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Chapter 4

Chapter 4. The Modern Model of the Atom. 4.3 The Bohr Theory of Atomic Structure (Continued). The electrons are attracted to the nucleus due to charge Energy must be added for the electron to move away from the nucleus

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Chapter 4

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  1. Chapter 4 The Modern Model of the Atom

  2. 4.3 The Bohr Theory of Atomic Structure (Continued) • The electrons are attracted to the nucleus due to charge • Energy must be added for the electron to move away from the nucleus • When the electrons move back to their original position, energy is released.

  3. 4.1 Seeing the Light—A New Model of the Atom • This energy gets released as light.Different colors correspond to different energies

  4. 4.3 The Bohr Theory of Atomic Structure (Continued) • Electrons in orbits • The closer they are to the nucleus, the less energy there is. • The further away (outer orbits) they are, the higher energy there is.

  5. 4.3 The Bohr Theory of Atomic Structure (Continued) • Bohr’s model • States electrons can only move between orbits, never moving halfway. • Shells • Name of the orbits in the model • Labeled as principal quantum number (n)

  6. 4.3 The Bohr Theory of Atomic Structure (Continued) • Maximum number is defined by 2n2 for each shell, where n is the shell number.

  7. 4.3 The Bohr Theory of Atomic Structure (Continued) Lithium Atom

  8. 4.4 Periodicity and Line Spectra Explained • Periodic nature of the chemical properties • Valence shell • The outermost shell of the Bohr model • The number of electrons in a valence shell defines its chemical properties.

  9. 4.4 Periodicity and Line Spectra Explained (Continued) • Energy states of atoms • Ground state—lowest energy state • Electrons packed as low as possible in the Bohr model • Excited state—any state higher in energy than the ground state • One or more electrons move outward in the Bohr model.

  10. 4.4 Periodicity and Line Spectra Explained (Continued) Ground StateExcited State

  11. 4.5 Subshells and Electron Configuration • Updates to the Bohr model • The Bohr model fails when more than one electron is present. • The spectra contained many more lines than originally appeared. • Irwin Shroedingerroposed the existence of subshells within the principle quantum shell

  12. Wave or Quantum Mechanics • Irwin Schrödinger proposed that matter can be described as a wave. E. Schrodinger 1887-1961

  13. 4.5 Subshells and Electron Configuration (Continued) • Each principle quantum shell has nsubshells present. • Denoted s, p, d, and f

  14. 4.5 Subshells and Electron Configuration (Continued) • Subshells • s can hold up to 2 electrons. • p can hold up to 6 electrons. • d can hold up to 10 electrons. • f can hold up to 14 electrons. • Energy increases as you move up the list. • s is the smallest subshell and has the lowest energy. • f is the largest subshell and has the highest energy.

  15. 4.5 Subshells and Electron Configuration (Continued) • Energy of subshells • Some shells can overlap due to this. • 4s has lower energy than the 3dsubshell.

  16. Electron Configuration Chartsto remember the order for filling orbitals s holds 2 p holds 6 d holds 10 f holds 14 g is theoretical etc. 4f 5f 6f 7f 3d 4d 5d 6d 7d 2p 3p 4p 5p 6p 7p 1s 2s 3s 4s 5s 6s 7s

  17. Other ways……

  18. 4.5 Subshells and Electron Configuration (Continued) • Electron configuration • Arrangement of electrons in the shells and subshells • Lists the shells and subshells in order of energy and number of electrons in each Electron Configuration of K

  19. 4.5 Subshells and Electron Configuration (Continued) • Electron configuration of Br • Remember 4s fills before 3d.

  20. 4.5 Subshells and Electron Configuration (Continued) What is the electron configuration for the ground state of Si?

  21. 4.5 Subshells and Electron Configuration (Continued) What is the electron configuration for the ground state of Si? 1s22s22p63s23p2

  22. 4.5 Subshells and Electron Configuration ( • Using the periodic table • The periodic table will show the order of electron filling • Notice the s block has two blocks corresponding to the two electrons an ssubshell can hold.

  23. Orbital Filling of Electrons Electrons fill orbitals from the bottom up: Aufbau Principle

  24. 4.5 Subshells and Electron Configuration (Continued) Using the periodic table what would be the electron configuration for Fe?

  25. 4.5 Subshells and Electron Configuration (Continued) Using the periodic table what would be the electron configurationfor Fe? 1s22s22p63s23p64s23d6

  26. 4.5 Subshells and Electron Configuration (Continued) • Abbreviating electron configuration • Use the previous noble gas to denote the electron configuration to that point. Lu Electron Configuration 1s22s22p63s23p64s23d104p65s24d105p66s25d14f14 or [Xe]6s25d14f14

  27. 4.5 Subshells and Electron Configuration (Continued) • Using the electron configuration to find on the periodic table • Arsenic has an electron configuration of [Ar]4s23d104p3 What group should it be in? What period? • The valence shell is n = 4, so it will be found in the fourth period. • Has five valence electrons (4s24p3), so it will be found in group VA ( or 15 )

  28. 4.7 Compound Formation and the Octet Rule • Octet rule—elements react to place eight electrons in the valence shell • Mimics the noble gas configurations

  29. 4.7 Compound Formation and the Octet Rule (Continued) • Explaining MgF2 • Each Mg gives up two electrons and each F accepts one electron. • Forming Mg+2 and F-1 ions

  30. 4.7 Compound Formation and the Octet Rule (Continued) • Predicting charges by element family • Knowing the number of valence electrons, one can easily predict what the final charge will be.

  31. 4.7 Compound Formation and the Octet Rule (Continued) Predict the chemical formula NaxOy of the compound that results from the reaction between Na and O.

  32. 4.7 Compound Formation and the Octet Rule (Continued) Predict the chemical formula NaxOy of the compound that results from the reaction between Na and O. Na2O would be the predicted compound.

  33. 4.8 The Modern Quantum Mechanical Model of the Atom • Surprising results: • Tunneling—the ability of an electron to pass through a solid material • Uncertainty principle • You cannot know where an electron is or where it is going. • The more precisely you measure the current location, the less you will be able to predict the next location.

  34. 4.8 The Modern Quantum Mechanical Model of the Atom (Continued) • Electron orbital • Can be thought of as electron clouds • Defines electron location as probability not exact location

  35. 4.8 The Modern Quantum Mechanical Model of the Atom (Continued) • Electron orbital shapes • Exhibit wavelike motion, giving rise to various shapes

  36. 4.8 The Modern Quantum Mechanical Model of the Atom (Continued) • Comparing the Bohr model and the quantum model with Li • The third electron only sits in the second shell for the Bohr model. • In the quantum model, all three electrons can occupy the same 1s area, and the third electron will also spend time outside the 1s area.

  37. 4.8 The Modern Quantum Mechanical Model of the Atom (Continued) • Orbitals can overlap in space.

  38. 4.1 Seeing the Light—A New Model of the Atom (Continued) • In the 20th century we discovered that light has two natures: • Particles (packets or quanta ) called “photons” • Waves of energy. What are waves? Waves are traveling vibrations. • Strangely, light behaves as both a particle and wave at the same time

  39. 4.1 Seeing the Light—A New Model of the Atom (Continued) • Light as a wave • Wavelength ()—distance between identical points on a wave

  40. 4.1 Seeing the Light—A New Model of the Atom (Continued) • Light spreads by wavelength through a prism. • Visible light ranges from 380−750 nm. • White light is when all colors are present.

  41. 4.1 Seeing the Light—A New Model of the Atom (Continued) • Electromagnetic spectrum • Range of wavelengths including both visible and invisiblewavelengths

  42. 4.1 Seeing the Light—A New Model of the Atom (Continued) • To calculate the energy of electromagnetic radiation: E = hc/ or E = hf E = energy of the electromagnetic radiation h = 6.626 x 10-34 J . S c = 3.00 x 108 m/s (speed of light) • = wavelength of electromagnetic radiation f = frequency of electromagnetic radiation

  43. 4.1 Seeing the Light—A New Model of the Atom (Continued) What is the energy of blue light that has a wavelength of 450.0 nm?

  44. 4.1 Seeing the Light—A New Model of the Atom (Continued) What is the energy of blue light that has a wavelength of 450.0 nm? What is the energy of red light that vibrates with a frequency of 4.5 x 1014 Hertz? E = hf = (6.6 x 10-34 Js )(4.5 x 1014 waves/sec) = 2.97 x 10-19 Joules

  45. 4.1 Seeing the Light—A New Model of the Atom (Continued) • What does all this mean? • Each color corresponds to a certain energy. • Each of the four colors seen in a hydrogen spectra are related to four specific energies. • Dark areas represent energies not emitted by hydrogen.

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