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Chapter 4
Atomic Structure
Theories about matter were based on
the ideas of Greek philosophers:
Democritus (400 B.C. ) – coins the
term “atom” saying matter can
be subdivided only as small as
an elemental particle.
 from Greek “a” means not and
“tomos” means cutting
 atom means “indivisible”
Aristotle thought that all
matter consisted of 4 basic
elements:
◦earth, air, fire, water
*Their views were not
backed by any evidence.
Dalton’s Atomic Theory (1808)
1.
2.
3.
All matter is composed of
small particles called atoms.
Atoms of a given element are
identical in size, mass, and
other properties.
Atoms cannot be subdivided,
created, or destroyed.
4.
5.
Atoms of different elements combine in
simple-whole number ratios to form
chemical compounds.
In chemical reactions, atoms are
combined, separated, or rearranged.
In 1981 we see atoms!
scanning electron
microscope
OR
(STM – scanning
tunneling microscope)
See page. 100 (Ni atoms)
What instrument is used to
observe individual atoms?
A: scanning electron
microscope (or STM)
B: particle accelerator
C: graduated cylinder
Please make your selection
Democritus, who lived in Greece during
the fourth century B.C., suggested that
matter is made up of tiny particles that
cannot be divided. He called these
particles
.
A) Electrons
B) Ions
C) atoms
Please make your selection
Suppose you could grind a sample of the
element copper into smaller and smaller
particles. The smallest particle that could no
longer be divided, yet still has the
chemical properties of copper, is
___________.
a) A copper atom
b) An ion
c) A proton
Please make your selection
Dalton’s atomic theory included the
idea that the atoms of different
elements can chemically combine in
_______ ratios.
A) Decimal
B) Fraction
C) Partially whole numbers
D) Whole numbers
Please make your selection
The modern process of discovery about
atoms began with the theories of an
English
schoolteacher named
.
A) Aristotle
B) Albert Einstein
C) John Dalton
D) Democritus
Please make your selection
Which of the following is not a part of
Dalton’s atomic theory?
a. All elements are composed of atoms.
b. Atoms of the same element are alike.
c. Atoms are always in motion.
d. Atoms that combine do so in simple
whole-number ratios.
Please make your selection
Dalton theorized that atoms are indivisible
and that all atoms of an element are
identical. Scientists now know that
a. Dalton’s theories are completely correct.
b. atoms of an element can have different
numbers of protons.
c. atoms are all divisible.
d. all atoms of an element are not identical
but they all have the same mass.
Please make your selection
Which of these statements is included in
Dalton’s atomic theory?
a. Chemical reactions occur when atoms are
separated, joined, or rearranged.
b. Some but not all elements are composed of
atoms.
c. Atoms of the same element can have
different numbers of protons.
d. Atoms are divisible.
Please make your selection
The identity of an element can be determined
on the basis of which of the following?
a. the number of protons in an atom of the
element
b. the number of neutrons in an atom of the
element
c. the mass number of the element
d. the atomic mass of the element
What instrument is used to
observe individual atoms?
A: scanning electron
microscope (or STM)
B: particle accelerator
C: graduated cylinder
Section 2: The Structure of the
Atom
consists of:
symbols
Protons
p+
Neutrons
no
location: in nucleus in nucleus
gain/loss: not gained
or lost
Electrons
eoutside nucleus
can be
can be gained
different number or lost
(isotopes)
(forms ions)
consists of:
Protons
mass number: 1 amu
*important
info:
# p+
determines
type of element
(atomic number)
Neutrons
1 amu
*holds nucleus
together
Electrons
0 amu
*involved in
chemical
reactions
Please make your selection
What is the atomic number and the mass
number of an atom with 11 protons and 12
neutrons?
a. atomic number = 11 and mass number = 12
b. atomic number = 12 and mass number = 11
c. atomic number = 11 and mass number = 23
d. atomic number = 23 and mass number = 12
Nuclear Force:
Very strong, short range force that holds
the particles of the nucleus together
Question: What is it about the composition
of the nucleus that requires the concept of
nuclear forces?
A: Like charges do repel each other, so protons
would not be expected to be close to other
protons. Nuclear forces prevent the protons
from repelling each other.
Question: What are the attraction
relationships in the nucleus
(application of nuclear forces)?
Answer:
proton-proton,
neutron-proton,
neutron-neutron
Section 4.2
Structure of the Nuclear Atom
 One
change to Dalton’s atomic
theory is that atoms are divisible
into subatomic particles (p+ n0 e-)
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 (pg. 106)
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
1916 – Robert Millikan determines the mass
of the electron: 1/1840 the mass of a
hydrogen atom; has one unit of negative
charge
Conclusions from the Study of the
Electron:
a) Cathode rays have identical properties
regardless of the element used to
produce them. All elements must contain
identically charged electrons.
b) Atoms are neutral, so there must be
positive particles in the atom to balance
the negative charge of the electrons
c) Electrons have so little mass that atoms
must contain other particles that account
for most of the mass
Conclusions from the Study of the
Electron:
 Eugen Goldstein in 1886 observed
what is now called the “proton” particles with a positive charge, and
a relative mass of 1 (or 1840 times
that of an electron)
 1932 – James Chadwick confirmed
the existence of the “neutron” – a
particle with no charge, but a mass
nearly equal to a proton
Subatomic Particles
Particle
Charge
Mass (g)
Location
Electron
(e-)
-1
9.11 x 10-28
Electron
cloud
Proton
(p+)
+1
1.67 x 10-24
Nucleus
Neutron
(no)
0
1.67 x 10-24
Nucleus
Ernest Rutherford’s
Gold Foil Experiment - 1911
Alpha particles are helium nuclei The alpha particles were fired at a thin
sheet of gold foil
 Particles that hit on the detecting
screen (film) are recorded

Rutherford’s Findings
Most of the particles passed right through
 A few particles were deflected
 VERY FEW were greatly deflected

“Like howitzer shells bouncing
off of tissue paper!”
Conclusions:
a) The nucleus is small
b) The nucleus is dense
c) The nucleus is positively
charged
The Rutherford Atomic Model
 Based
on his experimental evidence:
◦ The atom is mostly empty space
◦ All the positive charge, and almost all the
mass is concentrated in a small area in
the center. He called this a “nucleus”
◦ The nucleus is composed of protons and
neutrons (they make the nucleus!)
◦ The electrons distributed around the
nucleus, and occupy most of the volume
◦ His model was called a “nuclear model”
Rutherford’s atom
Section 4.3
Isotopes
For each element there are different
kinds
 They are chemically alike & react the
same in chemical reactions
 Differ in mass because there are different
numbers of neutrons in the nucleus.

Isotope hyphen notation:
Element name-dash mass number
Example: carbon-14
where14 is the mass number.
(NOTE: mass number = #p+ and #n0)
________ protons
________electrons
________ neutrons
How many protons, neutrons,
and electrons are there in
cobalt-60?
_______ protons
_______ electrons
_______ neutrons
isotope symbols
 Contain
the symbol of the element,
the mass number and the atomic
number.
Mass
Superscript →
number
Subscript →
Atomic
number
X
Symbols

Find each of these:
a) number of protons
b) number of
neutrons
c) number of
electrons
d) Atomic number
e) Mass Number
80
35
Br
Symbols
 If an element has 78
electrons and 117 neutrons
what is the
a) Atomic number
b) Mass number
c) number of protons
d) complete symbol
Symbols
 If an element has 91
protons and 140 neutrons
what is the
a) Atomic number
b) Mass number
c) number of electrons
d) complete symbol
*Another word for isotope is nuclide.
Mass # = p+ + n0
Nuclide
p+
n0
e- Mass #
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
15
16
15
31
Phosphorus - 31
Hydrogen isotopes have names. All other
isotopes are named with the mass number.
Isotope
Protons Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
1
1
2
Hydrogen-3
(tritium)
Nucleus
Measuring Atomic Mass
 Instead
of grams, the unit we use
is the Atomic Mass Unit (amu)
 It is defined as one-twelfth the
mass of a carbon-12 atom.
◦ Carbon-12 chosen because of its isotope purity.
 Each
isotope has its own atomic
mass, thus we determine the
average from percent abundance.
Atomic Masses pg. 116
Atomic mass is the average of all the
naturally occurring isotopes of that element.
Isotope
Symbol
Carbon-12
12C
Carbon-13
13C
Carbon-14
14C
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
Carbon = 12.011
% in nature
98.89%
1.11%
<0.01%
To calculate the average: p. 118
Multiply
the atomic mass of
each isotope by it’s
abundance (expressed as a
decimal), then add the
results.
 If
not told otherwise, the mass of the
isotope is expressed in atomic mass
units (amu)
Example problem:
Element X has two natural isotopes. The
isotope with a mass of 10.012 amu (10X)
has a relative abundance of 19.91%. The
isotope with a mass of 11.009 amu (11X)
has a relative abundance of 80.09%.
Calculate the atomic mass of this
element.
See page 119
Example problem:

The relative abundances and atomic
masses are 0.337% (mass = 35.978
amu), 0.063% (mass = 37.963 amu),
and 99.600% (mass = 39.962 amu),
respectively. Calculate the average
atomic mass of argon.