2. Light, a form of electronic radiation, hasLight, a form of electronic radiation, has
characteristics of both a wave and a particlecharacteristics of both a wave and a particle
Wavelike properties of electrons help relateWavelike properties of electrons help relate
atomic emission spectra, energy states ofatomic emission spectra, energy states of
atoms, and atomic orbitals.atoms, and atomic orbitals.
A set of three rules determines theA set of three rules determines the
arrangement in an atom.arrangement in an atom.
Main IdeasMain Ideas
4. • 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 QuantizedLight and Quantized
EnergyEnergy
Objectives:
5. • 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 andThe Atom and
Unanswered QuestionsUnanswered Questions
6. The Atom andThe Atom and
Unanswered QuestionsUnanswered 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.
7. Wave Nature of LightWave Nature of Light
Electromagnetic radiation, a form of energy
that exhibits wave-like behavior as it travels
through space.
• Visible light
• Microwaves
• X-rays
• Radio waves
8. Wave Nature of LightWave 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.
9. Wave Nature of LightWave 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 LightWave Nature of Light
The speed of light (3.00 × 108
m/s) is the
product of it’s wavelength and frequency
c = λν.
12. Wave Nature of LightWave Nature of Light
• All electromagnetic waves, including visible
light travels at 3.00 x 108
m/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.
13. Wave Nature of LightWave Nature of Light
• The electromagnetic spectrum includes all
forms of electromagnetic radiation.
• Not just visible light.
16. Particle Nature of LightParticle 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 LightParticle Nature of Light
Max Planck (1858-1947) – matter canMax Planck (1858-1947) – matter can
gain or lose energy only in smallgain or lose energy only in small
amounts.amounts.
• E=hv
• Planck’s constant has a value of
6.626 × 10–34
J ● s.
• Energy can only be emitted or absorbed in
whole number multiples of h.
18. Particle Nature of LightParticle 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 LightParticle 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 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.
21. Atomic Emission SpectrumAtomic 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.
22. Atomic Emission SpectrumAtomic Emission Spectrum
• Light in a neon sign is produced when
electricity is passed through a tube filled with
neon gas and excites the neon atoms. The
excited atoms emit light to release energy.
26. ObjectivesObjectives
Define a quantum of energy, andDefine a quantum of energy, and
explain how it is related to anexplain how it is related to an
energy change of matter.energy change of matter.
28. Question?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
29. What is a particle of electromagnetic
radiation with no mass called?
A. beta particle
B. alpha particle
C. quanta
D. photon
Question?Question?
31. Quantum Theory of theQuantum Theory of the
AtomAtom
• 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:
32. Bohr’s Model of the AtomBohr’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.
33. Bohr’s Model of the AtomBohr’s Model of the Atom
• Bohr suggested that an electron moves around the
nucleus only in certain allowed circular orbits.
•The smaller
the electrons
orbit the lower
the atoms
energy state
or level
34. Bohr’s Model of the AtomBohr’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
atoms energy
state or level.
35. Bohr’s Model of the AtomBohr’s Model of the Atom
• Each orbit was given a number, called the quantum
number. The orbit closed to the nucleus is n=1
36. Bohr’s Model of the AtomBohr’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. (Ya, know
the other kind of excited.)
• When electron moves from an excited state to
ground state, a photon is emitted.
37. Bohr’s Model of the AtomBohr’s Model of the Atom
• Change in Energy =
E (higher energy orbit) – E (lower energy orbital)
Ephoton = hv
40. Quantum Mechanical ModelQuantum 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 – odd number of
wavelengths.
42. Quantum Mechanical ModelQuantum 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.
ν represents frequency.
43. Quantum Mechanical ModelQuantum 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 ModelQuantum 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.
46. Quantum Mechanical ModelQuantum Mechanical Model
Schrödinger 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.
• 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.
47. Quantum Mechanical ModelQuantum Mechanical Model
Schrödinger treated electrons as waves in a
model called the quantum mechanical
model of the atom.
• 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.
48. Quantum Mechanical ModelQuantum Mechanical Model
• The wave function predicts a three-dimensional
region around the nucleus called the atomic orbital.
49. Quantum Numbers andQuantum Numbers and
the Revised Modelthe Revised Model
The revised model defines theThe revised model defines the
relationship between an electron’srelationship between an electron’s
energy level, sublevel and atomicenergy level, sublevel and atomic
orbitals.orbitals.
Four quantum numbers make up theFour quantum numbers make up the
identification of each electron in anidentification of each electron in an
atom.atom.
50. Atomic OrbitalsAtomic 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.
57. ObjectivesObjectives
Explain the impact of de Broglie'sExplain the impact of de Broglie's
wave article duality and thewave article duality and the
Heisenberg uncertainty principleHeisenberg uncertainty principle
on the current view of electrons inon the current view of electrons in
atoms.atoms.
58. ObjectivesObjectives
Identify the relationships among aIdentify the relationships among a
hydrogen atom's energy levels,hydrogen atom's energy levels,
sublevels, and atomic orbitals.sublevels, and atomic orbitals.
63. Electron ConfigurationElectron Configuration
Objectives: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.
65. Electron ConfigurationElectron 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.
67. Electron ConfigurationElectron 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 ConfigurationElectron 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.
69. Electron ArrangementElectron Arrangement
-Electron arrangement can be represented-Electron arrangement can be represented
by two common different methods.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. Electron ArrangementElectron Arrangement
-Electron arrangement can be represented-Electron arrangement can be represented
by two common different methods.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 ConfigurationElectron Configuration
• 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.
• Some energy levels overlap. Exceptions for
this start at Vandium, atomic #23.
75. Valence ElectronsValence 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.
76. Electron Dot StructureElectron Dot Structure
Electrons are placed one at a time on theElectrons are placed one at a time on the
four sides of the symbol and then pairedfour sides of the symbol and then paired
until used up. Side order doesn’t matter.until used up. Side order doesn’t matter.
Example:Example:
NaNa
ClCl
78. ObjectivesObjectives
Apply the Pauli exclusion principle,Apply the Pauli exclusion principle,
the aufbau principle, and Hund'sthe aufbau principle, and Hund's
rule to write electron configurationsrule to write electron configurations
using orbital diagrams and electronusing orbital diagrams and electron
configuration notation.configuration notation.
79. ObjectivesObjectives
Define valence electrons, and drawDefine valence electrons, and draw
electron-dot structures representingelectron-dot structures representing
an atom's valence electrons.an atom's valence electrons.
80. Question?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
83. Study GuideStudy 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
84. Study GuideStudy Guide
Key Concepts
• White light produces a continuous spectrum. An
element’s emission spectrum consists of a
series of lines of individual colors.
85. Study GuideStudy 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.
86. Study GuideStudy 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.
87. Chapter QuestionsChapter Questions
The shortest distance from equivalent
points on a continuous wave is the:
A. frequency
B. wavelength
C. amplitude
D. crest
88. Chapter QuestionsChapter Questions
The energy of a wave increases as ____.
A. frequency decreases
B. wavelength decreases
C. wavelength increases
D. distance increases
89. Chapter QuestionChapter 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
90. Chapter QuestionChapter 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
94. Chapter QuestionsChapter 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
96. Chapter QuestionsChapter 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
122. CIMCIM
.11 Balmer Series
.12 Electron Transitions
4 Electron Configurations and Orbital
Diagrams for Elements 1–10
6 Electron Configurations and
Dot Structures
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