoms, we can begin to look at how the frequencies of emitted light are related to changes in electron energies

  20. Light • Light waves properties: • Amplitude – the wave’s height from zero to crest • Wavelength – the distance between crests • Frequency – the number of wave cycles to pass a given point per unit of time (Usually Hz = 1/s)

  21. Light • Wavelength has the symbol (λ) lambda. • Frequency has the symbol (ν) nu. • The speed of light is a constant (c) = 3x108 m/s • c = λν

  22. Light • How are wavelength and frequency related? • They are inversely related. As one increases, the other decreases • How long are the wavelengths that correspond to visible light? • 700-380 nanometers

  23. Electromagnetic Spectrum • Visible light is only a tiny portion of the electromagnetic spectrum which also includes radio waves, microwaves, infrared, visible light, ultra violet, X-rays, and gamma rays. • If the entire electromagnetic spectrum was a strip of professional 16 mm movie film stretching from Los Angeles to Seattle, the portion of visible light would be only ONE frame of film.

  24. Atomic Spectra • When atoms absorb energy, electrons move to higher energy levels • Electrons then lose energy by emitting light as they return to lower energy levels • Atoms emit only specific frequencies of light that correspond to the energy levels in the atom • The frequencies of light emitted by an element separate into discrete lines to give the atomic emission spectrum of the element

  25. Atomic Spectra • An electron with its lowest possible energy is in its ground state • The light emitted by an electron is directly proportional to the energy change of the electron. • E = hν • Atomic spectra are like fingerprints: no two are alike!

More Related