Chapter 4: Atoms
The Building Blocks of Matter
An atom is the smallest particle of an element that
retains the chemical properties of that element.
Section 1
The Atom: From Philosophical
Idea to Scientific Theory

Page 67
The Early Atom


As early as 400 B.C., Democritus
called nature’s basic particle the
“atomon” based on the Greek word
meaning “indivisible”.
Aristotle succeeded Democritus
and did not believe in atoms.
Instead, he thought that all matter
was continuous. It was his theory
that was accepted for the next 2000
years. (Read page 43 of your
textbook.)
Three Basic Laws of Matter:
 Law
of Conservation of Mass
 Law of Definite Proportions
 Law of Multiple Proportions
Basic Laws of Matter

Law of Conservation of Mass- mass is
neither created nor destroyed during
ordinary chemical reactions or physical
changes.
CH4 + 2O2 → 2H2O + CO2
16g + 64g → 36g + 44g
Antoine Lavoisier
stated this about 1785
Antoine Lavoisier and his wife, Marie-Anne
"It took them only an instant to cut off that head, and a hundred years may not
produce another like it." Joseph-Louis Lagrange
Alka Seltzer in Water
Ziploc bag
 Alka seltzer tablet
 Water

Using the reaction
between the tablet
and the water, prove
that the Law of
Conservation of
matter is true.
HOMEWORK
Read Section 1
 Complete Questions 1-3 of the Section 1
Review (page 71) on a separate sheet of
paper to be collected.

Basic Laws of Matter

Law of Definite Proportions – no matter how much
salt you have, it is always 39.34% Na and 60.66% Cl by
mass.
Example: Sodium chloride always contains
39.34% Na and 60.66% Cl by mass.
2NaCl
100g
116.88g
→ 2Na + Cl2
→ 39.34g + 60.66g
→
? + ?
Joseph Louis Proust
stated this in 1794.
Basic Laws of Matter

Law of Multiple Proportions- Two or more
elements can combine to form different
compounds in whole-number ratios.
Example
John Dalton
proposed this
in 1803.
John Dalton’s Elements
Dalton’s Atomic Theory

In 1808, Dalton proposed a theory to
summarize and explain the laws of
conservation of mass, definite proportions,
& multiple proportions.
I was a school
teacher at the
age of 12!
Dalton’s Atomic Theory
John Dalton - 1808
1.
2.
3.
4.
5.
All matter is composed of extremely small particles
called atoms.
Atoms of a given element are identical 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.
**Today, we know these parts to have flaws.
Flaws of Dalton’s Theory…
2. Atoms of a given element are
identical in size, mass, and other
properties. Isotopes – atoms with the same
number of protons but a different
number of neutrons
3. Atoms cannot be subdivided,
created, or destroyed.
Subatomic particles – electrons,
protons, neutrons, and more
Section 2
 The

Structure of the Atom
Page 72
The Atom

Atom - the smallest particle of an element
that retains the chemical properties of that
element.
CARBON
The Structure of the Atom

The atom is composed of two main
regions, the nucleus & the electron cloud.
Nucleus of an Atom

Nucleus- very small region located at the center of the
atom. The nucleus accounts for most of an atoms
mass but very little volume, making it a very dense
region.
M
D=
V

The nucleus contains protons, neutrons, and more.
proton = p+
neutron = no
others – neutral, too
Electron Cloud of an Atom

The electron cloud is the negatively
charged region of the atom that accounts
for most of the atom’s volume but very
little of the atom’s mass.
The electron cloud is
composed of a number of
electrons, of which depends
the element.
electron = e-
Checking for Understanding

Does an electron from gold, act like gold?
NO, an electron is like any other electron, no matter the source.

What are the two main regions of the
atom?
The nucleus and the electron cloud are the two main regions.

What is the charge on the nucleus?
The nucleus is positive since it holds protons (+), neutrons (0)
and other neutral particles.
Subatomic Particles




Protons- positively charged particles found in
the nucleus of an atom.
Neutrons- neutral particles found in the nucleus
of an atom.
Electrons- negatively charged particles found in
the electron cloud.
Others – photon, boson, gluon, lepton, muon,
quark, tau, neutrino, meson, …
Properties of Subatomic
Particles
Particle Symbol Charge Mass #
Relative
Mass
(amu)
Actual
Mass (g)
Electron
e-
-1
0
0.0005486 9.109 X 10-28
Proton
p+
+1
1
1.007276
1.673 X 10-24
Neutron
no
0
1
1.008665
1.675 X 10-24
1 amu (atomic mass unit) = 1.660540 x 10-27 kg or
exactly 1/12 the mass of a carbon-12 atom
Discovery of the Subatomic
Particles


The discovery of the subatomic particles came
about from the study of electricity & matter.
Benjamin Franklin’s kite experiment in 1752
demonstrated that lightning was electrical.
Charged Particles
 In
1832, Michael
Faraday proposed
that objects are
made of positive
and negative
charges.
Discovery of the Electron
In the late 1870’s many experiments were
performed in which electric current was
passed through gases at low pressures due
to the fact that gases at atmospheric
pressure don’t conduct electricity well.
 These experiments were carried out in glass
tubes called cathode-ray tubes or Crookes
tubes.
 Sir William Crookes developed
these tubes.

Crookes Tube
CRT
Discovery of the Electron
When current was passed through the
cathode ray tube, the surface of the tube,
directly opposite the cathode, glowed.
 It was thought that this glow was caused
by a stream of particles called cathode
rays.
 The rays traveled from cathode (negative)
to anode (positive).

Discovery of the Electron
Negatively charged objects deflected the
rays away.
 Therefore, it was determined that the
particles making up the cathode rays were
negatively charged.

Joseph John Thomson

In 1897 the English physicist Joseph John
Thomson was able to measure the ratio of
charge of the cathode ray particles to their mass.

He found that the ratio was always the same
regardless of the metal used to make the
cathode or the nature of the gas inside the
cathode ray tube.

Thomson concluded that cathode rays were
composed of identical, negatively charged
particles called electrons.
Joseph John Thomson
Thomson’s experiments revealed that the
electron has a very large charge-to-mass
ratio.
 Thomson determined that electrons were
present in all elements because he noted
that cathode rays had identical properties
regardless of the element used to produce
them.

Cathode Ray Tube Experiment
Accomplishments
Proved that the atom was divisible and that
all atoms contain electrons.
 This contradicted Dalton’s Atomic Theory.
 This allowed a new model of the atom.

Plum-Pudding Model of the Atom
Checking for Understanding
Cathode Ray Tube

Why were the cathode rays deflected?
They were negatively charged, so they were repelled from the
negative plate and attracted to the positive plate.

Why did they assume there was a positive
portion to the atom?
They knew the atom was neutral, so by default, there must be a
positive portion if there are negative particles.

How did this contradict Dalton’s model of
the atom?
Dalton stated that atoms cannot be subdivided. Electrons are
subatomic particles.
Robert A. Millikan

In 1909, Robert Millikan performed an Oil
Drop Experiment & calculated the charge
of the electron.
Oil Drop Experiment



Millikan dropped negatively charged microscopic
oil particles into a chamber containing metallic
plates and viewed them with a microscope.
By applying voltage to the metallic plates,
Millikan created an electric field.
He was able to suspend the
oil droplets by adjusting the
electric field to the
appropriate strength and
direction to overcome
gravity.
Oil Drop Experiment



Knowing the mass of the droplets and the
strength of the electric field necessary to
suspend them, he was able to calculate the
charge of the electron.
He noticed that the charge was always a wholenumber multiple of 1.602 X10-19 Coulombs.
He determined that the charge of the electron to
be 1.602 X 10-19 C.
Checking for Understanding
Oil Drop Experiment

What year did Millikan perform this
experiment? 1909

How did he view the oil droplets?
He viewed them with a microscope.

He did NOT measure the charge on the
electron; he calculated it. What did he
measure?
He knew the mass of the droplets and the strength of the
electric field.
Discovery of X-Rays

In 1895 William Conrad Roentgen
discovered X-rays, a form of radiation.
Radioactivity

In 1896, the French scientist
Henri Becquerel was studying a
Uranium mineral. He discovered
it was spontaneously emitting
high-energy radiation.

In 1898, Marie and Pierre Curie
attempted to isolate radioactive
components of the mineral.
Radioactivity

In 1899, Ernest Rutherford, a British
scientist, began to classify radiation: alpha
(a), beta (b), and gamma (g).
Radiation

Look closely at the paths of radiation. Do
you notice something about the amount of
deflection of each type of particles?
Radiation
Discovery of the Nucleus

In 1911, Ernest
Rutherford performed a
Gold Foil Experiment.

He and his colleagues
bombarded a thin piece
of gold foil with fast
moving, positively
charged alpha particles.
Alpha Particles
Alpha (a) particles are Helium-4 nuclei.
 This means they are two protons and two
neutrons (with no electrons).
 Thus, they are positive.

4
2
He
+2
Gold Foil Experiment
Gold Foil Experiment
As expected, most of the alpha particles
passed straight through with little or no
deflection.
 However, 1/8000 of the positively charged
alpha particles were deflected, some back
at the source.

(Po)
Gold Foil
Experiment
Gold Foil Experiment
From this experiment, Rutherford
discovered that there must be a very
densely packed positively charged bundle
of matter within the atom which caused the
deflections.
 He called this positive bundle the nucleus.
 He tried this experiment with other metals
and found the same results.

Gold Foil Experiment
The volume of the
nucleus was very small
compared to the volume
of the atom.
 Therefore, most of the
atom was composed of
empty space. Niels Bohr
later found that this empty
space was where the
electrons were located.

Checking for Understanding
Gold Foil Experiment

Why did some of the alpha particles come
straight back to the source or deflect away
from the nucleus?

Why did he conclude that the nucleus must
be positive?

What things did Rutherford conclude from
the gold foil experiment?
Checking for Understanding
Gold Foil Experiment

If gold atoms were solid spheres stacked
together with no space between them, what
would you expect would happen to particles shot
at them?

What year did Ernest Rutherford perform
this experiment?

Rutherford experimented with many kinds of
metal foil as the target. The results were always
similar. Why was it important to do this?
“It was about as believable
as if you had fired a 15inch shell at a piece of
tissue paper, and it came
back and hit you.”
-Ernest Rutherford
Discovery of the Neutron
In 1932, James Chadwick discovered the
neutron.
 Rutherford predicted that there were
massive, neutrally charged particles in the
nucleus, but it was Chadwick who proved
their existence.

Bohr’s Model of The Atom
Forces in the Nucleus
REMEMBER:
The nucleus is positive (p+ and no)
 Like charges repel each other…so
shouldn’t the p+ in the nucleus repel each
other?
 But…when 2 p+ are close together in the
nucleus there is a strong attraction
between them. The same holds true for
neutrons.

Forces in the Nucleus
no act like the “glue” that holds the nucleus
together. They help to stabilize the
nucleus.
 Nuclear forces are the short-range p+-no,
p+- p+ , and no-no forces that hold the
nuclear particles together.

Atomic Number
atomic number (Z) - the number of protons
in the nucleus of each atom of a given
element.
The number of p+ identifies the element.
Atomic Number increases from left to right
on the periodic table.
Electrons
The number of electrons in a neutral atom is
equal to the number of protons in that atom.
e- = p+
•Electrons can be lost or gained.
• When electrons are lost or gained, ions are
formed.
Ions
ion- an atom with a positive or negative
charge.
cation- an atom with a positive charge
 Cations are formed when an atom loses
negatively charged electrons.
 Ca+2 is formed when calcium loses 2
electrons.
 Ca+2 has 2 less electrons than protons.
the lithium
atom
the lithium
ion
Ions
anion- an atom with a negative charge
 Anions are formed when an atom gains
negatively charged electrons.
 N-3 is formed when nitrogen gains 3
electrons.
 N-3 has 3 more electrons than protons.
Noble gases
are very stable
and don’t
react.
Every
element on
the periodic
table will try to
react to be
stable, like the
noble gases.
Metals vs. Nonmetals

Metals form cations.
Na  Na+ + 1e-

Nonmetals form anions.
Cl + 1 e-  Cl-
Charge determination
(WITH SOME EXCEPTIONS!)
Group 1 – forms +1
 Group 2 – forms +2
 Group 13(B and Al) – forms +3
 Group 15 – forms -3
 Group 16 – forms -2
 Group 17 – forms -1
 Group 18 – doesn’t forms ions easily!

Mass Number

mass number (A)- the number of p+ & no in
the nucleus of an atom.
# of neutrons = mass number – atomic number
Why aren’t
electrons included
when determining
the mass number
of the atom?
Isotopes
isotope- two or more atoms having the same
atomic number (same #p+) , but different
mass numbers (due to different #no).
nuclide- general term for a specific isotope
of an element.
Isotopes
Isotope Notation
Nuclear Notation
Hyphen Notation
Uses the elements symbol followed by a hyphen & the mass
number.
C-12
How many protons, neutrons &
electrons are there in the following?
Cl-38
35Cl-1
Br-80
32S-2
N-14
56Fe+3
You Try It.
Do the Subatomic Particles Table
Worksheet.
Changes in the Nucleus
Nuclear Reaction- changes that occur in the
atom’s nucleus.

Nuclear reactions can change the
composition of an atom’s nucleus
permanently.
Types of Radiation Produced in
Nuclear Reactions

Alpha (a)

Beta (b)

Gamma (g)
Nuclear Stability
Atoms with unstable nuclei are radioactive.
 Most atoms have stable nuclei and are,
therefore, not radioactive.
Nuclear Stability
Neutrons help to stabilize the nucleus.
Elements 1-20 have  p+ = no
 Above element 20, increasingly more no
are needed than p+ to maintain nuclear
stability.
 Element 84 and up, all atoms are
radioactive so the nucleus cannot be
stabilized regardless of the number of no.

Types of Radioactive Decay
Alpha Radiation (a)- stream of high energy alpha particles.
 Consists of 2 protons & 2 neutrons making it identical to
a He-4 nucleus.
 Alpha particles can be represented by:
a
4
2

4
2
He
+2
4
2
He
Most alpha particles are able to travel only a few
centimeters through air and are easily stopped by
clothing etc.
Alpha Decay
239
94
Pu 
234
92
U
parent
235
92
4
2
U + He
230
90
4
2
Th + He
daughter
Types of Radioactive Decay
Beta Radiation (b) – consists of a stream of high
speed electrons. These electrons are not
electrons that are in motion around the atom’s
nucleus.
 Beta particles can be represented by:
0 -1
1
e

0
1
e
0
1
b
Can penetrate through clothing and damage
skin.
Beta Decay
6
2
He  Li + b
24
11
6
3
Na 
parent
24
12
0
-1
Mg + b
0
-1
daughter
Types of Radioactive Decay
Gamma Rays (g)- energetic form of light that
cannot be seen.
 Does not contain particles.
 Gamma particles can be represented by:
g
0
0

Can penetrate heavy material including skin.
Can only be stopped by lead or concrete.
Checking for Understanding
alpha
decay
beta
decay
238
210
84
Po  He +
14
6
a
U 

4
2
C
234
14
7
206
82
Pb
N+ b
0
-1
b
b
a
a
Th 
 234 Pa 
 234 U 
 230 Th 
?
226
88
Ra
Other Types of Nuclear
Reactions



0
1
positron –
proton neutron -
1
1
H
1
0
n
e
Half Life
half life- the time required for half of the atoms in
any given quantity of a radioactive isotope to
decay
Each particular isotope has its own half-life.
Half Life
p.689 Sample Problem B
Phosphorus-32 has a half-life of 14.3 days.
How many milligrams of phosphorus-32
remain after 57.2 days if you start with 4.0
mg of the isotope?
Ans: 0.25 mg
You Try It.
Do the Nuclear Reactions Worksheet.
The Mole
mole (mol)- SI Unit for
the amount of a
substance that
contains as many
particles as there are
atoms in exactly 12g
of carbon-12.
 A unit of counting, like
the dozen.
Avogadro’s Number
Avogadro’s Number - the number of
particles in exactly one mole of a pure
substance.
1 mole = 6.0221415 X 1023
1 mol = 6.02 x
23
10
Amedeo Avogadro
Atomic Mass
atomic mass - the mass of one mole of an
atom
 Atomic mass is expressed in atomic mass
units (amu) or (u) or g/mol.
 Can be found on the periodic table.
 All atomic masses are based on the
atomic mass of carbon-12 being 12 amu.
Molar Mass
molar mass - the mass of one
mole of a pure substance.
 Molar mass is written in units
of amu or g/mol.
Atomic mass vs. Molar mass
 atomic
mass - the mass of one
mole of an atom.
 molar mass - the mass of one
mole of a pure substance.
Atomic Mass vs. Molar Mass
Example
Atomic Mass
Na
22.99 g/mol
Ag
107.87 g/mol
C
12.01 g/mol
O
16.00 g/mol
Molar Mass of Compounds
Compound
H2O
C6H12O6
Molar Mass
18.02 g/mol
180.18 g/mol
NaCl
58.44 g/mol
Cl2
70.90 g/mol
(NH4)3PO4
149.12 g/mol
CuSO4·5H2O
249.72 g/mol
Introduction to Molar Conversions
Amount
Mass
1 mol O2
32.00 g
½ mol O2
16.00 g
2 mol O2
64.00 g
3 mol O2
96.00 g
Cheer
I say grams, you say molar mass.
grams – molar mass
grams – molar mass

Grams to Moles
Converting grams to moles: divide by molar mass.
1. How many moles of Ca are in 5.00g of Ca?
1 mol Ca
5.00g Ca x
= 0.125 mol Ca
40.08 g Ca
2. How many moles of H2O are in 36.0g of H2O?
1 mol H 2 O
36.0 g H 2 O x
= 2.00 mol H 2O
18.02 g H 2 O
3. How many moles of AgNO3 are in 124.5g of AgNO3?
1 mol AgNO3
0.7329
124.5 g AgNO3 x
=
169.88 g AgNO3
mol AgNO3
Moles to Grams
Converting from moles to grams: multiply by molar mass
1.
What is the mass in grams of 2.25 moles of Fe?
55.85 g Fe
2.25 mol Fe x
= 126 g Fe
1 mole Fe
2.
What is the mass in grams of 0.896 moles of BaCl2?
208.23 g BaCl2
0.896 mol BaCl2 x
=
1 mole BaCl2
187 g BaCl2
Types of Particles

Atoms – C, Cu, He

Molecules – O2, C12H22O11, CO2 (all
nonmetals in the formula)

Formula units – NaCl, CaCl2, Mg(NO3)2
(includes a metal in the formula)
1 mole = 6.02 x 1023 particles
Particles to Moles
Converting particles to moles: divide by
Avogadro's Number.
1.
How many moles of Pb are in 1.50 X
1025 atoms of Pb?
2.49 x 101 moles Pb
2.
How many moles of CO2 are in 6.78 X
1021 molecules of CO2?
1.13 x 10-2 moles CO2
Moles to Particles
Converting moles to atoms: multiply by
Avogadro's Number.
1.
How many molecules of NO are in 0.87
moles of NO?
5.2 x 1023 molecules NO
2.
How many formula units of NaI are in
2.50 moles of NaI?
1.51 x 1024 formula units NaI
Grams to Moles to Particles
Example: How many molecules of N2
are in 57.1g of N2?
23
1 mol N 2
6.02
x
10
molecules
N
2
57.1 g N2 x
x
28.02 g N 2
1 mol N 2
= 1.23 x 1024 molecules N 2
Particles to Moles to Grams
Example: How many grams of NaF are
in 7.89 X 1024 formula units of NaF?
24
7.89 x 10 f.un. NaF x
1 mol NaF
6.02 x 1023 f.un. NaF
41.99 g NaF
x
1 mol NaF
= 550. g NaF
Atoms to Moles to Grams
Tough Example:
How many total atoms are in 235 g of CO2?
9.64 x 1024 total atoms
The Mole Bridge
Atomic Mass Determination
Average Atomic Mass - the weighted
average of atomic masses of the naturally
occurring isotopes of an element.
 The atomic mass is expressed relative to
the value of exactly 12u for a carbon-12
atom.
 Atomic Mass Unit – amu or u
What is a weighted average?
Example: Your grade in math might be 75%
tests and 25% homework. What would your
grade be if you had a test average of 80% and
a homework average of 100%?
Normally, you would average 80% & 100%
to get 90%.
However, with a weighted average, it is 85%.
80%(.75)  100%(.25)  85%
Calculating Average Atomic Mass
Average atomic mass =
atomic mass of each isotope X percent
natural abundance (in decimal form)
average atomic mass =
 (atomic mass of each isotope x % abundance of each isotope in decimal form)
Percent Natural Abundance
Percent Natural Abundance- the
relative proportions expressed as
percentages, in which isotopes of
an element are found in nature.
Calculating Average Atomic Mass
Ex: Uranium has two isotopes, uranium-235 and
uranium-238. Uranium-235 has an atomic mass of
235.043amu and a percent natural abundance of
0.720%. Uranium-238 has an atomic mass of
238.050amu and a percent natural abundance of
99.280%. Calculate the average atomic mass of
naturally occurring uranium.