1 / 22

Understanding Electrons: Light, Waves, and Quantum Mechanics

This chapter explores the nature of electrons and their behavior as light and waves within the framework of quantum mechanics. It defines essential concepts such as wavelength, frequency, amplitude, and the speed of electromagnetic waves. The influence of prominent scientists like Max Planck, Niels Bohr, Louis de Broglie, and Werner Heisenberg is examined. Key principles like the uncertainty principle and electron configurations in atomic structure are discussed, illustrating the quantum model of atoms and the significance of valence electrons in chemical behavior.

fionnuala-russell
Télécharger la présentation

Understanding Electrons: Light, Waves, and Quantum Mechanics

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 The Electron

  2. Light • A form of electromagnetic radiation • Characteristics of both wave & particle

  3. Waves

  4. Wavelength(λ)- shortest distance between equivalent points on a continuous wave • Measured in m, cm, or nm • Crest- top of wavelength • Trough- bottom of wavelength • Frequency(v)- number of waves that pass a certain point per second • Measured in waves/second • Hertz(Hz)- SI unit for frequency • Equal to 1 wave/second or 1/second • Amplitude- the wave’s height from it’s origin to it’s crest or trough • λ & v do not affect amplitude

  5. Waves (cont.) • All electromagnetic waves travel at a speed of 3.00 X 108 m/s in a vacuum • The symbol for the speed of light is c. • c=λv

  6. Quantum- minimum amount of energy that can be gained or lost by an atom • Energy of quantum= hv where h is Planck’s constant (6.626 x 10-34 Js, and v is the frequency

  7. Photon- a massless particle that carries a quantum of energy. • Energy of photon= hv, or hc/λ

  8. Max Planck • 1858-1947 • Studied the light emitted by heated objects and that this energy is quantized. • Matter absorbs energy in whole number multiples of hv

  9. Neils Bohr • Developed the quantum model for the hydrogen atom • Proposed that the H atom had only certain energy states • Ground state- the lowest allowable energy state of an atom (1st energy level) • Excited state- when an atom gains energy • Stated that the electron moves around the nucleus in only certain allowed circular motion  proven incorrect • The smaller the orbital- the lower the energy level

  10. Quantum Mechanical Model • Principal Quantum Number (n)- number assigned to each orbit (page 147) • n=1 • Closest orbital to the nucleus • n=2 • The next orbital

  11. The movement of electrons around the nucleus is not completely understood now. Evidence indicates that they do not travel in circular orbitals.

  12. Louis de Broglie • Questioned whether particles of matter can behave like waves • de Broglie equation- predicts that all moving particles have wave like characteristics • λ =(h/mv) • m=mass, h=Planck’s constant, v=velocity, λ= wavelength

  13. Werner Heisenburg • 1901-1976 • Showed that it was impossible to take any measurement of an object without disturbing it

  14. Heisenburg Uncertainty Principle- it is impossible to know precisely both the velocity and position of a particle at the same time • You can’t assign fixed paths for electrons • The only thing that can be known is the probability that an electron is in a certain region around the nucleus

  15. Atomic orbital- 3D area around the nucleus that indicates an electron’s probable location • Principle Quantum Number(n)- indicates the size and energy of atomic orbitals

  16. Energy Sublevels- the number of sublevels increase as ‘n’ increases • Sublevels are labeled s,p,d,f. • S is spherical • P looks like a dumbbell • D & F don’t always have the same shape • n=1 (1 sublevel) • n=2 (2 sublevels) • n=3 (3 sublevels) and so forth

  17. s holds 2 electrons • p holds 6 electrons • d holds 10 electrons • f holds 14 electrons

  18. Draw diagonal through to show the sublevels fill

  19. Ground State Electron Configuration • the most stable, lowest energy arrangement of electrons • Electron configuration- the arrangement of electrons in an atom • Lower energy levels are more stable than higher energy levels • Electrons typically arrange themselves so the atom is in the lowest energy state

  20. Aufbau Principle- each electron occupies the lowest energy level available • Pauli Exclusion Principle- only 2 electrons can occupy a single orbital and they must have opposite spins • Hund’s Rule- single electrons with the same spin must occupy each equal energy level before adding electrons with opposite spins

  21. Electron Configuration Notation- 1s22s22p3 • A noble gas can be used in brackets and then the rest of the electron configuration used • [Ne]3s2 • There are exceptions to the aufbau principle where there are only partial filled shells such as Cu (page 160)

  22. Valence Electrons • Determine the chemical properties of an atom • The electrons in the outermost shell • Involved in forming chemical bonds • S is [Ne]3s23p4 • It has 6 electrons in it’s outer shell • Electron dot structure- contains the symbol for the atom surrounded by the valence electrons

More Related