Ch. 3
Sizing up the Atom
 Elements are able to be subdivided into
smaller and smaller particles – these are
the atoms, and they still have properties
of that element
If you could line up 100,000,000
copper atoms in a single file, they
would be approximately 1 cm long
Despite their small size, individual
atoms are observable with instruments
such as scanning tunneling (electron)
microscopes
An STM image of nickel
atoms placed on a copper
surface.
Source: IBM Research
Red ridge is a series of Cesium atoms
Image of a ring of cobalt
atoms placed on a copper
surface.
Source: IBM Research
Atom - smallest particle making up
elements
One teaspoon of water has 3 times as
many atoms as the Atlantic Ocean has
teaspoons of water!
Think about the technological advances of
the past 100 years! They have been nothing
short of miraculous!
Radios
Calculators
Televisions
Computers
Automobiles
Cell phones
Jet airplanes
Ipods
Plastic
Velcro
Refrigerators
Internet (thanks, Al Gore)
Penicillin
CD’s & DVD’s
Insulin
and, of course Electric guitars
Sliced Bread!
Where did it all begin?
The word “atom”
comes from the
Greek word “atomos”
which means
indivisible.
The idea that all
matter is made up of
atoms was first
proposed by the
Greek philosopher
Democritus in the 5th
century B.C.
Then came the idea of “The 4 Basic Elements”
Earth, Air, Fire, & Water
After that came Alchemy.
The change to “real”
Chemistry didn’t occur
until the first true
element was discovered!
(1774)
The first element
discovered was
The discovery of oxygen is attributed to 3
scientists (working independently)
 Karl Scheele (1771) (German)
 first to prepare and describe oxygen
 Joseph Priestley (1774) (British)
 isolated oxygen gas from mercuric oxide.
 observed accelerated burning
 Antoine Lavoisier (1784) (French)
 made accurate measurements and
interpreted Priestley’s results

Carl Wilhelm Scheele beat Priestley to the
discovery but published afterwards.
Too bad! – So sad!
Priestley Medal
Source: Roald Hoffman, Cornell University
Priestley gets the
main credit for
discovering oxygen!
2HgO(s) → 2Hg(l) + O2 (g)

Priestley produced a gas (oxygen) by using sunlight to
heat mercuric oxide kept in a closed container. The
oxygen forced some of the mercury out of the jar as it
was produced, increasing the volume about five times.
Priestley: Scientific
Contributions
DISCOVERY OF 8 GASES
 Oxygen
 Nitrogen
 Carbon Dioxide
 Carbon Monoxide
 Sulfur Dioxide
 Nitrous Oxide
 Nitric Oxide
 Hydrogen Chloride
Priestley: Additional Scientific
Contributions
Discovered the interconnection between
photosynthesis and respiration
 Discovered carbonated water
 Discovered that India rubber removed
graphite pencil marks - the first rubber
eraser

Now we
can make
mistakes!!
Lavoisier: the Founder of Modern Chemistry
•Lavoisier continued the
investigations of Priestly
•Quantitative experiments led to:
Law of Conservation of Matter.
•He systematized the language of
chemistry, its nomenclature and
rhetoric.
Antoine-Laurent Lavoisier
•He was beheaded during the
Reign of Terror for his role as a
tax “farmer” prior to the
Revolution (Priestley escaped to
America!)
2Hg(l) + O2 (g) → 2HgO(s)

Lavoisier heated a measured amount of
mercury to form the red mercuric oxide. He
measured the amount of oxygen removed from
the jar and the amount of red oxide formed.
When the reaction was reversed, he found the
original amounts of mercury and oxygen.
Properties of Oxygen
P
P
P
P
P
C
C
Colorless
 Odorless
 Tasteless
 Gas at room temperature
 Slightly soluble in water
 Inflammable (does NOT burn)
 Only part of air that supports
combustion

Physical Property or Chemical Property?
These properties of oxygen were later used
to determine the properties of other
substances.
By the late 18th century, scientists finally came
to the conclusion that Oxygen was truly an
element (can’t be broken down into simpler
forms without losing its properties)
Scientists began to search for & test other
new elements.
Sometimes, when they tried to react
substances together, nothing happened!
Substances that DO NOT react are Inert
They found that most materials will react
to form new substances. These elements
are said to be chemically active (reactive)
Oxygen is very reactive, so is hydrogen
which we will look at next!
inert
Increasing chemical reactivity
Oxygen
hydrogen
Discovery of

Henry Cavendish (1766)
Reacted various metals with acids
producing a salt and hydrogen gas
Acid + metal → hydrogen gas + salt

Zinc + sulfuric acid → Hydrogen + zinc sulfate

Zn(s) + H2SO4(aq) → H2 (g) + ZnSO4 (aq)




(1731 – 1810)
Word
Equation
Chemical
equation
While testing the properties of Hydrogen he
found that water is a compound
Hydrogen + Oxygen
2H2 + O2
Water
2H2O
Antoine Lavoisier
Named Priestly’s newly discovered gas “oxygen” - meaning “acid former”
 Named Cavendish’s new gas “hydrogen” meaning “water former”

Dalton’s Atomic Theory
John Dalton (1766-1844)
While his theory was not completely
correct, it revolutionized how
chemists looked at matter and
brought about chemistry as we know
it today (instead of alchemy)
So, it’s an important landmark in
the history of science.
Dalton’s Modern Atomic Theory
(experiment based!)
1) All elements are composed of tiny indivisible
particles called atoms
2) Atoms of the same element are identical.
Atoms of any one element are different from
those of any other element.
3) Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds
4) In chemical reactions, atoms are combined,
separated, or rearranged – but never
changed into atoms of another element.
Law of Definite Proportions

Each compound has a specific ratio
of elements by mass.

Ex: Water is always 8 grams of oxygen
for each gram of hydrogen.
Discovery of the Electron
Began with the invention of the Crooke’s Tube
(cathode ray tube) c. 1875
Cathode Ray Tube
Voltage source
gas
+
Electric current sent through gases sealed
in tube at low pressure
 Anode- positive electrode
 Cathode- negative electrode

Metal Disks - electrodes
-
Modern Cathode Ray Tubes
Television
Computer Monitor
Cathode ray tubes pass electricity
through a gas that is contained at a
very low pressure.
In 1897, J.J. Thomson used a
cathode ray tube to study gases.
Thomson’s Experiment
Voltage source

+
Passing an electric current makes a
beam appear to move from the
negative to the positive end – so the
‘beam’ was called a “Cathode Ray”
Thomson’s Experiment
Voltage source
 Thomson found that cathode rays
were deflected from a negativelycharged plate.
Thomson’s Experiment
Voltage source
+


and that cathode rays were attracted to
plates with a positive charge
Does light bend like this?
Light doesn’t ‘bend’ so the cathode ray
must be made of particles rather than
Light!
Since they are attracted to a positive
plate & repelled by a negative one the
particles aren’t neutral –
What charge must they have?
That’s right! NEGATIVE!!
Thomson called these negative particles –
ELECTRONS
Mass of the Electron
Mass of the
electron is
9.11 x 10-28 g
The oil drop apparatus
1916 – Robert Millikan determined the mass
of the electron: 1/1840 the mass of a
hydrogen atom; and, 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. Therefore, all elements
must contain identically charged
electrons.
b) Atoms are neutral, so there must be a
positive substance 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 their mass
Thomson’s Atomic Model
J. J. Thomson
Thomson believed that the electrons
were like plums embedded in a
positively charged “pudding,” thus it
was called the “plum pudding” model.
In 1903, An important
discovery leading to
further understandings
of atomic structure
happened by accident.
Henri Becquerel
discovered radioactivity
Radioactivity is the spontaneous emission
of energy from an object
1903: Shared a Nobel Prize with Pierre and
Marie Curie for discovering radioactivity.
Ernest Rutherford (1871-1937)
The Nobel Prize in Chemistry 1908
Studied under J. J. Thomson
3 Types of Radiation discovered by
Ernest Rutherford
4
• Alpha (ά) – a positively charged
2
4
+2
He
helium nucleus
2 He
•Beta (β) – fast-moving0 electrons
-e
1 e
•Gamma (γ) – like high-energy
0
0
x-rays

Ernest Rutherford’s
Gold Foil Experiment - 1911
Shot alpha particles at a thin sheet of
gold foil
 Particles that hit on a detecting
screen (film) were recorded
Lead
block
Polonium
Flourescent
Screen
Gold Foil
He Expected:
 The
alpha particles to pass through
the foil without changing direction
very much.
 Because…
 The positive charges were spread
out evenly (according to Thomson’s
atomic theory). Alone they were not
enough to stop the alpha particles.
What he expected
Again, because he thought the mass
was evenly distributed in the atom
What he got
“Like howitzer
shells bouncing
Rutherford’s
Observations
off of tissue paper!”
Most of the particles went straight
through the foil (what he expected)
 A few particles were slightly deflected
 Still fewer actually bounced back
towards the source!
 Astonishing!!!
 Rutherford said it was like firing a
Howitzer shell at a piece of tissue paper
& having it bounce back & hit you!

+
Rutherford’s Conclusions




Since most of the particles went through the
foil - atoms are mostly empty space.
Because a few + particles were deflected
+ close to a positively
they must have come
charged core.
Since a very few particles were deflected
straight back, the positively-charged core
must be very dense.
This small dense positive area is the nucleus.
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 electrons are distributed around the
nucleus, and occupy most of the volume
 His model was called a “nuclear model”
Discovery of Protons
 Eugen Goldstein in 1886 observed
particles with a positive charge
passing through a perforated
cathode.
In 1920, Rutherford studied these
particles & called them protons.
They have a charge of positive 1
and a mass of 1.7 x 10-24 grams.
This is not a ‘handy’ number to work
with so we use a mass of 1 amu.
Amu stands for “atomic mass unit”
Discovery of the Neutron
Rutherford predicted the existence of
the neutron in 1920.
Twelve years later, his assistant found it!
1932 – James Chadwick confirmed
the existence of the “neutron”
– a particle with no charge,
but a mass nearly equal to a proton
(1 amu).
So now we have a more complete picture of an atom!
Subatomic
Particles
Particle
Electron
(e-)
Proton
(p+)
(H+)
Neutron
(no)
Charge
-1
+1
0
Mass (g)
9.11 x 10-28g
(virtually 0)
Location
outside
nucleus
1 amu
in nucleus
(1.7 x 10-24g)
1 amu
(1.67 x 10-24g) in nucleus
Elements are the new building blocks
Nitrogen-7
Hydrogen
Carbon-6
Oxygen-8
Henry Moseley
(1887 – 1915)
Between 1912 and 1914, the
physicist H.G.J. Moseley
conducted a series of experiments
where he bombarded targets
made out of different kinds of
metals with cathode rays. Each
metal he studied emitted X-rays
of a characteristic frequency,
almost like a set of "fingerprints".
The pattern that emerged when the
observed X-rays were organized in
order of increasing frequency
suggested to Moseley a regular
increase in the positive charge on
the nuclei of the atoms.
He called this positive nuclear chargethe Atomic Number of the element
Atomic Number
Henry Moseley – used x-ray spectra
& came up with the idea of the
Atomic Number
 Elements
are different because they
contain different numbers of PROTONS
 The “atomic number” of an element is
the number of protons in the nucleus
 Since
all atoms are neutral - the
# protons in an atom = # electrons
Atomic Number, Z
All atoms of the same element
have the same number of
protons in the nucleus, Z
13
Al
26.981
Atomic number
Atom symbol
AVERAGE Atomic Mass
Mass Number
Mass number is the number of protons
and neutrons in the nucleus of an
isotope:
Mass # = # protons + # neutrons
Subatomic Particles
ATOM
ATOM
NUCLEUS
NUCLEUS
ELECTRONS
ELECTRONS
PROTONS
PROTONS
NEUTRONS
NEUTRONS
POSITIVE
CHARGE
NEUTRAL
CHARGE
NEGATIVE
CHARGE
NEGATIVE CHARGE
equal in a
Atomic
Most Number
of the atom’s mass.
neutral atom
equals the # of...
Isotopes
 Frederick
Soddy (1877-1956)
proposed the idea of isotopes in
1912 (worked with Rutherford)
 Isotopes
are atoms of the same
element having different mass
numbers, due to varying numbers of
neutrons.

Soddy won the Nobel Prize in Chemistry in
1921 for his work with isotopes and
radioactive materials.
Isotopes
 Atoms
of the same element
(same Z) but different mass
number (A).
 Boron-10 (B-10) has 5 p and 5 n
Boron-11 (B-11) has 5 p and 6 n
11B
10B
Isotopes


Radioisotopes (radioactive isotopes) unstable isotopes that spontaneously decay
emitting radiation
They play an important part in the
technologies that provide us with food, water
and good health.
 Radio-carbon dating of fossils
 In medicine, diagnosis, treatment, and
research
 Sterilization of meat
Disinfestation of grain and spices
Increasing shelf life (eg, fruits)
Nuclear Symbols
 Contain
the symbol of the element,
the mass number and the atomic
number (represent isotopes of
elements)
Mass
Superscript →
number
Element
symbol
Atomic
Subscript →
number
X
REMEMBER! number of electrons = number of protons
So all atoms are neutral!
Rhenium
186
75
Re
Protons: 75
Neutrons: 111
Electrons: 75
Nuclear Symbols

Find each of these:
a) number of protons
b) number of neutrons
c) number of electrons
d) Atomic number
e) Mass Number
80
Br
35
Nuclear Symbols

If an element has an atomic
number of 34 and a mass
number of 78, what is the:
a) number of protons 34 78
Se
b) number of neutrons 44 34
c) number of electrons 34
d) Write the complete symbol
Naming Isotopes
We
can name isotopes by
placing the mass number
after the name of the
element:
carbon-12
carbon-14
uranium-235
Mass numbers
ISOTOPES
Isotope
Oxygen - 18
Arsenic
- 75
Phosphorus - 31
p+
n0
e- Mass #
8
10
8
18
33
42
33
75
15
16
15
31
The element hydrogen has 3 isotopes
Isotope
Hydrogen–1
(protium)
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
Protons Electrons
Neutrons
1
1
0
1
1
1
1
1
2
Nucleus
Examples of Isotopes
Learning Check – Counting
Naturally occurring carbon consists of three
isotopes, 12C, 13C, and 14C. State the number
of protons, neutrons, and electrons in each of
these carbon atoms.
12C
6
13C
6
14C
6
#p+ _______
_______
_______
#no _______
_______
_______
#e- _______
_______
Answers
12C
13C
14C
6
6
6
#p+
6
6
6
#no
6
7
8
#e-
6
6
6
Learning Check
An atom has 14 protons and 20 neutrons.
A. Its atomic number is
1) 14
2) 16
3) 34
B. Its mass number is
1) 14
2) 16
3) 34
C. The element is
1) Si
2) Ca
3) Se
D. Another isotope of this element is
1) 34X
2) 34X
3) 36X
16
14
14
Atomic Mass



How heavy is an atom of oxygen?
 It depends, because there are different
kinds of oxygen atoms.
We are more concerned with the
average atomic mass.
This is based on the abundance
(percentage) of each variety (isotope)
of that element in nature.

We don’t use grams for this mass because
the numbers would be too small –
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 mass
number, so we determine the
average atomic mass from the
element’s percent abundance.
To calculate the average atomic mass:
 Multiply
the mass of each
isotope by it’s abundance, then
add the results.
 Abundance may be expressed as
a decimal or a %, (Divide by 100
if using %’s)
Avg.
Atomic
Mass

(mass)(% )  (mass )(% )

100
Atomic Mass is the weighted average of all
the naturally occurring isotopes of an
element. on the Periodic Table
Isotope
Symbol
Carbon-12
C-12
Carbon-13
C-13
Carbon-14
C-14
Composition of
the nucleus
6 protons
6 neutrons
6 protons
7 neutrons
6 protons
8 neutrons
(98.89 x 12) + (1.11 x 13) + (0.01 x 14)
100
% in nature
98.89%
1.11%
<0.01%
= 12.011
D. Average Atomic Mass

EX: Calculate the avg. atomic mass of oxygen if
its abundance in nature is 99.76% 16O, 0.04%
17O, and 0.20% 18O.
Avg.
(16)(99.76 )  (17)(0.04)  (18)(0.20)
 16.00
Atomic 
100
amu
Mass
Sub-atomic Particles - Summary
Protons p+ positive charge, in
nucleus , mass of 1
amu
Neutrons n0 – no
charge, in nucleus,
mass of 1 amu
Electrons - e- negative
charge, orbiting nucleus, “no mass”