Pearson Chemistry Chap_03_Atoms_Sec_1-3 (1).ppt

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Table of Contents
Atoms and Moles
Section 1 Substances Are Made of Atoms
Section 2 Structure of Atoms
Section 3 Electron Configuration
Section 4 Counting Atoms
Chapter 3

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Objectives
• State the three laws that support the existence of
atoms.
• List the five principles of John Dalton’s atomic
theory.
Chapter 3
Section 1 Substances Are Made
of Atoms

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Atomic Theory
• The idea of an atomic theory is more than 2000 years
old.
• 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.
Chapter 3
of Atoms

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Atomic Theory, continued
Chapter 3
of Atoms
• The figure on the right is a more accurate
representation of an atom than the figure on the left.

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The Law of Definite Proportions
• 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.
Chapter 3
of Atoms

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The Law of Conservation of Mass
• 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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of Atoms

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Law of Conservation of Mass
of Atoms
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Law of Conservation of Mass, continued
of Atoms
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The Law of Multiple Proportions
• 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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of Atoms

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Law of Multiple Proportions
of Atoms
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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.
of Atoms
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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.
Chapter 3
of Atoms

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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.
Chapter 3
of Atoms

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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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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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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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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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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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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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An Electron Has a Negative Charge, continued
• Properties of Electrons
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Rutherford Discovered the Nucleus
• 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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Rutherford Discovers the Nucleus, continued
• Ernest Rutherford performed the gold foil experiment,
which disproved the plum-pudding model of the atom.
• 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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Gold Foil Experiment
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• 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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Gold Foil Experiment on the Atomic Level
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• 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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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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Proton and Neutrons Compose the Nucleus, continued
• Properties of a Proton and a Neutron
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Protons and Neutrons Can Form a Stable Nucleus
• Coulomb’s law states that the closer two charges are,
the greater the force between them.
Chapter 3
• The repulsive force between two protons is large
when two protons are close together.

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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.
• All atoms that have more than one proton also
have neutrons.
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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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Atomic Number and Mass Number, continued
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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Nucleus, continued
• The atomic number for oxygen tells you that the
oxygen atom has 8 protons and 8 electrons.
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Visual Concepts
Atomic Number
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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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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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Nucleus, continued
• The neon atom has 10 protons, 10 electrons, and
10 neutrons. The mass number is 20.
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Visual Concepts
Mass Number
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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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• The atomic number always appears on the lower left
side of the symbol.
1 2 3 4 5
H He Li Be B
1 2 3 4 6 7 9 10 11
H H He He Li Li Be B B
• Mass numbers are written on the upper left side of the
symbol.
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• Both numbers may be written with the symbol.
1 4 3
1 2 1
H H H
1 4 7 9 11
1 2 3 4 5
H He Li Be B
• An element may be represented by more than one
notation.
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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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Number, continued
• The second method of identifying isotopes shows the
composition of a nucleus as the isotope’s nuclear
symbol.
3 4
2 2
or
He He
Chapter 3
• 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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Number, continued
• The two stable helium isotopes are helium-3 and
helium-4.
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Number, continued
• The Stable Isotopes of Lead
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Visual Concepts
Isotopes and Nuclides
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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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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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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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Bohr’s Model Confines Electrons to Energy Levels
• Rutherford’s model of an atom
• electrons orbit the nucleus
just as planets orbit the
sun
• Bohr’s model of an atom
• electrons travel
around the nucleus in
specific energy levels
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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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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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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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Visual Concepts
Orbital
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Electrons and Light
• By 1900, scientists knew that light could be thought of
as moving waves that have given frequencies,
speeds, and wavelengths.
• In empty space, light waves travel at 2.998  108
m/s.
• The wavelength is the distance between two
consecutive peaks or troughs of a wave.
• The distance of a wavelength is usually measured in
meters.
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Electromagnetic Spectrum
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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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• 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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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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Wavelength and Frequency
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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 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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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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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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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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Hydrogen’s Line-Emission Spectrum
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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.
• 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.
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Quantum Numbers, continued
• 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,
and 4.
• As n increases, the electron’s distance from the
nucleus and the electron’s energy increases.
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Visual Concepts
Principal Quantum Number
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• 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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• 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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• The spin quantum number, indicates the orientation of
an electron’s magnetic field relative to an outside
magnetic field.
• The spin quantum number is represented by:
or or (
1
or
1
2
)
+
2
  
Chapter 3
• A single orbital can hold a maximum of two
electrons, which must have opposite spins.

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• Quantum Numbers of the First 30 Atomic Orbitals
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Electron Configurations
• 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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Electron Configurations, continued
• 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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• The arrangement of electrons in an atom is usually
shown by writing an electron configuration.
• 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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Shapes of s, p, and d Orbitals
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Visual Concepts
s Orbitals
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Visual Concepts
p Orbitals
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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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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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Available, continued
• This diagrams shows how
the energy of the orbitals
can overlap.
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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
1s2
2s2
2p6
3s2
3p4
.
• 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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An Electron Configuration Is 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 1s2
2s2
2p6
, so the
electron configuration of sulfur is
[Ne] 3s2
3p4
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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
s
1
  s
2
  p
2
     
p
3
   
s
3
 
Chapter 3

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