1. Quiz Review

10. Light and quantized Energy Section 1
Electrons in Atoms
Chapter 5

11. Unanswered Questions What have we learned? Subatomic particles exist
The nucleus of an atom contains all of the atoms positive charge (Protons) and mass (Neutrons + Protons)
Atoms are surrounded by fast, moving electrons

12. But there were some unanswered questions:
How the atoms electrons were arranged in space
Why the electrons were not pulled into the positively charged nucleus.
Why there were differences and similarities in the chemical behavior among various elements

13. Unraveling of the puzzle:
Unanswered Questions
Unraveling of the puzzle:
Scientists observed that certain elements emitted visible light when heated in a flame.
Analysis of this light revealed that an elements chemical behavior is related to the arrangement of the electrons in its atoms.
To understand this we have to first understand the wave nature of light

14. The Wave Nature of Light
Electromagnetic Radiation: Is a form of energy that exhibits wave like behavior as it travels through space
Visible light is a type of electromagnetic radiation
Other examples include those from:
Microwaves
X-rays
Signals sent from the Radio & TV station to your home

15. The Wave Nature of Light

16. The Wave Nature of Light
Characteristics of Waves
Wavelength (λ): the shortest distance between equivalent points on a continuous wave.
Usually expressed in meters (m)
Frequency (ν): the number of waves that pass a given point per second.
Expressed as:
Hertz (Hz) = One wave per second
Waves = Waves/ second (1/s) or (s-1)
Example: 652 Hz = 652 waves/second = 652/s = 652 s-1

17. The Wave Nature of Light
Characteristics of Waves
Amplitude: Wave’s height from the origin to a crest or origin to trough
Electromagnetic Wave Relationship:
Speed of Electromagnetic waves = 3.00x108 m/s
Speed of light (c) = 3.00x108 m/s
Equation: c = λv where:
c = 3.00x108 m/s
λ = wave length
ν = frequency

18. The Wave Nature of Light
Waves can have different wavelengths & frequencies
Wavelengths and frequency are inversely related

19. Electromagnetic Spectrum

20. Practice Problems 1—4 page 140

21. Practice Problems 1—4 page 140

22. Practice Problems 1—4 page 140

23. Practice Problems 1—4 page 140

24. The Particle Nature of Light
The wave nature of light does not explain certain phenomenon
Light emission of certain metals
Photoelectric effect
Studied by German physicist Max Planck ( ) in 1900

25. The Particle Nature of Light
Planck’s conclusions:
Matter can gain or lose energy only in small, specific amounts called quanta.
Quantum: minimum amount of energy that can be gained or lost by an atom
Planck’s constant (h) = × J ● s

26. The Photoelectric Effect
Photoelectric effect: electrons are emitted from the surface of a metal when light of a certain frequency shines on the metal

27. The Photoelectric Effect
Albert Einstein proposed in 1905 that light has a dual nature.
“A beam of light has wavelike and particle-like properties.”
Light particle: photon
a particle of electromagnetic radiation with no mass that carries a quantum of energy

28. The Photoelectric Effect

29. Practice Problems 5—7 page 143

30. Atomic Emission Spectrum

31. Atomic Emission Spectrum
Atomic emission spectrum of an element:
the set of frequencies of the electromagnetic waves emitted by the atoms of the element

32. Atomic Emission Spectrum Cont’d
Unique for each element and can be used to identify an element or determine whether an element is part of an unknown compound.

33. Atomic Emission Spectrum Cont’d
The atomic emission spectrum is not continuous (made up of certain frequencies of light)

34. Quantum Theory and the Atom Section 2
Electrons in Atoms
Chapter 5

35. Bohr’s model of the atom
Einstein’s theory of light’s dual nature accounted for several unexplainable phenomena (photoelectric effect) but not why atomic emission spectra of elements were discontinuous rather continuous.
In 1913, Niels Bohr, a Danish physicist working in Rutherford’s laboratory, proposed a quantum model for the hydrogen atom that seemed to answer this question.

36. Bohr’s model of the atom
Elements have certain allowable energy states
Ground State: Lowest energy state of an atom
Excited Sate: The state of an atom when it gains energy

37. Bohr’s model of the atom Cont’d
An atoms energy state is related to the location of the electron.
Electron moves around the nucleus only in certain allowed circular orbits
The smaller the electrons orbit the lower the atoms energy state or energy level
The larger the electrons orbit the higher the atoms energy state or energy level

38. Bohr’s model of the atom Cont’d
The energy state of an electron is related to its quantum number
The quantum number is number of the orbit in an atom

39. Ex.Hydrogen has one electron which exists in the ground state or quantum # of n= 1
When energy is absorbed the electron moves to a higher energy orbit e.g. n=2 (the excited state)

40. The electron of the atom can return from:
The excited state n=2 to its
Ground state n=1 by emitting (releasing) energy as a photon

41. This energy difference can be calculated using the formula
Δ E = E higher-energy orbit – E lower-energy orbit = E photon = hv

42. The Quantum Mechanical Model of the Atom
Bohr’s model was flawed: Electrons do not move in a circular orbit
Heisenberg uncertainty principle: It is fundamentally impossible to know precisely both the velocity and position of a particle at the same time

43. The Quantum Mechanical Model of the Atom
The quantum mechanical model makes no attempt to describe the electrons path around nucleus
Boundary encompasses the 95% probability that the electron is located there at any given time

44. The Quantum Mechanical Model of the Atom
Uses 4 numbers to “address” an electron in an atom. We will only work with the principal quantum #
Principal quantum number, n
Energy Sublevels is equal to the number of the principal quantum number

45. The Quantum Mechanical Model of the Atom
Shape of Orbitals
Sublevels are labelled s, p, d or f depending on the shape of the orbital
s- Spherical
p – Dumbbell shaped
d – Same shape different planes
f – Same shape different planes

46. The Quantum Mechanical Model of the Atom
s Orbitals
Have only 1 orbital
Increase in size with increasing quantum number

47. The Quantum Mechanical Model of the Atom
p Orbitals
Have 3 orbitals
Oriented along three axes: x, y & z

48. The Quantum Mechanical Model of the Atom
d Orbitals
Have 5orbitals
Oriented along different planes: xy, xz, yz, x2-y2
Z2 – shaped differently oriented differently

49. The Quantum Mechanical Model of the Atom
f Orbitals
Have 7 orbitals
Complex multilobed shapes

50. The Quantum Mechanical Model of the Atom
Shape of Orbitals
Each orbital contains at most 2 electrons
Principal
Quantum #
Sublevels (type of Orbitals)
# of orbitals related to sublevel
n = 1 1s 1
n=2 2s 2p 3
n=3 3s 3p 3d 5
n=4 3f 7

51. The Quantum Mechanical Model of the Atom
Uses four numbers to “address” an electron in an atom
Principal quantum number, n
Energy Sublevels is equal to the number of the principal quantum number
Angular momentum quantum number, l
Magnetic quantum number, m
Spin quantum number, + ½ or – ½

52. Electron Configuration Section 3
Electrons in Atoms
Chapter 5

53. Ground state electron configuration
Electron Configuration: Arrangement of electrons in an atom
The Aufbau principle: Each electron occupies the lowest energy orbital available because they are more stable at this level.
Hund’s rule: 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

54. Ground state electron configuration
The Pauli exclusion principle: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

55. Using the Aufbau principle: Each electron occupies the lowest energy orbital available because they are more stable at this level.
Order of filling orbitals

56. Order of filling orbitals

57. The Pauli exclusion principle: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

58. Hunts Rule: A maximum of two electrons can occupy a single orbital, but only if the electrons have opposite spins

59. Ground state electron configuration

60. Ground state electron configuration

62. Writing the electron configuration
Neon
Electron Configuration: 1s2 2s2 2p6

63. Ground state electron configuration
Only works perfectly for all elements up to and including Vanadium, atomic number 23

64. Practice Problem p 160

65. Ground state electron configuration
Noble Gas notation: A method of representing electron configurations of noble gases.

66. Ground state electron configuration
Noble Gases:
Have eight electrons in their outermost orbital
Usually unstable
Noble Gas Notation:
Uses bracketed symbols
For example: Helium (He)

67. Ground state electron configuration
For example:

68. Ground state electron configuration
For example:

69. Ground state electron configuration
Valence electrons
Defined as atom’s outermost orbitals
Atoms in the outermost orbitals are associated with the atom’s highest principal energy level
For example:

70. Ground state electron configuration
Electron-dot structure
Devised by American Chemist G.N. Lewis
Aka Lewis Dot structures
Used by chemists to represent Valence Electrons
Writing the Lewis Dot Structure:
Element’s Symbol: Represents the atomic nucleus and inner level electrons.

71. Ground state electron configuration
Dots:
Represent Valence electrons
Placed one at a time around the four sides of the symbol and then paired up until all are shown
Example:

72. Practice Problems
Pg. 162 #26-28

73. Practice Problems p 162

74. Practice Problems p 162

75. Practice Problems p 162

76. Practice Problems p 162

77. Practice Problems p 162