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Atomic Structure
Chemistry
Lesson 1
• Explain the law of conservation of mass,
the law of definite proportions, and the law
of multiple proportions.
• Summarize the five essential points of
Dalton’s atomic theory.
• Explain the relationship between Dalton’s
atomic theory and the law of conservation
of mass, the law of definite proportions,
and the law of multiple proportions.
Write these definitions in your
journal.
• Law of definite proportions: a chemical
compound contains the same elements in
exactly the same proportions by mass
regardless of the size of the sample or source of
the compound
• Law of multiple proportions: if two or more
different compounds are composed of the same
two elements, then the ratio of the masses of the
second element combined with a certain mass of
the first element is always a ratio of small whole
numbers
Dalton’s Atomic Theory
• All matter is composed of extremely small
particles called atoms.
• Atoms of a given element are identical in
size, mass, and other properties; atoms of
different elements differ in size, mass, and
other properties.
• Atoms cannot be subdivided, created, or
destroyed.
• Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds.
• In chemical reactions, atoms are
combined, separated, or rearranged.
Not all aspects of Dalton's theory have proven
to be correct.
We now know that:
•Atoms are divisible into even smaller particles.
•A given element can have atoms with different
masses.
Some important concepts remain unchanged.
• All matter is composed of atoms.
• Atoms of any one element differ in
properties from atoms of another element.
Lesson 2
• Summarize the observed properties of
cathode rays that led to the discovery of
the electron.
• Summarize the experiment carried out by
Rutherford and his co-workers that led to
the discovery of the nucleus.
Discovery of the Electron
•Experiments in the late 1800s showed that
cathode rays were composed of negatively
charged particles.
•These particles were named electrons.
Charge and Mass of the Electron
•Joseph John Thomson’s cathode-ray tube
experiments measured the charge-tomass ratio of an electron.
Thomson's Experiment
•Robert A. Millikan’s oil drop experiment
measured the charge of an electron.
Millikan's Experiment
•With this information, scientists were able to
determine the mass of an electron.
Discovery of the Atomic Nucleus
•More detail of the atom’s structure was provided
in 1911 by Ernest Rutherford and his associates
Hans Geiger and Ernest Marsden.
•The results of their gold foil experiment led to the
discovery of a very densely packed bundle of
matter with a positive electric charge.
•Rutherford called this positive bundle of matter
the nucleus.
Gold Foil Experiment
The Structure of the Atom
•An atom is the smallest particle of an
element that retains the chemical
properties of that element.
•The nucleus is a very small region located
at the center of an atom.
•The nucleus is made up of at least one
positively charged particle called a proton
and usually one or more neutral particles
called neutrons.
•Surrounding the nucleus is a region
occupied by negatively charged particles
called electrons.
•Protons, neutrons, and electrons are often
referred to as subatomic particles.
Composition of the Atomic Nucleus
• Except for the nucleus of the simplest type of
hydrogen atom, all atomic nuclei are made of
protons and neutrons.
• A proton has a positive charge equal in
magnitude to the negative charge of an electron.
• Atoms are electrically neutral because they
contain equal numbers of protons and electrons.
• A neutron is electrically neutral.
• The nuclei of atoms of different elements
differ in their number of protons and
therefore in the amount of positive charge
they possess.
• Thus, the number of protons determines
that atom’s identity.
Lesson 3
• Discuss the significance of the lineemission spectrum of hydrogen to the
development of the atomic model.
• Describe the Bohr model of the hydrogen
atom.
The Particle Description of Light
• A quantum of energy is the minimum quantity of
energy that can be lost or gained by an atom.
• German physicist Max Planck proposed the
following relationship between a quantum of
energy and the frequency of radiation:
E = hv
E is the energy, in joules, of a quantum of radiation, v is the
frequency, in s−1, of the radiation emitted, and h is a fundamental
physical constant now known as Planck’s constant; h = 6.626 
1034 J• s.
• A photon is a particle of electromagnetic
radiation having zero mass and carrying a
quantum of energy.
• The energy of a particular photon depends
on the frequency of the radiation.
Ephoton = hv
• Quantization of energy
• Energy of a photon
The Hydrogen-Atom Line-Emission
Spectrum
• The lowest energy state of an atom is its
ground state.
• A state in which an atom has a higher
potential energy than it has in its ground
state is an excited state.
• When investigators passed electric current
through a vacuum tube containing hydrogen
gas at low pressure, they observed the
emission of a characteristic pinkish glow.
• When a narrow beam of the emitted light
was shined through a prism, it was
separated into four specific colors of the
visible spectrum.
• The four bands of light were part of what is
known as hydrogen’s line-emission
spectrum.
Bohr Model of the Hydrogen Atom
Absorption and Emission Spectra
• Niels Bohr proposed a hydrogen-atom model
that linked the atom’s electron to photon
emission.
• According to the model, the electron can circle
the nucleus only in allowed paths, or orbits.
• The energy of the electron is higher when the
electron is in orbits that are successively farther
from the nucleus.
Bohr Model of the Atom
• When an electron falls to a lower energy
level, a photon is emitted, and the process
is called emission.
• Energy must be added to an atom in order
to move an electron from a lower energy
level to a higher energy level. This process
is called absorption.
Photon Emission and Absorption
Bohr model of a helium atom and a beryllium atom
Helium, He
Berylium, Be
– List three similarities and three
differences.
– How do you think a gold atom is
different from a copper atom?
Definitions
• Atomic number is the number of
protons in the nucleus of an atom.
• Mass number is the mass of an
individual atom.
• Atomic mass is the mass of a single
atom. It is the decimal number on the
periodic table.
• Atomic mass units (amu) are
“invented” measurement units of the
atomic mass.
element
chemical
symbol
atomic
# of
number protons
# of
electrons
# of
neutrons
beryllium
5
fluorine
10
6
18
lead
126
19
tin
35.45
39
70
tungsten
184
29
gold
atomic
weight
12
chlorine
potassium
mass
number
36
118
183.85
• What is different about the two atoms?
• What is the atomic number of each
atom?
• What is the mass number of each atom?
• Do you think they are both lithium
atoms? Why or why not?
• Atoms of the same element that have
different numbers of neutrons are
called isotopes.
boron
atom
# protons
# neutrons
#
electrons
1
2
3
4
5
6
7
8
9 10
Element
Boron
Chemical Atomic
Symbol Number
Atomic
Weight
# of
# of
electron neutrons
s
B
5 or 6
Chlorine
17
Lithium
Vanadium
# of
proton
s
6.94
V
23
Nitrogen
7
Magnesium
Argon
Ar
39.9
18, 20,
or 22
Comparing Atomic Models
Lesson 4
• Compare and contrast the Bohr model and the
quantum model of the atom.
• Explain how the Heisenberg uncertainty
principle and the Schrödinger wave equation led
to the idea of atomic orbitals.
• List the four quantum numbers and describe
their significance.
• Relate the number of sublevels corresponding
to each of an atom’s main energy levels, the
number of orbitals per sublevel, and the number
of orbitals per main energy level.
Electrons as Waves
• Scientists knew that any wave confined to
a space could have only certain
frequencies.
• Louis DeBroglie pointed out that in many
ways the behavior of electrons in Bohr's
orbits was similar to the behavior of
waves.
• He suggested that electrons be
considered waves in a confined space
around an atomic nucleus.
The Heisenberg Uncertainty
Principle
• German physicist Werner Heisenberg
proposed that any attempt to locate a
specific electron with a photon knocks the
electron off its course.
• The Heisenberg uncertainty principle
states that it is impossible to determine
simultaneously both the position and
velocity of an electron or any other
particle.
Heisenberg Uncertainty Principle
The Schrödinger Wave Equation
• In 1926, Austrian physicist Erwin Schrödinger
developed an equation that treated electrons in
atoms as waves.
• Together with the Heisenberg uncertainty
principle, the Schrödinger wave equation laid the
foundation for modern quantum theory.
• Quantum theory describes mathematically the
wave properties of electrons and other very
small particles.
• Electrons do not travel around the nucleus
in neat orbits, as Bohr had postulated.
• Instead, they exist in certain regions called
orbitals.
• An orbital is a three-dimensional region
around the nucleus that indicates the
probable location of an electron.
Electron cloud
Atomic Orbitals and Quantum
Numbers
• Quantum numbers specify the properties of
atomic orbitals and the properties of electrons in
orbitals.
• The principal quantum number, symbolized by
n, indicates the main energy level occupied by
the electron.
• The angular momentum quantum number,
symbolized by l, indicates the shape of the
orbital.
Orbitals and Quantum Numbers
• The magnetic quantum number,
symbolized by m, indicates the orientation
of an orbital around the nucleus.
• The spin quantum number has only two
possible values—(+1/2 , 1/2)—which
indicate the two fundamental spin states of
an electron in an orbital.
Quantum Numbers of the First 30
Atomic Orbitals
Electrons Accommodated in
Energy Levels and Sublevels