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Chapter 9 Lecture
Conceptual
Integrated Science
Second Edition
Atoms and the
Periodic Table
© 2013 Pearson Education, Inc.
This lecture will help you understand:
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Atoms Are Ancient and Empty
The Elements
Protons and Neutrons
Isotopes and Atomic Mass
The Periodic Table
Physical and Conceptual Models
Identifying Atoms Using the Spectroscope
The Quantum Hypothesis
Electron Waves
The Shell Model
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Atoms Are Ancient and Empty
• Atoms are
– ancient.
• The origin of most atoms goes back to the birth of
the universe.
– mostly empty space.
• Elements heavier than hydrogen and much of
the helium were produced in the interiors of
stars.
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Atoms Are Ancient and Empty
CHECK YOUR NEIGHBOR
Which of these statements about the atom are
incorrect?
A. Atoms have been around since the beginning
of the universe.
B. Atoms are mostly empty space.
C. Atoms are perpetually moving.
D. Atoms are manufactured in plants and in
humans during pregnancy.
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Elements
• An element is a material made of only one kind
of atom. For example, pure gold is an element
because it is made of only gold atoms.
• An atom is the fundamental unit of an element.
The term "element" is used when referring to macroscopic
quantities.
The term "atom" is used when discussing the submicroscopic.
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The Elements
• Atoms:
– Atoms make up all matter around us.
– To date, 115 distinct kinds of atoms are
known—90 are found in nature, and the rest
are synthesized.
• An element is any material that consists of only
one type of atom.
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Protons and Neutrons
• Protons
– carry a positive charge—the same quantity of
charge as electrons.
– are about 1800 times as massive as
electrons.
– have the same number of protons in the
nucleus as electrons that surround the
nucleus of an electrically neutral atom.
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Protons and Neutrons
• Positively charged
central core –
nucleus
– Protons – carry
positive charge. Each
proton has a charge
equal to the charge
of an electron.
– Neutrons – no
electrical charge.
Mass slightly more
than protons.
Protons and Neutrons
• Electrons
– are identical.
– repel the electrons of neighboring atoms.
– have electrical repulsion that prevents atomic
closeness.
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Protons and Neutrons
• Electrons
– move in the space around the
nucleus.
– Negative charge attracts to
the nucleus’ positive charge.
• In a neutral atom –
number of protons =
number of electrons.
• Protons have a much
greater mass than
electrons.
Protons and Neutrons
• The atomic number is the number of protons in
each element listed in the periodic table.
• Neutrons
– accompany protons in the nucleus.
– have about the same mass as protons but no
charge, so they are electrically neutral.
• Both protons and neutrons are nucleons.
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The structure of the Atom
• Charge
– Proton = +1
– Electron = -1
– Neutron = 0
• Mass
– Proton = 1 amu
– Electron = 0 amu
– Neutron = 1 amu
Atomic Numbers
• Henry Moseley (1887 – 1915)
– Found that atoms of each element contain a unique positive
charge in their nucleus.
– Atom’s identity comes from the number of protons in its nucleus.
• Atomic number – number of protons
– Unique to every element.
– Example : Nitrogen = 7 protons. Hydrogen = 1 proton.
– In a neutral atom the number of protons always equals number
of electrons.
Sample Problem #1
• How many protons and electrons are present in
a Sodium atom?
Ions
• An atom with an electrical charge – caused by
loss or gain of an electron.
– An atom with more electrons than protons = negative
charge
– An atom with less electrons than protons = positive
charge
Ions
• Ion charge= # protons - # electrons
Ions
Sample Problem #2
• What is the chemical symbol for the ion with 18
protons and 15 electrons?
Isotopes
• All atoms of a given element have the same
number of protons
• However, all atoms of an element do not
necessarily have the same number of neutrons.
Isotopes
• Isotopes – atoms that have the same number of
protons but different numbers of neutrons.
Isotopes
• In nature elements are found as mixtures of
isotopes, but always in the same percentages.
– Example: hydrogen found as 99.9844% atoms with no
neutrons and 0.0156% atoms with one neutron.
• Isotopes are almost indistinguishable.
• Chemical properties depend primarily on electrons
and protons.
• Difference in mass – “heavy,” “light”
Isotopes
Isotopes
Isotopes
• Identifying isotopes:
– Number added after the element’s
name.
– Isotope’s mass number – the sum of the
isotope’s number of protons and
neutrons.
– Example: chlorine-37, 17 protons and
20 neutrons.
Isotopes
Isotopes
Sample Problem #3
Isotopes and Atomic Mass
• Isotopes
– are atoms of the same element that contain the same
number of protons but different numbers of neutrons
in the nucleus.
– are identified by mass number, which is the total
number of protons and neutrons in the nucleus.
– differ only in mass, not electric charge; therefore,
isotopes share many characteristics.
• Total number of neutrons in isotope:
Mass number  atomic number
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Isotopes and Atomic Mass
• Atomic mass is
– the total mass of the atom(s) [protons,
neutrons, and electrons].
– listed in the periodic table in atomic mass
units.
• One atomic mass unit is equal to
1.661  10–24 gram or 1.661  10–27 kg.
© 2013 Pearson Education, Inc.
Isotopes and Atomic Mass
CHECK YOUR NEIGHBOR
The atomic number of an element matches the
number of
A.
B.
C.
D.
protons in the nucleus of an atom.
electrons in a neutral atom.
both of the above
none of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
Isotopes and Atomic Mass
CHECK YOUR NEIGHBOR
A nucleus with an atomic number of 44 and a
mass number of 100 must have
A.
B.
C.
D.
44 neutrons.
56 neutrons.
100 neutrons.
none of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Periodic Table
• The periodic table is a listing of all the known
elements.
• It is NOT something to be memorized.
• Instead, we learn how to READ the periodic
table.
• A chemist uses the periodic table much as a
writer uses a dictionary. NEITHER needs be
memorized!
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The Periodic Table
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The Periodic Table
• The elements are highly organized within the
periodic table.
• Each vertical column is called a group/family.
• Each horizontal row is called a period.
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The Periodic Table
© 2013 Pearson Education, Inc.
The Periodic Table
© 2013 Pearson Education, Inc.
The Periodic Table
© 2013 Pearson Education, Inc.
The Periodic Table
© 2013 Pearson Education, Inc.
The Periodic Table
CHECK YOUR NEIGHBOR
Which is larger: a lithium atom or a fluorine atom?
Li
A.
B.
C.
F
A lithium atom
A fluorine atom
There is no way to tell without memorizing the periodic
table.
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Periodic Table
CHECK YOUR NEIGHBOR
Which is larger: an arsenic atom or a sulfur atom?
S
As
A.
B.
C.
An arsenic atom
A sulfur atom
There is no way to tell without memorizing the periodic table.
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
Physical and Conceptual Models
• A physical model replicates an object at a
convenient scale.
– Model trains, cars, planes, buildings
• A conceptual model describes a system.
– An atom is best described by a conceptual
model.
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Identifying Atoms Using the Spectroscope
• A spectroscope
– is an instrument that separates and spreads
light into its component frequencies.
– allows the analysis of light emitted by
elements when they are made to glow. It
identifies each element by its characteristic
pattern.
– Each element emits a distinctive glow when
energized and displays a distinctive
spectrum.
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Identifying Atoms Using the Spectroscope
• An atomic spectrum is an element's fingerprint—
a pattern of discrete (distinct) frequencies of
light.
• Discoveries of the atomic spectrum of hydrogen:
– A researcher in the 1800s noted that
hydrogen has a more orderly atomic
spectrum than others.
– Johann Balmer expressed line positions by
a mathematical formula.
– Johannes Rydberg noted that the sum of
the frequencies of two lines often equals the
frequency of a third line.
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Identifying Atoms Using the Spectroscope
• Spectral lines of
various elements
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Identifying Atoms Using the Spectroscope
• Atomic excitation
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Identifying Atoms Using the Spectroscope
• There are three transitions in an atom. The sum
of the energies (and frequencies) for jumps A
and B equals the energy (and frequency) for
jump C.
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Identifying Atoms Using the Spectroscope
CHECK YOUR NEIGHBOR
Each spectral line in an atomic spectrum represents
A. a specific frequency of light emitted by an
element.
B. one of the many colors of an element.
C. a pattern characteristic of the element.
D. all of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
Identifying Atoms Using the Spectroscope
CHECK YOUR NEIGHBOR
The hydrogen spectrum consists of many spectral lines.
How can this simple element have so many lines?
A. One electron can be boosted to many different
energy levels.
B. The electron can move at a variety of speeds.
C. The electron can vibrate at a variety of frequencies.
D. Many standing electron waves can fit in the shell of the
hydrogen atom.
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
Identifying Atoms Using the Spectroscope
CHECK YOUR NEIGHBOR
When an atom is excited, its
A. electrons are boosted to higher energy
levels.
B. atoms are charged with light energy.
C. atoms are made to shake, rattle, and roll.
D. none of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
Identifying Atoms Using the Spectroscope
CHECK YOUR NEIGHBOR
The frequencies of light emitted by an atom often add up to
A. a higher frequency of light emitted by the same
atom.
B. a lower frequency of light emitted by the same
atom.
C. both of the above
D. none of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
© 2013 Pearson Education, Inc.
The Quantum Hypothesis
• Max Planck, a German physicist, hypothesized
that warm bodies emit radiant energy in discrete
bundles called quanta. The energy in each
energy bundle is proportional to the frequency of
the radiation.
• Einstein stated that light itself is quantized. A
beam of light is not a continuous stream of
energy but consists of countless small discrete
quanta of energy, with each quantum called a
photon.
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The Quantum Hypothesis
• Is light a wave or a stream of particles?
• Light can be described by both models: It
exhibits properties of both a wave and a particle,
depending on the experiment.
• The amount of energy in a photon is directly
proportional to the frequency of light:
E
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The Quantum Hypothesis
CHECK YOUR NEIGHBOR
In the relationship E  , the symbol  stands for
the frequency of emitted light, and E stands for the
A.
B.
C.
D.
potential energy of the electron emitting the light.
energy of the photon.
kinetic energy of the photon.
all of the above
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Quantum Hypothesis
CHECK YOUR NEIGHBOR
Which of these has the most energy per photon?
A.
B.
C.
D.
Red light
Green light
Blue light
All have the same.
Explain your answer to your neighbor.
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The Quantum Hypothesis
CHECK YOUR NEIGHBOR
Which of these photons has the least energy?
A.
B.
C.
D.
Infrared
Visible
Ultraviolet
All have the same.
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Quantum Hypothesis
• Using the quantum hypothesis
– Danish physicist Niels Bohr explained the
formation of atomic spectra as follows:
• The potential energy of an electron depends on its
distance from the nucleus.
• When an atom absorbs a photon of light, it absorbs
energy. Then a low-potential-energy electron is
boosted to become a high-potential-energy
electron.
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The Quantum Hypothesis
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The Quantum Hypothesis
• Using the quantum hypothesis (continued):
– When an electron in any energy level drops closer to
the nucleus, it emits a photon of light.
– Bohr reasoned that there must be a number of distinct
energy levels within the atom. Each energy level has
a principal quantum number n, where n is always an
integer. The lowest level is n = 1 and is closest to the
nucleus.
– Electrons release energy in discrete amounts that
form discrete lines in the atom's spectrum.
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The Quantum Hypothesis
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The Quantum Hypothesis
CHECK YOUR ANSWER
Which of the following is a quantum number?
A.
B.
C.
D.
0.02
0.2
2
2.5
Explain your answer to your neighbor.
© 2013 Pearson Education, Inc.
The Quantum Hypothesis
• Bohr's model explains why atoms don't collapse
– Electrons can lose only specific amounts of
energy equivalent to transitions between
levels.
– An atom reaches the lowest energy level
called the ground state, where the electron
can't lose more energy and can't move closer
to the nucleus.
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The Quantum Hypothesis
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The Bohr Model of the Hydrogen Atom
• Niels Bohr (1885 – 1962) – first saw the
connection of how each element is capable of
emitting its own characteristic wavelength of
radiation.
• Bohr applied the planetary model of the atom to
explain the line spectrum.
The Bohr Model of the Hydrogen Atom
• Bohr postulate that the electron is allowed to
have only certain orbits corresponding to
different amounts of energy.
• Each energy level/orbit was labeled by a
quantum number, n
The Bohr Model of the Hydrogen Atom
• Ground state – lowest energy level, n = 1
– Orbit closes to the nucleus.
• Excited state – level above n = 1 that the
electron jumps to when it absorbs the
appropriate amount of energy.
– Larger orbits, farther from the nucleus.
The Bohr Model of the Hydrogen Atom
• When radiation is absorbed an electron jumps
from the ground state to an excited state.
• Radiation is emitted when the electron falls back
from the higher energy level to the lower energy
level.
• The energy of the emitted radiation equals the
difference between the two energy levels.
The Quantum Hypothesis
• Planetary model of the atom:
– Photons are emitted by atoms as electrons
move from higher-energy outer levels to
lower-energy inner levels. The energy of an
emitted photon is equal to the difference in
energy between the two levels. Because an
electron is restricted to discrete levels, only
lights of distinct frequencies are emitted.
– The move between levels is instantaneous!
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Electron Waves
• An electron's wave nature explains why
electrons in an atom are restricted to particular
energy levels. The permitted energy levels are a
natural consequence of standing electron waves
closing in on themselves in a synchronized
manner.
• The orbit for n = 1 consists of a single
wavelength, n = 2 of two wavelengths, and so
on.
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Orbitals and Energy
• Bohr said that energies of electrons
in atoms are quantized.
– Electrons = energy levels
• Principal energy levels – designated
by the quantum number, n.
Orbitals and Energy
• The energy of the electron
increases as n increases.
– Each principal energy level
is divided into one or more
sublevels.
– The number of sublevels in
each principal energy level
equals the quantum
number n for that energy
level.
– n=1, 1 sublevel. n=2, 2
sublevels…
Orbitals and Energy
• The sublevels are labeled with a number that is the
value of n, and a letter (s,p,d,f,) which identifies the
sublevel.
• *Each principal energy level (n=1, n=2..) consists of
one or more sublevels
– *Each sublevel consists of one or more orbitals.
• * Orbital – a region in which an electron with a
particular energy is likely to be found.
• We can use the periodic table to predict which sublevel is
being filled by a particular element.
Electron Waves
• Electrons or any particle, can show itself as a
wave or as a particle, depending on how you
examine it.
• This is the wave-particle duality
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Matter Waves
• We can’t observe both the particle and wave
properties of an electron by using the same
experiment.
– Depending on what property the experiment is
designed to bring out then that’s what the
electron is going to show.
Matter Waves
• Scientists aren’t always concerned with duality.
– When scientists study the motion of the space
shuttle, they are concerned more with the
shuttle as a particle than a wave.
– In small particles like the electron however,
scientists have to be concerned with both
wave and particle properties.
Heisenberg’s Uncertainty Principle
• Werner Heisenberg (1901-1976) – proposed that the
position and the momentum of a moving object cannot
simultaneously be measured and known exactly.
• Significant when dealing with electrons.
– Only way to locate an electron is to strike it with a
photon – because of its small mass measurement
changes its position.
– No way to observe or to measure the orbit of an
electron in an atom.
– Arrangement of electrons in atoms are discussed in
terms of the probability of finding an electron in
certain locations within the atom.
Heisenberg’s Uncertainty Principle
• For us to see an object it must be hit by a photon
of radiant energy
• A photon hitting you or any object you can see
has a negligible effect on it.
• However, a collision between a photon and an
electron results in a large change in the energy
of the electron
Heisenberg’s Uncertainty Principle
• When you “see” an electron using some sort of
radiant energy as “illumination” you have found
the exact position of the electron.
• However, the collision between the photon and
electron caused its velocity to change.
• Therefore, we know the electrons position but
not its velocity
Heisenberg’s Uncertainty Principle
• On the other hand, if we measure an electron’s
velocity, we will change the electron’s position
Heisenberg’s Uncertainty Principle
• Heisenberg stated that there is always some
uncertainty about the position and momentum of
an electron
• This became known as the Heisenberg
Uncertainty Principle.
Heisenberg’s Uncertainty Principle
• Computers can now calculate the probabilities of
the location of the electron for thousands of
points (places) in space.
Heisenberg’s Uncertainty Principle
• There will be many points of equal probability.
• If all points of highest probability are connected,
a 3D shape is formed.
• This shape is just a mental model and does not
exist.
Heisenberg’s Uncertainty Principle
• Electrons are constantly in motion at high
speeds. They look like a cloud.
– Ex = like fan blades when a fan is turned on.
• You can try to fit something in between the rotating
fan blades but you will find that you can’t.
• The fan blades are moving so fast that they are
taking up all the space.
Heisenberg’s Uncertainty Principle
• The volume occupied by an electron is
somewhat vague.
• This is why we refer to the electrons as an
electron cloud rather than to a specific position.
Electron Waves
• For a fixed circumference, only an integral
number of standing waves can occur, and
likewise for the paths of electrons about the
nucleus.
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The Shell Model
• Cutaway view of shells in the shell model of the
atom
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Octet Rule
• Sodium chloride is much more stable than each
element on its own.
• Atoms are stable when their valence shell is full
– 8 electrons.
• Electron configurations:
– Chlorine: [Ne]3s23p5
– Sodium: [Ne]3s1
• Chlorine has 7 valence electrons = needs to
gain one
• Sodium has 1 valence electrons = needs to lose
one
Octet Rule
• Atoms tend to gain, lose, or share electrons
in order to acquire a full set of valence
electrons.
• Sodium readily loses an electron
• Chlorine readily gains an electron
The Shell Model
• Shell model showing the first three periods of the
periodic table
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