Chapter 4
Atomic Structure
OBJECTIVES:
• Understand the history of atom
• Differentiate between different models of
the atom
• Identify the no. of protons and neutrons in
a neutral atom based on atomic number Z
and mass number A
KEY TERMS
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Atom
Atomic mass
Atomic mass unit
Atomic number
Cathode ray
Dalton’s atomic theory
Electron
Group
Isotopes
Mass number
Neutron
Nucleus
Period
Periodic table
Proton
SECTIONS
• 4.1 Defining the Atom
• 4.2 Structure of the Nuclear Atom
• 4.3 Distinguishing Among Atoms
4.1
4.1 Defining the Atom
Early Models of the Atom
Democritus’s Atomic Philosophy 400 BC (430?)
“Everything is made up of a few simple parts called
atomos.” Atomos means “uncuttable” in Greek.
He envisioned atomos as small, solid particles of
many different sizes and shapes.
Democritus’s ideas were limited because they didn’t
explain chemical behavior and they lacked
experimental support. His ideas were rejected
because Aristotle supported *the “earth, air, water,
and fire” concept of matter.
Democritus believed that
atoms were solid particles
that are indivisible and
indestructible.
1.ALL MATTER IS COMPOSED OF
EXTREMELY SMALL PARTICLES
CALLED ATOMS
2. ATOMS OF A GIVEN ELEMENT
ARE IDENTICAL IN SIZE, MASS,
AND OTHER PROPERTIES;
ATOMS OF DIFFERENT
ELEMENTS DIFFER IN SIZE,
MASS, & OTHER PROPERTIES
ELEMENT
2
ELEMENT
3
ELEMENT
4
3. ATOMS CANNOT BE
SUBDIVIDED, CREATED, OR
DESTROYED
4. ATOMS OF DIFFERENT
ELEMENTS COMBINE IN
SIMPLE WHOLE # RATIOS
TO FORM CHEM COMPDS
5. IN CHEMICAL RXNS, ATOMS
ARE COMBINED,
SEPARATED, OR
REARRANGED
+
+
Sizing up the Atom
• The radii of most atoms fall within the
range of 5x10-11m to 2 x 10-10m. Individual
atoms are observable with instruments
such as scanning tunneling microscopes.
Discovery of the Electron
In 1897, J.J. Thomson used a cathode ray
tube to deduce the presence of a negatively
charged particle: the electron
4.2
Subatomic Particles
Cathode Ray Tube
A cathode ray is deflected by a
magnet.
A cathode ray is deflected by electrically
charged plates.
Thomson concluded that a
cathode ray is a stream of
electrons. Electrons are parts of
the atoms of all elements.
4.2
Subatomic Particles
In 1886, Eugen Goldstein (1850–1930) observed a
cathode-ray tube and found rays traveling in the
direction opposite to that of the cathode rays. He
concluded that they were composed of positive
particles.
Such positively charged subatomic particles are called
protons.
In 1932, the English physicist James Chadwick (1891–1974)
confirmed the existence of yet another subatomic
particle: the neutron.
Neutrons are subatomic particles with no charge but with a
mass nearly equal to that of a proton.
• Plum Pudding Model
– JJ Thomson – 1897
• Most of the atom was pos. charged
• Atoms had negative charges embedded in pos.
charged material
• Rutherford’s Model
– Ernest Rutherford – 1911
• Gold Foil Experiment – Rutherford fired positively
charged particles at a sheet of gold foil
• Rutherford discovered the nucleus as a result of
his experiment
4.2
The Atomic Nucleus
Rutherford’s Gold-Foil Experiment
In 1911, Rutherford and his coworkers at the University
of Manchester, England, directed a narrow beam of
alpha particles at a very thin sheet of gold foil.
Alpha particles scatter from the gold foil.
4.2
The Atomic Nucleus
The Rutherford Atomic Model
Ernest Rutherford concluded that the atom is mostly
empty space. All the positive charge and almost all of
the mass are concentrated in a small region called
the nucleus.
The nucleus is the tiny central core of an atom and is
composed of protons and neutrons.
In the nuclear atom, the protons and neutrons are
located in the nucleus.
The electrons are distributed around the nucleus and
occupy
almost all the volume of the atom.
4.3
4.3 Distinguishing Among Atoms
Just as apples
come in different
varieties, a
chemical element
can come in
different “varieties”
called isotopes.
4.3
The Periodic Table—A Preview
A periodic table is an arrangement of
elements in which the elements are
separated into groups based on a set of
repeating properties.
A periodic table allows you to easily compare the properties
of one element (or a group of elements) to another
element (or group of elements).
Each horizontal row of the periodic table is called a period. Within a
given period, the properties of the elements vary as you move
across it from element to element.
Each vertical column of the periodic table is called a group, or family.
Elements within a group have similar chemical and physical
properties.
4.3
The Periodic Table—A Preview
• A Period
Period goes across
4.3
Isotopes
Isotopes are atoms that have the same
number of protons but different numbers of
neutrons.
Because isotopes of an element have different numbers of
neutrons, they also have different mass numbers.
Despite these differences,
isotopes are chemically alike
because they have identical
numbers of protons and
electrons.
Isotopic notation
Mass number = p + n (in the nucleus)
Element
symbol
A
Z
E
• Atomic number = # p
– Elements are put in order of atomic number
on the periodic table.
Ex:
An atom of carbon with 7 neutrons:
13C
6
An atom of lead with 125 neutrons:
207Pb
82
4.3
Mass Number
Au is the chemical symbol for gold.
How many protons,
electrons, and neutrons
does a gold atom have?
The atomic number is 79. Therefore,
there are 79 protons and 79 electrons.
The mass number is 197, which is the
total number of protons and neutrons.
Therefore, 197-79= 118 neutrons.
Practice #2 (on worksheet)
4.3
Atomic Mass
An atomic mass unit
(amu) is defined as
one twelfth of the
mass of a carbon12 atom.
It is useful to to compare the
relative masses of atoms to
a standard reference
isotope. Carbon-12 is the
standard reference isotope.
Carbon-12 has a mass of
exactly 12 atomic mass
units.
Some elements and
their isotopes
4.3
Atomic Mass
The atomic mass of an element is a
weighted average mass of the atoms in a
naturally occurring sample of the element.
A weighted
average mass
reflects both the
mass and the
relative
abundance of
the isotopes as
they occur in
nature.
4.3
Atomic Mass
To calculate the atomic mass of an element,
multiply the mass of each isotope by its natural
abundance, expressed as a decimal, and then
add the products.
For example, carbon has two stable isotopes:
Carbon-12, which has a natural abundance of 98.89% and
a mass of 12.000 amu
Carbon-13, which has a natural abundance of 1.11% and a
mass of 13.003 amu.
Silver is found in two isotopes with atomic
masses 106.9041 and 108.9047 amu,
respectively. The first isotope represents
51.82% and the second 48.18%.
Determine the average atomic mass of
silver.
(106.9041)(.5182)= 55.398
(108.9047)(.4818)= 52.470
55.398 + 52.470 = 107.868 amu
• Planetary Model – Bohr Model
– Neils Bohr – 1913
• Electrons move in definite orbits
• Orbits are referred to as energy levels
4.2 The Structure of the Atom
• Atoms consist of subatomic particles
– Protons
• Positive charge
• Located in the nucleus
– Neutrons
• No charge
• Located in the nucleus
– Electrons
• Negative Charge
• Located in the electron cloud
Atomic Number and Mass Number
• Atomic Number Z
– # of protons in the nucleus of an atom
– Different elements have different numbers of
protons
• Mass Number A
– Sum of an atoms protons and neutrons
– n0 = A - Z
Isotopes
• Isotopes are atoms
of the same element
with different
numbers of neutrons
and different mass
numbers.