Brief Timeline of Atomic Theory
Democritus
• 400BC
• Greek philosopher
Hard Particle (Cannonball)Theory
• Proposed that they world was made up of tiny,
indivisible particles moving through a void of
empty space
• “atom” comes from the Greek word “atomos”,
meaning indivisible (cannot be divided)
John Dalton
• 1808 AD
• First modern atomic theory
Daltons Atomic Theory
1. All matter is composed of tiny, indivisible
particles called atoms
2. All atoms of an element are identical
3. Atoms of different elements are all different
4. Atoms combine in simple ratios to form
compounds
J.J. Thomson
• 1897-1904
• “Plum Pudding Model”
• Cathode Ray tube experiment
• demo
demo
Cathode Ray Tube
• Thompson showed that cathode rays
(electrons) were composed of negatively
charged particles that separated from the gas
atoms inside the tube
• Significant because: this meant that atoms
are not hard, indivisible particles. Atoms are
composed of smaller “subatomic” particles
Thomson’s Plum Pudding Model
• The atom was a hard
sphere that was
positively charged with
negatively charged
electrons that “dotted”
the atom like raisins in
plum pudding
The discovery of radioactivity
• Henri Becquerel
– 1896
– Discovered that uranium
ore released rays that could
expose photographic film
The discovery of radioactivity
• Marie & Pierre Curie
– Extracted 2 new elements
from uranium (U)ore:
radium (Ra) and
polonium (Po)
Marie Curie
Ernest Rutherford
Magnetic Field Experiment
• Was able to separate
radioactive rays into 2
types: alpha (a) & beta (B)
• Determined that a rays
were composed of helium
nuclei (He +2 charge)
Gold Foil Experiment (1911)
• Lead to discovery of the
nucleus, as a positively
charged center of atom,
containing the mass
• Most of the atom is
negatively charged empty
space, electrons are outside
the nucleus
Magnetic Field Experiment
Gold Foil Experiment
Gold Foil Experiment
Gold Foil Experiment
Gold Foil Experiment
Rutherford’s Atomic Model
Rutherford’s “Nuclear Model”
• Most of the atom is negatively charged empty
space, surrounding a small, positively charged
nucleus, containing most of the mass of the
atom
Modern Theory of Atomic Structure
• Developed by Niels Bohr, based on the science
of nuclear physics
• Bohr determined that an element's position
on the periodic table was related to its
electron configuration.
Electron configuration
Electron configuration – shows how many
electrons are in each energy level or “ring”
• Ex: Carbon 2-4
Bohr’s Planetary Atomic Model
• Niels Bohr (1922)
• Determined that electrons
rotate around the nucleus
in discrete paths or rings
Planetary Model of Atomic Structure
Wave-Mechanical Model
• Current (modern) theory of atomic structure
• Moseley used x-ray analysis to calculate an
integer for each element: these integers are
the atomic numbers
Wave-Mechanical Model
• There is a tiny, dense positively charged
nucleus at the center of a huge negatively
charged electron cloud
Wave-Mechanical Model
Orbital
• Region of probability of finding an electron
“The whole point:”
• The modern model of the atom is the result of
many investigations that have been revised
over a long period of time by many scientists
• Atomic theory song
Place the models of atomic structure in
order from earliest to the modern theory:
Basic Atomic Structure
• The nucleus occupies less than 0.01% of the
total volume of an atom but accounts for
99.97% of its mass. Thus most of an atom is
EMPTY SPACE where the ELECTRONS are
found, this is called an ELECTRON CLOUD.
• One atomic mass unit is 1/12TH THE MASS OF
A CARBON-12 ATOM. This is the standard by
which the masses of all other elements are
determined. It is abbreviated “u”.
Subatomic Particles
Use your Periodic Table to complete the following:
Atomic Number
27 Al
13
35 Cl
17
1 H
1
207 Pb
82
Mass Number
Nuclear
# of
Charge
Protons
# of Neutrons
# of electrons
The only number that never changes
for an element is
ATOMIC
NUMBER
!!
Atomic Structure 1
Name
Symbol
Atomic
Mass
Number
Number
19 F
9
Charge
# of
# of Neutrons # of electrons
Protons
0
Helium-4
0
11
23
Nitrogen-14
0
0
0
32
64 Cu
29
14
14
0
16
0
25
53
131
0
35
0
81
0
0
53
74
56
Atomic Structure 1
Name
Symbol
Atomic
Mass
Number
Number
19 F
9
Charge
# of Neutrons # of electrons
Protons
0
Helium-4
0
11
23
Nitrogen-14
0
0
0
32
64 Cu
29
Manganese-60
# of
14
14
0
16
0
60 Mn
25
25
60
0
25
35
25
137 Ba
56
56
137
0
56
81
56
Iodine-131
131 I
53
53
131
0
53
78
53
Iodine-127
127 I
53
53
127
0
53
74
53
Barium-137
Phosphorus-32
0
14 C
6
0
Potassium-39
0
16
56 Fe
26
0
8
0
18
40
0
0
79
24 Mg
12
197
29
0
0
**Shade the columns representing the nucleons light blue
35
Changes in number of subatomic
particles
Isotopes
• Change in number of
neutrons
• Same atomic number,
different mass
• Same number protons,
different number neutrons
Ions
• Change in number of
electrons
• A cation is positive ion,
results from loss of
electrons, reducing radius
• An anion is negative ion,
results from gain of
electrons, increasing radius
ISOTOPE
• Forms of the same element having different
mass due to different number of neutrons.
• Indicated by “element name-mass”
15
8O
16
8O
Name: _______________
Name: _______________
Mass: ________________
Mass: ________________
Protons: ______________
Protons: ______________
Neutrons: _____________
Neutrons: _____________
Practice:
Name
Symbol
235U
238U
Carbon-12
Carbon-13
Atomic #
Mass #
# Protons
# Neutrons
# Electrons
The mass number on the periodic table indicates the
weighted average of all the naturally occurring
isotopes of an element
To calculate a weighted average:
% X mass + % X mass
100
100
+ …..
Neon is naturally found in nature having 90.51% mass
of 20.00u, 0.24% mass of 21.00u and 9.22% mass of
22.00u. Calculate the weighted atomic mass of neon.
1.) Uranium is found naturally in nature as 3 isotopes: isotope
mass
% abundance
U-238
238.05g
99.28
U-235
235.04g
0.7110
U-234
234.04g
0.0054
Calculate the weighted average atomic mass of the elements
below. Show all work, round to the nearest hundredth.
a.)99.63%14N & 0.37%10N
b.)69.1%63Cu (actual mass of 63.93g) & 30.9%65Cu (actual mass
of 64.93g)
c.)78.9%24Mg, 10.00%25Mg & 11.01%26Mg
You can estimate which isotope is found in the highest
abundance as the one with a mass closest to the mass
listed on the periodic table
Example:
Chlorine-35 mass 34.969g
Chlorine-37 mass 36.966g
Look on the periodic table for the mass of chlorine
____________________________
The more abundant isotope has a mass closer to the
mass given on the periodic table:_____________
Practice: Which isotope of silicon would be
found in the highest percentage?
28 Si,
14
Why?
mass 27.977
29 Si,
14
mass 28.976
30 Si,
14
mass 29.974
Atomic Structure 2
Isotopic
Number of
Number of
Number of
Mass
protons
neutrons
electrons
number
4.
18
18
5.
16
1.Oxygen-16
Notation
O-16
16O
2.Oxygen-18
3.
Ar-40
6.
32
34S
7.
19
8.
19
20
41
9. Iron-
10.
57Fe
11.
12.
26
Ne-20
13.
10
14.Hydrogen15.
16.
32
22
1
H-2
3H
2.) Calculate the weighted average of the following naturally
occurring isotopes. SHOW ALL WORK!
a.) 95.50%7Li & 7.50% 6Li
d.) 99.63%14N & 0.37%15N
b.)80.20%11B & 19.80%10B
e.) 78.9%24Mg, 10.00%25Mg, & 11.01%26Mg
c.)95.02%32S, 0.75%33S, & 4.21%34S
f.) 92.23%28Si, 4.67%29Si, & 3.10%30Si
IONS
• A charged part of an atom, resulting from the
loss or gain of electrons
• VALENCE electrons: outermost electrons, the
last number in an electron configuration
• KERNEL electrons: all electrons except
valance electrons
Electron configuration
Electron configuration – shows how many
electrons are in each energy level or “ring”
• Ex: Carbon 2-4
Electron configuration of sodium:
2 diagrams of atomic structure:
Bohr diagrams
Lewis electron dot diagrams
Bohr realized that the rows on the periodic table corresponded
Lewis realized that the groups/families on the periodic table
to the number of shells of electrons
correspond to the number of valence electrons
This model shows the nucleus, indicating the number of protons
This model shows the element symbol surrounded by dots,
and neutrons, surrounded by rings, representing each energy
representing the valence electrons. You must place one dot at
level
each (3, 6,9,12 o’clock) location before “doubling up” (exception:
Helium)
18 F
9
electron configuration 2-7
F electron configuration 2-7
1
18
Bohr Atomic Structures
1
4
tables to fill in the electron
configurations, as shown, then
draw the Bohr Atomic Structure
for each element 1-20.
1
2
1
2
13
14
15
16
17
2
7
9
11
12
14
16
19
20
3
4
5
6
7
8
9
10
2-1
2-2
2-3
2-4
23
24
27
28
31
32
35
40
11
12
13
14
15
16
17
18
2-8-1
2-8-2
2-8-3
2-8-4
39
40
Rules:
19
20
2-8-8-1
2-8-8-2
1.) Show placement
2.) The nucleus is
3.) Indicate the number
of ALL electrons
represented by a center
of electrons in each
circle showing the
energy level, by writing
*use atomic #
# of protons & the
the number on each ring.
OR the entire
# of neutrons
electron configuration
** closest to nucleus is 1st
1
18
Directions:use your reference
LEWIS Electron Dot Structures
1
2
tables to fill in the electron
configurations, as shown, then
draw the Lewis Dot Structure
for each element 1-20.
1
2
13
14
3
4
5
6
2-1
2-2
2-3
2-4
11
12
13
14
2-8-1
2-8-2
2-8-3
2-8-4
19
20
Rules:
2-8-8-2
16
17
7
8
9
10
15
16
17
18
1.) Only show
2.) Electrons are represent- ex:
3.) You must place 1 dot(e -)
4.) Exception is row 1:
outermost(VALENCE)
ed as dots, placed at the
at each location before
for element #2, indicate
electrons
12
*use group #
12,3,6,9
or the last # in the
2-8-8-1
15
electron configuration
around the
element symbol.
you double up.
both electrons at the 12
o'clock location.
2 Main Types of Ions:
a
nion
A negative ion
Ex: Cl-, O-2
t
ca ion
A positive ion
Ex: Na+, Al+3
The octet rule
Atoms will gain or lose electrons in order to have
a full valence shell of 8 electrons.
Exception: Helium can have a maximum of 2
valance electrons
When an atom gains 1 or more
electrons
It becomes a negative ion and it’s radius
increases. A negative ion is an anion.
When an atom loses 1 or more
electrons
It becomes a positive ion and it’s radius
decreases. A positive ion is a cation.
CATION
ANION
Definition
Results from
Indicated by
What happens
to radius???
Na
Naming
Lewis Dot
Structure
Na+
CATION
ANION
positive ion
negative ion
Loss of electron(s)
Gain of electron(s)
(+) charge
(-) charge
Gets smaller
Gets bigger
Definition
Results from
Indicated by
What happens to
radius???
Na
Na+
“Element name-ion”
Change ending of element to “-ide”
Naming
Lewis Dot Structure
[Na]
+
..
[:.F.:]-
How to predict if an element will form
an anion or cation:
The “electron clock”:
8/0
7
1
6
2
5
3
4
# valance electrons
Atomic Structure 3: Predicting Ions
Atomic Structure 3: Predicting Ions
Atomic Structure 3: Predicting Ions
Radius
increase or
decrease?
Element
Electron
configuration
Lewis dot
structure of
atom
F
2-7
F
gain
1
-1
F
increase
Mg
2-8-2
Mg
lose
2
+2
Mg
decrease
O
Al
N
Fr
C
Lose or gain
electrons?
How many
electrons lost
or gained?
Ionic Charge
**
Lewis dot
structure of ion
Element
Electron
configuration
Lewis dot
structure of
atom
Lose or gain
electrons?
How many
electrons lost
or gained?
Ionic Charge
**
2-8-8-1
2-8-7
2-8-18-18-8-2
2-8-6
2-8-5
2-3
**In the “ionic charge” column only: shade the cation charges red and the anion charges blue
Lewis dot
structure of
ion
Radius
increase or
decrease?
Atomic Structure 4
ex
35 Cl
17
1
23 Na
11
2
9 Be
4
3
65 Zn
30
4
14 N
7
5
32 S
16
6
20 Ne
10
7
127 I
53
8
108 Ag
47
9
70 Ga
31
10
12 C
6
# of
Protons
# of
Neutrons
# of
Electrons
Nuclear
Charge
17
18
17
+17
Bohr Diagram of
Atom
Lewis Dot
of Atom
Predict
Ionic Charge
Cl
-1
Lewis Dot
of Ion
Name of Ion
Chloride
Atomic Spectra
Radiant Energy
• Energy that travels through space as
electromagnetic waves at the speed of light
Electromagnetic Spectrum
• Includes all types of radiant energy from
gamma rays (hi E) to radiowaves (lo E)
• Visible light is only a small portion of the
spectrum
1 photon = 1 quantum
Quanta: tiny packets of energy released or
absorbed by objects
*Einstein and Plank determined that energy is
released or absorbed in a continuous flow of
small packets or quantum/photons
Release or Absorption of Energy:
Higher energy levels
(excited state)
Electrons absorb energy
when jumping to
Electrons release energy
when falling to
Lower energy levels
(ground state)
Bohr used the emission spectrum as
proof of planetary model
But his model only works for hydrogen because
he didn’t account for electrons moving between
energy levels
Spectral Lines
Characteristic wavelengths (l) of photons of
energy released as electrons fall from hi to lo
energy
Spectral lines demo:
Salt of
Element
Color of Flame
Strontium Chloride
Barium Chloride
Copper (II) Chloride
Lithium Chloride
Potassium Chloride
Identity
Unknown Element
Unknown Mixture
Emission Spectrum:
Each element has it’s own
characteristic spectrum:
Compare H & He:
hydrogen
helium
Because electrons do move between
energy levels, emitting “spectral lines”,
we had to change our view of atomic
structure:
Excited State Electron Configurations
Occurs when elements absorb energy and jump
to a higher energy level.
** it will not look like it is written on periodic
table, be sure they add to the correct number!
Ground state: 2-8-1
Excited state : 2-7-2
“Crib Sheet”
• #p+ = atomic number *#n0 = mass-atomic number
• #e- = #p+ - charge (use the sign of the charge)
• Isotope: same #p+, different #no OR same atomic
number, different mass
• To calculate weighted average: (%/100 x atomic mass)
+ (%/100 X atomic mass) + …..
• *Ion: same # p+, different #e• Charge= #p+ - #e-
Atomic Structure Review p. 17
1.
2.
3.
4.
5.
6.
7.
8.
11
9
43
92
118
13
11
4
9. Br
10. C
11. Sn
12. Zn
13. Cl
14. 40
15. 16
Atomic Structure Review p. 17
16.)
17)
=(.925x7) + (.0750x6)
=(.789x24)+(.10x25)+(.1101x26)
=6.475 + .45
= 18.936 + 2.5 + 2.8626
=6.925
= 24.2986
=6.93g
= 24.30g
Atomic Structure Review p. 18
18.) 2-8-1
19.) Na
20.) 2-7-2
21.) 19
22.) 1
23.) Y
24.) Ar
25.) Not possible
27.) as electrons fall from excited
state to ground state energy is
released as radiant energy
(spectral lines).
28.) you can ID the gas element
using spectral line analysis.
29.) electrons are negatively
charged particles. B has 5 e-, its econfig. is 2-3, with 2 e- in the 1st
energy level and 3 e- in the 2nd
(valence) level
Atomic Structure Review MC?s
1.) 2
2.) 4
3.) 1
4.) 1
5.) 3
6.) 1
7.) 4
8.) 2
9.) 4
10.) 3
11.) 4
12.) 3
13.) 1
14.) 4
15.) 1
16.) 3
17.) 4
18.) 3
19.) 2
20.) 2
21.) 3
pg 19-20
1.) 4
2.) 3
3.) 2
4.) 3
5.) 2
6.) 3
7.) 1
8.) 2
9.) 3
10.) 1
11.) 3
12.) 1
13.) 2
14.) 4
15.) 3
16.) 3
17.) 2
18.) 1
19.) 1
20.) 4
21.) 4
22.) 2
23.) 3
pg 21-22
Atomic Structure Review p. 23
1.) 19p, 20n, 18e
3.) 5p,6n,2e
5.) 16p,16n,18e
7.) 7p,7n,10e
9.) 37p,48n,36e
11.) 30p,35n,28e
2.) 9p,10n,10e
4.) 15p,16n,18e
6.) 14p,14n,10e
8.) 20p,20n,20e
10.) 53p,75n,54e
12.) 6p,6n,10e