Chemistry: Chapter 4
Atomic Structure
The Atom
An atom is the
smallest particle of
an element that
retains the identity
of that element.
If we repeatedly cut
a piece of
aluminum, the
smallest possible
piece is an atom of
aluminum
Democritus (460 – 370 BCE)
– Greek Philosopher
Proposed that all matter
was made of tiny
indivisible particles.
He called these particles
atomos (meaning
indivisible).
We call them atoms.
the “laughing philosopher”
Aristotle (384 – 322 BCE)
– Greek Philosopher
Did not agree with
Democritus.
Believed matter was
made up of only one
substance called “hyle”
(greek for wood, or
materials)
It wasn’t until the
1700’s when his ideas
were reexamined.
Robert Boyle (1627 – 1691)
Isaac Newton (1642 – 1727)
Published articles stating their
belief in the atomic nature of
elements
They had no proof
only 13 elements were known:
antimony, arsenic, bismuth,
carbon, copper, gold, iron,
lead, mercury, silver, sulfur,
tin, and zinc.
Robert Boyle ↑
← Isaac Newton
Antoine Lavoisier (1743 – 1794)
– French Chemist
The “Father of Modern Chemistry”
Developed the “Law of Conservation
of Matter.”
Matter is neither created nor destroyed.
Lavoisier published a list of elements,
adding another 11: chlorine, cobalt,
hydrogen, manganese, molybdenum,
nickel, nitrogen, oxygen, phosphorus,
platinum, and tungsten.
Pioneered Stoichiometry
Guillotined on May 8th in Paris, 1794
Joseph Proust (1754 – 1826)
– French Chemist
French Chemist
Developed The Law
of Definite
Proportions
Compounds always
contain elements in
the same proportion
by mass.
Law of Definite Proportions
H20 (by mass is
always)
88.9% Oxygen,
11.1% Hydrogen
If we had an 80g
sample of H20 how
much is O?
 .889 x 80 = 71g
How much is H?
 .111 x 80 = 9g
John Dalton (1776 – 1844)
– English Scientist
Proposed the
atomic theory of
matter, which is
the basis for present
atomic theory
John Dalton, English schoolteacher
Atomic Theory of Matter
Each element is
composed of
extremely small
particles called
atoms.
All atoms of a given
element are
identical, but differ
from those of any
other element.
Which element is this?
Atomic Theory of Matter
When elements unite
to form compounds, they
do so in a ratio of small
whole numbers. This is
called the Law of
Multiple Proportions.
Ex: C and O can
combine to form CO or
CO2, but not CO1.8.
Dalton’s Model of an Atom
All matter is composed of tiny solid particles
Joseph Louis Gay-Lussac (1778 – 1850)
– French Chemist
 Observed that working with
gas reactions at constant
volume: temperature and
pressure are directly
related.
 He named the discovery of
this relationship Charles
Law, which is represented
by:
P1
P2
=
T1
T2
Amadeo Avogadro (1776 – 1856)
– Italian Physicist
Explained GayLussac’s work using
Dalton’s theory: Equal
volumes of gases at
the same temperature
and pressure have
the same number of
gas molecules.
Michael Faraday (1791 – 1867)
– English Physicist
Suggested that atomic
structure was related to
electricity.
Atoms contain particles that
have electrical charges.
 Positive (+)
 Negative (-)
Opposite charges attract
Like charges repel
William Crookes (1832 – 1919)
– English Physicist
Developed the cathode ray tube
to find evidence for the existence
of particles within the atom.
J.J. Thomson (1856 – 1940)
– English Physicist
Used a cathode ray tube to
identify negatively charged
particles, called electrons.
Determined the ratio of an
electron’s charge to its
mass.
Developed the “plum
pudding” model of an
atom.
Cathode Ray Tube Experiment
Plum Pudding Model
Atoms are composed
of a solid sphere with
charged particles
embedded
throughout
Robert Millikan(1868 – 1953)
– American Physicist
US Physicist
Used the oil drop
experiment to prove the
electron has a negative
charge
Was able to determine the
charge of the electron
Millikan’s Oil Drop Experiment
Ernest Rutherford (1871 – 1937)
– New Zealand Physicist
Used the gold foil
experiment to prove the
atom is mostly empty
space.
Rutherford concluded that all
of an atom’s positive charge,
and most of its mass is
located in the center, called
the nucleus.
Gold Foil Experiment
Gold Foil Experiment
98% of the particles passed
straight through
2% of the particles deflected
off at varying angles
0.01% of the particles
bounced straight back
Rutherford’s
Planetary Model of an atom
Positive charge
and majority of
mass located in
the nucleus.
Negatively
charged electrons
orbit the nucleus
like planets.
James Chadwick (1891 – 1974)
– English Physicist
Found high energy
particles with no charge
with the same mass as
the proton called
neutrons.
Max Planck (1858 – 1947)
– German Physicist
Planck demonstrated that the
electron in orbit around the
nucleus accelerates because it
changes direction.
Because the electron has charge,
acceleration means a changing
electric field, which means
photons should be emitted.
But, then the electron would
lose energy and fall into the
nucleus.
Niels Bohr (left) and Max Planck
Planck’s Dilemma – Why don’t the
electrons fall into the nucleus?
Planck determined that energy, at the
sub-atomic level, can only be transferred
in small units, called quanta.
Quanta are not divisible.
There are no “in between” energy levels.
Quantization limits the energy to be
transferred to photons and resolves the
dilemma of the electron falling into the
nucleus.
Components of the Atom
The smallest particle
of an element that
has the properties of
that element.
Make up of nucleus
consists of protons
and neutrons
Surrounded by an
electron cloud
Protons (p+)
Found in the nucleus
Positive (+) charge
Composed of 3 quarks
(2 up, 1 down)
Mass= 1.6726 x 10-27 kg
Atomic mass  1 amu (µ)
The number of protons in an atom
refers to the atomic number (Z)
Neutrons (n0)
Found in the nucleus
Neutral (0) charge
Composed of 3 quarks
(1 up, 2 down)
Mass= 1.6749 x 10-27 kg
Atomic mass  1 amu (µ)
Isotopes: atoms of the same
element that have a different
number of neutrons.
Electrons (e–)
Found in electron clouds
surrounding the nucleus.
Negative (–) charge
Elementary Particle
Mass =9.1094 x 10-31 kg
1800 times smaller than
protons & neutrons
Mass  0 amu (µ)
Electrons (e–)
Have negligible mass
(when compared to
protons and neutrons)
Orbit the nucleus at
very high speed in
energy levels
(electron clouds).
Atomic Number = Protons
 The atomic number of an
element is the number of
protons an element has.
 Located above the symbol
of the element
 The number of protons
determines the identity of
the element.
 Each element has a different
atomic number
How many electrons are in an atom?
For an atom to have an
overall neutral charge
the number of electrons
must equal the number
of protons.
p+ = e–
What element is this?
Mass number
 The Mass number of an atom
is the average of the mass of
the isotopes of that element
 Located below the symbol of
the element
 Atomic mass is measured in
amu’s, (atomic mass units)
 Based on Carbon having a mass
of 12
Mass = Protons + Neutrons
How many neutrons are in this
isotope of platinum?
 mass = p+ + n0
 196 = 78 + n0
 n0 = 196 – 78
 Platinum has 118 Neutrons
 Find the number of neutrons in:
Hydrogen - 1 Carbon - 13
Helium - 4
Potassium - 40
Boron - 11
Gold - 198
196
mass = p+ + n0
 Hydrogen–1 (11H)
hydrogen–1 has 0 neutrons
 Helium–4 (42He)
helium–4 has 2 neutrons
 Boron–11 (11
B)
5
boron–11 has 6 neutrons
 Carbon–13 (13
C)
6
carbon–13 has 7 neutrons
 Potassium–40 (40
K)
19
 Gold–198 (198
Au)
79
potassium–40 has 21 neutrons
gold–198 has 119 neutrons
Average Atomic Mass Number
 The average mass of all of the isotopes of an element.
 Also known as the average atomic mass number, or
atomic weight.
 Isotopes: atoms of the same element with different
masses.
Average Atomic Mass

20
Ne
10
22
Ne
10






90.00% of a sample of neon is neon–20
10.00% of a sample of neon is neon–22
Calculate the average atomic mass of neon
.90 x 19.992
=
17.9928
.10 x 21.991
=
2.1991
average mass =
20.1919 amu

has a mass of 19.992 amu
has a mass of 21.991 amu.
Ions
An atom that has gained or lost an electron.
It acquires a net electrical charge.
If an atom loses an electron (oxidation) it has more
protons than electrons and has a net positive charge.
(cation)
Ions
If an atom gains an electron (reduction) it has more
electrons than protons and has a net negative charge
(anion)
Ionic Charges
Charge of ion = p+ – e–
What is the charge of a magnesium atom that
loses 2 electrons?
 Number of p+
12
 Number of e–
10
charge of ion
+2
Mg2+ or Mg+2
Charge is written to the upper right of the
symbol.
Representations of atoms
General form:
(Elemental Notation)
X = Element Symbol
A = Atomic Mass
(p+ + n0)
Z = Atomic Number
(p+)
Ionic Charge
A charge
X
Z
What is the atomic structure?
Determine the
number of:
p+ =
0
n =
–
e =
23 +
Na
11
What is the atomic structure?
Determine the
number of:
p+ = 11
0
n = 12
–
e = 10
23 +
Na
11
What is the atomic structure?
Determine the
number of:
p+ =
0
n =
–
e =
127 −
I
53
What is the atomic structure?
Determine the
number of:
p+ = 53
0
n = 74
–
e = 54
127 −
I
53
Put into elemental notation
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2
?
How many electrons?
Atomic # = 29
Atomic Mass = 64
Ionic charge = +2
Number of e– =
64 2+
Cu
29
Put into elemental notation
37 p+
48 n0
36 e–
?
Put into elemental notation
37 p+
48 n0
36 e–
85 +
Rb
37
Alpha Decay (α)
2 protons and 2 neutrons
bound together
emitted from the nucleus
during radioactive decay
Occurs when nucleus is
unstable
4
 2He2+
Beta Decay (β)
An electron is emitted
from the nucleus
This occurs when the
neutron to proton
ratio is too large
A neutron decays into
a proton, an electron,
and a antineutrino
Gamma Decay (γ)
High energy
electromagnetic waves
(no mass) emitted from
the nucleus as it loses
energy
Very penetrating form
of radiation
High energy radiation,
stopped by lead plates.