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Chemistry Unit Three
Atomic Structure
A Brief History of the
Atomic Theory
Democritus
400 B.C. Greek
philosopher
Coined the term
“atomos” which means,
Indivisible.
All matter is made of
atoms.
Atoms are hard, solid
particles, made of the
same material but are
of different shapes
and sizes.
John Dalton
1803 English Chemist
Atoms are solid,
neutral spheres.
Atoms of same
element are the same.
Atoms of different
elements are
different.
Compounds form from
the joining of atoms
of two or more
elements.
J.J. Thomson
1897 English Chemist.
Atoms are made of even
smaller particles.
Called the Plum Pudding
Model (Chocolate Chip
Cookie Dough Model)
Positively charged material
through which negative
particles are scattered.
Since atoms are neutral,
therefore, there must be
(+) particles too, but
Thomson never found them.
Cathode Ray Experiments;
discovered the electron.
Ernest Rutherford
1911 British physicist
Gold Foil experiment.
Atom has a small,
dense positively
charged center called
the Nucleus.
Negative electrons are
scattered outside the
nucleus.
Most of the atom is
empty space
If an atom was the size of a
baseball stadium, the nucleus
would be the size of a marble.
Rutherford’s
Gold Foil Experiment
•A beam of + particles (alpha particles) shot
through a thin sheet of gold foil.
•Most particles passed straight through.
(Most of atom is empty space.)
•A few were deflected. (Positive core-similar
charges repel each other.)
•Very few bounced off. (Solid core is very
small.)
Neils Bohr
1913 Danish Scientist
Planetary model.
Electrons are held in
place by the attraction
between them and the
+ charged nucleus.
Each electron occupies
a specific energy level
and orbit the nucleus
like planets circling
the sun.
Labeled each energy level
by a quantum number.
Wave Model
Electrons are not
discreet particles
moving in discreet
orbits.
The probable location
of an electron depends
on how much energy it
has.
Electrons seem to be
everywhere at once,
like the moving blades
of a fan.
Electron Cloud Model
Positively charged
protons and neutral
neutrons are held
together with a huge
amount of energy
forming the nucleus
of the atom.
Negatively charged
electrons move rapidly
around the outside of
the nucleus forming
“clouds” of negative
charge.
Most of the mass of
the atom is in the
nucleus.
Quantum Model
Summary of Atomic Models
Atomic Structure
Protons
(p+)
Found in the
Nucleus
Has 1 amu of
mass
Neutrons
Electrons
(n)
(e-)
Found in the Found outside
Nucleus
the nucleus
Has 1 amu of
mass
Has 0 amu of
mass
Has a positive Has no charge Has a negative
(neutral)
charge
charge
Atomic Structure
Atom
Isotope
Ion
The number Atoms of the An atom that
of protons in same element has lost or
an atom never
that have
gained
changes.
different
electrons.
numbers of
neutrons.
Atomic Mass
Average Atomic Mass – the average
mass of all of the isotopes of an
element. Decimal number.
Mass Number – the total number of
protons and neutrons in the nucleus
of an atom. Whole number.
Atomic Structure
Atomic Number = the number of
protons. Equal to the number of
electrons.
Atomic Mass = the number of protons
and neutrons added together.
Atomic Mass – Atomic Number = the
number of neutrons.
Calculating Numbers of Protons,
Neutrons and Electrons.
Atomic Number = 6
6
C
12.011
6 protons = 6 electrons
6 p+ = 6 e(atom is neutral)
Atomic Mass = 12
12 p+ and n
-6 p+
6 neutrons
Practice Calculating p+, n, eElement Atomic
number
Silver
47
Mass
#
Number proton
108
47
#
#
electron neutron
47
61
*Atomic # is number of protons so protons = 47
*number (+) charges (p+) must equal (–) charges to
make the atom neutral so electrons = 47
*Mass Number is total particles with mass
(p+ and n) so 47 + 61 = 108
Practice Calculating p+, n, eElement Atomic
number
Copper
29
Mass
#
Number proton
64
29
#
#
electron neutron
29
35
*number of protons is the atomic # so atomic
number is 29
*number (+) charges (p+) must equal (–) charges to
make the atom neutral so electrons = 29
*Mass Number is total of all particles with mass
(p+ and n) so subtract away the atomic number
(#p+) and you will have just neutrons (64 – 29 = 35)
Practice Calculating p+, n, eElement Atomic
number
Mass
#
Number proton
#
#
electron neutron
Tin
50
69
50
119
50
+ + n
p
+ = - (neutral)
Is # p+
Helium
4
2
2
2
2
Mass # Is # p+
+ = - (neutral)
atomic #
Boron
11
6
5
5
5
Is # p+
Mass #+ =n- (neutral)
Remember e- = p+ (to make atom neutral)
 #p+ is atomic number
 p+ + n = mass number
 Mass number – atomic number = n
Bohr Models
p+ & n in nucleus
e- in energy
levels around
nucleus
3 energy levels
-1st has up to 2e-2nd has up to 8e-3rd has up to 8e-
e-- e-e- ee e
e-- e-ee
e-- e-ee
e-- e-e e
Bohr Model of Lithium
e- ee
3
Li
6.941
3
4
Bohr Model of Argon
e- ee
e
18
e e
Ar
39.948
e- ee- e-
18
22
e--e-ee
e- eee
Electron Configuration
Shows the distribution of electrons
among the orbitals of an atom.
Describes where the electrons are
located and how much energy each
one has.
Rules for Electron Configuration
Aufbau Principle - Electrons enter orbitals
of lowest energy level first.
Pauli Exclusion Principle – An orbital can
hold a maximum of 2 electrons. To occupy
the same orbital, the 2 electrons must spin
in opposite directions.
Hund’s Rule - one electron enters each
orbital until each orbital contain one
electron with parallel spins before a second
electron is added.
Determining Electron Configurations
Quantum Numbers describe the
amount of energy in that level. The
lower the number, the less energy it
has. (n = 1, 2, 3, 4, etc.)
Sublevels are divisions of the
principle energy levels. The main
sublevels are called s, p, d and f.
Each sublevel has a different shape
caused by the different energy
levels.
Number of Electrons per Sublevel
Sublevel
Number of
Orbitals
Maximum
# of e-
s
1
2
p
3
6
d
5
10
f
7
14
s and p Orbital Electrons
d Orbital Electrons
f Orbital Electrons
Periodic Table to remember order
Sublevels (s,p,d,f) by columns - Energy levels by rows
(1,2,3,4,5,6,7 except d(row-1) & f(row-2))
X – 1s22s22p63s23p4
(16 e-)
1
2
s
Y – 1s22s22p63s23p6 4s23d104p6
5s1
s
2
1 p2 p3 p4 p5
s
p
(36
e-)
1
2
3
4
5
6
d1 d2 d3 d4 d5 d6 d7 d8 d9d10
X
Y
7
f1 f2 f3 f4 f5 f6 f7 f8 f9 f10 f11 f12 f13 f14
p6
Example 1
He (atomic # = 2) (Which means 2 p+ = 2e-)
1s2
(1=energy level; s=sublevel; 2=electrons)
He
Example 2
Li (atomic #=3)
(means 3p+ = 3e-)
1s22s1
(1,2=energy levels; s=sublevel; 2+1=electrons)
Li
Example 3
Be (atomic #=4)
(means 4p+ = 4e-)
1s22s2
(1,2=energy levels; s=sublevel; 2+2=electrons)
Be
Example 4
Si (atomic #=14)
1s22s22p63s23p2
(means 14p+ = 14e-)
(1,2,3=energy levels; s,p=sublevels; 2+2+6+2+2=electrons)
Si
Lewis Dot Structures
One more type of atomic model…
(In addition to Bohr models and electron
configurations)
Consists of the element’s symbol and the
atom’s valence electrons.
Symbol = kernel (represents the protons,
neutrons and full electron shells).
Dots = valence electrons.
Lewis Dot Structures Con’t
B
B = Kernel
(The protons,
neutrons and full
electron shells.)
Valence shell
electrons
You can use the Electron
Configuration to get the Lewis Dot
Structure…
Ca
1s22s22p63s23p64s2
Locate the highest quantum number. (4)
Add the s and p orbital electrons, and place
them around the element symbol. (2)
Ca
One Final Example
Tin
1s22s22p63s23p64s23d104p65s24d105p2
Locate the highest quantum number (5)
Add the s and p orbital electrons (4)
Sn
How to place electrons on a
Lewis Dot
First two dots represent the s orbital
electrons and are placed at the top of
the element’s symbol.
Then the p orbital electrons are
placed in this order: right, bottom,
left, right, bottom, left.
So, it goes like this…
8
5
1 2
Ne
7 4
3
6
Percent Abundance
The percentage of how much one specific
isotope of an element is found in nature.
FORMULA:
% abundance = amount of one isotope
total amount of all isotopes
Average Atomic Mass
(How the number ends up on the periodic table!!)
1st  Mass of one isotope x % abundance
in decimal form (watch SIG FIGS!!)
2nd  Do this for each isotope of that
element
3rd  Then add all individual isotopes
together to get the average atomic
mass.
1. Calculate the average atomic mass of
potassium using the following data:
Isotope
Mass
% abundance
Potassium-39
38.964 amu
93.12%
Potassium-41
40.962 amu
6.88 %
Potassium-39 38.964 amu x 0.9312 = 36.28 amu
Potassium-41 40.962 amu x 0.0688 = 2.82 amu +
Average atomic mass for K = 39.10 amu
2. Calculate the average atomic mass of
magnesium using the following data:
Isotope
Magnesium-24
Magnesium-25
Magnesium-26
Mass
23.985 amu
24.986 amu
25.983 amu
% abundance
78.70%
10.13 %
11.17 %
Magnesium-24 23.985 amu x 0.7870 = 18.88 amu
Magnesium-25 24.986 amu x 0.1013 = 2.531 amu +
Magnesium-26 25.983 amu x 0.1117 = 2.902 amu +
Average atomic mass for K = 24.31 amu