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Atomic Theory
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The Evolution of the Atomic
Model
Atom
Basic building block of matter
Cannot be broken down chemically
A single unit of an element
Dalton (1803)
Known as the founder of the atomic theory
Dalton’s Postulates:
1. All matter is composed of indivisible particles called
atoms
2. all atoms of a given element are identical in mass and
properties. Atoms of different elements have different
masses and different properties
3. Compounds are formed by a combination of 2 or more
atoms
4. Atoms cannot be created, destroyed, or converted into
other kinds of atoms during chemical reactions
Cannonball
Theory/Model
1. Spherical
2. Uniform Density
J.J. Thomson (1897)
Used a cathode ray tube with charged particle field (+/-)
Cathode ray deflected by negative electrode toward
positive electrode
Discovered subatomic particle called the ELECTRON
Small
Negatively charged
J.J Thomson
Plum Pudding Theory/Model
Positive “pudding”
Negative electrons embedded (just like raisin bread)
Rutherford (1909)
Conducted the GOLD FOIL EXPERIMENT
where he bombarded a thin piece of gold foil
with a positive stream of alpha particles
Expected virtually all alpha particles to pass
straight through foil
Most passed through but some were severely
deflected
Conclusion led to…
Rutherford’s
Experiment
Rutherford
Nuclear Theory/Model
1. The atom is mostly empty space
2. At the center of the atom there is a DENSE, POSITIVE CORE
called the NUCLEUS
**Provided no information about electrons other than the fact
that they were located outside the nucleus**
Neils Bohr (1913)
Bohr Model or Planetary Model
1. Electrons travel AROUND the nucleus in well
defined paths calls ORBITS (like planets in the solar
system)
2. Electrons in different orbits possess different
amounts of energy
Neils Bohr
3. Absorbing/Gaining a certain amount of energy
causes electrons to jump to a higher energy level or an
excited state
When excited electrons emit/lose a certain amount of energy
which causes electrons to fall back to a lower energy level or
the Ground State
Wave-Mechanical/Cloud
Model
Modern, present day model
Electrons have distinct amounts of energy and move in areas
called orbitals
Orbital= an area of high probability for finding an electron (not
necessarily a circular path)
Developed after the famous discovery that energy can behave
as both waves & particles
Many scientists have contribute to this theory using X-Ray
diffraction
Vocabulary
(of the Periodic Table)
Atomic#= the number
of protons in every
atom of the element
(NEVER CHANGES!)
Atomic Mass= average
mass of all the isotopes
of an element
Element Symbol=
the letter(s) used
to identify an
element
**Nucleons= Protons and Neutrons
(any subatomic particle found within the nucleus)
Subatomic Charge Relative
Particle
Mass
Location
Symbol
How to
Calculate
or
Look at
atomic #
Proton
+1
1 amu
Nucleus
Neutron
0
1 amu
Nucleus
Mass # Atomic #
Electron
-1
1/1836
amu
Outisde
nucleus
P= e
(in neutral
atoms)
Determining Subatomic
Particles (p, n, e)
# of Protons = sum of the protons and neutrons in an atom of
an element (Atomic Number)
# of Neutrons = Mass # - Atomic # or Mass # - # Protons
# of Electrons = sum of the protons and neutrons in an atom of
an element ( equal to # protons in a neutral atom)
Determining Subatomic
Particles (p, n, e)
Atomic Number:
1. Look at the element symbol and locate on the periodic table
2. same as the # of protons or the nuclear charge
Mass Number = # Protons + # Neutrons
Nuclear Charge = # of protons or the atomic #
P = nuclear charge
Atoms vs. Ions
Vocabulary Term
Neutral Atom
Ion
Definition
An atom with the
same number of
protons and
electrons
P=e
Two Types
Anion and Cation
Protons and
electrons are not
the same
Example
(no charge indicated)
Atoms vs. Ions
Ions
Anion
aNion
Cation
ca+ion
An atom that has
An atom that has
GAINED one or more LOST one or more
electrons
electrons
e>p
p>e
NEGATIVE ION
POSITIVE ION
Isotope
Atoms of the same element with
different mass numbers; same
atomic number, same number of
protons, different number of
neutrons
Isotopes
Example 1: Carbon (12, 13, 14)
p=
e=
n=
p=
e=
n=
p=
e=
n=
Isotopes
Example 2: Uranium (238, 240)
U-240
p=
e=
n=
p=
e=
n=
Practice
1. Two different isotopes of the same
element must contain the same number of
*
a. protons
b. neutrons
c. electrons
2. Two different isotopes of the same
element must contain a different number of
a. protons
*
b. neutrons
c. electrons
Practice
3. Isotopes of a given element have
*
a. the same mass number and a different atomic
number
b. the same atomic number and a different mass
number
c. the same atomic number and the same mass
number
Calculating Atomic Mass
(for any element)
Atomic mass = the weighted average of an element’s
naturally occurring isotopes
% abundance of isotope 1 x (mass of isotope 1)
% abundance of isotope 2 x (mass of isotope 2)
+
% abundance of isotope 3 x (mass of isotope 3)
Average Atomic Mass of the Element
Calculating Atomic Mass
Example 1: The exact mass of each isotope is given
Chlorine has two naturally occurring isotopes, Cl-35
(isotopic mass 34.9689 amu) and Cl-37 (isotopic
mass 36.9659 amu). In the atmosphere, 32.51% of the
chlorine is Cl-37, and 67.49% is Cl-35. What is the
atomic mass of atmospheric chlorine?
Step 1: Multiply the mass of each separate isotope by its
percent abundance
CL-35 = 34.9689 amu x
(.6749) = 23.6005
CL-37 = 36.9659 amu x
(.3251)
=
12.0176
Calculating Atomic Mass
Step 2: Add the products of all the calculated
isotopes together from step 1
23.6005
+ 12.0176
35.6181
***This is your average atomic mass***
Calculating Atomic Mass
The element Carbon occurs in nature as two isotopes.
Calculate the average atomic mass for Carbon based
on the information below
C-12 = 98.89%
C-13 = 1.11%
**Since the mass numbers were not given for either
isotope, use the mass number instead**
C-12 = 12
C-13 = 13
Calculating Atomic Mass
Step 1: Multiply the mass by the percent abundance
C-12 = 12 x (.9889) = 11.8668
C-13 = 13 x (.0111) =
*These are weighted masses*
Step 2: Add the products together
11.8668
+ 0.1443
12.0111
0.1443
Calculating Atomic Mass
Practice: The element Boron occurs in nature as two
isotopes. Calculate the average atomic mass for Boron, using
the information below.
Isotope
Mass
% abundance
Boron-10
10.0130 amu
19.9 %
Boron-11
11.0093 amu
80.1 %
Average Mass of Boron = 10.8104
Practice: The element Hydrogen occurs in nature as
three isotopes. Calculate the average atomic mass
of Hydrogen.
Isotope
% abundance
Protium
99.0%
Deuterium
0.6%
Tritium
0.4%
Average Mass= 1.014
Mass Number
Atomic Mass
The MASS of ONE
isotope of a given
element
The AVERAGE MASS of
ALL isotopes of a given
element
Electron Configurations
The dashed chain of numbers found in the lower left
hand corner of an element box
Tells the number of energy levels as well as the
number of electrons in each level
(how the electrons are arranged around the nucleus)
Electron Configuration
Is the representation of the arrangement of electrons
distributed among the orbitals
Used to describe the orbitals of an atom in the
ground state
Can be used to describe ionized atoms (cations and
anions)
Many of the chemical and physical properties of
elements can be correlated to their electron
configuration
Electron Configurations
All electron configurations on the Periodic
Table are NUETRAL ( p = e)
For IONS, add or subtract electrons from the
LAST NUMBER in the electron configuration
only
Electron Configuration
Orbitals
There are four types (s, p, d, and f)
They have different shapes
Each orbital can hold a maximum of two
electrons
P, d and f orbitals have different sublevels
Electron Configuration
Is unique to an element’s position on the periodic
table
Energy level is determined by the period
Number of electrons is given by the atomic number
Electron Configuration
Orbitals on different energy levels are similar to
each other but they occupy different areas in space
ex. 1s and 2s orbitals both have s orbital
characteristics ( radial nodes, spherical volume, can
only hold two electrons) but since they are found in
different energy levels they occupy different spaces
around the nucleus
Electron Configuration
Electrons fill orbitals to minimize energy
Electrons fill the principal energy levels in order
of increasing energy
Electrons are getting further from the nucleus
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
Electron Configuration
A single orbital can hold a maximum of two electrons
These electrons must have opposing spins
One electron spins up and the other spins down
Ensures they have different quantum numbers
Pauli Exclusion Principal
Electron Configuration
S subshell 1 orbital that can hold 2 electrons
P subshell 3 orbitals that can hold up to 6 electrons
D subshell 5 orbitals that can hold up to 10 electrons
F subshell 7 orbitals that can hold up to 14 electrons
Principal Energy Level (n)
Electron energy levels consisting of orbitals which
designated s, p, d, or f.
Electron Configurations
Valence Electrons:
Electrons found in the OUTERMOST shell or orbital
Kernel Electrons:
INNER electrons (all non-valence electrons)
Orbital Notation
Orbital Notation
When filling in the orbital
Electrons fill the lowest vacancy levels first
When there’s more than one subshell at a particular
energy level (ex. 3p or 4d) only one electron fills
each subshell until each subshell has one electron.
Then electrons start pairing in each subshell
Hund’s Rule
Every orbital in a subshell is singly
occupied with one electron before
any one orbital is doubly occupied,
and all electrons in singly occupied
orbitals have the same spin.
Energy Level Diagram: Oxygen.
Oxygen is atomic number 8.
8 protons  How many electrons?
8 electrons
Bohr Diagrams
A method for showing electron location in an
atom/ion
All electrons must be drawn
Look up the electron configuration and the period
(row) the element is in on the periodic table
Draw a circle (nucleus) write in the number of
protons and neutrons
Draw the shells (energy levels) around the nucleus
Bohr Models
Add the electrons
The first shell only holds 2 electrons
Now add the rest of the electrons
Add one at a time starting on the right side and going
counter clockwise
Shell (energy level)
1
2
3
4
5
6
7
2
2n
n=
Maximum number of
electrons
2
8
18
32
50
72
98
Formula for maximum # of electrons
per quantum number (or energy
level)
Bohr Models
Example:
Draw the Bohr Model for Carbon
Bohr Models
Example
Draw the Bohr Model for Oxygen
Bohr Models
Example
Draw the Bohr Model for Na+
E- configuration 2-8
Lewis Dot Diagrams
Only illustrates valence electrons
Write the element symbol
Look up the electron configuration (use the last
number in the configuration- # of valence
electrons)
Use either an X or dot to represent the electrons
Place that many electrons around the symbol at
(12, 3, 6 and 9)
Lewis Dot Diagram
Example: Carbon
e- configuration 2-4
Practice
The number of unpaired electrons is equal to the
number of BONDS that an element can form with
other elements
When determining the number of bonds an element
can form, arrange the valence electrons so that you
have the MAXIMUM number UNPAIRED
Example: Carbon
How many bonds can carbon form?
4
Ground State vs. Excited State
Ground State electrons are in the lowest energy
configuration possible (the configuration found on
the periodic table)
Excited State electrons are found in a higher
energy configuration (any configuration not listed
on the periodic table)
Ground State
Excited State
EXCITED
EXCITED
GROUND
GROUND
EXCITED
GROUND
EXCITED
EXCITED
The greater the distance from the nucleus, the
greater the energy of the electron
When ground state electrons absorb energy they
jump to a higher energy level or an excited state
This is very unstable/temporary condition
Excited electrons fall rapidly to a lower energy level
When excited electrons fall from an excited state to a
lower energy level, they release energy in the form of
light
Ground  Excited
Energy is absorbed
Dark line spectrum is produced
Excited  Ground
Energy is released
Bright line spectrum is produced
Dark Lines
Absorbed
Bright Lines
Emitted
Balmer Series: electrons falling from an excited
state down to the second (2nd) energy level give off
visible light (Bright Line Spectrum or Visible Light
Spectrum)
Different elements produce different colors of
light or spectra
These spectra are unique for each element
Spectral lines are used to identify different
elements
*
Gas A and Gas D