Atomic History and
Structure of the Atom
Chapters 4 and 5
Early Theories of Matter
• The concept of the atom developed over a
few thousand years.
• Consider that there were no controlled
experiments and few tools for scientific
exploration.
•2
• The power of the mind and intellectual
thought were the primary ways to truth
about the universe.
• Curiosity drove discovery
• Philosophers, scholarly thinkers, speculated
about the nature of matter.
• Ideas came from their own life experiences.
•3
Early Greek philosophers thought that matter
was made of Earth, Wind, Water, and Fire.
•4
• Commonly accepted that matter could
be divided endlessly into smaller and
smaller pieces…..
– The Continuous Theory of Matter
• No methods to test the validity of these
ideas.
•5
Democritus
• a Greek Philosopher
• the first to propose
the idea that matter
was not infinitely
divisible.
discontinuous theory
• He proposed the idea
that matter was
made up of tiny
individual particles
called atomos.
•6
• The English word for atom comes from
this Greek word atomos.
• He believed that atoms could not be
created, destroyed, or further divided
(discontinuous theory).
•7
• His belief in atoms was amazingly
accurate.
• However, his ideas were met with
much criticism from other philosophers.
• The harshest criticism was from the
influential Greek philosopher, Aristotle.
•8
Aristotle
• Aristotle rejected the
theory of atoms
because it did not
agree with own ideas
on nature.
• He did not believe
that atoms could
move through empty
space.
• He did not believe
that the
“nothingness” of
empty space
could
•9
exist.
Democritus Rejected
• Because Democritus had no way to answer
the challenges to his ideas, his theory was
eventually rejected.
• It is important to realize that these ideas
were just that—ideas and not science.
• Incredibly, the ideas of Aristotle were so
great and the science so primitive that his
denial of the existence of the atom went
unchallenged for two thousand years
•10
John Dalton
• In the 19th century John Dalton
– an English School teacher
– revived and revised the
work of Democritus.
– This time it was based on
scientific research
conducted by Dalton.
•11
Dalton’s Atomic Theory
–
–
1808
marks the beginning of modern atomic
theory.
•12
Dalton was wrong about:
1.the atom being divisible (atoms are
divisible into subatomic particles)
2.all atoms of an element having
identical properties. (atoms of an
element can have slightly different
masses.)
•13
Dalton’s Model of the Atom
•14
Next Set of Questions
1. What is an atom like?
2. How are atoms shaped?
3. Is the atom composed of other
particles?
•15
The Discovery of the Electron
• British scientist named J. J.Thompson
• Used a glass tube, called a
cathode-ray tube, connected to
two electrodes.
•16
Cathode Ray Tube
•17
• One electrode was attached to one
end of the tube.
• It was negative in charge and was thus
called the cathode.
• At the other end was attached a
positive electrode, the anode.
•18
•19
Thomson’s Experiment
• Thomson placed a paddle wheel inside the
tube, between the cathode and the
anode.
• He passed an electric current through a
variety of gases in the tube. The gases
within the tube were changed with each
experiment.
•20
Thomson made the following
observations
1. The current passing through the tube creates
different colors in different gases.
2. The anode (positive end) glows where as the
cathode (negative end) does not.
3. An object placed within the cathode-ray casts a
shadow.
4. The paddle-wheel spins in the direction of the
anode.
•21
Conclusions:
1. The different colors that were created by using
different gases showed that atoms of different
elements possessed different energies.
2. The cast shadow was thought to be due to the
beam of light created by the cathode-ray.
However, the experiment made with the
spinning paddle-wheel showed that the
cathode-ray was composed of particles.
3. The fact that the anode glowed and the
cathode did not shows that the cathode-ray
travels from a negative potential to a positive
potential. This was one of the most important
findings of J. J. Thomson's experiment and was
repeated using another apparatus. •22
• To test the polarity of the cathode ray,
Thompson replaced the paddle-wheel with
another pair of electrodes.
• When the current was passed through this new
apparatus, Thompson found that the cathoderay was "bent" towards the positive electrode
and repelled by the negative electrode.
• Thompson attributed the "bending" of the
cathode-ray to charge-charge repulsion of the
second cathode and the negatively charged
particles in the cathode-ray
•23
•24
Electrons
• He named these negatively charged
particles electrons.
• Since Thompson could get these results
regardless of what element he used
within the tube, he concluded that all
atoms have electrons.
•25
Thomson’s Model of the Atom
• The Plum Pudding Model
•26
Oil Drop Experiment
• In 1909, American scientist, Robert Millikan,
determined the charge of an electron in his
famous Oil Drop Experiment.
•27
Discovery of the Nucleus
• 1911 Ernst Rutherford
• Rutherford’s Gold Foil Experiment
– determined that nearly all of the mass of
the atom consisted of a positively
charged nucleus
•28
Rutherford’s Gold Foil Experiment
• He did this by measuring the deflection of helium
nuclei by gold atoms.
•29
Predicted Results
•30
•
While most atoms were not deflected
much at all, a few were deflected by 180
degrees.
•31
How was this possible?
• Rutherford calculated that the only
way this was possible was if the gold
atoms consisted of a cloud of
electrons surrounding the very dense,
positively charged nucleus, for only
then could the gold atom transfer
enough momentum to the ions to turn
them around.
•32
Rutherford’s Conclusions
• Atoms are mostly empty space and
the nucleus is incredibly dense.
– If the nucleus were a marble, the atom
would be the size of a football field
•33
Rutherford ‘s Nuclear Model
• By 1919 Rutherford
refined his concept of
the nucleus and
concluded that it
contained protons.
• Protons have a
charge of +1.
• This model could not
account for the mass
of the atom.
• 20 more years before
this mystery was
solved.
•34
Discovery of the Neutron
• 1932
• James Chadwick, an English Physicist
• He showed that the nucleus also contained
another particle, the neutron
• The neutron is neutral, meaning it has no
charge.
• The neutron has a mass equal to the proton.
•35
Properties of Subatomic Particles
Properties of Subatomic Particles
Particle
Electron
Electrical
charge
Symbol

e
1
Relative
mass
1/1840
Actual
mass
(g)
9.111028
Proton
p
1
1
1.6731024
Neutron
n0
0
1
1.6751024
•36
Nuclear Model of the Atom
• Atom: an electrically neutral particle
composed of protons, neutrons, and
electrons.
• Atoms are spherical in shape
• tiny, dense nucleus of positive charge
surrounded by one or more negatively
charged electrons
• Nucleus contains 99.7 % of the mass of an
atom.
•37
Subatomic Particles
• An Element consists of atoms with the same
number positive charges.
• Each atom has a unique number of positive
charges in its nucleus. (Moseley)
• These positive charges are protons.
• The number of Protons in the nucleus
determines the identity of an atom.
• Atomic Number = the number of protons
•38
Atomic Number(Z): the number of protons
in the nucleus of each atom of the element
• Elements are arranged in the
periodic table from left to right in
order of increasing atomic number.
• Atomic number identifies the
element
•39
Atoms are Neutral
Atoms have no charge. So we know
that # of positive charges and the # of
negative charges must be equal.
# of protons = # of electrons
•40
Isotopes and Mass Number
• We also know that elements have
neutrons.
• Neutrons are neutral. (no charge)
• Neutrons contribute only mass to an
atom.
• Naturally occurring elements are a
mixture of atoms that have different
numbers of neutrons.
• Atoms with the same number of
protons, but different number of
neutrons are called Isotopes.•41
Isotopes and Mass Number
• Isotopes: atoms of the same element that
have different masses.
• Different number of neutrons so mass is
different
• Atoms of different isotopes have different
masses so identity is given by name and
mass #.
• Mass Number = the total number of protons
and neutrons in the nucleus of an isotope
•42
Calculating the number of neutrons
• # neutrons = Mass # - Atomic #
Isotope Symbols
Consists of three parts
1. the symbol of the element
2. the atomic number of the element
3. the mass number of the specific
isotope.
Isotope Names
Element Name-Mass Number
Helium-4
•45
What is the name of this isotope?
Counting the subatomic particles in an isotope
Atomic Number: 47
Mass Number 108
# of Protons 47
# of Electrons 47
# of Neutrons: 61
Counting the subatomic particles in an isotope
Atomic Number: 3
Mass Number 7
# of Protons 3
# of Electrons 3
# of Neutrons: 4
Counting the subatomic particles in an isotope
Atomic Number: 22
Mass Number 48
# of Protons 22
# of Electrons 22
# of Neutrons: 26
Draw the isotope symbol for
calcium with 21 neutrons.
Comparing Potassium Isotopes
Potassium-41
• protons = ________
• electrons = ______
• neutrons = _______
Potassium-40
• protons = ________
• electrons = ______
• neutrons = _______
Potassium-39
• Protons = __________
• Electrons = _________
• Neutrons = __________
Average Atomic Mass
• Naturally occurring elements are a
mixture of atoms that have different
numbers of neutrons
• The Average Atomic Mass of an
element is given on the periodic
table.
• Atomic masses measured in grams
are very small.
• Example: An atom of Oxygen-16
has a mass of 2.657 x 19-23grams.
•52
To set up a relative scale of
atomic masses:
1. We measure atomic masses in atomic mass
units.
2. One atom is chosen as a standard and
assigned a relative mass value
3. All the other masses are expressed in
relation to this defined standard.
4. Carbon-12 is the chosen atomic mass
standard.
5. A single atom of C-12 is assigned a mass of
exactly 12 atomic mass units(u).
An AMU
• One amu is exactly 1/12th of the mass
of a carbon-12 atom
– 1.660 5402 x 10-24grams.
• Atomic Mass: The mass of an atom
expressed in atomic mass units.
•54
Average Atomic Mass
• Average Atomic Mass
– given on the periodic table
– weighted average for the naturally occurring mixtures of
isotopes in each element
• Average Atomic Mass depends on
– mass and relative abundance of the isotopes.
– How big and how much
•55
Calculating average atomic mass:
Step 1:
Isotope 1: Atomic mass x relative abundance = mass contribution 1
Isotope 2: Atomic mass x relative abundance = mass contribution 2
Step 2: Add the mass contributions results.
•56
Example Problem:
Naturally occurring copper consists of
69.17% Cu-63, mass of 62.939 598u and
30.83% Cu-65, mass of 64.927 793u
Cu-63
Cu-65
0.6917 x 62.939 598u = 43.535u
0.3083 x 64.927 793u = 20.017
63.552u
•57
Example Problem 2
Gallium occurs in nature as a mixture of two
isotopes. They are Ga-69 with a 60.11% abundance
and a mass of 68.93 amu and Ga-71 with a 39.89%
abundance and an atomic mass of 70.93. Calculate
the atomic mass of gallium.
Example 3
The element chlorine occurs in nature as a mixture of two
isotopes. Chlorine-35 has an atomic mass of 34.7 amu and
makes up 75.8% of chlorine atoms. Chlorine-37 atoms make
up the remaining 24.2% of all chlorine. Use the average
atomic mass of chlorine from the periodic table to calculate
the atomic mass of Cl-37 atoms.
•59
Example 4
The atomic mass of bromine given in the periodic
table is 79.904 amu, which is very close to 80 amu.
Use a reference book to find the percent of Br-80 in
naturally occurring bromine. Explain the value of the
atomic mass of bromine from the data you find.
•Naturally Occurring Isotopes
Mass Number
Natural Abundance
Half-life
79
50.69%
STABLE
81
49.31%
STABLE
Part 2 Atomic Theory
ELECTRON STRUCTURE
The Atom and Unanswered Questions
• Although three subatomic particles had been
discovered by the early 1900s, the quest to
understand the atom and its structure had just
begun.
1. How are electrons arranged in an atom?
2. How does that arrangement play a role
in chemical behavior?
Comparing Element Properties
Chlorine
Argon
Atomic #
Protons
Electrons
Location on Table
Group/Family
State
Properties
Atomic #
Protons
Electrons
Location on Table
Group/Family
State
Properties
Potassium
Atomic #
Protons
Electrons
Location on Table
Group/Family
State
Properties
Why do the elements behave
differently?
• Electron Structure
• The arrangement of the electrons in the atom
The Atom and Unanswered Questions
• Rutherford’s Model
• all of an atom’s positive
charge and virtually all of
its mass are concentrated in
a nucleus that is surrounded
by fast-moving electrons.
• A major scientific development
• Won Nobel Prize 1908
• Many scientists in the early twentieth century found
Rutherford’s nuclear atomic model to be fundamentally
incomplete.
What Did Rutherford’s Model Lack?
To Physicists:
• Did not explain how the
atom’s electrons are
arranged in the space around
the nucleus.
• Did not address why the
negatively charged electrons
are not pulled into the
atom’s positively charged
nucleus.
To Chemists:
• Could not account for the
differences in chemical
behavior among the various
elements.
The Atom and Unanswered Questions
• In the early 1900s, scientists began to unravel the
puzzle of chemical behavior.
• They had observed that 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 better understand this relationship
and the nature of atomic structure, it will be
helpful to first understand the nature of light.
Wave Nature of Light
• Visible Light is a form of Electromagnetic
Radiation.
• All waves can be described by the following
characteristics
Wavelength
• Wavelength (represented by λ, the Greek
letter lambda) is the shortest distance between
equivalent points on a continuous wave.
Wavelength
• Measured from crest to crest or from trough
to trough.
• Expressed in meters, centimeters, or
nanometers (1nm = 1 x 10–9 m).
Frequency
• Frequency (represented by ν, the Greek
letter nu) is the number of “waves” that pass
a given point per second.
• One hertz (Hz), the SI unit of frequency,
equals one wave per second.
Frequency
• In calculations, frequency is expressed with
units of “waves per second,”
(
)
or (s–1) where the term “waves” is understood.
Amplitude
• The amplitude of a wave is the wave’s
height from the origin to a crest, or from
the origin to a trough.
Speed of Light
• All electromagnetic waves travel at a speed of
3.00 x 108 m/s in a vacuum.
• Symbol is c
• speed of light is the product of its
wavelength (λ) and its frequency (ν).
Wave nature of Light
• Although the speed of all electromagnetic
waves is the same, waves may have
different wavelengths and frequencies.
• As you can see from the equation,
wavelength and frequency are inversely
related; in other words, as one quantity
increases, the other decreases.
Wave Nature of Light
Calculating Wavelength of an EM Wave
• Microwaves are used to transmit information.
• What is the wavelength of a microwave
having a frequency of 3.44 x 109 Hz?
• Solve the equation relating the speed,
frequency, and wavelength of an
electromagnetic wave for wavelength (λ).
Light Passing through a Prism
The Continuous Spectrum also
called
The Electromagnetic Spectrum
Particle Nature of Light
• While considering light as
a wave does explain much
of its everyday behavior, it
fails to adequately describe
important aspects of light’s
interactions with matter.
Particle Nature of Light
• The wave model of light cannot
• explain why heated objects emit only
certain frequencies of light at a given
temperature
• or why some metals emit electrons when
colored light of a specific frequency shines on
them (The photoelectric effect)
• Obviously, a totally new model or a revision
of the current model of light was needed to
address these phenomena.
The quantum concept
• In 1900, the German physicist Max Planck
(1858–1947) began searching for an
explanation as he studied the light emitted
from heated objects.
• His study of the phenomenon led him to
a startling conclusion:
• Matter can gain or lose energy only in
small, specific amounts called quanta.
The quantum concept
• 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.
• Planck’s constant has a value of
6.626  10–34 J ● s.
The quantum concept
• Planck’s constant has a value of
6.626 x 10–34 J · s, where J is the symbol
for the joule, the SI unit of energy.
• Looking at the equation, you can see that
the energy of radiation increases as the
radiation’s frequency, v, increases.
The quantum concept
• According to Planck’s theory, for a given
frequency, ν, matter can emit or absorb
energy only in whole-number multiples of
hν; that is, 1hν, 2hν, 3hν, and so on.
• Matter can have only certain amounts of
energy—quantities of energy between
these values do not exist.
The photoelectric effect
• Scientists knew that the wave model (still
very popular in spite of Planck’s proposal)
could not explain a phenomenon called the
photoelectric effect.
The photoelectric effect
• In the photoelectric effect, electrons, called
photoelectrons, are emitted from a metal’s
surface when light of a certain frequency
shines on the surface.
The photoelectric effect
• Albert Einstein proposed in 1905 that
light has a dual nature.
• A beam of light has wavelike and
particle-like properties.
• A photon is a particle of electromagnetic
radiation with no mass that carries a
quantum of energy.
Calculating the Energy of a Photon
• Tiny water drops in
the air disperse the
white light of the
Sun into a rainbow.
• What is the energy
of a photon from the
violet portion of the
rainbow if it has a
frequency of 7.23 x
1014 s–1?
Calculating the Energy of a Photon
• Substitute the known values for frequency
and Planck’s constant into the equation
relating energy of a photon and frequency.
Multiply the known values and cancel units.
Atomic Emission Spectra
• The atomic emission spectrum of an element
is the set of frequencies of the electromagnetic
waves emitted by atoms of the element.
• Light is produced when electricity is
passed through a tube filled with
hydrogen gas and excites the hydrogen
atoms.
•• The
Theexcited
excitedatoms
atomsemit
emitlight
lighttotorelease
release
energy.
energy.
Hydrogen’s Emission Spectrum
Atomic Emission Spectra
• Hydrogen’s atomic emission spectrum
consists of several individual lines of color,
not a continuous range of colors as seen in the
visible spectrum.
• Each element’s atomic emission spectrum is
unique and can be used to determine if that
element is part of an unknown compound.
Atomic Emission Spectra
• An atomic emission spectrum is characteristic
of the element being examined and can be
used to identify that element.
• The fact that only certain colors appear in an
element’s atomic emission spectrum means
that only certain specific frequencies of light
are emitted.
• http://www.trschools.com/staff/g/cgirtain/Web
labs/spectrolab.htm
Atomic Emission Spectra
• And because those emitted frequencies of
light are related to energy by the formula
Ephoton = hν, it can be concluded that only
photons having certain specific energies
are emitted.
Atomic Emission Spectra
• Scientists found atomic emission spectra
puzzling because they had expected to
observe the emission of a continuous series
of colors and energies as excited electrons
lost energy and spiraled toward the nucleus.
Why are elements’
atomic emission spectra
discontinuous rather than
continuous?
The Bohr Model of the Atom
• Niels Bohr, a young Danish physicist
working in Rutherford’s laboratory in 1913,
proposed a quantum model for the hydrogen
atom that seemed to answer this question.
• Impressively, Bohr’s model also correctly
predicted the frequencies of the lines in
hydrogen’s atomic emission spectrum.
Energy States of Hydrogen
• Building on Planck’s and Einstein’s
concepts of quantized energy (quantized
means that only certain values are allowed),
Bohr proposed that the hydrogen atom has
only certain allowable energy states.
• The lowest allowable energy state of an
atom is called its ground state.
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.
• Bohr suggested that an electron moves
around the nucleus only in certain
allowed circular orbits.
The Bohr Model
• Each orbit was given a number, called the
quantum number.
The Bohr Model
• 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.
Bohr's Model of the Atom
An explanation of
hydrogen’s line spectrum
• The four electron
transitions that
account for visible
lines in hydrogen’s
atomic emission
spectrum are shown.
Bohr's Model of the Atom
• 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.
The Quantum Mechanical
Model of the Atom
• Scientists in the mid-1920s, by then
convinced that the Bohr atomic model was
incorrect, formulated new and innovative
explanations of how electrons are arranged
in atoms.
• In 1924, a young French graduate student in
physics named Louis de Broglie (1892–1987)
proposed an idea that eventually accounted
for the fixed energy levels of Bohr’s model.
Electrons as waves
• De Broglie had been
thinking that Bohr’s
quantized electron orbits
had characteristics similar
to those of waves.
• Louis de Broglie introduces
the wave/particle duality of
matter (1921)
• Traditional (classical)
physics had assumed that
particles were particles and
waves were waves and
that’s that. However, de
Broglie suggested that
particles could sometimes
behave as waves and waves
could sometimes behave as
particles
The Quantum Mechanical Model of the Atom
• The de Broglie equation predicts that all
moving particles have wave
characteristics.
 represents wavelengths
h is Planck's constant.
m represents mass of the particle.
 represents frequency.
The Quantum Mechanical Model of the Atom
• The figure illustrates that electrons orbit
the nucleus only in whole-number
wavelengths.
Electrons as waves
• Step by step, scientists such as Rutherford,
Bohr, and de Broglie had been unraveling
the mysteries of the atom.
• However, a conclusion reached by the
German theoretical physicist Werner
Heisenberg (1901–1976), a contemporary
of de Broglie, proved to have profound
implications for atomic models.
The Heisenberg Uncertainty Principle
• Heisenberg concluded that it is impossible to
make any measurement on an object without
disturbing the object—at least a little.
• The act of observing the electron produces a
significant, unavoidable uncertainty in the
position and motion of the electron.
The Heisenberg Uncertainty Principle
• Heisenberg’s analysis of interactions such
as those between photons and electrons led
him to his historic conclusion.
• 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.
• Classical physics had always assumed that
precise location and velocity of objects was
always possible.
• Heisenberg, however discovered that this was
not necessarily the case at the atomic level.
• In particular, he stated that the act of
observation interfered with the location and
velocity of small particles such as electrons.
• This is the case because observation requires
light and light has momentum. When light
bounces off an electron momentum exchange
can occur between light and the electron which
means the electrons location and velocity have
been altered by the act of measurement. This
scenario has important implications to what we
can measure at the atomic level.
The Schrödinger wave equation
• In 1926, Austrian physicist Erwin
Schrödinger (1887–1961) furthered the waveparticle theory proposed by de Broglie.
• Schrödinger derived an equation that treated
the hydrogen atom’s electron as a wave.
• Remarkably, Schrödinger’s new model for
the hydrogen atom seemed to apply equally
well to atoms of other elements—an area in
which Bohr’s model failed.
The Schrödinger wave equation
• The atomic model in which electrons are
treated as waves is called the wave
mechanical model of the atom or, more
commonly, the quantum mechanical
model of the atom.
The Schrödinger wave equation
• Like Bohr’s model, the quantum mechanical
model limits an electron’s energy to certain
values.
• However, unlike Bohr’s model, the
quantum mechanical model makes no
attempt to describe the electron’s path
around the nucleus.
The Schrödinger wave equation
• A three-dimensional region around the
nucleus called an atomic orbital describes
the electron’s probable location.
• You can picture an atomic orbital as a fuzzy
cloud in which the density of the cloud at a
given point is proportional to the probability
of finding the electron at that point.
The Schrödinger wave equation
• This electron density
diagram for a
hydrogen atom
represents the
likelihood of finding
an electron at a
particular point in the
atom.
• The greater the density of the dots, the greater
the likelihood of finding hydrogen’s electron.
The Schrödinger wave equation
• The boundary of
an atom is
defined as the
volume that
encloses a 90%
probability of
containing its
electrons.
Hydrogen’s Atomic Orbitals
• Because the boundary of an atomic orbital is
fuzzy, the orbital does not have an exactly
defined size.
• To overcome the inherent uncertainty about
the electron’s location, chemists arbitrarily
draw an orbital’s surface to contain 90% of
the electron’s total probability distribution.
Hydrogen’s Atomic Orbitals
• In other words, the electron spends 90% of
the time within the volume defined by the
surface, and 10% of the time somewhere
outside the surface.
Hydrogen’s Atomic Orbitals
• Recall that the Bohr atomic model assigns
quantum numbers to electron orbits.
• In a similar manner, the quantum mechanical
model assigns principal quantum numbers
(n) that indicate the relative sizes and
energies of atomic orbitals.
Hydrogen’s Atomic Orbitals
• That is, as n increases, the orbital becomes
larger, the electron spends more time
farther from the nucleus, and the atom’s
energy level increases.
• Therefore, n specifies the atom’s major
energy levels, called principal energy
levels.
Hydrogen’s Atomic Orbitals
• An atom’s lowest principal energy level is
assigned a principal quantum number of one.
• When the hydrogen atom’s single electron
occupies an orbital with n = 1, the atom is in
its ground state.
• Up to seven energy levels have been detected
for the hydrogen atom, giving n values
ranging from 1 to 7.
Hydrogen’s Atomic Orbitals
• Principal energy levels contain energy
sublevels.
• Principal energy level 1 consists of a single
sublevel, principal energy level 2 consists of
two sublevels, principal energy level 3
consists of three sublevels, and so on.
Hydrogen’s Atomic Orbitals
• To better understand the relationship between
the atom’s energy levels and sublevels, picture
the seats in a
wedgeshaped
section of a
theater.
Hydrogen’s Atomic Orbitals
• As you move away from the stage, the rows
become higher and contain more seats.
• Similarly,
the number
of energy
sublevels in
a principal
energy level
increases as
n increases.
Hydrogen’s Atomic Orbitals
• Sublevels are labeled s,
p, d, or f according to
the shapes of the
atom’s orbitals.
• All s orbitals are
spherical and all p
orbitals are dumbbell
shaped; however, not
all d or f orbitals have
the same shape.
Hydrogen’s Atomic Orbitals
• Each orbital may
contain at most
two electrons.
• The single
sublevel in
principal energy
level 1 consists of
a spherical orbital
called the 1s
orbital.
Hydrogen’s Atomic Orbitals
• The two sublevels in principal energy level 2
are designated 2s and 2p.
• The 2s sublevel
consists of the 2s
orbital, which is
spherical like the 1s
orbital but larger in
size.
Hydrogen’s Atomic Orbitals
• The 2p sublevel consists of three dumbbellshaped p orbitals of equal energy designated
2px, 2py, and 2pz.
• The subscripts x,
y, and z merely
designate the
orientations of p
orbitals along
the x, y, and z
coordinate axes.
•
•
•
•
Hydrogen’s Atomic Orbitals
Principal energy level 3 consists of three
sublevels designated 3s, 3p, and 3d.
Each d sublevel
consists of five
orbitals of equal
energy.
Four d orbitals have identical shapes but
different orientations.
However, the fifth, dz2 orbital is shaped and
oriented differently from the other four.
Scientist conclude that
• Electrons occupy energy levels.
• That is they must have certain amounts of
energy and no others
• Energy is said to be Quantized
• Quantized: to have a certain specific
quantity
• Within energy levels are sublevels
Energy States
• Excited State: An atom has one or more of
its electrons in a higher energy level than
the lowest state.
• Ground State: All electrons are in the
lowest energy levels possible
Electron Configurations
•
•
•
•
•
•
Energy levels are numbered:
n=1
n=2
n=3
n=4
These are known as the Principal
Quantum Number
Sublevels
•
•
•
•
When n = 1 there is 1 sublevel
When n = 2 there are 2 sublevels
When n = 3 there are 3 sublevels
When n = 4 there are 4 sublevels
Sublevel Designations
•
•
•
•
n = 1 the sublevel is denoted by the letter s
n = 2 the sublevels are s and p
n = 3 the sublevels are s and p and d
n = 4 the sublevels are s and p and d and f
Numbers of electrons
•
•
•
•
s sublevels can hold 2 electrons
p sublevels can hold 6 electrons
d sublevels can hold 10 electrons
f sublevels can hold 14 electrons
The electron structure of the atom can
be described as
• A number (1,2,3) denotes the quantum shell
• A letter (s,p,d,f) denotes the sublevel
• A superscript indicates the number of
electrons in the sub level.
• 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f
Orbitals
• Regions that enclose the electron cloud
• Regions of the charge cloud where there is
the highest probability of there being an
electron
• Orbitals can hold a maximum of two
electrons
• Two electrons in the same orbital will spin
in opposite directions
Orbital Shapes
• s orbitals are spherical in shape
• p orbitals are dumb-bell shaped
each p sub-shell has 3 p orbitals
• d and f orbitals are very complex
each d sublevel has 5 d orbitals
each f sublevel has 7 orbitals
Summary Chart of the QMM
Principle
Energy
Level (n)
Sublevels
Number of
orbitals
Shape
Max # of Electrons
n=1
1 Sublevel
s
1
Spherical
2
n=2
2 sublevels
s and p
1 and 3
Spherical
and
Dumbbell
2 and 6
n=3
3 sublevels
s, p, and d
1,3,5
spherical
dumbbell
complex
2, 6, 10
n=4
4 sublevels
s, p, d, f
1, 3, 5, 7
spherical
dumbbell
complex
complex
2,6, 10, 14
Electron Configurations
• The arrangement of electrons in an atom
• Determines chemical reactivity
Rules for Determining Electron Configurations
• AufbauPrinciple
Electrons will occupy the
lowest energy levels first.
• Hund’s Rule
Orbitals of equal energy
are each occupied by
one electron before any
one orbital gets two.
• Pauli Exclusion
Principle
Electrons in the same
orbital will spin in
opposite directions.
The electron filling order
• 1s22s22p63s23p64s23d104p65s24d105p66s24f14
5d106p67s25f146d107p6
• Electrons will occupy the lowest energy levels
first.
Why do we see 4s appear in the
filling order before 3d??
• Scientist have determined that the s sublevel
is slightly lower in energy than the d
sublevel
• This is called Overlap.
• Seen at 4s and 3d 5s and 4d
6s and
4f and 5d
Three Methods to Represent
Electron Configurations
• Electron Configuration Notation
• Orbital Notation
• Nobel Gas Configurations
Electron Configuration Notation
•
•
•
•
•
•
•
1s2s2p3s3p4s3d4p5s4d
Write the configuration for:
H
He
Li
Be
B
Orbital Notation
• Line or box represents the orbital and up
and down arrows represent a pair of
electrons of opposite spin
• Draw an orbital diagram for:
• H
• He
• Li
Orbital Diagrams
•
•
•
•
•
•
•
Be
B
C
N
O
F
Ne
•http://www.iun.edu/~cpanhd/C101w
ebnotes/modern-atomictheory/aufbau-principle.html
Short Hand Notation
• Use the noble gas that precedes the atom you’re
writing
• Put the noble Gas in brackets
• Finish the configuration
• K
• Ca
• Y
• Al
Valence Electrons
• The outermost electrons
• Electrons in the highest principal energy
level
• Determine the chemical reactivity of the
element
• Involved in forming chemical bonds
Representing valence electron
structure visually
• Electron Dot
Diagrams
– Also called Lewis
Structures
• The symbol
represents the Kernel
(non valence electrons
and the nucleus)
• Surrounding Dots
represent the valence
electrons
Examples Lewis Dot Structures