Atomos: Not to Be Cut
The History of Atomic Theory
Atomic Models

A model uses familiar
ideas to explain
unfamiliar facts observed
in nature.

A model can be changed
as new information is
collected.
 The
atomic
model has
changed
throughout the
centuries,
starting in 400
BC, when it
looked like a
billiard ball →
Democritus

This is the Greek
philosopher Democritus -

His theory: Matter could
not be divided into
smaller and smaller
pieces forever, eventually
the smallest possible
piece would be obtained.
400 BC
Atomos


This piece would be
indivisible.
He named the smallest
piece of matter “atomos,”
meaning “not to be cut.”
This theory was ignored and
forgotten for more than 2000
years!
Why?

The eminent
philosophers of
the time,
Aristotle and
Plato, had a
more respected,
(and ultimately
wrong) theory.
Aristotle and Plato favored the earth, fire, air
and water approach to the nature of matter.
Their ideas held sway because of their
eminence as philosophers. The atomos idea
was buried for approximately 2000 years.
Dalton’s Model

In the early 1800s,
the English
Chemist John
Dalton performed a
number of
experiments that
eventually led to
the acceptance of
the idea of atoms.
Dalton’s Theory




All elements are made of tiny,
indivisible particles called atoms.
All atoms of the same element are
identical. All atoms of different
elements are different (unique).
Atoms combine with different
atoms to form compounds. Atoms
unite in whole number ratios
(H2O, CO2).
Atoms are not created or
destroyed, simply rearranged in
chemical reactions.
.
 This
theory
became one
of the
foundations
of modern
chemistry.
Thomson’s Plum Pudding
Model
 In
1897, the
English scientist
J.J. Thomson
provided the first
hint that an atom
is made of even
smaller particles.
Thomson Model
 Thomson
studied
the passage of an
electric current
through a gas.
 As the current
passed through
the gas, it gave off
rays of negatively
charged particles.
Thomson Model

He proposed -“Plum
Pudding” model.

Atoms were made
from a positively
charged substance
with negatively
charged electrons
scattered about, like
raisins in a pudding.
Thomson Model
 This
surprised
Thomson,
because the
atoms of the gas
were uncharged.
Where had the
negative charges
come from?
Where did
they come
from?
Rutherford’s Gold Foil
Experiment

In 1908, the English physicist Ernest
Rutherford.

Rutherford’s experiment Involved firing a
stream of tiny positively charged particles
at a thin sheet of gold foil (2000 atoms
thick)
Gold Foil Experiment
Gold Foil Experiment


Most of the positively
charged “bullets” passed
right through the gold
atoms in the sheet of
gold foil without changing
course at all.
Some of the positively
charged “bullets,”
however, did bounce
away from the gold sheet
as if they had hit
something solid. He
knew that positive
charges repel positive
charges.

This could only mean atoms were mostly open
space.

Rutherford concluded that an atom had a small,
dense, positively charged center that repelled
his positively charged “bullets.”

He called the center of the atom the “nucleus”

The nucleus is tiny compared to the atom as a
whole.
Bohr Model
 In
1913, the
Danish scientist
Niels Bohr
proposed that
electrons were in
a specific energy
level.
Bohr Model

electrons move in
definite orbits around
the nucleus, much like
planets circle the sun.

These orbits, or energy
levels, are located at
certain distances from
the nucleus.
Wave Model
The Wave Model

Today’s atomic model is based on the
principles of wave mechanics.

electrons do not move about an atom in a
definite path, like the planets around the
sun.

In fact, it is impossible to determine the
exact location of an electron. The probable
location of an electron is based on how
much energy the electron has.
The Wave Model

According to the modern atomic model, at atom
has a small positively charged nucleus
surrounded by a large region (electron cloud) in
which there are enough electrons to make an
atom neutral.
Electron Cloud:

Depending on their energy they are locked into a
certain area in the cloud.

Electrons with the lowest energy are found in
the energy level closest to the nucleus

Electrons with the highest energy are found
in the outermost energy levels, farther from
the nucleus.
The Atom
CHAPTER 4, PART I

http://www.gpb.org/chemistryphysics/chemistry/302
Relative Sizes
What is the relative distance from the
nucleus of an atom to the electrons? If the
nucleus of an atom were the size of a
marble how far away would the first
electron be?
Answer
Subatomic Particles
Proton
 Neutron
 Electron

FYI:

Quarks
 Up
/ Down
 Charm / Strange
 Top / Bottom
Proton
Symbol = P+
 Charge = Positive
 Mass = 1 atomic mass unit (amu)
 Location = In the nucleus

Neutron
Symbol = n0
 Charge = Neutral
 Mass = 1 atomic mass unit (amu)
 Location = In the nucleus

Electron
Symbol = e Charge = Negative
 Mass = 0 atomic mass unit (amu)
 Location = Orbiting the nucleus

Atomic Number, (Z) = # of P+ in an
atom’s nucleus
 Mass Number, (A) = Sum of the # of N0
and # of P+ in a given nucleus.
A
Z
A = mass #, Z = atomic #, X = symbol

X
Atomic Calculations
1) Atomic Number = Number of Protons
Z = Atomic Number = # p+
2) If neutral, then protons must equal
electrons
3) Atomic Mass = Number of Protons +
Neutrons
A = Atomic Mass
A = Z + n0
Atomic Calculations Game
http://education.jlab.org/elementmath/index.
html

Isotopes – atoms of the same element
with the same # of P+, but different # of N0
and different mass numbers.


Nuclide – a particular atom containing a
definite number of protons and neutrons.


Ex.
Ex. Carbon-12 Iron-56
Nucleons – particles that make up the
atomic nucleus

Ex. protons and neutrons.
Atoms vs Ions
C
B-1
C+1
Ion – an atom that has lost or gained
electrons and now has a charge
Number of protons = Z (atomic number)
Number of neutrons = A – Z (atomic mass –
atomic number)
Number of electrons = Number of
protons - charge
Average Atomic Mass
AMU – atomic mass unit, 1/12 the mass of
a carbon-12 atom.
 Mass on periodic table is a weighted
average mass of all isotopes of an
element.

Example
Neon-12 relative abundance – 98%
Carbon-14, relative abundance – 2%
0.98 (12 amu) + 0.02 (14 amu) = 12.04 amu
Example 2

What is the atomic mass of silicon if
92.21% of its atoms have mass of 27.977
amu, 4.70% have mass of 28.976 amu,
and 3.09% have a mass of 29.974 amu?

28.09 amu