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Resources
Chapter Presentation
Bellringer
Transparencies
Sample Problems
Visual Concepts
Standardized Test Prep
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Chapter 3
Atoms and Moles
Table of Contents
Section 1 Substances Are Made of Atoms
Section 2 Structure of Atoms
Section 3 Electron Configuration
Section 4 Counting Atoms
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Chapter 3
Section 1 Substances Are Made
of Atoms
Bellringer
• Make a list of inferences about any properties of
objects in the box.
• How could you learn more about the objects in the
box without opening the box?
• Scientist face these same questions as they try to
learn more about atoms.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Objectives
• State the three laws that support the existence of
atoms.
• List the five principles of John Dalton’s atomic
theory.
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Section 1 Substances Are Made
of Atoms
Chapter 3
Atomic Theory
• The idea of an atomic theory is more than 2000 years
old.
Aristotle modified an earlier theory
that matter was made of four
“elements”: earth, fire, water, air.
Aristotle
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Democritus
Democritus (430-370 BC), a Greek philosopher, believed that all
Matter consists of very small, indivisible particles. He called
these particles atomos (Greek for “uncuttable” or “indivisible”).
Democritus’ idea, supported by experimental evidence nearly
2500 years later, laid the foundations for our modern
understanding of the nature of elements and compounds.
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• Until recently, scientists had never seen evidence of
atoms.
• The law of definite proportions, the law of
conservation of mass and the law of multiple
proportions support the current atomic theory.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Atomic Theory, continued
• The figure on the right is a more accurate
representation of an atom than the figure on the left.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Atomic Theory, continued
The Law of Definite Proportions
Joseph Proust (1754-1826), in 1799. Proust’s
Law of Definite Proportions states:
• The law of definite proportions states that a
chemical compound always contains the same
elements in exactly the same proportions by weight or
mass.
• The law of definite proportions also states that every
molecule of a substance is made of the same number
and types of atoms.
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Chapter 3
Visual Concepts
Law of Definite Proportions PLAY
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Chapter 3
Section 1 Substances Are Made
of Atoms
Atomic Theory, continued
The Law of Conservation of Mass
Law of Conservation of Mass: Antoine Lavoisier
consider “the father of modern chemistry”
• The law of conservation of mass states that mass
cannot be created or destroyed in ordinary chemical
and physical changes.
• The mass of the reactants is equal to the mass of the
products.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Law of Conservation of Mass
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Chapter 3
Section 1 Substances Are Made
of Atoms
Law of Conservation of Mass, continued
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Law of Conservation of Mass PLAY
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Chapter 3
Section 1 Substances Are Made
of Atoms
Atomic Theory, continued
The Law of Multiple Proportions
• Proposed by John Dalton
• The law of multiple proportions states that when
two elements combine to form two or more
compounds, the mass of one element that combines
with a given mass of the other is in the ratio of small
whole numbers.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Law of Multiple Proportions
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Chapter 3
Visual Concepts
Law of Multiple Proportions PLAY
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Chapter 3
Section 1 Substances Are Made
of Atoms
Dalton’s Atomic Theory
• In 1808, John Dalton developed an atomic theory.
• Dalton believed that a few kinds of atoms made up all
matter.
• According to Dalton, elements are composed of only
one kind of atom and compounds are made from two
or more kinds of atoms.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Dalton’s Atomic Theory , continued
Dalton’s Theory Contains Five Principles
1. All matter is composed of extremely small particles
called atoms, which cannot be subdivided, created,
or destroyed.
2. Atoms of a given element are identical in their
physical and chemical properties.
3. Atoms of different elements differ in their physical
and chemical properties.
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Chapter 3
Section 1 Substances Are Made
of Atoms
Dalton’s Atomic Theory , continued
Dalton’s Theory Contains Five Principles, continued
4. Atoms of different elements combine in simple,
whole-number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged but never created,
destroyed, or changed.
•
Data gathered since Dalton’s time shows that the
first two principles are not true in all cases.
•
Hwk page 78 Que. 1,2,3,4,5,6 Quiz to Follow
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Chapter 3
Section 2 Structure of Atoms
Bellringer
• Look at the following terms:
electron, nucleus, proton, neutron, atomic number,
mass number, isotope
• Chapter 3 Sec 2 Video Intro
• Make a list of the terms that are unfamiliar to you?
• After completing this section, look over your list to
check that you are familiar with and understand all of
the terms.
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Chapter 3
Section 2 Structure of Atoms
Objectives
• Describe the evidence for the existence of electrons,
protons, neutrons, and describe the properties of
these subatomic particles.
• Discuss atoms of different elements in terms of their
numbers of electrons, protons, neutrons, and define
the terms atomic number and atomic mass.
• Define isotope, and determine the number of
particles in the nucleus of an isotope.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles
• Experiments by several scientists in the mid-1800s
led to the first change to Dalton’s atomic theory.
Scientists discovered that atoms can be broken into
pieces after all.
• The smaller parts that make up atoms are called
subatomic particles.
• The three subatomic particles that are most important
for chemistry are the electron, the proton, and the
neutron.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Electrons Were Discovered Using Cathode Rays
• To study current, J. J. Thomson pumped most of the
air out of a glass tube. He applied a voltage to two
metal plates, called electrodes, which were placed at
either end of the tube.
• One electrode, called the anode, was attached to the
positive terminal of the voltage source, so it had a
positive charge.
• The other electrode, called a cathode, had a negative
charge because it was attached to the negative
terminal of the voltage source.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Electrons Were Discovered Using Cathode Rays,
continued
• Thomson observed a glowing beam that came out of
the cathode and struck the anode and the nearby
glass walls of the tube.
• He called these rays cathode rays.
• The glass tube Thomson used is known as a
cathode-ray tube (CRT).
• CRTs are used in television sets, computer monitors,
and radar displays.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
An Electron Has a Negative Charge
• Because the cathode ray came from the negatively
charged cathode, Thomson reasoned that the ray was
negatively charged.
• Thomson confirmed this prediction by seeing how electric and
magnetic fields affected the cathode ray.
• Thomson also observed that when a small paddle
wheel was placed in the path of the rays, the wheel
would turn.
• This suggested that the cathode rays consisted of tiny
particles that were hitting the paddles of the wheel.
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Chapter 3
Visual Concepts
Thompson’s Cathode Ray Tube Experiment
PLAY
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United Streaming (Discovery)
Search Cathode rays to xrays Program overview
VIDEO
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
An Electron Has a Negative Charge, continued
• Thomson’s experiments showed that a cathode ray
consists of particles that have mass and a negative
charge.
• These particles are called electrons.
• An electron is a subatomic particle that has a negative
electric charge.
• Electrons are negatively charged, but atoms have no
charge.
• Atoms contain some positive charges that
balance the negative charges of the electrons.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
An Electron Has a Negative Charge, continued
• Properties of Electrons
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
• Thomson proposed that the electrons of an atom
were embedded in a positively charged ball of matter.
His model of an atom was named the plum-pudding
model.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Rutherford Discovers the Nucleus, continued
• Ernest Rutherford performed the gold foil experiment,
which disproved the plum-pudding model of the atom.
•
Rutherford’s Expt.
• A beam of small, positively charged particles, called alpha
particles, was directed at a thin gold foil.
• Rutherford’s team measured the angles at which the
particles were deflected from their former straight-line paths
as they came out of the foil.
• Rutherford found that most of the alpha particles shot
at the foil passed straight through the foil. But very
few were deflected, in some cases backward.
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Chapter 3
Gold Foil Experiment
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Rutherford Discovers the Nucleus, continued
• Rutherford reasoned that only a very concentrated
positive charge in a tiny space within the gold atom
could possibly repel the fast-moving, alpha particles
enough to reverse the alpha particles’ direction.
• Rutherford also hypothesized that the mass of this
positive-charge containing region, called the nucleus,
must be larger than the mass of the alpha particle.
• Rutherford argued that the reason most of the alpha
particles were undeflected, was that most parts of the
atoms in the gold foil were empty space.
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Chapter 3
Gold Foil Experiment on the Atomic Level
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Rutherford Discovers the Nucleus, continued
• The nucleus is the dense, central portion of the atom.
• The nucleus is made up of protons and neutrons.
• The nucleus has all of the positive charge, nearly all
of the mass, but only a very small fraction of the
volume of the atom.
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Chapter 3
Visual Concepts
Rutherford’s Gold Foil Experiment
PLAY
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Proton and Neutrons Compose the Nucleus
• Protons are the subatomic particles that have a
positive charge and that is found in the nucleus of an
atom.
• The number of protons of the nucleus is the atomic number,
which determines the identity of an element.
• Because protons and electrons have equal but opposite
charges, a neutral atom must contain equal numbers of
protons and electrons.
• Neutrons are the subatomic particles that have no
charge and that is found in the nucleus of an atom.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Proton and Neutrons Compose the Nucleus, continued
• Properties of a Proton and a Neutron
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Symbol
Charge
Mass
Relative
Charge
Relative
Mass
Electron
e-
-1.602 x 10-19
C
9.109 x 10-31
kg
1-
1
Proton
p+
+1.602 x10-19
C
1.673 x 10-27
kg
1+
1837
Neutron
n0
0C
1.675 x 10-27
kg
0
1839
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Chapter 3
Visual Concepts
Parts of an Atom PLAY
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Protons and Neutrons Can Form a Stable Nucleus
• Coulomb’s law states that the closer two charges are,
the greater the force between them.
• The repulsive force between two protons is large
when two protons are close together.
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Chapter 3
Section 2 Structure of Atoms
Subatomic Particles, continued
Protons and Neutrons Can Form a Stable Nucleus
• Protons form stable nuclei despite the repulsive force
between them.
• A strong attractive force between these protons
overcomes the repulsive force at small distances.
• Because neutrons also add attractive forces, some
neutrons can help stabilize a nucleus. (Act as glue)
• All atoms that have more than one proton also
have neutrons.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number
Atomic Number Is the Number of Protons of the
Nucleus
• The number of protons that an atom has is known as
the atom’s atomic number.
• The atomic number is the same for all atoms of an
element.
• Because each element has a unique number of
protons in its atoms, no two elements have the
same atomic number.
• Example: the atomic number of hydrogen is 1 because
the nucleus of each hydrogen atom has one proton.
The atomic number of oxygen is 8.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Atomic Number Is the Number of Protons of the
Nucleus, continued
• Atomic numbers are always whole numbers.
• The atomic number also reveals the number of
electrons in an atom of an element.
• For atoms to be neutral, the number of negatively charged
electrons must equal the number of positively charged
protons.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Atomic Number Is the Number of Protons of the
Nucleus, continued
• The atomic number for oxygen tells you that the
oxygen atom has 8 protons and 8 electrons.
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Chapter 3
Visual Concepts
Atomic Number
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Mass Number Is the Number of Particles of the
Nucleus, continued
• The mass number is the sum of the number of
protons and neutrons in the nucleus of an atom.
• You can calculate the number of neutrons in an atom
by subtracting the atomic number (the number of
protons) from the mass number (the number of
protons and neutrons).
mass number – atomic number = number of neutrons
• Unlike the atomic number, the mass number can vary
among atoms of a single element.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Mass Number Is the Number of Particles of the
Nucleus, continued
• Example: a particular atom of neon has a mass
number of 20.
• Because the atomic number for an atom of neon is
10, neon has 10 protons.
number of protons and neutrons (mass number) = 20
 number of protons (atomic number) = 10
number of neutrons = 10
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Mass Number Is the Number of Particles of the
Nucleus, continued
• The neon atom has 10 protons, 10 electrons, and
10 neutrons. The mass number is 20.
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Chapter 3
Visual Concepts
Mass Number
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Chapter 3
Section 2 Structure of Atoms
Determining the Number of Particle In An
Atom
Sample Problem A
How many protons, electrons, and neutrons are
present in an atom of copper whose atomic number is
29 and whose mass number is 64?
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Chapter 3
Section 2 Structure of Atoms
Sample Problem A Solution
• The atomic number indicates the number of protons in
the nucleus of a copper atom.
atomic number (29) = number of protons = 29
• A copper atom must be electrically neutral, so the
number of electrons equals the number of protons.
number of protons = number of electrons = 29
• The mass number indicates the total number of protons
and neutrons
mass number (64) - atomic number (29) =
number of neutrons = 35
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Atomic Structures Can Be Represented by Symbols
• Each element has a name, and the same name is
given to all atoms of an element.
• Example: sulfur is composed of sulfur atoms.
• Each element has a symbol, and the same symbol is
used to represent one of the element’s atoms.
• Atomic number and mass number are sometimes
written with an element’s symbol.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Isotopes of an Element Have the Same Atomic
Number
• All atoms of an element have the same atomic
number and the same number of protons. Atoms do
not necessarily have the same number of neutrons.
• Atoms of the same element that have different
numbers of neutrons are called isotopes.
• One standard method of identifying isotopes is to
write the mass number with a hyphen after the name
of an element.
helium-3 or helium-4
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Atomic Structures Can Be Represented by Symbols
A
X
Z
A is the mass number
Z is the atomic number
X is the element symbol
or
Ex)
Element Name-Mass Number
239
U
92
Uranium-239
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Section 2 Structure of Atoms
Chapter 3
Atomic Number and Mass Number, continued
Atomic Structures Can Be Represented by Symbols
• The atomic number always appears on the lower left
side of the symbol.
H
1
He 3Li
Be
2
4
B
5
• Mass numbers are written on the upper left side of the
symbol.
1
H
2
H
3
He
4
He
6
Li
7
Li
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9
Be
10
B
11
B
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Section 2 Structure of Atoms
Chapter 3
Atomic Number and Mass Number, continued
Atomic Structures Can Be Represented by Symbols
• Both numbers may be written with the symbol.
1
1
H
4
2
He
7
3
Li
9
4
Be
11
5
B
• An element may be represented by more than one
notation.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Isotopes of an Element Have the Same Atomic
Number, continued
• The second method of identifying isotopes shows the
composition of a nucleus as the isotope’s nuclear
symbol.
3
2
He or 24He
• All isotopes of an element have the same atomic
number. However, their atomic masses are not the
same because the number of neutrons of the atomic
nucleus of each isotope varies.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Isotopes of an Element Have the Same Atomic
Number, continued
• The two stable helium isotopes are helium-3 and
helium-4.
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Chapter 3
Section 2 Structure of Atoms
Atomic Number and Mass Number, continued
Isotopes of an Element Have the Same Atomic
Number, continued
• The Stable Isotopes of Lead
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Chapter 3
Visual Concepts
Isotopes and Nuclides
PLAY
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Chapter 3
Section 2 Structure of Atoms
Determining the Number of Particle In An
Isotope
Sample Problem B
Calculate the numbers of protons, electrons, and
neutrons in oxygen-17 and in oxygen-18.
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Chapter 3
Section 2 Structure of Atoms
Sample Problem B Solution
• atomic number = number of protons =
number of electrons = 8
• mass number - atomic number = number of neutrons
For oxygen-17, 17 - 8 = 9 neutrons
For oxygen-18, 18 - 8 = 10 neutrons
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Chapter 3
Ex)
Section 2 Structure of Atoms
32
S
16
Ex)
108
Ag2+
47
Ex)
80
Br1-
35
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Chapter 3
Section 2 Structure of Atoms
Ex) Fluorine-19
Ex) Calcium-44
Ex) Write the symbol that has 34 protons, 34
electrons, and 46 neutrons.
Ex) Write the symbol that has 17 protons, 18
electrons, and 17 neutrons.
Ex) Write the symbol that has 79 protons, 77
electrons, and 118 neutrons.
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Homework  Worksheet
Quiz to Follow
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Chapter 3
Section 3 Electron Configuration
Objectives
• Compare the Rutherford, Bohr, and quantum models
of an atom.
• Explain how the wavelengths of light emitted by an
atom provide information about electron energy
levels.
• List the four quantum numbers, and describe their
significance.
• Write the electron configuration of an atom by using
the Pauli exclusion principle and the the aufbau
principle.
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Evolution of the Atom or Models of the Atom
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1) John Dalton Model  Billard Ball Model
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2) J.J. Thomson Model or Plum Pudding Model
1904
the electrons are like raisins dispersed in a pudding
(positive charge cloud)
seeds in a watermelon
blue-berry muffin
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Chapter 3
Section 3 Electron Configuration
Atomic Models
Rutherford’s Model Proposed Electron Orbits
• The experiments of Rutherford’s team led to the
replacement of the plum pudding model of the atom
with a nuclear model of the atom.
• Rutherford suggested that electrons, like planets orbiting the
sun, revolve around the nucleus in circular or elliptical orbits.
• Rutherford’s model could not explain why electrons did not
crash into the nucleus.
• The Rutherford model of the atom was replaced only
two years later by a model developed by Niels Bohr.
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Evolution of the Atom
3) Rutherford Model or Nuclear Atom Model
1911
positive dense center called the nucleus
rest of atom empty space (electrons reside)
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Evolution of the Atom
4) Bohr Model or the Shell Model
1913
pictured the atom as a small positive nucleus
with electrons orbiting around it in curved
circular pathways
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Chapter 3
Section 3 Electron Configuration
Atomic Models, continued
Bohr’s Model Confines Electrons to Energy Levels
• According to Bohr’s model, electrons can be only certain
distances from the nucleus. Each distance corresponds
to a certain quantity of energy that an electron can have.
• An electron that is as close to the nucleus as it can be is in
its lowest energy level.
• The farther an electron is from the nucleus, the higher the
energy level that the electron occupies.
• The difference in energy between two energy levels
is known as a quantum of energy.
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Niels Bohr
(Born in Denmark 1885-1962)
• Student of Rutherford
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Niels Bohr’s Model (1913)
• Electrons orbit
the nucleus in
circular paths of
fixed energy
(energy levels).
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Max Plank
E=hn
E=energy
n=frequency
h=Plank’s constant 6.7x10-34Js
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Energy of Emitted Photon
Energy of the emitted photon =
Difference in energy between two states
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• Energy emitted by the electron as it leaps from the
higher to the lower energy level is proportional to the
frequency of the light wave.
• Frequency define the color of visible light.
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Niels Bohr’s Atom Cont’d
• Electrons can jump from energy level
to energy level.
• Electrons absorb or emit light energy
when they jump from one energy level
to another.
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Quantum
• A quantum of energy is the amount
of energy required to move an
electron from one energy level to
another.
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The energy levels are like the
rungs of a ladder but are not
equally spaced.
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Photons
• Photons are bundles of light energy
that is emitted by electrons as they go
from higher energy levels to lower
levels.
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Excited State and Ground State
• Ground state: the lowest possible energy level an
electron can be at.
• Excited state: an energy level higher than the ground
state.
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Emission Spectrum
• Light emitted produces a unique
emission spectrum.
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Hydrogen Emission Spectrum
Violet
Blue
Red
Balmer
Series
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Bohr Model for Hydrogen
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• The Bohr model explained the
emission spectrum of the hydrogen
atom but did not always explain
those of other elements.
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Evolution of the Atom
5) Wave Mechanical Model or the Electron Cloud Model
1920’s
Louis Victor de Broglie and Erwin Schrodinger
suggested that because light seems to behave
both as a wave and as a stream of particles, then
the electron should exhibit both of these
characteristics
Orbitals (electron states) are nothing like orbits
Similar to a cloud or firefly analogy (highest
probability)
Heisenberg Uncertainty Principle  we will never
know simultaneously the exact momentum and
position of an electron
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Chapter 3
Section 3 Electron Configuration
Atomic Models, continued
Electrons Act Like Both Particles and Waves
• Thomson’s experiments demonstrated that electrons
act like particles that have mass.
• In 1924, Louis de Broglie pointed out that the
behavior of electrons according to Bohr’s model was
similar to the behavior of waves.
• De Broglie suggested that electrons could be
considered waves confined to the space around a
nucleus.
• As waves, electrons could have only certain frequencies
which correspond to the specific energy levels.
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Chapter 3
Visual Concepts
De Broglie and the Wave-Particle Nature of
Electrons PLAY
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Chapter 3
Section 3 Electron Configuration
Atomic Models, continued
Electrons Act Like Both Particles and Waves,
continued
• The present-day model of the atom takes into account
both the particle and wave properties of electrons.
• In this model, electrons are located in orbitals,
regions around a nucleus that correspond to specific
energy levels.
• Orbitals are regions where electrons are likely to be found.
• Orbitals are sometimes called electron clouds because they
do not have sharp boundaries. Because electrons can be in
other places, the orbital has a fuzzy boundary like a cloud.
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Chapter 3
Section 3 Electron Configuration
Atomic Models, continued
Electrons Act Like Both Particles and Waves,
continued
• According to the current model of an atom, electrons
are found in orbitals.
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Quantum Mechanical Model
•
•
•
•
1920’s
Werner Heisenberg (Uncertainty Principle)
Louis de Broglie (electron has wave properties)
Erwin Schrodinger (mathematical equations using
probability, quantum numbers)
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Werner Heisenberg: Uncertainty
Principle
• We can not know both the
position and momentum of a
particle at a given time.
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Louis de Broglie, (France, 18921987)
Wave Properties of Matter (1923)
•Since light waves have a particle
behavior (as shown by Einstein in the
Photoelectric Effect), then particles
could have a wave behavior.
•de Broglie wavelength
l= h
mv
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Electron Motion Around Atom
Shown as a de Broglie Wave
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Davisson and Germer (USA,
1927) confirmed de Broglie’s
hypothesis for electrons.
Electrons produced a diffraction
pattern similar to x-rays.
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Erwin Schrodinger, 1925
Quantum (wave) Mechanical
Model of the Atom
• Four quantum numbers are
required to describe the state
of the hydrogen atom.
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Atomic Orbital:
A region in space in which there is
high probability of finding an
electron.
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Chapter 3
Visual Concepts
Comparing Models of Atoms PLAY
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Chapter 3
Section 3 Electron Configuration
Electrons and Light
• By 1900, scientists knew that light could be thought of
as moving waves that have given frequencies,
speeds, and wavelengths. Light waves are
electromagnetic waves and light is a form of
electromagnetic radiation (A form of energy called
radiant energy that travels at the speed of light with
wave-like behavior). All waves, whether they are
water waves or electromagnetic waves, can be
described in terms of four characteristics-amplitude,
wavelength, frequency, and speed.
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The Nature of Waves
• What is a wave?
• A wave is a repeating disturbance or
movement that transfers energy
through matter or space
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Waves transfer energy not matter. The water waves
below are carrying energy but are not moving.
Waves can only exist as they have energy to carry.
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What are the parts of a wave?
The crest is the highest
point on a wave. The trough
is the lowest point on a
wave.
The rest position of the
wave is called the node or
nodal line.
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Electrons and Light
• Characteristics of a wave.
1) The amplitude of a wave is the height of the wave
measured from the origin to its crest, or peak.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light
• Characteristics of a wave.
2) The wavelength, λ (lambda) is the distance between
successive crests of the wave. It is the distance that
the wave travels as it completes one full cycle of
upward and downward motion.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light
• Characteristics of a wave.
3) The frequency, v (nu) of a wave tells how fast the
wave oscillates up and down. The frequency of light
is measured by the number of times a light wave
completes a cycle of upward and downward motion
in one second. Thus, the unit for frequency is cycles
per second. Because it is understood that cycles are
involved, frequency is commonly expressed simply
as "per second," which is written as s-1 or 1/s. A
cycle per second is also called a hertz (Hz): 1 Hz =
1 s-1.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light
4) Light, regardless of its wavelength, moves through
space at a constant speed of 3.00 x 108 meters per
second (m/s), which is the speed of light, c
Because light moves at a constant speed, there is a relationship
between its wavelength and its frequency. The shorter the
distance between the crests of the wave, the faster the wave
oscillates up and down. That is, the shorter the wavelength,
the greater the frequency. This relationship can be expressed
in a simple equation. Using the symbol λ (the Greek letter
lambda) for wavelength, v (the Greek letter nu) for frequency,
and c for the speed of light, the relationship between
wavelength and frequency is C=λv
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C=λν
λ is inversly proportional to the ν
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Chapter 3
Visual Concepts
Characteristics of a Wave
PLAY
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Chapter 3
Section 3 Electron Configuration
Electromagnetic Spectrum
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Chapter 3
Visual Concepts
Electromagnetic Spectrum
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
• The electromagnetic spectrum is all of the
frequencies or wavelengths of electromagnetic
radiation.
• The wavelength of light can vary from 105 m to less
than 10–13 m.
• In 1905, Albert Einstein proposed that light also has
some properties of particles.
• His theory would explain a phenomenon known as the
photoelectric effect.
• This effect happens when light strikes a metal and electrons
are released.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
• Einstein proposed that light has the properties of both
waves and particles.
• Light can be described as a stream of particles,
the energy of which is determined by the light’s
frequency.
Light is an electromagnetic wave.
• Red light has a low frequency and a long wavelength.
• Violet light has a high frequency and a short wavelength.
• The frequency and wavelength of a wave are
inversely related.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
Light is an electromagnetic Wave, continued
• The frequency and wavelength of a wave are
inversely related.
• As frequency increases, wavelength decreases.
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Chapter 3
Section 3 Electron Configuration
Wavelength and Frequency
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
Light Emission
• When a high-voltage current is passed through a tube
of hydrogen gas at low pressure, lavender-colored
light is seen. When this light passes through a prism,
you can see that the light is made of only a few
colors. This spectrum of a few colors on a black
background that is related to electron transitions from
higher energy levels to lower energy level s is called
a line-emission spectrum.
• Experiments with other gaseous elements show that
each element has a line-emission spectrum that is
made of a different pattern of colors.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
Light Emission, continued
• In 1913, Bohr showed that hydrogen’s line-emission
spectrum could be explained by assuming that the
hydrogen atom’s electron can be in any one of a
number of distinct energy levels.
• An electron can move from a low energy level to a high
energy level by absorbing energy.
• Electrons at a higher energy level are unstable and can
move to a lower energy level by releasing energy. This
energy is released as light that has a specific wavelength.
• Each different move from a particular energy level to a lower
energy level will release light of a different wavelength.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
Light Provides Information About Electrons
• An electron in a state of its lowest possible energy, is
in a ground state.
• The ground state is the lowest energy state of a quantized
system
• If an electron gains energy, it moves to an excited
state.
• An excited state is a state in which an atom has more
energy than it does at its ground state
• An electron in an excited state will release a specific
quantity of energy as it quickly “falls” back to its
ground state.
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Chapter 3
Section 3 Electron Configuration
Electrons and Light, continued
Light Provides Information About Electrons, continued
• An electron in a hydrogen
atom can move between
only certain energy states,
shown as n = 1 to n = 7.
• In dropping from a higher
energy state to a lower
energy state, an electron
emits a characteristic
wavelength of light.
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Chapter 3
Section 3 Electron Configuration
Hydrogen’s Line-Emission Spectrum
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Chapter 3
Visual Concepts
Absorption and Emission Spectra
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers
• The present-day model of the atom is also known as the
quantum model.
• According to this model, electrons within an energy level are
located in orbitals, regions of high probability for finding a
particular electron.
• The model does not explain how the electrons move about
the nucleus to create these regions. Heinsenberg
Uncertainity Principle
• To define the region in which electrons can be found,
scientists have assigned four quantum numbers that
specify the properties of the electrons.
• A quantum number is a number that specifies the
properties of electrons. Chapter menu
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers, continued
1) The principal quantum number, symbolized by n,
indicates the main energy level occupied by the
electron.
• Values of n are positive integers, such as 1, 2, 3,
4, ∞
• As n increases, the electron’s distance from the
nucleus and the electron’s energy increases.
• Max # e- is 2n2
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Chapter 3
Visual Concepts
Principal Quantum Number
PLAY
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers, continued
2) The main energy levels can be divided into
sublevels. These sublevels are represented by the
angular momentum quantum number, l.
• This quantum number indicates the shape or type
of orbital that corresponds to a particular sublevel.
• A letter code is used for this quantum number.
• l = 0 corresponds to an s orbital
• l = 1 to a p orbital
• l = 2 to a d orbital
• l = 3 to an f orbital
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers, continued
3) The magnetic quantum number, symbolized by m, is a
subset of the l quantum number.
• It also indicates the numbers and orientations of
orbitals around the nucleus.
• The value of m takes whole-number values,
depending on the value of l.
• The number of orbitals includes
•
•
•
•
one s orbital
three p orbitals
five d orbitals
seven f orbitals
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers, continued
4) The spin quantum number, ms indicates the
orientation of an electron’s magnetic field relative to
an outside magnetic field.
• The spin quantum number is represented by:
1
1
+ or 
or (  or  )
2
2
• A single orbital can hold a maximum of two
electrons, which must have opposite spins.
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Chapter 3
Visual Concepts
Quantum Numbers and Orbitals PLAY
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Chapter 3
Section 3 Electron Configuration
Quantum Numbers, continued
• Quantum Numbers of the First 30 Atomic Orbitals
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Chapter 3
Section 3 Electron Configuration
• In 1925 the German chemist Wolfgang Pauli
established a rule is known as the Pauli exclusion
principle.
• The Pauli exclusion principle states that two
particles of a certain class cannot be in the exact
same energy state.
• This means that that no two electrons in the same
atom can have the same four quantum numbers.
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Chapter 3
Visual Concepts
Pauli Exclusion Principle PLAY
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Chapter 3
Section 3 Electron Configuration
• Two electrons can have the same value of n by
being in the same main energy level.
• These two electrons can also have the same value
of l by being in orbitals that have the same shape.
• These two electrons may also have the same value
of m by being in the same orbital.
• But these two electrons cannot have the same spin
quantum number.
• If one electron has the value of 1/2, then the
other electron must have the value of –1/2.
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Chapter 3
Section 3 Electron Configuration
Shapes of s, p, and d Orbitals
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Homework: Worksheet on Quantum Numbers
QUIZ on Section 3 to this point next class
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Chapter 3
Section 3 Electron Configuration
Electron Configurations
• The arrangement of electrons in the ground state of
an atom is usually shown by writing an electron
configuration or an orbital notation.
• Like all systems in nature, electrons in atoms tend to
assume arrangements that have the lowest possible
energies.
• An electron configuration of an atom shows the
lowest-energy arrangement of the electrons for the
element.
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Electron configurations vs Orbital Notation
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Chapter 3
Visual Concepts
Orbital Notation PLAY
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Chapter 3
Section 3 Electron Configuration
Electron Configurations, continued
An Electron Occupies the Lowest Energy Level
Available
• The aufbau principle states that electrons fill orbitals
that have the lowest energy first.
• Aufbau is the German word for “building up.”
• The smaller the principal quantum number, the lower
the energy. Within an energy level, the smaller the l
quantum number, the lower the energy.
• So, the order in which the orbitals are filled
matches the order of energies.
1s < 2s < 2p < 3s < 3p
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Chapter 3
Section 3 Electron Configuration
Electron Configurations, continued
An Electron Occupies the Lowest Energy Level
Available, continued
• The energy of the 3d orbitals is slightly higher than
the energy of the 4s orbitals.
• As a result, the order in which the orbitals are
filled is as follows:
1s < 2s < 2p < 3s < 3p < 4s < 3d
• Additional irregularities occur at higher energy levels.
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Chapter 3
Section 3 Electron Configuration
Electron Configurations, continued
An Electron Occupies the Lowest Energy Level
Available, continued
• This diagrams shows how
the energy of the orbitals
can overlap.
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Chapter 3
Visual Concepts
Aufbau Principle PLAY
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See Handout for notes
Aufbau Principle
a) Memorize
b) Periodic table
c) Diagonal Rule
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Chapter 4
Section 1 How Are Elements
Organized?
Blocks of the Periodic Table
4f
5f
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Chapter 3
Visual Concepts
s Orbitals
PLAY
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Chapter 3
Visual Concepts
p Orbitals
PLAY
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Chapter 3
d Orbitals
Visual Concepts
PLAY
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Chapter 3
Visual Concepts
Electron Configuration PLAY
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Section 3 Electron Configuration
Chapter 3
Electron Configurations, continued
An Electron Configuration Is a Shorthand Notation,
continued
• Electron orbitals are filled according to Hund’s Rule.
• Hund’s rule states that orbitals of the same n and l
quantum numbers are each occupied by one electron
before any pairing occurs.
• Orbital diagram for sulfur
  

3p
   3s
2p

 2s
1s
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Chapter 3
Section 3 Electron Configuration
Electron Configurations, continued
An Electron Configuration Is a Shorthand Notation
• Based on the quantum model of the atom, the
arrangement of the electrons around the nucleus can
be shown by the nucleus’s electron configuration.
• Example: sulfur has sixteen electrons.
Its electron configuration is written as
1s22s22p63s23p4.
• Two electrons are in the 1s orbital, two electrons
are in the 2s orbital, six electrons are in the 2p
orbitals, two electrons are in the 3s orbital, and
four electrons are in the 3p orbitals.
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Examples of both Electron
Configuration and Orbital Notation
2 worksheets for homework.
QUIZ to follow.
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Chapter 3
Section 3 Electron Configuration
Electron Configurations, continued
An Electron Configuration As a Shorthand Notation,
continued
• Each element’s configuration builds on the previous
elements’ configurations.
• To save space, one can write this configuration by
using a configuration of a noble gas.
• neon, argon, krypton, and xenon
• The neon atom’s configuration is 1s22s22p6, so the
electron configuration of sulfur is
[Ne] 3s23p4
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Chapter 3
Visual Concepts
Noble Gas Notation PLAY
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Chapter 3
Section 3 Electron Configuration
Sample Problem C
Writing Electron Configurations
Write the electron configuration for an atom whose
atomic number is 20.
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Chapter 3
Section 3 Electron Configuration
Sample Problem C Solution
• atomic number = number of protons =
number of electrons = 20
• According to the aufbau principle, the order of orbital
filling is 1s,2s, 2p, 3s, 3p, 4s, 3d, 4p, and so on.
• The electron configuration for an atom of this element
is written as follows:
1s22s22p63s23p64s2
• This electron configuration can be abbreviated as
follows:
[Ar]4s2
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Chapter 3
• The electrons in the outer shell are called valence
electrons.
• Valence electrons are found in the outermost
shell of an atom and that determines the atom’s
chemical properties.
• Elements with the same number of valence
electrons tend to react in similar ways.
• Because s and p electrons fill sequentially, the
number of valence electrons in s- and p-block
elements are predictable.
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Valence Electrons
• valence electrons are the electrons in
the OUTERMOST energy level…
that’s why we did all those electron
configurations!
• B is 1s2 2s2 2p1; so the outer energy
level is 2, and there are 2+1 = 3
electrons in level 2. These are the
valence electrons!
• Br is [Ar] 4s2 3d10 4p5
How many valence electrons are
present?
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Valence Electrons
Number of valence electrons of a main (A)
group atom = Group number
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Electron Dot Structures
Symbols of atoms with dots to represent the
valence-shell electrons
1
2
13
14
15
16
17
H
Li
18
He:



Be
B 


C


Na Mg


Al


N


 Si 


O



P
S


Chapter menu


: F  :Ne :




:Cl  :Ar :


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Worksheet for Homework
Quiz to Follow
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Chapter 3
Section 4 Counting Atoms
Bellringer
• A penny has 2.97  1022 copper atoms.
• On a sheet of paper, write out this number in regular
notation with all the zeros.
• What does this tell you about the size of an atom?
Chapter 3
Section 4 Counting Atoms
Objectives
• Compare the mass quantities and units for atomic
mass with those for molar mass.
• Define mole, and explain why this unit is used to
count atoms.
• Calculate either mass with molar mass or number
with Avogadro’s number given an amount in moles.
List some common Counting Units
A counting unit
Similar to a dozen, except instead of 12, it’s
602 billion trillion
602,200,000,000,000,000,000,000
6.022 X 1023 (in scientific notation)
This number is named in honor of Amedeo
Avogadro(1776 – 1856)
How we determine quantity: Counting, Weighing or Volume Measurement
We will first consider how people normally keep
track of quantity in everyday life. There are
generally three ways to do this:
•
•
•
Counting  by number of units.
Example: oranges priced by number.
If a single unit is too small, we devise a lump
sum.
Example: eggs are priced by the dozen,
which is a lump sum of 12 units.
Weighing  by weight or mass.
Example: meat is priced by the pound.
Measuring  by volume.
This is easier to use with liquids or gases.
Example: gasoline is priced by the gallon.
Is it surprising to you that we
employ similar ways in chemistry to
keep track of atoms and molecules as
described above? But, indeed we do.
The next two slides compare two
parallel ways of counting by number
of units: As we count small items by
the “dozen” in everyday life, we count
atoms and molecules by the “mole” in
chemistry.
How many eggs are in 15 dozen eggs?
How many dozens of donuts are in 26
donuts?
How many atoms are in 57 dozen of Al
atoms?
How many dozen of atoms are in 1.8 * 105
Fe?
How many atoms are in 57 mole of Al
atoms?
How many moles of atoms are in 1.8 * 105
Fe?
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Previous slides showed how counting
eggs in dozens is similar to counting
atoms by the mole. We use similar math
equations for both mole <--> number
conversions and dozen <--> number
conversions.
Since everyday items contain
astronomical numbers of atoms and
molecules, it is easier to count them by a
huge lump sum:
1 mole = 6.022 x 1023
But why do we use Avogadro’s number,
6.022 x 1023 , for one mole? Did he invent
the number? No, neither he nor anyone
else did. The number is defined by how
the mass units, amu and gram, relate to
each other:
amu is the mass of one atom
units are atomic mass unit
1 amu = 1.66058 x 10-24 grams
Experimental data:
1 amu = 1.66058 x 10-24 grams leads to this
equality:
6.022 x 1023 amu = 1 gram
because: 1 / (1.66058 x 10-24) = 6.022 x 1023
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Just How Big is a Mole?
• If you had Avogadro's number of soft
drink cans to cover the surface of
the earth to a depth of over 200
miles.
• If you had Avogadro's number of
unpopped popcorn kernels, and
spread them across the United
States of America, the country would
be covered in popcorn to a depth of
over 9 miles.
• If we were able to count atoms at the
rate of 10 million per second, it
would take about 2 billion years to
count the atoms in one mole.
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Mole Trivia
• If there were a mole of rice grains, all the
land area in the whole world would be
covered with rice to a depth of about 75
meters.
• One mole of rice grains is more grains
than all the grain that has been grown
since the beginning of time. (1)
• One mole of rice would occupy a cube
about 120 miles on an edge! (1)
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Mole Trivia
• A mole of marshmallows would cover the
United States to a depth of 600 miles (3)
• In order to put a mole of rain drops in a 30
meter (about 100 feet) diameter tank, the
sides of the tank would have to be 280
times the distance from the Earth to the
Sun. (4)
• A mole of hockey pucks would be equal to
the mass of the Moon.
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Mole Trivia
• Assuming that each human being has 60
trillion body cells (6.0 x 1013) and the
Earth's population is 6 billion (6 x 109), the
total number of living human body cells on
the Earth at the present time is 3.6 x 1023
or a little over half of a mole.
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Mole Trivia
• If one mole of pennies were divided up
among the Earth's population, each
person would receive 1 x 1014 pennies.
• Personal spending at the rate of one
million dollars a day would use up each
persons wealth in about three thousand
years.
• Life would not be comfortable because the
surface of the Earth would be covered in
copper coins to a depth of at least 400
meters.
Mole Trivia
The mole is a REALLY BIG number!!!!
Just how big is it?
Mole Trivia
• If you had a mole of pennies and wanted to buy
kite string at the rate of a million dollars per inch,
you would get your money's worth.
• After stretching your string around the Earth one
million times, and to the Moon and back twentyfive times, you would have enough string left
over to sell back at a dollar an inch (a decided
loss) to gain enough money to buy every man,
woman and child in the US a $50,000
automobile and enough gasoline to run it at 55
mph for a year.
• After those purchases, you would still have
enough money left over to give every man,
woman, and child in the whole world about
$5000.
Mole Trivia
• Basis for calculations:
–Earth's circumference = 25,000
miles
–Distance to moon = 240,000 miles
–Cost of gasoline = $2.50 per gallon
–Gasoline mileage = 20 miles per
gallon
–U.S. population =220,000,000
–World population = 6,000,000,000
• Given the chemical composition of
an everyday item, we can easily
determine the number of atoms or
molecules present in it by applying
the conversions we have practiced in
the previous slides
Avogadro’s Number
• The standard laboratory unit of mass is
the gram.
• We would like to choose a number of
atoms which would have a mass in grams
equivalent to the mass of one atom in
atomic mass units.
• The same number would fit all elements
since equal numbers of different atoms
always have the same mass ratio.
• There is one problem in using the molecular
and formula masses of substances.
• These masses are in atomic mass units,
which is only 1.66 x 10-24 g.
• The mass of a single molecule is so small
that it is impossible to measure it in the
laboratory.
• For everyday use in chemistry, a larger unit,
such as a gram, is needed.
http://www.youtube.com/watch?v=TqDqLmwWx3A
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Chapter 3
Section 4 Counting Atoms
Atomic Mass, continued
Masses of Atoms Are Expressed in Atomic Mass
Units
• A special mass unit is used to express atomic mass.
• This unit has two names—the atomic mass unit
(amu) and the Dalton (Da).
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Chapter 3
Section 4 Counting Atoms
Introduction to the Mole
• Most samples of elements have great numbers of
atoms.
• A mole is defined as the number of atoms in
exactly 12 grams of carbon-12.The mole is the SI
unit for the amount of a substance.
• The molar mass of an element is the mass in
grams of one mole of the element. Molar mass has
the unit grams per mol (g/mol).
• The mass in grams of 1 mol of an element is
numerically equal to the element’s atomic mass
from the periodic table in atomic mass units.
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Other Names Related to Molar
Mass
• Molecular Mass/Molecular Weight: If you have a single
molecule, mass is measured in amu’s instead of grams. But,
the molecular mass/weight is the same numerical value as 1
mole of molecules. Only the units are different. (This is the
beauty of Avogadro’s Number!)
• Formula Mass/Formula Weight: Same goes for
compounds. But again, the numerical value is the same.
Only the units are different.
• THE POINT: You may hear all of these terms
which mean the SAME NUMBER… just different units
Learning Check!
Find the molar mass
(usually we round to the hundreths place)
A. 1 mole of Br atoms =
B. 1 mole of Sn atoms =
79.90 g/mole
118.69 g/mole
Chapter 3
Visual Concepts
Molar Mass
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Calculations with Molar Mass
molar mass
Grams
Moles
What is in a Mole?
It is important to note that;
one mole of atoms = 6.022 x 1023 atoms
one mole of atoms= molar mass in grams
1 mole of Cu = 6.022 x 1023 atoms of Cu
1 mole of Cu = 63.55 g of Cu
mol= ___________atoms
atoms=__________mol
g=_____________mol
mol=____________g
g=____________atoms
atoms=__________g
These problems can be solved in TWO ways:
1) Factor Label Method (conversion factors)(proportions)
We know that:
1 mole of a element = 6.022 x 1023 atoms of that element
= molar mass (atomic weight) in grams
Given (units given)
1
x Conversion ( units of unknown)
factor
(units of given)
= desired answer
Avogadro’s Number as
Conversion Factor
6.022 x 1023 particles
1 mole
or
1 mole
6.02 x 1023 particles
Note that a particle could be an atom OR a molecule!
Using Formulas
Molar Mass = mass ÷ moles
Mass (grams of the subs.) = (moles) (Molar Mass)
Moles = Mass (grams of the subs.) ÷ Molar Mass
Number of Particles = (Moles) (Avogadro’s #)
Ex) What is the mass in grams of 0.586 mol Zn
atoms?
Converting Moles and Grams
Aluminum is often used for the structure
of light-weight bicycle frames. How
many grams of Al are in 3.00 moles of
Al?
3.00 moles Al
? g Al
1. Molar mass of Al
1 mole Al = 27.0 g Al
2. Conversion factors for Al
27.0g Al
1 mol Al
or
1 mol Al
27.0 g Al
3. Setup 3.00 moles Al
Answer
x
27.0 g Al
1 mole Al
= 81.0 g Al
Atoms/Molecules and Grams
• Since 6.02 X 1023 particles = 1 mole
AND
1 mole = molar mass (grams)
• You can convert atoms/molecules to
moles and then moles to grams! (Two step
process)
• You can’t go directly from atoms to
grams!!!! You MUST go thru MOLES.
• That’s like asking 2 dozen cookies weigh
how many ounces if 1 cookie weighs 4 oz?
You have to convert to dozen first!
Calculations
molar mass
Grams
Avogadro’s number
Moles
particles
Everything must go through
Moles!!!
Atoms/Molecules and Grams
How many atoms of Cu are present in 35.4 g
of Cu?
35.4 g Cu
Cu
1 mol Cu
6.02 X 1023 atoms
63.5 g Cu
1 mol Cu
= 3.4 X 1023 atoms Cu
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Learning Check!
How many atoms of K are present in 78.4 g of K?
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Learning Check!
How many atoms of O are present in 78.1 g of oxygen?
78.1 g O2 1 mol O2 6.02 X 1023 molecules O2 2 atoms O
32.0 g O2 1 mol O2
1 molecule O2
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Chapter 3
Section 4 Counting Atoms
Converting from Amount in Moles to Mass
Sample Problem D
Determine the mass in grams of 3.50 mol of copper.
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Chapter 3
Section 4 Counting Atoms
Sample Problem D Solution
• First, make a set-up that shows what is given and
what is desired.
3.50 mol Cu  ? = ? g Cu
• Use a conversion factor that has g Cu in the numerator
and mol Cu in the denominator.
? g Cu
3.50 mol Cu 
= ? g Cu
1 mol
• The correct conversion factor is the molar mass of
Cu, 63.55 g/mol.
63.55 g Cu
3.50 mol Cu 
= 222 g Cu
1 mol
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Chapter 3
Section 4 Counting Atoms
Converting from Amount in Moles to
Number of Atoms
Sample Problem E
Determine the number of atoms in 0.30 mol of fluorine
atoms.
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Chapter 3
Section 4 Counting Atoms
Sample Problem E Answer
• To determine the number of atoms, select the
conversion factor that will take you from the amount
in moles to the number of atoms.
amount (mol)  6.022  1023 atoms/mol =
number of atoms
6.022  1023 F atoms
3.50 mol F 
=
1 mol F
1.8  1023 F atoms
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Ex) How many moles are in 2.83 * x 1025 atoms Cr?
Ex) How many moles are in 6.195 g of K?
Ex) How many grams are in 9.76 x 1024 atoms of
Ba?
Determining Moles of Atoms
Aluminum (Al) is a metal with a high
strength-to-mass ratio and a high
resistance to corrosion: thus it is often
used for structural purposes.
1. Compute the number of moles in a 10.0
gram sample of Aluminum
2. Compute the number of atoms in the
same sample.
Answers
1. 0.371 mol Al atoms
2. 2.23 x 1023 atoms
Calculating Numbers of Atoms
A silicon chip used in an integrated circuit
of a microcomputer has a mass of 5.68
mg. How many silicon (Si) atoms are
present in the chip?
Answer
1.22 x 1020 atoms
Calculating Number of Moles
and Mass
Cobalt (Co) is a metal that is added to
steel to improve its resistance to
corrosion. Calculate both the number of
moles and the number of grams in a
sample of cobalt containing 5.00 x 1020
atoms of cobalt.
Answers
8.30 x 10-4 mol Co
4.89 x 10-2 g Co
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