History of the Development of the
Atomic Model, Part 1
AS Aim #1
Where did the idea
of the atom come
from?
Atomic Model of the Atom
• Models help us describe
several things about an
atom:
– ____________ – what its
made of
– __________________ – lets us determine
how atoms interact with each other
• We will look at several models of the atom
that build upon previous models
• The modern model of the atom is based
on the work of many scientists, not just
one!
Early Greek
Theories
• Democritus
• 400 B.C.
• His theory:
– everything is composed of "atoms", which are
physically ___________________________;
– are always ________________;
– come in there are an __________________
of types, variety, and shapes
• Based his theory on _____________________
Early Greek Theories
• Aristotle
• 350 B.C
• Presented a modified earlier
theory that matter was made
of four “elements”; ______, ____,
______, and ________
• Based his theory on reason, ______________
• But…
• ____________________!
• __________________________ must be done
to ensure theories hold up under scrutiny!
John Dalton’s Atomic Theory of the Atom
• Early 1800’s English teacher
John Dalton
– proposed a modern atomic
model of structure ________
_______________________
_______________________
– Described elements as being
composed of particles called
_______________________
– _______________________
to a given element
Components of Dalton’s Model
• All matter is __________________
•
Atoms of an element _________________
•
Each element has ___________________
4. Atoms of different elements combine in
__________________ to form compounds.
• Think of H20 vs H2O2
5. Atoms are __________________________
• 2 H2 +O2  2 H2O
Dalton’s Theory accounts for:
• The Law of Conservation of Mass
– Mass cannot be ____________________
– Hydrogen + Oxygen = Water
– 2 grams + 16 grams = ______________
• The Law of Constant Composition
– elements combine in fixed ratios.
– 2 hydrogen + 1 oxygen = _______________
– 2 hydrogen + 2 oxygen = _______________
• PROBLEM – no clues in his model as to the
_________________________ of the atom
Cathode Rays and Electrons
• By 1897, experiments suggested atoms are
composed of subatomic particles
– Subatomic particles - ________________
• British physicist J.J. Thomson
– Used a ________________________ and
discovered particles he called ___________
– ____________ are
negatively charged
– Mass of one electron
only 1/1836 of a
________________
The Thomson Atomic Model
• “Plum pudding” model
– an atom is a positively charged, jellylike
mass with electrons “stuck” in it
• Did not _____________
Dalton’s model
• Built upon Dalton’s
model
– __________ positive
and negative
charges
– _________________
to the charges
History of the Development of the Atomic
Model, Part 2
AS Aim #2 – What is significance of the
Gold Foil Experiment?
Ernest Rutherford’s Gold Foil Experiment
• In his experiment, he bombarded (hit)
extremely thin gold foil with _____________
– a helium nucleus only (_______________)
– has a __________________
– _______________________ (more later)
• Based on Thomson’s Theory:
– Particles should bounce off the
___________________ or
– Particles should stick to the
negative electrons made
up of Thomson’s “plums”
• Experimental Results
– Some alpha particles _______________ or
__________________
– But some particles ____________________
the atom
– Why?
• Rutherford’s Conclusion
– Most of the mass of the atom is concentrated
in a ________________________________
– Electrons are present in the space
________________________________
– The typical model of the
atom represented in the
media is Rutherford’s
model at the right
– So most of the volume
of an atom is
____________________
Bohr’s Atomic Model
• Rutherford didn’t say
where electrons were
__________________
• Bohr said that electrons
exist in _________ or
_________________
• n represents the energy level
– Energy level n = 1 , holds up to ___________
– Energy level n = 2 , hold up to ___________
– Energy level n = 3 , hold up to ____________
– Energy level n = 4 , holds up to ___________
Bohr’s Experiment:
– Bohr used hydrogen gas which he heated to
___________________
– He analyzed the light patterns using a device
called a ___________________________
which separated the colors of light produced
– In hydrogen,
he found four
specific
____________
of color
Bohr’s Theory:
• Bohr found that electrons
moved from one energy level to
another when they gained
energy
• They released the energy as
light (photons)
• In the lowest levels, or the
ground state, to the excited
state ____________________
________________________
• When electrons moved from
the excited state back to the
ground state, ______________
Bohr’s Theory:
• This energy
appears as
wavelengths
(______________)
_______________
• Each element
produces its own
pattern of
______________ or _________________
• This is because each has different numbers of
electrons
Summary of History of the Atomic Model
• Democritus – came up with the word _____
• Dalton – his original atom had no ________,
neutrons, or ________ in it
• Thompson – used _______________ tubes
to discover electrons and their charge
• Rutherford – his experiment shot
_____________ at a piece of ____________
• Bohr – used excited hydrogen atoms to
produce ___________ in various colors, and
proved electrons existed in different energy
levels
History of the
Development of
the Atomic Model,
Part 3
AS Aim #3:
Where does the
Modern Model of
the Atom place
electrons?
Modern Atomic or Wave
Mechanical Model
• Bohr’s shell model at the right
is not quite right either!
• Electrons actually exist in
________________________
around the nucleus, not in
orbits like planets around the
Sun
• As per the Modern Atomic
Model
• Also known as the Wave
Mechanical Model of the
Atom
Modern Atomic or Wave
Mechanical Model
• These locations are based on
where they are most likely
found, ________________
• We call this arrangement an
_____________________
• ______________ are a three
dimensional representation of
principal energy levels
• Each energy level (n)
contains smaller areas
called ____________
• In the Periodic Table,
each block represents
a __________________
with electrons
• This is what gives the Periodic Table its
_____________________
• There are _______________
• each labeled “s”, “p”, “d”, and “f”
•
•
Sublevels are further broken down into
areas called ____________
Each orbital only holds two (2) electrons
each maximum and has a _____________
Number of
Sublevel
orbitals
s
1
p
3
d
5
f
7
Shape of
Maximum #
orbitals
of electrons
Single round
sphere
Three dumbbell
shapes
Five dumbbell
shapes
Seven dumbbell
shapes
•
Orbital shapes affect
– how the Periodic Table _____________
– how _______________ with each other
• Look at the Periodic
Table
• Count the elements
across each block.
How many elements
are there in each?
• s sublevel block = ____
• p sublevel block = ____
• d sublevel block = ____
• f sublevel block = ____
• Each ____________________ in the Periodic
Table represents 1 more electron being added
• Simplified Electron Configurations:
– ____________________ surround an atom
– ____________________ are in each energy
level
– ____________________ they exist in
Atomic mass
12.011
-4
+2
+4
Atomic Number
6
C
Symbol
2-4
Electron
Configuration (2 e- in 1st level, 4 e- in 2nd level)
Basic Electron Configuration
• Electrons fill the lowest energy levels first
(_______________)
• The electron configuration is a “code” for
showing _______________ around an atom
1st
Element
Shell
2nd
Shell
3rd
Shell
4th
Shell
Electron
Config
He
2
2
Na
2
8
1
Br
2
8
18
7
2-8-18-7
Ca
2
8
8
2
2-8-8-2
2-8-1
The Octet Rule of Electron Configs
• Why is calcium’s 3rd shell not filled?
• It should take up to __________ in the 3rd shell
1st
Element
Shell
Ca
2nd
Shell
3rd
Shell
4th
Shell
Electron
Config
8
8
2
2-8-8-2
2
• The Octet Rule
– no atom can have more than 8 electrons in the
_____________________ energy level
– If more than 8 electrons in an energy level occurs,
we push two up to the next energy level
The Octet Rule of Electron Configs
• Having 8 electrons in the valence shell also
makes the ___________
• This occurs in the last column of the Periodic
Table, a group of elements called Noble _____
• Very _____________ with other elements
1st
Element
Shell
2nd
Shell
3rd
Shell
4th
Shell
Electron
Config
Ne
2
8
-
-
2-8
Ar
2
8
8
-
2-8-8
Kr
2
8
18
8
2-8-18-8
Excited State Electron Configurations
• Excited state electrons can be shown by not
filling the _________________
• Electrons have ___________ to higher shells
Element
He
Na
Br
Ca
Ground state
electron
configuration
2
2-8-1
2-8-18-7
2-8-8-2
Excited state
electron
configurations
1-1, 1-0-1
2-7-2, 2-7-1-1
2-8-18-6-1, 2-8-17-8
2-8-7-3, 2-8-8-1-1
• Fill in the chart below for each element’s
ground state electron configuration and one
excited state configuration:
Element
K
Mg
O
S
N
P
Ar
Ground state
configuration
Excited state
configuration
Atoms + Electrons = Ions
AS Aim #4 – Why are most elements
“wannabees”?
HAIL THE MIGHTY VALENCE ELECTRONS!
• Most of chemistry is really all about electrons
and where they go and stay
• All elements in the Periodic Table are
“__________________________” (Group 18)
– If an atom can gain or lose electrons, it can have
the electron configuration as the noble gases
• These elements become stable when they
form _________:
– a gain or loss of electrons gives an _____________
– a __________ electrons creates a negative ion
– a __________ electrons creates a positive ion
HAIL THE MIGHTY VALENCE ELECTRONS!
• Ions are atoms with a charge, or an unequal
number of protons and electrons
• What is the charge on a proton?
______
• What is the charge on an electron? ______
• What is the charge on each of the following
atoms:
– 5 protons and 5 electrons
– 5 protons and 4 electrons
– 5 protons and 6 electrons
– 19 protons and 21 electrons
= ______________
= ______________
= ______________
= ______________
HAIL THE MIGHTY VALENCE ELECTRONS!
• Ions of opposite charge can thus form
compounds
• Positive ions ________________ negative ions
(___________________!)
• In compounds, as in atoms, charges must add
_________________
• Therefore:
– A +1 ion bonds with a -1 ion (+1 + -1 = ____)
– A +2 ion bonds with a -2 ion (+2 +-2 = _____)
– A +2 ion bonds with two -1 ions (+2 +(-1x2))= ___)
• Each of the atoms below want to be ions with a
stable electron configuration of eight
• Determine how many electrons are gained or lost
• Write the new electron configuration
Element
Electron
Config of
atom
Gained or
lost e-
Electron
Config of ion
K
2-8-8-1
1 lost
2-8-8
Mg
2-8-2
2 lost
2-8
O
2-6
2 gained
2-8
N
2-5
3 gained
2-8
Li
2-1
1 lost
2
Lewis Electron-Dot Diagrams
• Another way to represent _____________
• Lewis Dot Diagrams shows the number of
______________________
• Procedure
– Write the symbol first
– Use the Periodic Table to find the number of
valence (outermost) electrons
– Place two dots to represent the first electrons
on top
– Place the rest evenly around the atom
Lewis Electron-Dot Diagrams
• Example: draw the Lewis Dot for
chlorine and for sodium
• Notice: chlorine __________ to be like
a noble gas, sodium ________________
Name,
Symbols,
and Atomic
Numbers
AS Aim #5 – What
does a chemical
symbol tell us about
an element?
Identifying Elements – Names, Symbols,
and Atomic Numbers
• Every element (as well as its atoms) is
associated with three unique identifiers
– Names
– Symbols
– Atomic numbers (number of protons in an atom)
– obtained from the periodic table
Element
Element Element Atomic
Name
Symbol
Number
Hydrogen
H
1
Sodium
Na
11
Gold
Au
79
Identifying Elements – Names,
Symbols, and Atomic Numbers
• Names
– are based on
– _______– like Einsteinium
– ________ – like Francium
– ________________ - like
chlorine (comes from the
Greek work chloros, or
“yellow green”)
Identifying Elements
– Names, Symbols,
and Atomic Numbers
Symbols
• Each element with
permanent names have unique letters
associated with them
– First letter ____________capitalized
– Second letter ______________capitalized
• New elements have three letter symbols
__________________ to them
• Some elements originally had ___________
• Example – Mercury, or hydragyras (Hg)
Identifying Elements – Names, Symbols,
and Atomic Numbers
• Give either the name or the symbol of each
(you may need to use Table S to do this!)
Name
Symbol
Name
Symbol
Nickel
Ni
Magnesium
Mg
Tungsten
W
Radium
Ra
Radon
Rn
Uranium
U
Bromine
Br
Arsenic
As
Identifying Elements – Names, Symbols,
and Atomic Numbers
• Atomic number - this represents the number of
____________ in the atom’s ______________
• Each element has its own atomic number
• Therefore,
• the _______________ gives you the element
• Examples:
Atomic number = 2 = 2 protons = _______
Atomic number = 8 = 8 protons = _______
Atomic number = 79 = 79 protons = _____
Identifying Elements – Names, Symbols,
and Atomic Numbers
• Give the name of the element based on the atomic
number (use Table S and the Periodic Table)
Atomic
#
4
Name
Beryllium
Atomic
#
25
Name
Manganese
8
Oxygen
50
Tin
12
Magnesium
75
Rhenium
16
Sulfur
100
Fermium
Identifying Elements – Names, Symbols,
and Atomic Numbers
• The Atomic Number also gives us
the number of electrons in an atom
• Remember,
– Protons = charge of _____
– Electrons = charge of ______
– Neutrons = charge of ___________
– ATOMS are always electrically ___________
(charge = 0)
– Therefore, in an atom,
the # of __________= the # of ___________
but not the number of neutrons (that changes!)
Masses and Isotopes
AS Aim #6:
What makes
something an
isotope?
Using the Periodic Table to find Names,
Symbols, and Atomic Numbers
• The Periodic Table of the elements is the
master chart of chemistry
• Contains various pieces of information
including:
Atomic mass
12.011
(Total protons+neutrons)
Atomic Number
(number of protons, only)
6
-4
+2
+4
C
Oxidation
states
Symbol
2-4
Electron
Configuration (arrangement of electrons in energy levels)
Using the Periodic Table to find Names,
Symbols, and Atomic Numbers
• Determine the following information for each
element from the Periodic Table:
Name
Hydrogen
Boron
Chlorine
Argon
Calcium
Atomic Atomic Electron
Symbol
Number Mass
Config
H
B
Cl
Ar
Ca
1
5
17
18
20
1.0079
1
10.81
2-3
35.45 2-8-7
39.95 2-8-8
40.08 2-8-8-2
Isotopes, Mass Numbers, and Neutrons
• Mass number – a measure of the number of
____________________________ in an atom
• Why not electrons too?
• _________________________!!!
• Masses of subatomic particles are measured in
units called _____________________ or amu’s
– Mass of 1 neutron = 1 amu
– Mass of 1 proton = 1 amu
– Mass of 1 electron = 0.0005 amu
• Therefore,
Mass number = # of protons + # of neutrons
Isotopes, Mass Numbers, and Neutrons
• Does this mean that atoms of the same
element all have the same mass numbers?
• ________!
• Atoms of the same element are actually a bit
different
• We call them ISOTOPES
– All atoms of an element have the
___________________ (# of protons)
– But atoms of the same element can have
____________________
(different #s of neutrons)
Isotopes, Mass Numbers, and Neutrons
• Examples of Isotopes - hydrogen
– There are three forms of hydrogen isotope
Form of
Hydrogen
Isotope
Atomic
Number
(# of protons)
Mass Number
(# of protons +
neutrons)
# of
Neutrons
Protium
Deuterium
Tritium
1
1
1
1
2
3
0
1
2
Neutrons, Isotopes, and Mass Numbers
• Isotopes of hydrogen
Isotopes, Mass Numbers, and Neutrons
• Question 1 - an atom has an atomic number of 6,
and a mass number of 12 amu.
– What element is it?
______________
– How many protons does it have?
______________
– How many neutrons does it have?
mass of 12 – 6 protons = 6 neutrons
– How many electrons does it have?
________________________________
– IMPORTANT! 1 amu = 1/12 the mass of a
typical carbon atom
Isotopes, Mass Numbers, and Neutrons
• Question 2 - an atom has an atomic number of 6,
and a mass number of 14 amu.
– What element is it?
____________
– How many protons does it have?
____________
– How many neutrons does it have?
_______________________________
– How many electrons does it have?
_______________________________
– This is an ISOTOPE of carbon
Neutrons, Isotopes, and Mass Numbers
• Isotopes of carbon
Representing isotopes
Isotopes can be represented in several ways
• As the element with it _______________
– Ex: carbon-12, carbon-13,
carbon-14
• As the element’s symbol with
its ________________
– Ex: C-12, C-13, C-14
• As the symbol with both the
__________ and the atomic
number represented
Neutrons, Isotopes, and Mass Numbers
• Problem – how many protons, neutrons,
and electrons are contained in a neutral
atom of uranium-238, if the atomic number
of uranium is 92?
• Mass number
238 = _______
• Protons
92 = _______
• Electrons
92 = ___ = ___
• Neutrons
146 = _____________
- _____________
Average
Atomic
Masses
AS Aim #7:
Why do atomic
mass numbers
contain
decimals?
Calculating Grade Averages
• You are in the class from H… the teacher
has decided that your grade for the quarter
will be based on the following weighting:
– Exams
– Homeworks
– Labs
60%
30%
10%
• You score 50% average on your exams, a
70% average on your homework, and an
90% average on your labs.
• Do you pass the course the first quarter?
Calculating Grade Averages
• If the teacher averaged the 3 grades, you would
simply add your 3 grades and divide by 3
(50 + 70 + 90 ) / 3 = ___________
= you pass and your family is happy!
• BUT = the grade is weighted, so:
– Exams
60% x 50% avg =
– Homeworks
30% x 70% avg =
– Labs
10% x 100% avg = ____
– The total comes out to be
• You fail, and now you get to attend extra help
FOREVER
Calculating Grade Averages
• Problem #1 - Evil Mr. Foley decides your
second quarter exams will be 80% of your
Test grade, HW will be 10%, and Labs will be
10%. If you score a 60 avg on exams, an 80
avg on labs, and a 100 avg on labs, do you
pass?
Calculating Grade Averages
• Problem #2 – Mr. Foley’s good twin decides
in his class that the weighting will be quite
different. For the second quarter, exams will
be 50% of your grade, HW will be 30%, and
Labs will be 20%. If you still score a 60 avg
on exams, an 80 avg on homework, and a 100
avg on labs, do you pass?
Calculating Atomic Weights
• Determining the atomic weights of elements is the
same
• _______________________________________
_______________________________________
• So we need to calculate the atomic weight based
_______________________________
• Example 1 – a sample of hydrogen isotopes:
– Hydrogen-1 has an abundance of 95%
– Hydrogen-2 has an abundance of 3%
– Hydrogen-3 has an abundance of 2%
• What is the average atomic mass of this sample?
Calculating Atomic Weights
• Hydrogen-1 has an abundance of 95%
• Hydrogen-2 has an abundance of 3%
• Hydrogen-3 has an abundance of 2%
Hydrogen 1 = 1 amu x 95% =
Hydrogen 2 = 2 amu x 3% =
Hydrogen 3 = 3 amu x 2% = ______
Calculating Atomic Weights
• Example 2: A sample of sulfur has the
following isotopes in it
– Sulfur-30 with an abundance of
60%
– Sulfur-32 with an abundance of
30%
– Sulfur-34 with an abundance of
10%
• What is the average atomic mass of this
sulfur sample?
Sulfur-30 = 30 amu x 60% =
Sulfur-32 = 32 amu x 30% =
Sulfur-34 = 34 amu x 10% = ________