Unit 2
Dalton’s Atomic Theory

It was only 200 years ago that John
Dalton proposed his atomic theory

Several postulates to this accepted
theory
Dalton’s Atomic Theory
1.
2.
3.
4.
All matter is made of atoms. Atoms are
indivisible and indestructible.
All atoms of a given element are
identical in mass and properties
Compounds are formed by a
combination of two or more different
kinds of atoms.
A chemical reaction is a
rearrangement of atoms.
Atoms are TINY

Atoms, particles on the atomic scale, are
almost unconceivably small

Think about picking up ONE grain of
sand from the beach. There are more
atoms in that one grain than can be
counted in a lifetime
Building Blocks

Despite this size, atoms are the main
building blocks that make up everything

These atoms, which make up the
elements, are the bridging unit that
connects the microscopic and the
macroscopic world
Early Ideas…

Leucippus and his student Democritus
were the first people to propose matter
was composed of small, indivisible
particles

5th century B.C.

Called these particles atomos
Critics

Democritus wrote “nothing exists except
atoms and empty space; everything else is
opinion”

Some disagreed with this idea…
 Plato and Aristotle said this was incorrect.
 They thought there was no “tiniest” indivisible
particle and that everything was ultimately made
up of the fire, air, earth, and water

Because Plato and Aristotle had such
influential contributions, this idea
prevailed

You could not easy
test either ideas
so any actual
intelligent thought
took a back seat
for nearly
2,000 years
Then we move forward…

In the sixteenth century modern science
began to emerge

Nicolaus Copernicus proposed the sun,
not the earth, was the center of the
universe… the turning point in the
scientific revolution

This scientific revolution led to the work of
Francis Bacon (1561-1626)
Johannes Kepler (1571-1630)
Galileo Galilei (1564-1642)
Robert Boyle (1627-1691)
Isaac Newton (1642-1727)
And then…
John Dalton(1766-1844) – which offered
convincing evidence of the early atomic ideas of
Leucippus and Democritus
The Law of Conservation of Mass

In 1789, Antoine Lavoiser formulated the
Law of Conservation of Mass.. Which
states the following:
In a chemical reaction, matter is
neither created or destroyed.

In other words, the total mass of the
substances involved in the reaction does
not change.
Concept Check

When a small log completely burns in a
campfire, the mass of the ash is much
less than the mass of the log. What
happened to the original matter that
composed the log??
The Law of Definite Proportions

1797. A French scientist named Joseph
Proust made an observation on the
composition of compounds

Elements that made up a compound are
found in fixed (or definite) proportions in
all samples
The Law of Definite Proportions

States:
All samples of a given compound,
regardless of their source or how
they were prepared, have the
same proportions of their
constituent elements.
Example

The decomposition of 18.0 grams of water
results in 16.0 grams of oxygen and 2.0
grams of hydrogen, or an oxygen-tohydrogen mass ratio of:
16.0 𝑔 𝑂
𝑀𝑎𝑠𝑠 𝑅𝑎𝑡𝑖𝑜 =
= 8.0 𝑜𝑟 8: 1
2.0 𝑔 𝐻

This ratio holds true for any sample of
water, regardless of its origin
Practice

Consider Ammonia

Ammonia consists of 14.0 grams of
nitrogen for every 3.0 grams of
hydrogen. What is the ratio?
14.0 𝑔 𝑁
𝑀𝑎𝑠𝑠 𝑅𝑎𝑡𝑖𝑜 =
= 4.7 𝑜𝑟 4.7: 1
3.0 𝑔 𝐻
The Law of Multiple Proportions

In 1804, John Dalton published his law of
multiple proportions:
When two elements (call them A and B)
form two different compounds, the
masses of element B that combine with 1
g of element A can be expressed as a ratio
of small whole numbers
i.e. AB1, AB2, AB3 …

Think about carbon monoxide and
carbon dioxide..
CO

CO2
Two compounds with the same
elements but in different proportions or
ratios
Subatomic Particles

By the 19th century, scientists accepted
the idea of the atom and it being the
fundamental unit of matter

However, they then realized it too could
be broken up into smaller particles
 “subatomic particles”
The Discovery of the Electron

J.J. Thompson (1856-1940)

Worked with a Cathode Ray Tube –
determined that the traveling particles
carried a negative charge

Discovered the negatively charged, low
mass particle (the electron)
The Structure of the Atom

The discovery of the (-) charged
electrons brought new questions… the
atom is overall neutral… so something
must carry a (+) charge to cancel out

J.J. Thompson proposed his “PlumPudding Model”
The Plum Pudding Model

Proposes the negatively charged
electrons are held within a positively
charged sphere (the nucleus)
Rutherford’s Gold Foil Experiment

Tried to prove J.J.’s Plum Pudding
model… disproved it instead

Shot alpha particles at a very small
piece of gold foil.. Very, very thin piece

Instead of passing straight through,
particles were deflected and even
bounced right back at him
Rutherford

Concluded that the atom was much
smaller than the plum pudding implied

Rutherford proposed the “Nuclear
Theory” of the atom
The Nuclear Theory
1.
Most of the atom’s mass and all (+) charge
is in a small core; the nucleus
2.
Most of the volume of the atom is empty
space with tiny negative electrons
dispersed
3.
There are as many negative charged
electrons outside the nucleus as there are
positive charged protons inside
Rutherford’s Model

Rutherford’s Model was great but still
incomplete.

DID NOT account for neutrons yet

They knew this because there was missing
mass (mass unaccounted for) in the atoms
with considering the protons and mass
number
James Chadwick

Rutherford and his student, James
Chadwick, later determined the missing
mass was due to the neutrons

Neutrons are neutrally charged particles
similar in size to the protons

They are literally there to stabilize the
atom – keep the protons (+) apart
Research Project



Read section 2.1 – 2.6
Select a contributor to the atomic theory progression
Use your textbooks, notebooks, and the internet to supplement your
research

Must include:
1.
2.
3.
4.
5.
A brief biography of the individual you’re researching
The contributions he/she made to the atomic theory and why they were
important/helped. Include who they worked with or the work they built on
Timeframe/timeline of their work
A graphic depiction of the experiment/work he or she performed (must be
clearly and articulately labeled and explained)
Any flaws, criticisms, or necessary improvements to their ideas/work that were
later required
Specifics: you may choose any form of media you desire. You and
ONE partner will present this work to the class. Presentations must be a
minimum of 4 minutes and must be accompanied with a 2-3 page
research paper addressing each of the 5 points above. 12 point font,
double-spaced, Times New Roman, 1-inche margins. Papers will be
graded based on grammar, correctness, and completeness.
Subatomic Particles

Protons and neutrons have nearly
identical masses

However, the electrons are much
smaller – Negligible

If the protons and neutrons had the
mass of a baseball then the electrons
would be the size of 1 grain of rice
Elements

Elements are defined by their number of
protons

The Atomic Number = the number of
protons … aka their social security
number

The mass number = #N + #P
 No electrons !!! Why not???
Isotopes

“When the number of neutrons varies”

All elements of the same atom must have
the same number of __________.

We can, however, vary in the number of
neutrons

Same atoms, different neutrons = isotopes
Masses

Elements may have 2 different isotopes
or 3… or more

Ex: All neon atoms have 10 protons but
there are three different isotopes: one
has 10 neutrons, one has 11, and one
has 12 – which gives them all three
different atomic masses
Natural Abundance

The mass number (a decimal value from
the table) is not a whole number because it
is an average of all occurring isotopes

Neon’s 3 isotopes are:
 90.48% Ne-10
 0.27% Ne-11
 9.25% Ne-12
**Take the average
of all three of these
to determine the
average atomic mass
Symbol
No. of
Protons
No. of
Neutrons
A (mass
number)
Natural
Abundance
(%)
Ne-20
or
20 Ne
10
10
10
20
90.48
Ne-21
or
21 Ne
10
10
11
21
0.27
Ne-22
or
22 Ne
10
10
12
22
9.25
Practice

What are the atomic number, mass
number, and symbol of a chlorine
isotope with 18 neutrons?

How many protons, electrons, and
neutrons are present in an atom of
52 Cr ?
24
Most Elements are Mixtures of
Isotopes

The periodic table reports Average
Atomic Mass, a weighted average of
the atomic mass of an element’s
isotopes.

Example: If I said there was a larger
abundance of C-12 than C-13, which
value would the average atomic mass
be closer to?

If you know the abundance of each
isotope, you can calculate the average
atomic mass of an element.
Calculate the Average Atomic
Mass
90.48% Ne-20
0.27% Ne-21
9.25% Ne-22
10p + 10N = 20AMU
0.9048
20 x 0.9048 =
18.096
10p + 11N = 21AMU
0.0027
21 x 0.0027 =
0.0567
+
10p + 12N = 22AMU
0.0925
22 x 0.0925 =
2.035
+
= 20.1877 g/mol
 What is the mass of each??
 What is the decimal of each percent?
 Multiply the mass by the abundance and add it all up
Calculating Average Atomic Mass

The mass of a Cu-63 atom is 62.94
amu, and that of a Cu-65 atom is 64.93
amu. Abundance of Cu-63 is 69.17%
and Cu-65 is 30.83%. Find the average
atomic mass of Cu.
More Practice

Calculate the average atomic mass for
gallium if 60.00% of its atoms have a
mass of 68.926 amu and 40.00% have a
mass of 70.925 amu.

Calculate the average atomic mass of
oxygen. Its composition is 99.76% of
atoms with a mass of 15.99 amu,
0.038% with a mass of 17.00 amu, and
0.20% with a mass of 18.00 amu.
How do we calculate these
masses??

Mass Spectrometry – instrument used to
measure the masses of atoms and the
percent abundances of isotopes of
elements

Atoms are converted to positively
charged ions, accelerated, and passed
through a magnetic field that deflects
their path. The heaviest ions undergo
the least deflection
Ions (losing or gaining e-)

If the atom is neutral, the #P = #e-

But, if it is not neutral, than the number
of electrons has been changed…
increased (gained) or decreased (lost)
Li  Li+ + 1 eF + 1 e-  F -
Ions

If an atom gains electrons, than it will be
___________ charged
 We can this an anion

If an atoms loses electrons, than it will
be _____________ charged
 We call this a cation
Summary

You can change the number of electrons
(ions)

YOU CANNOT change the number of
protons (SS#)

You CAN change the number of
neutrons – we make isotopes
The Periodic Table

Dmitri Mendeleev (1834 – 1907)
 Worked on developing the Table
 Only 65 different elements at this time
 Knew the elements’ relative masses,
chemical activity, and many physical
properties through experimentation … but
there was not real way of organizing them
The Periodic Table

Dmitri Mendeleev (1834 – 1907)
 Worked on developing the Table
 Only 65 different elements at this time
 Knew the elements’ relative masses,
chemical activity, and many physical
properties through experimentation … but
there was not real way of organizing them
Periodic Law

“When the elements are arranged in
order of increasing mass, certain sets of
properties occur periodically”
Periodic Law

Mendeleev summarized these
observations and arranged them in a
way that elements with similar
characteristics are in the same group /
column
Predictions

Since not all the elements had been
discovered yet, Mendeleev’s Table had
some gaps in it

This also led to predictions of the
existence and behaviors of unknown
elements – found to hold truth as they
were discovered
Periodic Tables

Color code your periodic tables
illustrating the following terms:
Arrangement









Metals
Nonmetals
Metalloids
Main-group elements (define)
Transition elements (transition metals)
Noble gases
Alkali metals
Alkaline earth metals
Halogens
Ions and the Periodic Law

A main-group metal tends to lose
electrons, forming a cation with the
same number of electrons as the
nearest noble gas

A main-group nonmetal tends to gain
electrons, forming an anion with the
number of electrons as the nearest
noble gas
Practice

Predict the charges of the monoatomic
ions formed by the following
Al
S
O
Li
Mg
N
Molar Mass

A molar mass is the mass of one mole

The mole (mol) is a “chemist’s dozen”
 A unit to quantify or count atoms, molecules,
compounds, ions …. Little things
1 mol (of anything) = 6.022x1023
Avogadro’s Number

Named after Amedeo Avogadro

The numerical value of the mole is
defined as being equal to the number of
atoms in exactly 12 grams of pure
carbon-12
(12 grams of C = 1 mol of C atoms =
6.022x1023 C atoms)
(12 grams of C = 1 mol of C atoms =
6.022x1023 C atoms)

A CONVERSION FACTOR!!!
Conversions
1 mol atoms
6.022x1023
or
6.022x1023
1 mol atoms
Use these conversions just as you would use
1 foot = 12 inches
And
12 inches = 1 foot
Practice

How many atoms are in 2.5 mol of Cu?

How many water molecules are in 3.25
mol of water?
Practice

How many moles are in 4.22x1026
sodium atoms

How many moles are in 5.2x1022 carbon
dioxide molecules?
Molar Mass

“the mass of one mole”

The mass numbers on the table are the
mass of one mole of that substance
meaning the mass of 6.022x1023 atoms
Sodium = 22.98977 g/mol
 Hydrogen = 1.00794 g/mol
 Nitrogen = 14.0067 g/mol

Conversions!!
Sodium = 22.98977 g/mol
 Hydrogen = 1.00794 g/mol
 Nitrogen = 14.0067 g/mol


Look at those units…. What can we do
with them??
Practice

How many moles are in 17.45 grams V?

Convert: 104.5 grams of Ag to mol Ag

Convert: 0.098 mol Na to grams
Multi STEP

Convert:
17.88 grams Ti to atoms of Ti
1.009 grams Na + to Na+ ions
5.99x1024 Ne atoms to grams
Challenge yourselves
An Al sphere contains 8.55x1022 aluminum atoms.
What is the radius of the sphere in centimeters? The
density of aluminum is 2.70g/cm3
Volume of a sphere:
4
𝑉 = 𝜋 𝑟3
3
Tip: Use your density equation like another
conversion factor… remember, if the units work, its
WORKS 
Practice with a buddy

A Ti cube contains 2.86x1023 atoms. What is the edge
length of the cube? The density of titanium is
4.50g/cm3 (V = LxWxH = Side3 (b/c it’s a cube))

Find the number of atoms in a copper rod with a length
of 9.85 cm and a radius of 1.05 cm. The density of
3
2
copper if 8.96 g/cm (V = 𝜋 r h)