1 / 78

Chapter 6 Electronic Structure of Atoms

Chapter 6 Electronic Structure of Atoms. or “How I Learned to Stop Worrying and Love the Electron”. Problems with the Rutherford Model. Classical physics says atoms should emit light and destroy themselves - they don’t

halil
Download Presentation

Chapter 6 Electronic Structure of Atoms

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 6Electronic Structureof Atoms or “How I Learned to Stop Worrying and Love the Electron”

  2. Problems with the Rutherford Model • Classical physics says atoms should emit light and destroy themselves - they don’t • Atoms can be induced to emit light, but they give off a line spectrum, rather than a continuous spectrum. • Every atom gives off different colours of light. • No explanation of why different atoms have different properties, or the same properties.

  3. Let’s Talk About Light

  4. What is Light? • light is radiant energy. • a better term is electromagnetic energy (EMR) • includes not only visible light, but many forms of EMR we cannot see directly: • heat • microwaves • x-rays • light has properties of both waves and particles.

  5. Wave model of Light • The distance between corresponding points on adjacent waves is the wavelength • The symbol is the Greek letter lamda(). • The height of the wave is the amplitude. It corresponds to the intensity of the light

  6. Waves • The number of waves passing a given point per unit of time is the frequency, • The symbol is the Greek letter nu (). • The longer the wavelength, the smaller the frequency.

  7. Electromagnetic Radiation • EMR is a continuous spectrum of wavelengths.

  8. Using λ and ν to determine “colour” • “Colour” is a term used to describe visible light. • Visible light is a very small part of the electromagnetic spectrum: • radio waves -- ultraviolet • microwaves -- x-rays • infrared -- gamma rays • visible

  9. Colour of Light • we can identify light by its wavelength or frequency: • a wave 5.00 x 10-7 m is • green • a wave 6.80 x 10-7 m is • red • a wave 2.60 x 10-5 m is • infrared • a wave 7.80 x 10-10 m is • in the x-ray region.

  10. Speed of Light • is the one thing that is constant in the universe (sort of). • in a vacuum, the speed of light is 3.00 x 108 m/s • there is a relationship between wavelength and frequency: • where • c = speed of light • λ = wavelength • ν = frequency

  11. Complete questions 6.13 to 6.18, even

  12. Quantized Energy and Photons • The wave model of light is good, but it does not explain: • how an object can glow when its temperature increases. (Blackbody radiation) • emission of electrons by shining light on the surface of a metal. (photoelectric effect) • emission of light from electrons of excited gas atoms. (emission spectra)

  13. Blackbody Radiation • Heated solids emit radiation (blackbody radiation) • The wavelength distribution depends on the temperature (i.e., “red hot” objects are cooler than “white hot” objects). • Why does wavelength or frequency depend on temperature? • Max Planck suggested a way out by assuming that energy comes in packets called quanta.

  14. Planck proposed a relationship between energy and the frequency of light quanta: • where: • E = energy in Joules • h = Planck’s constant (6.6262 x 10-34 J·s) • ν = frequency, in Hertz (1/s, s-1) • as energy increases, so does frequency.

  15. Complete questions 6.21 to 6.28

  16. The PhotoelectricEffect

  17. Photons • Einstein used the quantum to explain the photoelectric effect: • Light comes in particles, called photons. • The energy of each photon is determined by Planck’s equation. • Light shining on the surface of a metal can cause electrons to be ejected from the metal.

  18. The electrons will only be ejected if the photons have sufficient energy: • Below the threshold frequency no electrons are ejected. • Above the threshold frequency, the excess energy appears as the kinetic energy of the ejected electrons. • Light has wave-like AND particle-like properties. • Complete question 6.30

  19. Line Spectra and the Bohr Model Another mystery involved the emission spectra observed from energy emitted by atoms and molecules.

  20. The Nature of Energy • A white light source produces a continuous spectrum, like a rainbow. • When elements are excited, only a line spectrum of discrete wavelengths is observed.

  21. Hydrogen Spectrum • has line spectra in 3 regions of the EM Spectrum: • Ultraviolet – Lymann Series • Visible – Balmer Series • Infrared - Paschen Series

  22. The Nature of Energy • Niels Bohr adopted Planck’s ideas about the quantum and applied them to the electrons around a nucleus.

  23. The Nature of Energy • Bohr’s model is based on three postulates: • Electrons in an atom can only occupy certain permitted orbits, or quanta. • Electrons in permitted orbits have specific, “allowed” energies; these energies will not be radiated from the atom. • Energy is only absorbed or emitted to move an electron from one “allowed” energy state to another; the energy is emitted by a photon: E = h

  24. The Nature of Energy • An electron in its lowest permissible energy is at ground state • If an electron accepts a quantum of energy it will move to a higher energy level, or excited state. • When the electron moves back down to ground state it emits a photon of light of a frequency which correlates to the energy of the quantum.

  25. 1 nf2 ( ) - E = −RH 1 ni2 The Nature of Energy The energy absorbed or emitted from the process of electron promotion or demotion can be calculated by the equation: where RH is the Rydberg constant, 2.18  10−18 J, and ni and nf are the initial and final energy levels of the electron.

  26. this explains the line spectra of hydrogen: • Lyman series. High energy, in the UV range. Represents energy transition from higher quanta to ground state. • Balmer series. Intermediate energy, in the visible range. Represents energy transition from higher quanta to quantum 2. • Paschen series. Low energy, in the IR range. Represents energy transition from higher quanta to quantum 3.

  27. Limitations of Bohr Model • cannot explain the spectra of atoms other than hydrogen. • However, the model introduces two important ideas: • The energy of an electron is quantized: electrons exist only in certain energy levels described by quantum numbers. • Energy gain or loss is involved in moving an electron from one energy level to another.

  28. h mv  = The Wave Nature of Matter • Louis de Broglie posited that if light can have material properties, matter should exhibit wave properties. • He demonstrated that the relationship between mass and wavelength was

  29. What does this mean? • electrons have wave properties. • the orbits of electrons are multiples of the electron wavelength • the first orbit has a circumference of 1 λ, the second is 2 λ , etc.

  30. The Uncertainty Principle • Heisenberg showed that the more precisely the momentum of a particle is known, the less precisely is its position known. • Our uncertainty of the whereabouts of an electron is greater than the size of the atom itself!

  31. Quantum Mechanics • Erwin Schrödinger developed a mathematical treatment into which both the wave and particle nature of matter could be incorporated. • It is known as quantum mechanics.

  32. The wave function describes the electron’s matter wave; it gives the probability of finding the electron. • Electron density is another way of expressing probability. • A region of high electron density is one where there is a high probability of finding an electron.

  33. Orbitals and Quantum Numbers • Solving the wave equation gives a set of wave functions, or orbitals, and their corresponding energies. • Each orbital describes a spatial distribution of electron density. • An orbital is described by a set of three quantum numbers.

  34. Principal Quantum Number, n • describes the energy level on which the orbital resides. • The values of n are integers ≥ 0. • correspond to the periods on the Periodic Table.

  35. Azimuthal Quantum Number, l • defines the shape of the orbital. • Allowed values of lare integers ranging from 0 to n − 1. • letter designations communicate the different values of land, therefore, the shapes and types of orbitals.

  36. Azimuthal Quantum Number, l • Theoretical g, h, i, etc. orbitals exist, but no atoms have been created to use them.

  37. Magnetic Quantum Number, ml • three-dimensional orientation of the orbital. • Values are integers ranging from -l to l : −l ≤ ml≤ l • on any given energy level, there can be up to: • 1 s orbital • 3 p orbitals • 5 d orbitals • 7 f orbitals

  38. Magnetic Quantum Number, ml • Orbitals with the same value of n form a shell. • Different orbital types within a shell are subshells. • Each subshell is designated by a number and a letter. • For example, 3p orbitals have n = 3 and l = 1.

  39. In Summary

  40. s Orbitals • Value of l = 0. • Spherical in shape. • Radius of sphere increases with increasing value of n.

  41. p Orbitals • Value of l = 1. • Have two lobes with a node between them. • The letters correspond to allowed the values of ml of –1, 0, and +1.

  42. d Orbitals • Value of l is 2. • Four of the five orbitals have 4 lobes; the other resembles a p orbital with a doughnut around the center. • The letters correspond to allowed the values of ml of -2, –1, 0, +1 and +2.

  43. forbitals • Value of l is 3. • 7 possible shapes, including 8 lobes and 2 doughnuts. • The letters correspond to allowed the values of mlof -3, -2, –1, 0, +1, +2 and +3.

  44. Energies of Orbitals • Orbitals of the same energy are degenerate.

  45. Spin Quantum Number, ms • two electrons in the same orbital do not have exactly the same energy. • The “spin” of an electron describes its magnetic field, which affects its energy.

  46. Spin Quantum Number, ms • There is a spin quantum number, ms. • has only 2 allowed values: +1/2 and −1/2.

More Related