Chapter 5
Electrons in Atoms
Section 1
Light and Quantized Energy
Section 1: Light and Quantized Energy
Light, a form of electromagnetic radiation, has characteristics of both a
wave and a particle.
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The Atom and Unanswered Questions
•
Recall that in _____________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.
•
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.
Light and Quantized Energy
The Wave Nature of Light
•
Visible light is a type of _________________________, a form of energy that
exhibits wave-like behavior as it travels through space.
•
All waves can be described by several characteristics.
•
The ____________(λ) is the shortest distance between equivalent points on a
continuous wave.
•
The frequency (_) is the number of waves that pass a given point per second.
•
The ____________is the wave’s height from the origin to a crest.
Light and Quantized Energy
The Wave Nature of Light
•
The speed of light (________________) is the product of it’s wavelength and
frequency c = λν.
Light and Quantized Energy
The Wave Nature of Light
•
Sunlight contains a continuous range of wavelengths and frequencies.
•
A prism separates sunlight into a continuous spectrum of colors.
•
The __________________________includes all forms of electromagnetic
radiation.
Light and Quantized Energy
Properties of Waves
• Wave length:
– The length of each wave
- _________________
Constant: speed of light (c)
3.00 × 108 m/s
Frequency (v)
Hz
Wavelength (λ)
m
Calculations
Microwaves are used to cook food and
transmit information. What is the wavelength
of a microwave that has a frequency of 3.44
× 109 Hz?
KNOWN
UNKNOWN
ν = 3.44 × 109 Hz
c = 3.00 × 108 m/s
λ=?m
Calculations
• What is the frequency of a microwave that
has a wavelength of 6.72 × 10-9 m?
KNOWN
UNKNOWN
λ=
ν=
c=
The Particle Nature of Light
•
The wave model of light cannot explain all of light’s characteristics. Some
examples include:
– Why heated objects emit only certain frequencies of light at a given temperature.
– Why some metals emit electrons when light of a specific frequency shines on them.
Light and Quantized Energy
The Particle Nature of Light
The Quantum Concept
•
In 1900, German physicist Max Planck (1858-1947) began searching for an
explanation of this phenomenon as he studied the light emitted by heated
objects.
•
Planck’s study led him to a startling conclusion:
– Matter can gain or lose _______only in small, specific amounts called ________.
– A quantum is the __________amount of energy that can be gained or lost by an
atom.
– Planck’s constant has a value of ______________________________
Light and Quantized Energy
The Particle Nature of Light
The Photoelectric Effect
•
The _____________________is when electrons are emitted from a metal’s
surface when light of a certain frequency shines on it.
Light and Quantized Energy
The Particle Nature of Light
Light’s Dual Nature
•
Albert Einstein proposed in 1905 that light has a dual nature.
– A beam of light has wavelike and particle-like properties.
– A __________is a particle of electromagnetic radiation with no mass that carries a
quantum of energy.
Light and Quantized Energy
Energy of a Photon
- ________________________
Ephoton
energy
Frequency (v)
Hz
H (Planck’s contstant)
6.626 × 10–34 J ● s.
Calculations
Every object gets its color by reflecting a certain portion of incident
light. The color is determined by the wavelength of the reflected
photons, thus by their energy. What is the energy of a photon from the
violet portion of the Sun’s light if it has a frequency of 7.230 × 1014 s-1?
KNOWN
UNKNOWN
ν = 7.230 × 1014 s-1
Ephoton = ? J
h = 6.626 × 10-34 J•s
Atomic Emission Spectra
•
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 return to their __________state by emitting light to release
energy.
Light and Quantized Energy
Atomic Emission Spectra
•
The atomic __________ spectrum of an element is the set of frequencies of
the electromagnetic waves emitted by the atoms of the element.
•
Each element’s atomic emission spectrum is unique.
Light and Quantized Energy
Essential Questions
• How do the wave and particle natures of light compare?
• How is a quantum of energy related to an energy change of matter?
• How do continuous electromagnetic spectra and atomic emission
spectra compare and contrast?
Light and Quantized Energy
Section 2
Quantum Theory and the Atom
Section 2: Quantum Theory and the Atom
Wavelike properties of electrons help relate atomic emission
spectra, energy states of atoms, and atomic orbitals.
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Bohr’s Model of the Atom
•
Einstein’s theory of light’s dual nature accounted for several unexplainable
phenomena 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.
•
Bohr correctly predicted the frequency lines in hydrogen’s atomic emission
spectrum.
•
The lowest allowable energy state of an atom is called its ______________.
•
When an atom gains energy, it is in an _____________state.
Quantum Theory and the Atom
Bohr’s Model of the Atom
•
Bohr suggested that an electron moves around the __________only in certain
allowed circular orbits.
Quantum Theory and the Atom
Bohr’s Model of the Atom
•
Each orbit was given a number, called the _________________ number.
•
Hydrogen’s single electron is in the n = 1 orbit in the ground state.
•
When energy is added, the electron moves to the n = 2 orbit.
Quantum Theory and the Atom
Atomic orbitals
• ___________is a region of space around the nucleus where an
electron is likely to be found.
• An electron cloud is a good approximation of how electrons
behave in their orbitals
• The level in which an electron has the least energy—the lowest
energy level—has only one orbital. Higher energy levels have
more than one orbital
Electron Configuration
Energy Levels, Orbitals, and Electrons
• The most stable Energy
Level
electron
configuration is
___
the one in which
the electrons are
2
in orbitals with the
3
lowest possible
energies.
4
Number
of
Orbitals
1
Max. # of
Electrons
____
8
9
____
_____
32
2
DRAWING BOHR MODELS OF ATOMS
• Draw the nucleus and label with #p &
#n
• Draw electron orbitals
– 1st orbital can have ____ electrons only
– 2nd and 3rd ring can each have
_____electrons
– Fill – _________________________
– Inner rings must be filled first before any
electron enters a higher ring!!!!!
– Let’s Practice:
Bohr Model - Hydrogen
Bohr Model - Oxygen
Bohr Model - Sulfur
Bohr’s Model of the Atom
The limits of Bohr’s model
•
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 substantial
evidence indicates they do not move around the nucleus in circular orbits.
Quantum Theory and the Atom
The Quantum Mechanical Model of the Atom
•
Heisenberg showed it is impossible to take any measurement of an object
without disturbing it.
•
The _____________ 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.
Quantum Theory and the Atom
The Quantum Mechanical Model of the Atom
•
Schrödinger treated electrons as waves in a model called the quantum
_______________ model of the atom.
•
Schrödinger’s equation applied equally well to elements other than
hydrogen.
•
The wave function predicts a three-dimensional region around the nucleus
called the ______________.
Quantum Theory and the Atom
Hydrogen Atomic Orbitals
•
Principal _________ number (__) indicates the relative size and energy of
atomic orbitals.
•
n specifies the atom’s major energy levels, called the __________ energy
levels.
•
______________________are contained within the principal energy levels.
Quantum Theory and the Atom
Hydrogen Atomic Orbitals
•
Each energy sublevel relates to orbitals of different ______________.
Quantum Theory and the Atom
Hydrogen Atomic Orbitals
Quantum Theory and the Atom
Essential Questions
• How do the Bohr and quantum mechanical models of the atom
compare?
• What is the impact of de Broglie’s wave-particle duality and the
Heisenberg uncertainty principle on the current view of electrons in
atoms?
• What are the relationships among a hydrogen atom’s energy levels,
sublevels, and atomic orbitals?
Quantum Theory and the Atom
Section 3
Electron Configuration
Section 3: Electron Configuration
Three rules are used to determine electron arrangement in an atom.
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Ground-State Electron Configuration
•
The arrangement of electrons in the atom is called the ____________
______________________
•
The ________ principle states that each electron occupies the lowest
energy orbital available.
Electron Configuration
Ground-State Electron Configuration
Electron Configuration
Ground-State Electron Configuration
•
The aufbau diagram can be used to write correct ground-state electron
configurations for all elements up to and including Vanadium, atomic
number 23.
•
The electron configurations for certain transition metals, like chromium and
copper, do not follow the aufbau diagram due to increased stability of halffilled and filled sets of s and d orbitals.
Electron Configuration
Electron Configuration
• _________Configuration - a representation of the arrangement of
electrons in an atom
Electron Configuration
• Examples of electron Configuration
– 1. Li 1s22s1
– 2. C 1s22s22p6
________
_________
# of e- in that
________
Electron Configuration
• Take note that after 4s is filled, 3d is than filled
before 4p.
• …… 6s than 4f than 5d than 6p
• When writing out the electron configuration,
always write your numbers in numerical order
– Y 1s22s22p63s23p64s23d104p65s24d1 – NO!
– Y 1s22s22p63s23p63d104s24p64d15s2
Electron Configuration
• Examples:
• Be
• O
• Ca
• Mn
Electron Configuration
• Examples
• Pb
• Os
Ground-State Electron
Configuration
• _________ Rule (better known as the Bus Rule)
– Before any second electron can be placed in a sub level, all the
orbitals of that sub level must contain at least one electron –
spread out the e- before pairing them up.
• Pauli _________ Principle - electrons occupying the same
orbital must have ______________spin.
• See a good online illustration at
http://www.avogadro.co.uk/light/aufbau/aufbau.htm
Orbital Notation
H
1s
F
1s
2s
2p
Orbital Notation
• Examples:
• Li
F
• Na
Sc
Valence Electrons
•
________electrons are defined
as electrons in the atom’s
outermost orbitals—those
associated with the atom’s
highest principal energy level.
•
___________structure consists
of the element’s symbol
representing the nucleus,
surrounded by dots
representing the element’s
valence electrons.
Electron Configuration
Significance of Electron Configurations
•
•
•
•
Valence shell electrons - ______________electrons involved with bonding
no atom has more than _______________________
Noble gases - ___valence electrons – _____reactive of all elements
Lewis Dot structures: NSEW (cheating) also show correct way, count to 8
Lewis Dot Structures
Essential Questions
• How are the Pauli exclusion principle, the aufbau principle, and
Hund’s rule used to write electron configurations?
• How do electron-dot structures represent an atom’s valence
electrons?
Electron Configuration