General Chemistry

1. Why are Electrons
Important?
Chemistry Unit 4

2. Why are Electrons
Important?

3.  Light, a form of electronic radiation, has
characteristics of both a wave and a particle
 Wavelike properties of electrons help relate atomic
emission spectra, energy states of atoms, and atomic
orbitals.
 Atomic emission spectra corresponds to the release of
energy from an electron changing atomic energy
levels.
 A set of three rules determines the arrangement in an
atom.
Main Ideas

4. 4:1 Light and Quantized
Energy

5. • Compare the wave and particle natures of
light.
• 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.
Light and Quantized
Energy
Objectives:

6. • 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
were arranged around the nucleus.
• The model doesn’t explain why negatively
charged electrons aren’t pulled into the
positively charged nucleus.
The Atom and
Unanswered Questions

7. The Atom and
Unanswered Questions
• In the early 1900s, scientists observed certain
elements emitted visible light when heated in a
flame.
• Analysis of the emitted light revealed that an
element’s chemical behavior is related to the
arrangement of the electrons in its atoms.
• In order to understand this relationship and the
nature of atomic structure, it will be helpful to
first understand the nature of light.

8. Wave Nature of Light
Electromagnetic radiation is a form of energy
that exhibits wave-like behavior as it travels
through space.
• Visible light
• Microwaves
• X-rays
• Radio waves

9. Wave Nature of Light
• The wavelength (λ) is the shortest
distance between equivalent points on a
continuous wave. (crest to crest, trough to
trough)
• The frequency (ν) is the number of waves that
pass a given point per second.
• Hertz- SI unit for frequency= one wave/sec
• Energy increases with increasing frequency
All waves can be described by several
characteristics.

10. Wave Nature of Light
• The amplitude is the wave’s height from the
origin to a crest.
• Independent of wavelength and frequency
All waves can be described by several
characteristics.

11. Wave Nature of Light

12. Wave Nature of Light
The speed of light (3.00 x 108 m/s) is the
product of it’s wavelength and frequency
c = λν.

13. Wave Nature of Light
• All electromagnetic waves, including visible
light travels at 3.00 x 108m/s in a vacuum.
• Speed is constant but wavelengths and
frequencies vary.
• Sunlight contains a continuous range of
wavelengths and frequencies.
• A prism separates sunlight into a continuous
spectrum of colors.

14. Wave Nature of Light
• The electromagnetic spectrum includes all
forms of electromagnetic radiation.
• Not just visible light.

15. Wave Nature of Light

16. Particle Nature of Light
The wave model of light cannot explain all of
light’s characteristics.
• 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.

17. Particle Nature of Light
• E=hv
• Planck’s constant has a value of
6.626 x 10–34 J ● s.
• Energy can only be emitted or absorbed in
whole number multiples of h.

18. Particle Nature of Light
• The photoelectric effect is when electrons
are emitted from a metal’s surface when
light of a certain frequency shines on it.

19. Particle Nature of Light
• Albert Einstein proposed in 1905 that light
has a dual nature. Nobel prize in 1921.
• A beam of light has wavelike and particlelike
properties.
• A photon is a mass-less particle of
electromagnetic radiation with no mass that
carries a quantum of energy.
Ephoton = hν Ephoton represents energy.
h is Planck's constant.
ν represents frequency.

20. Example Problem 1
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?

21. Example Problem 2
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?
What is the amount of energy in this wavelength?

22. Example Problem 2
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?

23. Example Problem 2
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?
What is the associated energy for each of these
broadcasts?

24. Example Problem 3
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?

25. Atomic Emission Spectrum
The atomic emission spectrum of an element
is the set of frequencies of the electromagnetic
waves emitted by the atoms of the element.
• Emission lines are specific to an element and
can be used for identification.

26. Atomic Emission Spectrum

27. Absorption Spectra
The absorption spectra of an element is
the set of frequencies of the
electromagnetic waves absorbed by
the atoms of the element.
 Absorption lines are specific to an
element and can be used for
identification.

28. Emission vs. Absorption

29. Question?
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

30. What is a particle of electromagnetic
radiation with no mass called?
A. beta particle
B. alpha particle
C. quanta
D. photon
Question?

31. 4:2 Quantum Theory of the
Atom

32. Quantum Theory of the
Atom
• Compare the Bohr and quantum mechanical
models of the atom.
• 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.
Objectives:

33. Bohr’s Model of the Atom
Bohr correctly predicted the frequency lines in
hydrogen’s atomic emission spectrum.
• The lowest allowable energy state of an atom
is called its ground state.
• When an atom gains energy, it is in an
excited state.

34. Bohr’s Model of the Atom
• Bohr suggested that an electron moves around the
nucleus only in certain allowed circular orbits.
•The smaller
the electron’s
orbit the lower
the atom’s
energy state
or level

35. Bohr’s Model of the Atom
• Bohr suggested that an electron moves around the
nucleus only in certain allowed circular orbits.
•The larger
the electron’s
orbit the
higher the
atom’s
energy state
or level.

36. Bohr’s Model of the Atom
• Each orbit was given a number, called the quantum
number. The orbit closest to the nucleus is n=1

37. Bohr’s Model of the Atom
• Example: Hydrogen’s single electron is in
the n = 1 orbit in the ground state. Atom
does not radiate energy.
• When energy is added, the electron moves to
the n = 2 orbit. Atom is excited.
• When electron moves from an excited state to
ground state, a photon is emitted.

38. Electron States

39. Quantum Mechanical Model
The Quantum Mechanical Model of the Atom –
this model progressed through a series of
scientific findings:
• Louis de Broglie (1892–1987) hypothesized that
particles, including electrons, could also have
wavelike behaviors.
• Like vibrating guitar strings – multiples of half
waves.
• Orbiting electron – whole number of
wavelengths.

40. Quantum Mechanical Model

41. Quantum Mechanical Model
• The de Broglie equation predicts that all
moving particles have wave characteristics.
λ represents wavelengths
h is Planck's constant.
m represents mass of the particle.
v represents velocity.

42. Example Problem
Why do we not notice the wavelengths of
moving objects such as automobiles?

43. Quantum Mechanical Model
Heisenberg showed it is impossible to take
any measurement of an object without
disturbing it.
• The Heisenberg uncertainty principle
states that it is fundamentally impossible to
know precisely both the velocity and position
of a particle at the same time.
• Means that it is impossible to assign fixed
paths for electrons like the circular orbits
as previously thought.

44. Quantum Mechanical Model
Heisenberg showed it is impossible to take
any measurement of an object without
disturbing it.
• The Heisenberg uncertainty principle
states that it is fundamentally impossible to
know precisely both the velocity and position
of a particle at the same time.
• The only quantity that can be known is the
probability for an electron to occupy a
certain region around the nucleus.

45. Quantum Mechanical Model
Schrödinger expanded on the de Broglie wave
particle theory and created the quantum
mechanical model that we know today.
• treated electrons as waves in a model called the
quantum mechanical model of the atom.
• Schrödinger’s equation applied equally well
to elements other than hydrogen.

46. Quantum Mechanical Model
• 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.
• 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.

47. Quantum Mechanical Model
• The wave function predicts a three-dimensional
region around the nucleus called the atomic orbital.

48. Quantum Numbers and
the Revised Model
The revised model defines the
relationship between an electron’s
energy level, sublevel and atomic
orbitals.
 Four quantum numbers make up the
identification of each electron in an
atom.

49. Atomic Orbitals
• 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.

50. Atomic Orbitals
• Energy sublevels (s,p,d or f) are
contained within the principal energy
levels.

51. Atomic Orbitals
• n= # of sublevels per principal energy levels.

52. Atomic Orbitals
• Each energy sublevel relates to orbitals of different
shape.

53. Atomic Orbitals
• Each orbital can contain 2 electrons

54. Atomic Orbitals

55. Question?
Which atomic suborbitals have a
“dumbbell” shape?
A. s
B. f
C. p
D. d

56. Question
Who proposed that particles could also
exhibit wavelike behaviors?
A. Bohr
B. Einstein
C. Rutherford
D. de Broglie

57. 4:3 Energy and Wavelength
Relationships

58. Atomic Emission
Atomic emission spectra corresponds to the release
of energy from an electron changing atomic energy
levels.
 Hydrogen’s emission spectrum comprises three
series of lines.
 One series of lines are ultraviolet (Lyman series)
and one series is infrared (Paschen series).
 Visible wavelengths comprise the Balmer series.

59. Bohr’s Model of the Atom

60. Atomic Emission
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).
 The change in energy of this jump can be
calculated with:
 ΔE = -2.178 x 10-18 J (1/n2
final – 1/n2
initial)

61. Example Problems
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.

62. 4:4 Electron Configuration

63. Electron Configuration
Objectives:
• Apply the Pauli exclusion principle, the
aufbau principle, and Hund's rule to write
electron configurations using orbital diagrams
and electron configuration notation.
• Define valence electrons, and draw electron-
dot structures representing an atom's valence
electrons.

64. Electron Configuration
The arrangement of electrons in the atom is called
the electron configuration.

65. Electron Configuration
Three rules/ principals define how electrons can
be arranged in atom’s orbitals.
1. The aufbau principle
states that each
electron occupies the
lowest energy orbital
available.

66. Electron Configuration

67. Electron Configuration
2. The Pauli exclusion principle states that a
maximum of two electrons can occupy a single
orbital, but only if the electrons have opposite
spins.
• Electrons in orbitals can be
represented by arrows in
boxes and each electron
has an associated spin.

68. Electron Configuration
3. 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.
• Electrons don’t like
roommates.

69. Electron Arrangement
-Electron arrangement can be represented
by two common different methods.
•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.

70. Orbital Diagram

71. Orbital Diagram
 Examples:

72. Electron Arrangement
-Electron arrangement can be represented
by two common different methods.
•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.
•Noble Gas Notation – abbreviated electron
configuration by substituting noble gas
symbols for a long series of notation.

73. Electron Configuration
Notation
 Examples:

74. Electron Configuration

75. Electron Configuration

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

77. Electron Dot Structure
 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.
 Example:
Na
Cl

78. Valence Electrons

79. 4:5 Accumulating
Content

80. Accumulating Content
 Wavelengths are often measured in nm or Å.
How do measurements like these affect
calculations?

81. Accumulating Content
 Gamma rays can be very harmful radiation.
What did you learn about the magnetic
spectrum that helps you understand why?

82. Accumulating Content
 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?

83. Question?
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

84. Question?
The outermost electrons of an atom are
called what?
A. suborbitals
B. orbitals
C. ground state electrons
D. valence electrons

85. Study Guide
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.
Equantum = hν

86. Study Guide
Key Concepts
• White light produces a continuous spectrum. An
element’s emission spectrum consists of a
series of lines of individual colors.

87. Study Guide
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.

88. Study Guide
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.

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

90. Chapter Questions
The energy of a wave increases as ____.
A. frequency decreases
B. wavelength decreases
C. wavelength increases
D. distance increases

91. Chapter Question
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

92. Chapter Question
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

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

94. Chapter Questions
Spherical orbitals belong to which
sublevel?
A. s
B. p
C. d
D. f

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

96. Chapter Questions
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

97. Chapter Questions
How many valence electrons does boron
contain?
A. 1
B. 2
C. 3
D. 5

98. Chapter Questions
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

99. The End