•
Young people should not smoke.
•
Studies show that smoking at an early
age may make it more difficult to quit
smoking later.
•
Which of the above statements is an
opinion and which is a theory?
•
What is the difference?
Atomic Theory
What is it?
 Who figured it out?
 When did they do it?
 How did they do it?
 Why do we believe it?

What can you tell from this picture?
What is an ATOM?

An atom is the smallest particle of an
element that retains the identity of the
element.
ZZZ
ZZZ
ZZZ
ATOM
Atoms combine to form compounds
The Ancient Greeks

Democritus



Lived 450 B.C.
Proposed that all the
stuff in the world is
“atomos”
Tiny, indivisible
particles

Aristotle


Questioned theory of
Democritus
Rejected it for lack of
proof
Neither offered PROOF, but Aristotle
was respected as the smartest guy
around…the “Einstein” of his time! So
Democritus’ idea was ignored!
ALCHEMISTS




Tried to turn base
metals into precious
metals.
Developed
knowledge and
techniques…very
smart and knew a lot!
Not true scientists
Didn’t share!
Roger Bacon
Lived in 13th century
 Believed that science should be based on
experimental evidence.

Antoine Lavoisier




1743-1794
French
Father of Chemistry
Sadly, beheaded in French Revolution
Lavoisier

LAW OF CONSERVATION OF MATTER


Matter, like energy, is neither created nor
destroyed in a chemical reaction.
This concept established modern chemistry.
Law of Conservation of Mass
What was different?

Experimentation

He used a balance to
study the role of
oxygen in rusting and
burning.
Other people did it, too!



Priestly did similar
experiments.
He believed a false
theory of Phlogiston.
Kept his mind closed
to new idea.
PROUST

HYDROGEN
11%
OXYGEN
89%
1799
 Law
of
Constant
Composition

A given compound
always contains
the same elements
in the same
proportion by
mass.
Law of Multiple Proportions
(stated by Dalton)
The ratios of the masses of
elements in compounds will be
ratios of small whole
numbers…multiple proportions.
JOHN DALTON
LAVOISIER
PROUST
ATOMIC THEORY OF MATTER
DEMOCRITUS
BACON
DALTON’S ATOMIC THEORY OF
MATTER
1.
All matter is composed of
submicroscopic (extremely small) indivisible
particles called ATOMS.
2.
All atoms of a given element are
identical. The atoms of any one element are
different from those of any other element.
DALTON’S ATOMIC THEORY OF
MATTER
3. Atoms are neither created nor destroyed in a
chemical reaction. Chemical reactions occur
when atoms are separated, joined, or
rearranged. However, atoms of one element are
never changed into atoms of another element
as a result of a chemical reaction.
4. Atoms of different elements can mix
physically or can combine chemically with one
another in simple whole number ratios to form
compounds.
Dalton’s Theory

Based on
CAREFUL EXPERIMENTATION
DALTON’S MODEL OF THE ATOM
1803
Modern Atomic Theory

Not all aspects of Dalton’s atomic theory
have proven to be correct. We now know
that:
--Atoms are divisible into even smaller particles.
--A given element can have atoms with different
masses.
Some important concepts remain unchanged.
--All matter is composed of atoms.
--Atoms of any one element differ in properties
from atoms of another element.
What does this mean?
Not perfect...
but a workable theory to build on.
Atomic Theory of Matter





Scientists had been doing many
experiments with electricity since
Ben Franklin flew his kite.
Faraday suggested that electricity
might explain the atom
English physicist, J.J. Thomson
1856-1940
Discovered electrons in 1897
Thomson experimented with a
“cathode ray tube”
Thomson’s Cathode Ray Tube
Experiment

Thomson used
the atom
ELECTRICITY To probe
Cathode Ray Tube
Flow of electric current through gases.
Sealed gas in glass tube with metal plates at
the end.
Connected plates to high voltage source:
Cathode –
and
Anode +
CRT
+
This ray could be deflected
toward a positive
charge…
It has a negative charge.
This ray could move
things…
It was made of particles of
matter.
This ray acted the same no
matter what materials
were used…
It was not atoms…it must be
part of all atoms
Cathode Ray Tube
Cathode ray is composed of very small
negatively charged particles that are part of
atoms
ELECTRONS
Thomson’s Plum Pudding Model
1896
Do you like
Chocolate
chip cookies
better?
Chocolate
Chip Cookie
Dough
Millikan’s Oil Drop Experiment
1909
Determined the
charge of the
electron which
along with
Thomson’s
experiment
determined the
mass of the
electron
Atomic Theory
Millikan devised experiments to determine
the mass and charge of the electron.
 Protons : discovered in 1886
Positively charged particles
Also discovered with cathode ray
tube…these particles went the other
direction.

The Nuclear Atom

What’s missing?
The Nucleus!
 Rutherford did an experiment using
RADIOACTIVITY as a tool to probe the
atom
 Identified three types of radiation

 Alpha…positive
particle
 Beta…negative particle
 Gamma…high energy

Devised the GOLD FOIL EXPERIMENT
A. The Structure of the Atom
Rutherford’s Experiment
Unexpected Results of the
Rutherford Experiment
(a) The results that the metal foil
experiment would have yielded if the
plum pudding model had been correct
(b) Actual results
Rutherford’s Nuclear Atom
1911

Most of the mass of
an atom is in the
center…the
nucleus…with
electrons moving
around it.
Nuclear atom
By 1932, neutron was discovered, too.
 The nucleus is the central core of the
atom, composed of protons and neutrons.
Because, protons and neutrons have
much greater mass than electrons, almost
all of the mass of the atom is concentrated
in a tiny nucleus…a dime in a football
stadium!

Forces
in the Nucleus
• When two protons are extremely close to
each other, there is a strong attraction
between them.
• A similar attraction exists when neutrons are
very close to each other or when protons and
neutrons are very close together.
• The short-range proton-neutron, proton-proton,
and neutron-neutron forces that hold the nuclear
particles together are referred to as nuclear
forces.
What next?
We know the parts of the atom.
 We know about the nucleus.
 What about the electrons?

Bohr’s Model
LIGHT
provided the next
clues for the
structure of the
atom
Light has a Dual Nature
Like this
picture…
Young lady
or
Old lady?
Light is a wave
Electromagnetic radiation travels in waves.

As the wavelength increases, the frequency
decreases.

The greater the frequency the greater the
energy.
Light acts like Particles

Electromagnetic radiation also has the
properties of particles.

Bohr suggested that energy is emitted and
absorbed in discrete quantities called
Quanta or Quantum
packets or pieces of energy
“Jumps”
Dual Nature of Light is the next tool
for understanding the atom
Energy is directly proportional to
frequency…wave nature.
 Example: Light diffuses through small slits
 Einstein proposed that light consists of
quanta of energy that behave like particles
of light…he called these photons.
 Example: photoelectric effect…garage
door openers, film
This is the DUAL NATURE OF LIGHT

Demonstrate Quantum
Continuous Spectrum vs.
Line Spectrum
Emission Spectra



Line Spectra of elements
show discrete or quantized
energy levels in the atom.
These energy levels are
different for each element.
The color or wavelength of
light shows the energy level
because energy can be
calculated from frequency of
light emitted.
Quantum Leaps!
Electrons exist at low energy…
ground state
 Add energy to go to higher energy…
excited state
• Electron drops back down to ground
state as soon as it can
• It releases the exact energy it needed to
jump up as a photon of a frequency or
color.

Bohr’s Model of the Atom
1913




Bohr applied Quantum
Theory to the structure of
the hydrogen atom.
Quantum theory means
that electrons jump from
level to level.
Bohr pictured the
electrons spinning around
set orbits, like planets
around the sun.
Planetary Model
Schroedinger’s
Quantum Mechanical
(Today’s) Model
 We still have energy levels, like Bohr
 MATH is used as the tool to probe the
atom
 Heisenberg Uncertainty Principle
It is impossible to know the velocity and
position of a particle at the same time
 We can not know the exact location of an
electron.
 Any effort to do so, will change the position.
 We can only figure the probability of finding it
in a certain region.

Orbitals
Area where electrons are most likely to be
found.
 Example:
If your mom wants you to do some work,
she can find you in your room most of the
time, or the house, or the yard…but you
could be at the mall.
And sometimes her efforts to find you,
make you move!

Today’s Atom
1926
Dense,
small,nucleus (with
protons and
neutrons) is
surrounded by a
fuzzy cloud shapes
where electrons
(that act more like
waves) are most
likely to be found…
orbitals.
Today’s Model of the Atom
aka
Schroedinger’s
or
Quantum
Mechanical
Model
The Crazy World of Quantum Mechanics
“If you aren’t shocked by it, you don’t
understand it.” Bohr
An atom is the smallest particle of an element that retains
the chemical identity of the element.
Three Parts of Atom
Particle
Proton
Neutron
Electron
Charge
+1
0
-1
Mass
1.67x 10-24 g
1.67x 10-24 g
9.11x 10-28 g
Atomic mass units
1 amu
1 amu
0 amu
Subatomic particles
B. Introduction to the Modern
Concept of Atomic Structure
Comparing the Parts of an Atom
Where are they?
Nucleus is made up of protons and
neutrons.
 Electrons move around the nucleus.
 Mass is concentrated in the nucleus.
 99.97% of the mass is in the nucleus.

A
NEUTRON
WALKED INTO A BAR AND ASKED…
HOW MUCH FOR A
ROOT BEER?
THE BARTENDER SAID
FOR YOU, BUDDY…
NO CHARGE!
How big are atoms?
If nucleus is size of a golf ball, the
electrons are how far away?

1 mile! Lots of empty space
 How big are they?
 6.02 x 1023 atoms in 12 grams of carbon.

How many protons?

The atomic number is the number of
protons in the nucleus of the atom.
Look at the periodic table for this.
The number of protons identifies the
element.
CARBON
Carbon has
 6 protons
 6 neutrons
 6 electrons

6 amu
6 amu
0 amu
Total atomic mass = 12 amu
How do we write this?







Carbon-12
C-12
12C
12 C0
6
The periodic table tells us that carbon has 6
protons.
The mass is 12 amu and equals protons plus
neutrons, so it has 6 neutrons.
It is neutral (has no charge), so it has 6
electrons.
How many neutrons?





The number of neutrons may vary.
An isotope is an atom that has the same number
of protons as other atoms of the same element,
but different number of neutrons.
Same atomic number = same element
Protons + Neutrons = Atomic Mass
Isotopes have different atomic masses, because
they have different number of neutrons.
Isotopes

Isotopes are atoms with the same number
of protons but different numbers of
neutrons.
Isotope of carbon







Carbon-14
C-14
14C
14 C0
6
The periodic table tells us that carbon has 6
protons.
The mass is 14 amu and equals protons plus
neutrons, so it has 8 neutrons.
It is neutral (has no charge), so it has 6
electrons.
Carbon-14 isotope
said to Carbon-12
isotope,
“Do these neutrons
make my mass look
big?
What’s the
safe answer
here?
???
How many electrons?
The number of electrons equals the
number of protons in a neutral atom
 An ion is a charged atom in which the
number of electrons can increase or
decrease.
 More electrons than protons, charge is
negative
 Fewer electrons than protons, charge is
positive

Ion of carbon







Carbon-12
C-12
12C
12 C+4
6
The periodic table tells us that carbon has 6
protons.
The mass is 12 amu and equals protons plus
neutrons, so it has 6 neutrons.
It is positive (it lost 4 electrons), so it has only 2
electrons.
OH NO!!!
I think LOST an
ELECTRON!!
Are you
POSITIVE?
A
n (- or +)
E
Z
E = element symbol
Z = atomic number = # protons (identifies element)
A = atomic mass = # protons + # neutrons
n = charge
(+ if fewer electrons than protons)
(- if more electrons than protons)
AVERAGE ATOMIC MASS
is a
WEIGHTED AVERAGE
•
Imagine that your semester grade depends 60%
on exam scores and 40% on laboratory
explorations.
•
Your exam scores would count more heavily
toward your final grade.
•
In this section, you will learn that the atomic
mass of an element is a weighted average of the
masses of the naturally occurring isotopes of that
element.
Relative Atomic Masses
•
The standard used by scientists to compare
units of atomic mass is the carbon-12 atom,
which has been arbitrarily assigned a mass of
exactly 12 atomic mass units, or 12 amu.
•
One atomic mass unit, or 1 amu, is exactly
1/12 the mass of a carbon-12 atom.
•
The atomic mass of any atom is determined by
comparing it with the mass of the carbon-12
atom.
2 examples

Test scores for a class
2 examples

Easy atoms
Average Atomic Masses of
Elements
Average atomic mass is the weighted
average of the atomic masses of the
naturally occurring isotopes of an
element. It accounts for the relative
abundance of each isotope
Calculating Average Atomic Mass
• The average atomic mass of an element
depends on both the mass and the
relative abundance of each of the
element’s isotopes.
•
Calculating
Average Atomic Mass
• The average atomic mass of copper can be
calculated by multiplying the atomic mass of
each isotope by its relative abundance
(expressed in decimal form) and adding the
results.
• Copper consists of 69.09% copper-63,
which has an atomic mass of 62.9298
amu, and 30.91% copper-65, which has
an atomic mass of 64.9278 amu.
Calculating
Average Atomic Mass
• The average atomic mass of copper can be
calculated by multiplying the atomic mass of
each isotope by its relative abundance
(expressed in decimal form) and adding the
results.
• Copper consists of 69.09% copper-63,
which has an atomic mass of 62.9298
amu, and 30.91% copper-65, which has
an atomic mass of 64.9278 amu.
• (0.6909×
62.9298 amu) +
(0.3091 × 64.9278 amu) =
63.55 amu
• The calculated average atomic mass of
naturally occurring copper is 63.55
amu.
Solve

Calculate the average atomic mass for
element X, the element that is a “goofy
prisoner!” Then identify the element.
X-28
X-29
X-30
27.977 amu
28.976 amu
29.974 amu
92.21%
4.70%
3.09%
Counting Atoms to Determine
Molar Mass



We need to be able to count atoms to determine molar
mass
FIRST:
See handout to practice counting ALL atoms





Subscripts…# to right and below multiply what goes before it
Parentheses…atoms within act as a group, subscript applies to
all of the elements in parentheses
Coefficients…multiply everything
SECOND:
IGNORING COEFFICIENTS, add masses of all atoms to
determine molar mass of each substance. Use masses
to the hundredths place.
THE MOLE MAP
Everything goes through
the MOLE.
VOLUME OF A GAS (L)
(in Liters at STP)
Use mole conversion factors
and factor-label to solve
Molar volume of a gas
22.4 L of a gas / mole of gas (at STP)
Molar mass from PT
grams / mole
MASS (g)
(mass in grams)
MOLES (mol)
Avogadro’s #
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)
Relating Mass and Volume to
Numbers of Atoms
 The
mole (mol) is the SI unit for the AMOUNTof
a substance. It relates mass or volume of a gas
(things we can easily measure ) to the # of
particles (things too small to measure)
• The number of PARTICLES of a substance.
A mole of anything contains as many
particles as there are atoms in exactly 12 g
of carbon-12.
• “particles” can be atoms, ions,
molecules, or formula units
• Avogadro’s number—6.022 1415 ×
1023—is the number of particles in exactly
one mole of a pure substance.
Conversions
•
•
with Avogadro’s
Number
Avogadro’s number can be used to find
the number of atoms of an element
from the amount in moles or to find the
amount of an element in moles from
the number of atoms.
In these calculations, Avogadro’s
number is expressed in units of atoms
per mole.
Sample
Problem A
How many moles of silver, Ag, are in
3.01  1023 atoms of silver?
THE MOLE MAP
VOLUME OF A GAS (L)
(in Liters at STP)
Everything goes through
the MOLE.
Use mole conversion factors
and factor-label to solve
22.4 L of a gas/ mole of a gas at STP
Molar mass from PT
g/mole
MASS (g)
(mass in grams)
MOLES (mol)
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)

Sample Problem A Solution

Given: 3.01 × 1023 atoms of Ag
Unknown: amount of Ag in moles
Solution:


moles Ag
Ag atoms 
= moles Ag
Avogadro's number of Ag atoms
3.01  10
23
1 mol Ag
Ag atoms 
=
23
6.022  10 Ag atoms
0.500 mol Ag
Gram/Mole
Conversions
• Chemists use molar mass as a conversion factor in
chemical calculations.
• For example, the molar mass of helium is 4.00 g
He/mol He.
• To find how many grams of helium there are in two
moles of helium, multiply by the molar mass.
4.00 g He
2.00 mol He 
= 8.00 g He
1 mol He
Molar
Mass
• The mass of one mole of a pure
substance is called the molar mass of
that substance.
• Molar mass is usually written in units of
g/mol.
• The molar mass of an element is
numerically equal to the atomic mass
of the element in atomic mass units.
Sample
Problem B
What is the mass in grams of 3.50 mol
of the element copper, Cu?
THE MOLE MAP
VOLUME OF A GAS (L)
(in Liters at STP)
Everything goes through
the MOLE.
Use mole conversion factors
and factor-label to solve
22.4 L of a gas/ mole of a gas at STP
Molar mass from PT
g/mole
MASS (g)
(mass in grams)
MOLES (mol)
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)

Sample Problem B Solution
Given: 3.50 mol Cu
 Unknown: mass of Cu in grams
 Solution: the mass of an element in
grams can be calculated by multiplying
the amount of the element in moles by
the element’s molar mass.

moles Cu ×
grams Cu
= grams Cu
moles Cu
Sample
Problem B Solution,
continued
The
molar mass of copper from the
periodic table is rounded to 63.55 g/mol.
63.55 g Cu
3.50 mol Cu ×
= 222 g Cu
1 mol Cu
Sample
Problem C
A chemist produced 11.9 g of
aluminum, Al. How many moles of
aluminum were produced?
THE MOLE MAP
VOLUME OF A GAS (L)
(in Liters at STP)
Everything goes through
the MOLE.
Use mole conversion factors
and factor-label to solve
22.4 L of a gas/ mole of a gas at STP
Molar mass from PT
g/mole
MASS (g)
(mass in grams)
MOLES (mol)
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)

Sample Problem C Solution


Given: 11.9 g Al
Unknown: amount of Al in moles

Solution:
moles Al
grams Al 
= moles Al
grams Al
The molar mass of aluminum from the
periodic table is rounded to 26.98 g/mol.
1 mol Al
11.9 g Al 
= 0.441 mol Al
26.98 g Al
Sample
Problem D
What is the mass in grams of
1.20  108 atoms of copper, Cu?
THE MOLE MAP
VOLUME OF A GAS (L)
(in Liters at STP)
Everything goes through
the MOLE.
Use mole conversion factors
and factor-label to solve
22.4 L of a gas/ mole of a gas at STP
Molar mass from PT
g/mole
MASS (g)
(mass in grams)
MOLES (mol)
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)

Sample Problem D Solution

Given: 1.20 × 108 atoms of Cu
Unknown: mass of Cu in grams
Solution:


Cu atoms 
moles Cu
grams Cu

= grams Cu
Avogadro's number of Cu atoms
moles Cu
The molar mass of copper from the periodic
table is rounded to 63.55 g/mol.
1 mol Cu
63.55 g Cu
1.20 10 Cu atoms 

=
23
6.022 10 Cu atoms
1 mol Cu
8
1.27  10–14 g Cu
The Mole and Volume
The mole also relates the amount of atoms
to the volume of a gas at STP, Standard
Temperature and Pressure.
 One mole of gas at STP is 22.4 Liters

THE MOLE MAP
VOLUME OF A GAS (L)
(in Liters at STP)
Everything goes through
the MOLE.
Use mole conversion factors
and factor-label to solve
22.4 L of a gas/ mole of a gas at STP
Molar mass from PT
g/mole
MASS (g)
(mass in grams)
MOLES (mol)
6.022 x 1023 particles / mole
NUMBER OF PARTICLES (#)
(atoms, ions, molecules, formula units)
Sample Problem E

How many moles of Helium are present in
44.8 L at STP? How many grams?
44.8 L He | 1 mole He = 2 moles He
22.4 L He
44.8 L He | 1 mole | 4.00 g He = 8 g He
22.4 L He | 1 mole He