Successfully reported this slideshow.
We use your LinkedIn profile and activity data to personalize ads and to show you more relevant ads. You can change your ad preferences anytime.

Chem 1 unit 4 presentation

2,655 views

Published on

Published in: Education
  • Be the first to comment

Chem 1 unit 4 presentation

  1. 1. Why are Electrons Important?Chemistry Unit 4<br />
  2. 2. Why are Electrons Important?<br />
  3. 3. Light, a form of electronic radiation, has characteristics of both a wave and a particle<br />Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals.<br />Atomic emission spectra corresponds to the release of energy from an electron changing atomic energy levels.<br />A set of three rules determines the arrangement in an atom.<br />Main Ideas<br />
  4. 4. 4:1 Light and Quantized Energy<br />
  5. 5. Light and Quantized Energy<br />Objectives:<br />Compare the wave and particle natures of light.<br />Define a quantum of energy, and explain how it is related to an energy change of matter.<br />Contrast continuous electromagnetic spectra and atomic emission spectra.<br />
  6. 6. The Atom and Unanswered Questions<br />Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it.<br />The model doesn’t explain how the electrons were arranged around the nucleus.<br />The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.<br />
  7. 7. The Atom and Unanswered Questions<br />In the early 1900s, scientists observed certain elements emitted visible light when heated in a flame.<br />Analysis of the emitted light revealed that an element’s chemical behavior is related to the arrangement of the electrons in its atoms.<br />In order to understand this relationship and the nature of atomic structure, it will be helpful to first understand the nature of light.<br />
  8. 8. Wave Nature of Light<br />Electromagnetic radiationis a form of energy that exhibits wave-like behavior as it travels through space.<br />Visible light<br />Microwaves<br />X-rays<br />Radio waves<br />
  9. 9. Wave Nature of Light<br />All waves can be described by several characteristics.<br />The wavelength(λ) is the shortest distance between equivalent points on a continuous wave. (crest to crest, trough to trough)<br />The frequency(ν) is the number of waves that pass a given point per second.<br />Hertz- SI unit for frequency= one wave/sec<br />Energy increases with increasing frequency<br />
  10. 10. Wave Nature of Light<br />All waves can be described by several characteristics.<br />The amplitude is the wave’s height from the origin to a crest.<br />Independent of wavelength and frequency<br />
  11. 11. Wave Nature of Light<br />
  12. 12. Wave Nature of Light<br />The speed of light (3.00 x108 m/s) is the product of it’s wavelength and frequency c = λν.<br />
  13. 13. Wave Nature of Light<br />All electromagnetic waves, including visible light travels at 3.00 x 108m/s in a vacuum.<br />Speed is constant but wavelengths and frequencies vary.<br />Sunlight contains a continuous range of wavelengths and frequencies.<br />A prism separates sunlight into a continuous spectrum of colors.<br />
  14. 14. Wave Nature of Light<br />The electromagnetic spectrum includes all forms of electromagnetic radiation.<br />Not just visible light. <br />
  15. 15. Wave Nature of Light<br />
  16. 16. Particle Nature of Light<br />The wave model of light cannot explain all of light’s characteristics.<br />Matter can gain or lose energy only in small, specific amounts called quanta.<br />A quantum is the minimum amount of energy that can be gained or lost by an atom.<br />
  17. 17. Particle Nature of Light<br />Max Planck (1858-1947) – matter can gain or lose energy only in small amounts.<br />E=hv<br />Planck’s constanthas a value of 6.626 x10–34 J ● s.<br />Energy can only be emitted or absorbed in whole number multiples of h.<br />
  18. 18. Particle Nature of Light<br />The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.<br />
  19. 19. Particle Nature of Light<br />Albert Einstein proposed in 1905 that light has a dual nature. Nobel prize in 1921.<br />A beam of light has wavelike and particlelike properties.<br />A photon is a mass-less particle of electromagnetic radiation with no mass that carries a quantum of energy.<br />Ephoton = hνEphoton represents energy.h is Planck's constant.νrepresents frequency. <br />
  20. 20. Example Problem 1<br />Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz?<br />
  21. 21. Example Problem 2<br />Microwaves are used to cook food and transmit information. What is the wavelength of a microwave that has a frequency of 3.44 x 109 Hz?<br />What is the amount of energy in this wavelength?<br />
  22. 22. Example Problem 2<br />While an FM radio station broadcasts at a frequency of 94.7 MHz, an AM station broadcasts at a frequency of 820 kHz. What are the wavelengths of the two broadcasts?<br />
  23. 23. Example Problem 2<br />While an FM radio station broadcasts at a frequency of 94.7 MHz, an AM station broadcasts at a frequency of 820 kHz. What are the wavelengths of the two broadcasts?<br />What is the associated energy for each of these broadcasts?<br />
  24. 24. Example Problem 3<br />Every object gets its color by reflecting a certain portion of visible light. The color is determined by the wavelength of the reflected photons, and therefore their energy. The blue color in some fireworks occurs when copper (I) chloride is heated to approximately 1500K and emits blue light of wavelength 4.50 x 102 nm. How much energy does one photon of this light carry?<br />
  25. 25. Atomic Emission Spectrum<br />The atomic emission spectrumof an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.<br />Emission lines are specific to an element and can be used for identification.<br />
  26. 26. Atomic Emission Spectrum<br />
  27. 27. Absorption Spectra<br />The absorption spectra of an element is the set of frequencies of the electromagnetic waves absorbed by the atoms of the element.<br />Absorption lines are specific to an element and can be used for identification.<br />
  28. 28. Emission vs. Absorption<br />
  29. 29. Question?<br />What is the smallest amount of energy that can be gained or lost by an atom? <br />A. electromagnetic photon<br />B. beta particle<br />C. quanta<br />D. wave-particle<br />
  30. 30. What is a particle of electromagnetic radiation with no mass called? <br />A. beta particle<br />B. alpha particle<br />C. quanta<br />D. photon<br />Question?<br />
  31. 31. 4:2 Quantum Theory of the Atom<br />
  32. 32. Quantum Theory of the Atom<br />Objectives:<br />Compare the Bohr and quantum mechanical models of the atom.<br />Explain the impact of de Broglie's wave article duality and the Heisenberg uncertainty principle on the current view of electrons in atoms.<br />Identifythe relationships among a hydrogen atom's energy levels, sublevels, and atomic orbitals.<br />
  33. 33. Bohr’s Model of the Atom<br />Bohr correctly predicted the frequency lines in hydrogen’s atomic emission spectrum.<br />The lowest allowable energy state of an atom is called its ground state.<br />When an atom gains energy, it is in an excited state.<br />
  34. 34. Bohr’s Model of the Atom<br />Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.<br />The smaller the electron’s orbit the lower the atom’s energy state or level<br />
  35. 35. Bohr’s Model of the Atom<br />Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.<br />The larger the electron’s orbit the higher the atom’s energy state or level.<br />
  36. 36. Bohr’s Model of the Atom<br />Each orbit was given a number, called the quantum number. The orbit closest to the nucleus is n=1<br />
  37. 37. Bohr’s Model of the Atom<br />Example: Hydrogen’s single electron is in the n = 1 orbit in the ground state. Atom does not radiate energy.<br />When energy is added, the electron moves to the n = 2 orbit. Atom is excited. <br />When electron moves from an excited state to ground state, a photon is emitted. <br />
  38. 38. Electron States <br />
  39. 39. Quantum Mechanical Model<br />The Quantum Mechanical Model of the Atom – this model progressed through a series of scientific findings:<br />Louis de Broglie (1892–1987) hypothesized that particles, including electrons, could also have wavelike behaviors.<br />Like vibrating guitar strings – multiples of half waves. <br />Orbiting electron – whole number of wavelengths.<br />
  40. 40. Quantum Mechanical Model<br />
  41. 41. λrepresents wavelengthsh is Planck's constant.m represents mass of the particle.vrepresents velocity.<br />Quantum Mechanical Model<br />The de Broglie equationpredicts that all moving particles have wave characteristics.<br />
  42. 42. Example Problem<br />Why do we not notice the wavelengths of moving objects such as automobiles?<br />
  43. 43. Quantum Mechanical Model<br />Heisenberg showed it is impossible to take any measurement of an object without disturbing it.<br />The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.<br />Means that it is impossible to assign fixed paths for electrons like the circular orbits as previously thought.<br />
  44. 44. Quantum Mechanical Model<br />Heisenberg showed it is impossible to take any measurement of an object without disturbing it.<br />The Heisenberg uncertainty principlestates that it is fundamentally impossible to know precisely both the velocity and position of a particle at the same time.<br />The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.<br />
  45. 45. Quantum Mechanical Model<br />Schrödinger expanded on the de Broglie wave particle theory and created the quantum mechanical model that we know today.<br />treated electrons as waves in a model called the quantum mechanical model of the atom.<br />Schrödinger’s equation applied equally well to elements other than hydrogen.<br />
  46. 46. Quantum Mechanical Model<br />Both models limit an electron’s energy to certain values. Unlike the Bohr model, the quantum mechanical model makes no attempt to describe the electron’s path around the nucleus.<br />Electrons are located around the nucleus at a position that can be described only by a probability map. A boundary surface is chosen to contain the region that the electron can be expected to occupy 90% of the time.<br />
  47. 47. Quantum Mechanical Model<br />The wave function predicts a three-dimensional region around the nucleus called the atomic orbital.<br />
  48. 48. Quantum Numbers and the Revised Model<br />The revised model defines the relationship between an electron’s energy level, sublevel and atomic orbitals.<br />Four quantum numbers make up the identification of each electron in an atom. <br />
  49. 49. Atomic Orbitals<br />Principal quantum number (n) indicates the relative size and energy of atomic orbitals. n specifies the atom’s major energy levels, called the principal energy levels.<br />
  50. 50. Atomic Orbitals<br />Energy sublevels (s,p,d or f) are contained within the principal energy levels.<br />
  51. 51. Atomic Orbitals<br />n= # of sublevels per principal energy levels.<br />
  52. 52. Atomic Orbitals<br />Each energy sublevel relates to orbitals of different shape.<br />
  53. 53. Atomic Orbitals<br />Each orbital can contain 2 electrons<br />
  54. 54. Atomic Orbitals<br />
  55. 55. Question?<br />Which atomic suborbitals have a “dumbbell” shape? <br />A. s <br />B. f<br />C. p<br />D.d <br />
  56. 56. Question<br />Who proposed that particles could also exhibit wavelike behaviors? <br />A. Bohr<br />B. Einstein<br />C. Rutherford<br />D. de Broglie<br />
  57. 57. 4:3 Energy and Wavelength Relationships<br />
  58. 58. Atomic Emission<br />Atomic emission spectra corresponds to the release of energy from an electron changing atomic energy levels.<br />Hydrogen’s emission spectrum comprises three series of lines. <br />One series of lines are ultraviolet (Lyman series) and one series is infrared (Paschen series). <br />Visible wavelengths comprise the Balmer series.<br />
  59. 59. Bohr’s Model of the Atom<br />
  60. 60. Atomic Emission<br />The Bohr atomic model attributes these spectral lines to transitions from higher-energy states with larger electron orbits (ni) to lower-energy states with smaller electron orbits (nf).<br />The change in energy of this jump can be calculated with:<br />ΔE = -2.178 x 10-18 J (1/n2final – 1/n2initial)<br />
  61. 61. Example Problems<br />For the Balmer series, electron orbit transition occur from larger orbits to the n=2 orbit (nfinal= 2). Calculate the change in energy and wavelengths for the following electron transitions. ninitail = 3,4,5, and 6.<br />
  62. 62. 4:4 Electron Configuration<br />
  63. 63. Electron Configuration<br />Objectives:<br />Applythe Pauli exclusion principle, the aufbau principle, and Hund's rule to write electron configurations using orbital diagrams and electron configuration notation.<br />Define valence electrons, and draw electron-dot structures representing an atom's valence electrons.<br />
  64. 64. Electron Configuration<br />The arrangement of electrons in the atom is called the electron configuration.<br />
  65. 65. Electron Configuration<br />Three rules/ principals define how electrons can be arranged in atom’s orbitals. <br />The aufbau principle states that each electron occupies the lowest energy orbital available.<br />
  66. 66. Electron Configuration<br />
  67. 67. Electron Configuration<br />The Pauli exclusion principle states that a maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins.<br />Electrons in orbitals can be represented by arrows in boxes and each electron has an associated spin.<br />
  68. 68. Electron Configuration<br />Hund’s rulestates that single electrons with the same spin must occupy each equal-energy orbital before additional electrons with opposite spins can occupy the same energy level orbitals.<br />Electrons don’t like roommates.<br />
  69. 69. Electron Arrangement<br />-Electron arrangement can be represented by two common different methods. <br />Orbital Diagram – boxes labeled with principle energy level and sublevel associated with each orbital. Arrows are drawn up and down in the box to represent electrons and their spins.<br />
  70. 70. Orbital Diagram<br />
  71. 71. Orbital Diagram<br />Examples:<br />
  72. 72. Electron Arrangement<br />-Electron arrangement can be represented by two common different methods. <br />Electron Configuration Notation- lists the following in order: Principle energy number, sublevel, superscript of number of electrons in the sublevel. Electron distribution follows the main three rules. <br />Noble Gas Notation – abbreviated electron configuration by substituting noble gas symbols for a long series of notation.<br />
  73. 73. Electron Configuration Notation<br />Examples:<br />
  74. 74. Electron Configuration<br />
  75. 75. Electron Configuration<br />
  76. 76. Valence Electrons<br />Valence electronsare defined as electrons in the atom’s outermost orbitals—those associated with the atom’s highest principal energy level.<br />Electron-dot structureconsists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.<br />
  77. 77. Electron Dot Structure<br />Electrons are placed one at a time on the four sides of the symbol and then paired until used up. Side order doesn’t matter.<br />Example: <br /> Na<br /> Cl<br />
  78. 78. Valence Electrons<br />
  79. 79. 4:5 Accumulating Content<br />
  80. 80. Accumulating Content<br />Wavelengths are often measured in nm or Å. How do measurements like these affect calculations?<br />
  81. 81. Accumulating Content<br />Gamma rays can be very harmful radiation. What did you learn about the magnetic spectrum that helps you understand why?<br />
  82. 82. Accumulating Content<br />Light travels slower in water than it does in air due to the change in densities of the medium; however its frequency remains the same. How does the wavelength of light change as it travels from air to water?<br />
  83. 83. Question?<br />In the ground state, which orbital does an atom’s electrons occupy? <br />A. the highest available<br />B. the lowest available<br />C. the n = 0 orbital<br />D. the d suborbital<br />
  84. 84. Question?<br />The outermost electrons of an atom are called what? <br />A. suborbitals<br />B. orbitals<br />C. ground state electrons<br />D. valence electrons<br />
  85. 85. Study Guide <br />Key Concepts<br />All waves are defined by their wavelengths, frequencies, amplitudes, and speeds. c = λν<br />In a vacuum, all electromagnetic waves travel at the speed of light.<br />All electromagnetic waves have both wave and particle properties.<br />Matter emits and absorbs energy in quanta.Equantum = hν<br />
  86. 86. Study Guide <br />Key Concepts<br />White light produces a continuous spectrum. An element’s emission spectrum consists of a series of lines of individual colors.<br />
  87. 87. Study Guide <br />Key Concepts<br />Bohr’s atomic model attributes hydrogen’s emission spectrum to electrons dropping from higher-energy to lower-energy orbits. ∆E = E higher-energy orbit - E lower-energy orbit = E photon = hν<br />The de Broglie equation relates a particle’s wavelength to its mass, its velocity, and Planck’s constant. λ = h / mν<br />The quantum mechanical model of the atom assumes that electrons have wave properties. <br />Electrons occupy three-dimensional regions of space called atomic orbitals.<br />
  88. 88. Study Guide <br />Key Concepts<br />The arrangement of electrons in an atom is called the atom’s electron configuration.<br />Electron configurations are defined by the aufbau principle, the Pauli exclusion principle, and Hund’s rule. <br />An element’s valence electrons determine the chemical properties of the element. <br />Electron configurations can be represented using orbital diagrams, electron configuration notation, and electron-dot structures.<br />
  89. 89. Chapter Questions<br />The shortest distance from equivalent points on a continuous wave is the: <br />A. frequency<br />B. wavelength<br />C. amplitude<br />D. crest<br />
  90. 90. Chapter Questions<br />The energy of a wave increases as ____. <br />A. frequency decreases<br />B. wavelength decreases<br />C. wavelength increases<br />D. distance increases<br />
  91. 91. Chapter Question<br />Atom’s move in circular orbits in which atomic model? <br />A. quantum mechanical model<br />B. Rutherford’s model<br />C. Bohr’s model<br />D. plum-pudding model<br />
  92. 92. Chapter Question<br />It is impossible to know precisely both the location and velocity of an electron at the same time because: <br />A. the Pauli exclusion principle<br />B. the dual nature of light<br />C. electrons travel in waves<br />D. the Heisenberg uncertainty principle<br />
  93. 93. Chapter Assessment 5<br />How many valence electrons does neon have? <br />A. 0 <br />B. 1<br />C. 2<br />D. 3 <br />
  94. 94. Chapter Questions<br />Spherical orbitals belong to which sublevel? <br />A. s <br />B. p<br />C. d<br />D. f <br />
  95. 95. Chapter Questions<br />What is the maximum number of electrons the 1s orbital can hold? <br />A. 10 <br />B. 2<br />C. 8<br />D. 1 <br />
  96. 96. Chapter Questions<br />In order for two electrons to occupy the same orbital, they must: <br />A. have opposite charges<br />B. have opposite spins<br />C. have the same spin<br />D. have the same spin and charge<br />
  97. 97. Chapter Questions<br />How many valence electrons does boron contain? <br />A. 1 <br />B. 2<br />C. 3<br />D. 5 <br />
  98. 98. Chapter Questions<br />What is a quantum? <br />A. another name for an atom<br />B. the smallest amount of energy that can be gained or lost by an atom<br />C. the ground state of an atom<br />D. the excited state of an atom<br />
  99. 99. The End<br />

×