3. Chapter Menu Electrons in Atoms Section 5.1Section 5.1Light and Quantized Energy Section 5.2Section 5.2 Quantum Theory and the Atom Section 5.3Section 5.3 Electron Configuration Exit Click a hyperlink or folder tab to view the corresponding slides.

4. Section 5-1 Section 5.1 Light and Quantized Energy Compare the wave and particle natures of light. radiation: the rays and particles —alpha particles, beta particles, and gamma rays—that are emitted by radioactive material Define a quantum of energy, and explain how it is related to an energy change of matter. Contrast continuous electromagnetic spectra and atomic emission spectra.

5. Section 5-1 Section 5.1 Light and Quantized Energy (cont.) electromagnetic radiation wavelength frequency amplitude electromagnetic spectrum Light, a form of electronic radiation, has characteristics of both a wave and a particle. quantum Planck's constant photoelectric effect photon atomic emission spectrum

6. Section 5-1 The Atom and Unanswered Questions Recall that in Rutherford's model, the atom’s mass is concentrated in the nucleus and electrons move around it. The model doesn’t explain how the electrons are arranged around the nucleus. The model doesn’t explain why negatively charged electrons aren’t pulled into the positively charged nucleus.

7. Section 5-1 The Atom and Unanswered Questions (cont.) In the early 1900s, scientists observed certain elements emitted characteristic visible light when heated in a flame. Analysis of the emitted light suggested a way to understand the behavior of electrons around a nucleus.

8. Section 5-1 The Wave Nature of Light Visible light is a type of electromagnetic radiation, a form of energy that exhibits wave-like behavior as it travels through space.electromagnetic radiation

9. Section 5-1 The Wave Nature of Light (cont.) The wavelength (λ) is the shortest distance between equivalent points on a continuous wave.wavelength The frequency (ν) is the number of waves that pass a given point per second.frequency The amplitude is the wave’s height from the origin to a crest.amplitude All waves can be described by several characteristics.

10. Section 5-1 The Wave Nature of Light (cont.)

11. Section 5-1 The Wave Nature of Light (cont.)

12. Section 5-1 The Wave Nature of Light (cont.) The speed of light c is the product of it’s wavelength λ and frequency ν c = λν In a vacuum, the speed of light is 3.00x10 8 m/s all forms of EM radiation travel at the same speed wavelength and frequency are inversely related if one increases, the other decreases

13. 1) What is the wavelength of light with a frequency of 6.40x10 14 Hz? Hz means ‘cycles per second’ (1/sec) or (sec -1 ) c = λν 3.00x10 8 m/s = (λ)(λ)( 6.40x10 14 Hz ) λ = 4.69x10 -7 m (or 469 nm)

14. 2) What is the frequency of EM radiation with a wavelength of 14.3 mm? c = λν 3.00x10 8 m/s = ( 14.3x10 -3 m )(ν)(ν) ν = 2.10x10 10 Hz

15. Section 5-1 The Wave Nature of Light (cont.) The electromagnetic spectrum includes all forms of electromagnetic radiation.electromagnetic spectrum

16. Section 5-1 The Wave Nature of Light (cont.) Sunlight contains a wide range of wavelengths and frequencies. A small section of the EM spectrum is the visible light we can see.

17. S or F 5.1

18. S or F 5.2

19. Section 5-1 The Particle Nature of Light The wave model of light cannot explain all of light’s characteristics. Max Plank proposed that matter can gain or lose energy only in small, specific amounts called quanta. A quantum is the minimum amount of energy that can be gained or lost by an atom All changes in energy must be multiples of a tiny increment called Plank’s constant (h)quantum h has a value of 6.626  10 –34 J ● s.

20. Section 5-1 The Particle Nature of Light (cont.) The photoelectric effect is when electrons are emitted from a metal’s surface when light of a certain frequency shines on it.photoelectric effect

22. Section 5-1 The Particle Nature of Light (cont.) Albert Einstein proposed in 1905 that light has a dual nature. A beam of light has both wavelike and particlelike properties.A beam of light has both wavelike and particlelike properties. A photon is a particle of electromagnetic radiation with no mass that carries a quantum of energy.photon E photon = h E photon represents energy. h is Planck's constant. represents frequency.

23. 1) What is the energy carried by a photon of light with a frequency of 6.40x10 14 Hz? E = hν E =(6.626x10 -34 Js) ( 6.40x10 14 Hz ) E = 4.24x10 -19 J

24. 2) What wavelength of EM radiation corresponds with a photon of 7.85x10 -11 J? Hint: First find the frequency and then find the wavelength. 7.85x10 -11 J = (6.626x10 -34 Js)( ν) ν = 1.18x10 23 Hz 3.00x10 8 m/s = λ ( 1.18x10 23 Hz ) c = λν λ = 2.54x10 -15 m E = hν

25. A.A B.B C.C D.D Section 5-1 Section 5.1 Assessment What is the smallest amount of energy that can be gained or lost by an atom? A.electromagnetic photon B.beta particle C.quanta D.wave-particle

26. A.A B.B C.C D.D Section 5-1 Section 5.1 Assessment What is a particle of electromagnetic radiation with no mass called? A.beta particle B.alpha particle C.quanta D.photon

27. End of Section 5-1

28. Section 5-2 Section 5.2 Quantum Theory and the Atom Compare the Bohr and quantum mechanical models of the atom. atom: the smallest particle of an element that retains all the properties of that element, is composed of electrons, protons, and neutrons. Explain the impact of de Broglie's wave article duality and the Heisenberg uncertainty principle on the current view of electrons in atoms. Identify the relationships among a hydrogen atom's energy levels, sublevels, and atomic orbitals.

29. Section 5-2 Section 5.2 Quantum Theory and the Atom (cont.) ground state quantum number de Broglie equation Heisenberg uncertainty principle Wavelike properties of electrons help relate atomic emission spectra, energy states of atoms, and atomic orbitals. quantum mechanical model of the atom atomic orbital principal quantum number principal energy level energy sublevel

30. Section 5-1 Atomic Emission Spectra (cont.) The atomic emission spectrum of an element is the set of frequencies of the electromagnetic waves emitted by the atoms of the element.atomic emission spectrum Each element’s atomic emission spectrum is unique.

31. Section 5-1 Atomic Emission Spectra (cont.)

32. Section 5-2 Bohr's Model of the Atom In 1913, Niels Bohr suggested a model of the atom that would explain the observed spectrum of a hydrogen atom. Bohr suggested that an electron moves around the nucleus only in certain allowed circular orbits.

33. Section 5-2 Bohr's Model of the Atom (cont.) The lowest allowable energy state of an atom is called its ground state.ground state When an atom gains energy, it is in an excited state.

34. Section 5-2 Bohr's Model of the Atom (cont.) Each orbit was given a number, called the quantum number, describing its relative size (energy).

35. Section 5-2 Bohr's Model of the Atom (cont.) Hydrogen’s single electron is in the n = 1 orbit in the ground state. When energy is added, the electron moves to a higher orbit.(n = 2 or n = 3 or n = 4, etc.) When the electrons drop back down to lower levels, energy is released as EM radiation.

36. Section 5-2 Bohr's Model of the Atom (cont.)

37. Section 5-2 Bohr's Model of the Atom (cont.) Different drops will release different amounts of energy, therefore releasing different wavelengths of radiation. So the radiation released gives us insight into the locations of the electrons in an atom.

38. Section 5-2 Bohr's Model of the Atom (cont.) Bohr’s model explained the hydrogen’s spectral lines, but failed to explain any other element’s lines. The behavior of electrons is still not fully understood, but it is known they do not move around the nucleus in circular orbits.

39. Section 5-2 The Quantum Mechanical Model of the Atom Louis de Broglie (1892– 1987) hypothesized that particles, including electrons, could also have wavelike behaviors.

40. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The figure illustrates that electrons orbit the nucleus only in whole-number wavelengths.

41. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The de Broglie equation predicts that all moving particles have wave characteristics.de Broglie equation represents wavelengths h is Planck's constant. m represents mass of the particle. represents velocity. When describing these extremely small particles, we must consider their wave- particle duality, just like when we describe how light behaves.

42. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) Heisenberg showed it is impossible to take any measurement of an object without disturbing it. Influence of the Observer for big objects the effect is insignificant, but for tiny fundamental particles, it can make a big difference

43. Section 5-2 The Quantum Mechanical Model of the Atom (cont.)

44. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The Heisenberg uncertainty principle states that it is fundamentally impossible to know exactly both the velocity and position of a particle at the same time.Heisenberg uncertainty principle The only quantity that can be known is the probability for an electron to occupy a certain region around the nucleus.

45. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) Schrödinger treated electrons as waves in a model called the quantum mechanical model of the atom.quantum mechanical model of the atom Schrödinger’s equation applied equally well to elements other than hydrogen. http://en.wikipedia.org/wiki/Schr%C3%B6dinger%27s_equation

46. Section 5-2 The Quantum Mechanical Model of the Atom (cont.) The wave function predicts a three- dimensional region around the nucleus called the atomic orbital.atomic orbital

47. Section 5-2 Hydrogen Atomic Orbitals Principal quantum number (n) indicates the relative size and energy of atomic orbitals.Principal quantum number n specifies the atom’s major energy levels, called the principal energy levels.principal energy levels

48. Section 5-2 Hydrogen Atomic Orbitals (cont.) Energy sublevels are contained within the principal energy levels.Energy sublevels

49. Section 5-2 Hydrogen Atomic Orbitals (cont.) Each energy sublevel relates to orbitals of different shape.

50. Section 5-2 Hydrogen Atomic Orbitals (cont.)

51. A.A B.B C.C D.D Section 5-2 Section 5.2 Assessment Which atomic suborbitals have a “dumbbell” shape? A.s B.f C.p D.d

52. A.A B.B C.C D.D Section 5-2 Section 5.2 Assessment Who proposed that particles could also exhibit wavelike behaviors? A.Bohr B.Einstein C.Rutherford D.de Broglie

53. Section 5-3 Section 5.3 Electron Configuration Apply the Pauli exclusion principle, the aufbau principle, and Hund's rule to write electron configurations using orbital diagrams and electron configuration notation. electron: a negatively charged, fast-moving particle with an extremely small mass that is found in all forms of matter and moves through the empty space surrounding an atom's nucleus Define valence electrons, and draw electron-dot structures representing an atom's valence electrons.

54. Section 5-3 Section 5.3 Electron Configuration (cont.) electron configuration aufbau principle Pauli exclusion principle Hund's rule valence electrons electron-dot structure A set of three rules determines the arrangement in an atom.

55. Section 5-3 Ground-State Electron Configuration The arrangement of electrons in the atom is called the electron configuration.electron configuration The aufbau principle states that each electron occupies the lowest energy orbital available.aufbau principle

56. Section 5-3 Ground-State Electron Configuration (cont.)

57. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 This is the order that the electrons fill orbitals according to the Aufbau principle.

58. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Write an electron configuration for sulfur. S = 16 electrons 1s 2 2s 2 2p 6 3s 2 3p 4 How many electrons are in the 3 rd energy level? 6

59. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Write an electron configuration for vanadium. V = 23 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 How many electrons are in the outermost energy level? 2

60. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Write an electron configuration for strontium Sr = 38 electrons 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 Here is a shortcut. Skip to the previous noble gas. Put the symbol in square brackets and continue to the end of the configuration. [Kr] 5s 2

61. Section 5-3 Ground-State Electron Configuration (cont.) Noble gas notation uses noble gas symbols in brackets to shorten inner electron configurations of other elements.

62. 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 14 5s 2 5p 6 5d 10 5f 14 6s 2 6p 6 6d 10 7s 2 7p 6 Write an electron configuration for tantalum. Ta = 73 [Xe] 6s 2 4f 14 5d 3 After helium (1s 2 ) all noble gas configurations end with p 6.

63. Section 5-3 Ground-State Electron Configuration (cont.) The Pauli exclusion principle states that a maximum of two electrons can occupy a single orbital, and only if the electrons have opposite spins.Pauli exclusion principle Hund’s rule states 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.Hund’s rule

64. Drawing an orbital diagram, means representing each orbital with a box and each electron as an arrow. The direction of the arrow represents the ‘spin’ quantum number. Ne – 10 electrons orbital diagram electron configuration = 1s 2 2s 2 2p 6 1s2s2p

65. The Pauli exclusion principle dictates that we may only have electrons with opposite spin within the same orbital. He – 2 electrons orbital diagram electron configuration = 1s 2 1s correct incorrect

66. Hund’s rule dictates that electrons will fill open orbitals of the same energy before they pair up. C – 6 electrons orbital diagram electron configuration = 1s 2 2s 2 2p 2 1s2s2p 1s2s2p correct incorrect

67. Section 5-3 Ground-State Electron Configuration (cont.)

68. 1s2s2p Draw an orbital diagram for sodium. 3s

69. 1s2s2p Draw an orbital diagram for sulfur. 3s3p

70. Section 5-3 Ground-State Electron Configuration (cont.) The electron configurations (for chromium, copper, and several other elements) reflect the increased stability of half-filled and filled sets of s and d orbitals.

71. Section 5-3 Valence Electrons Valence electrons are defined as electrons in the atom’s outermost orbitals— those associated with the atom’s highest principal energy level.Valence electrons Electron-dot structure consists of the element’s symbol representing the nucleus, surrounded by dots representing the element’s valence electrons.Electron-dot structure

72. Section 5-3 Valence Electrons (cont.)

73. Science or Fiction The element curium is named after Marie Curie and her husband Pierre. Curium is produced by bombarding plutonium with alpha particles (helium ions) and was named after Marie Skłodowska- Curie and her husband Pierre. plutoniumalpha particlesheliumions Marie Skłodowska- CuriePierre

74. Science or Fiction Marie Curie is the only person who has a won Nobel prizes in both chemistry and physics. Curie was the first person to win or share two Nobel Prizes. She is one of only two people who have been awarded a Nobel Prize in two different fields, the other person being Linus Pauling (for chemistry and for peace).Linus Pauling

75. Science or Fiction Marie and Pierre Curie are given credit for discovering 2 elements. The Curies are given credit for the discovery of both radium and polonium.

76. Science or Fiction The element polonium was discovered in Poland. Marie Curie discovered polonium in France. Poland at the time was under Russian, Prussian, and Austrian partition, and did not exist as an independent country. It was Curie's hope that naming the element after her native land would publicize its lack of independence.

77. Science or Fiction Americium was discovered in Illinois. Americium was first isolated by Glen Seaborg and his team in late 1944 at the wartime Metallurgical Laboratory at the University of Chicago (now known as Argonne National Laboratory).University of Chicago Argonne National Laboratory

78. A.A B.B C.C D.D Section 5-3 Section 5.3 Assessment In the ground state, which orbital does an atom’s electrons occupy? A.the highest available B.the lowest available C.the n = 0 orbital D.the d suborbital

79. A.A B.B C.C D.D Section 5-3 Section 5.3 Assessment The outermost electrons of an atom are called what? A.suborbitals B.orbitals C.ground state electrons D.valence electrons

80. End of Section 5-3

81. Resources Menu Chemistry Online Study Guide Chapter Assessment Standardized Test Practice Image Bank Concepts in Motion

82. Study Guide 1 Section 5.1 Light and Quantized Energy Key Concepts All waves are defined by their wavelengths, frequencies, amplitudes, and speeds. c = λν In a vacuum, all electromagnetic waves travel at the speed of light. All electromagnetic waves have both wave and particle properties. Matter emits and absorbs energy in quanta. E quantum = hν

83. Study Guide 1 Section 5.1 Light and Quantized Energy (cont.) Key Concepts White light produces a continuous spectrum. An element’s emission spectrum consists of a series of lines of individual colors.

84. Study Guide 2 Section 5.2 Quantum Theory and the Atom Key Concepts 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 The de Broglie equation relates a particle’s wavelength to its mass, its velocity, and Planck’s constant. λ = h / mν The quantum mechanical model of the atom assumes that electrons have wave properties. Electrons occupy three-dimensional regions of space called atomic orbitals.

85. Study Guide 3 Section 5.3 Electron Configuration Key Concepts The arrangement of electrons in an atom is called the atom’s electron configuration. Electron configurations are defined by the aufbau principle, the Pauli exclusion principle, and Hund’s rule. An element’s valence electrons determine the chemical properties of the element. Electron configurations can be represented using orbital diagrams, electron configuration notation, and electron-dot structures.

86. A.A B.B C.C D.D Chapter Assessment 1 The shortest distance from equivalent points on a continuous wave is the: A.frequency B.wavelength C.amplitude D.crest

87. A.A B.B C.C D.D Chapter Assessment 2 The energy of a wave increases as ____. A.frequency decreases B.wavelength decreases C.wavelength increases D.distance increases

88. A.A B.B C.C D.D Chapter Assessment 3 Atom’s move in circular orbits in which atomic model? A.quantum mechanical model B.Rutherford’s model C.Bohr’s model D.plum-pudding model

89. A.A B.B C.C D.D Chapter Assessment 4 It is impossible to know precisely both the location and velocity of an electron at the same time because: A.the Pauli exclusion principle B.the dual nature of light C.electrons travel in waves D.the Heisenberg uncertainty principle

90. A.A B.B C.C D.D Chapter Assessment 5 How many valence electrons does neon have? A.0 B.1 C.2 D.3

91. A.A B.B C.C D.D STP 1 Spherical orbitals belong to which sublevel? A.s B.p C.d D.f

92. A.A B.B C.C D.D STP 2 What is the maximum number of electrons the 1s orbital can hold? A.10 B.2 C.8 D.1

93. A.A B.B C.C D.D STP 3 In order for two electrons to occupy the same orbital, they must: A.have opposite charges B.have opposite spins C.have the same spin D.have the same spin and charge

94. A.A B.B C.C D.D STP 4 How many valence electrons does boron contain? A.1 B.2 C.3 D.5

95. A.A B.B C.C D.D STP 5 What is a quantum? A.another name for an atom B.the smallest amount of energy that can be gained or lost by an atom C.the ground state of an atom D.the excited state of an atom