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H/Chemistry
Atomic Theory
Atomic Theory
Essential Questions:

What are we made of?

How are scientific models developed?

Do atoms exist or are they just concepts invented by scientists? What evidence is there in
your everyday life for the existence of atoms?
I.

How did the understanding of the atom affect historical events?

How have historical events affected the model of the atom?

What do we think the atom “looks like” now?

If the atom is mostly empty space, why doesn’t my butt fall through the chair?

How are light and electrons related?

How do we “see” where electrons are located in the atom?

Why is the location of electrons so important?
Historical Background
A.
The Greek philosophers
Democritus (460-370 BC) – “atomism”
Aristotle (384-322 BC) and other philosophers –earth, air, fire and water, denied the
existence of atoms
No experiments
B.
On the way to chemistry – alchemy and phlogiston
Alchemy – transmutation of other metals into gold (~ Middle Ages and before)
Phlogiston – an imaginary element, believed to separate from every combustible
body in burning (1700s)
C.
The first experimental chemists – focus on gases
Henry Cavendish (1731-1810) – discovered hydrogen
Joseph Priestley (1733-1804) - oxygen
Antoine Lavoisier (1743-1794) - oxygen
Karl Wilhelm Scheele (1742-1786) - oxygen
Count Amedeo Avogadro (1776-1856) - gases  mole
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H/Chemistry
Atomic Theory
Inventions
Which of the following were invented
a) before 1800?
D.
b) between 1800 and 1900?
c) after 1900?
Three Laws
1.
Law of Conservation of Mass Antoine Lavoisier (1770s)
Mass of reactants = mass of products (relates to before and after reaction)
2.
Law of Definite Proportions
Joseph Proust (1799)
Compounds always contain elements in the same proportion by
mass, no matter how they are made or where they are found.
3.
Law of Multiple Proportions
John Dalton (1803)
Different compounds made from the same elements:
The ratio of mass of an element in the first compound relative to
the mass of the same element in a second compound is a fixed
whole number.
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H/Chemistry
E.
Atomic Theory
Dalton’s Atomic Theory
1. All matter is made of extremely small particles called atoms.
2. All atoms of a given element are identical (mass, physical and chemical
properties).
3. Atoms of different elements have different masses, and different physical and
chemical properties.
4. Different atoms combine in simple whole number ratios to form compounds.
5. In a chemical reaction, atoms are combined, separated, or rearranged in
chemical reactions.
6. Atoms cannot be created, divided into smaller particles or destroyed.
II.
Subatomic Particle Discovery
A.
Electron was the first discovered!
1897 – J.J. Thomson
B.
1.
Experiment: cathode rays
a. passed electricity through a partially-evacuated tube of gas
b. observed a ray of light crossing from one electrode to the other
c. the ray of light moved a paddle wheel inside the tube
2.
Conclusions:
a. Since ray moved toward (+) electrode, the ray must be (-) charged.
b. Since the paddle wheel moved, the ray must be made up of particles.
c. Since Thomson performed the experiment with different gases, with the
same results, the particles must be in all atoms.
Nucleons
1.
Protons: 1909/1910 – Rutherford (specialty = radioactivity)
(Geiger and Marsden performed the experiments; Rutherford interpreted
them)
Alpha particles (He nucleus, + charge) shot through gold foil were deflected in
peculiar ways, inconsistent with the plum pudding model of the atom
 area of concentration of positive charge
2.
Neutrons: 1932 – James Chadwick
(worked with Hans Geiger, then Ernest Rutherford)
Uncharged particles in the nucleus with mass were pushed out of beryllium
when bombarded with alpha particles. These particles accounted for the
“missing mass” in the nucleus.
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H/Chemistry
III.
Atomic Theory
Subatomic Particles
Location
Charge
Relative
mass
Symbol
A.Electrons
B.Nucleons
(1) Protons
(2) Neutrons
Atomic Mass
Atomic number = # protons (also = # electrons)
Mass number = # protons + # neutrons
# neutrons = mass number – atomic number
Isotopes
# protons
# neutrons
# electrons
Hydrogen-1
(protium)
1 H
1
1
0
1
Hydrogen-2
(deuterium)
2 H
1
1
1
1
Hydrogen-3
(tritium)
3 H
1
1
2
1
POGIL – Isotopes
POGIL - Ions
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H/Chemistry
Atomic Theory
Average Atomic Mass
1 atomic mass unit (amu) = 1/12 mass of a C-12 atom
Calculate average atomic mass by using weighted averages which take into account the relative
abundance of each isotope.
Note that what is on the periodic table is the average atomic mass, not the mass of a single
isotope.
POGIL – Average Atomic Mass
Atomic #, Mass # and Average Atomic Mass
IV.
Atomic #
Mass #
Average Atomic
Mass
Weighted average of
all isotopes
# p+
# p+ + #n0
whole #
Whole number
Decimal, limited sfs
found on PT
NOT on PT
Found on PT
Nuclear Chemistry and Radioactivity
Nuclear reaction  change in the identity of the element(s)
Radioactivity = radiation emitted spontaneously by certain elements whose nuclei contain an
unstable neutron:proton ratio
Smaller elements
Stable neutron:proton ratio = 1 n0:1 p+
for masses < 20 (i.e. mass # = 2 x atomic #)
Largest elements
Stable neutron:proton ratio = 1.5 n0:1 p+
for largest atoms
All elements with atomic #s > 83 are radioactive
Unstable nuclei  emit radiation and change their identities (“radioactive decay”)
Historical figures:
Wilhelm Roentgen (1845-1923) - discovered X-rays –1895
Henri Becquerel (1852-1908) - discovered radioactivity in U - late 1800s
Ernest Rutherford (1871-1937) – id’d different types of radiation (beg. 1898)
Pierre (1867-1934) and Marie Curie (1859-1906) - discovered radium and polonium – 1898;
first used the term “radioactivity”
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H/Chemistry
Atomic Theory
Types of radiation
1. Alpha radiation (most common in elements with atomic # > 83)
Alpha particles = 2 p+ + 2 n0 (He nucleus, 42He, ) with 2+ charge
226 Ra
88
e.g.

222 Rn
86
+
4 He
2
(+ energy)
2. Beta radiation (most common in elements with high n0:p+ ratio)
Beta particles = 1 e- (0-1) with 1- charge
Neutron  proton + beta particle
1 n
0
e.g.

1 p
1
14 C
6
+

0 
-1
14 N
7
0 (+
-1
+
energy)
Note: The sum of the mass #s and atomic #s on both sides of the equation are the same.
3. Gamma radiation
Gamma rays = high-energy radiation with no mass and no charge (00
- usually accompany alpha and beta radiation
e.g.
238 U
92

234 Th
90
+
4 He
2
+ 2 00
Nuclei with lower neutron:proton ratios than optimal:
4. Positron Emission
 more neutrons by converting a proton into a neutron
(most common in lighter elements with low n0:p+ ratio)
Positron = particle with same mass as an e-, but opposite charge
Proton  neutron + positron
1 p  1 n + 0 
1
0
1
e.g.
11 C
6

11 B
5
+
0 
1
5. Electron Capture
(most common in elements with a high n0:p+ ratio)
 more neutrons by pulling in an e- which combines with a proton to form a neutron
Proton + electron  neutron
1 p + 0 e  1 n
1
-1
0
e.g.
0 e
-1
+ 8137Rb 
81 Kr
36
+ X-ray photon
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H/Chemistry
Atomic Theory
Radioactive Particles WS
Positron
same mass as e-’s
Electron capture
electrons
(Added to the reactants side)
0 
+1
0 
-1
1/1840
1/1840
1+
1-
1. Which radioactive emission has the greatest mass? __________________________
Least mass? _______________________________
2. Why do you think gamma rays are drawn as wavy lines? __________________________
_______________________________________________________________________
3. To which charged plate are the alpha particles attracted? Explain.
______________________________________________________________________
4. To which charged plate are the beta particles attracted?
_______________________________________________________________________
Why do the beta particles have a greater curvature than the alpha particles?
______________________________________________________________________
5. Explain why the gamma rays do not bend toward one of the electrically charged plates.
______________________________________________________________________
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H/Chemistry
Atomic Theory
Nuclear Fission and Fusion
Nuclear fission = the splitting of a nucleus into smaller, more stable fragments, accompanied by a
large release of energy
e.g. Uranium-235:
235 U
92
+
1 n
0
 23692U 
(unstable)
92 Kr
36
+
141 Ba
56
+ 3 10n + energy
The new neutrons (10n)  fission of more U-235 (= chain reaction, a self-sustaining process)
Chain reaction requires a critical mass (= minimum amount of starting material to maintain a
chain reaction);
supercritical mass may  violent nuclear explosion
results in radioactive waste
Practical examples = nuclear power plant, atomic bomb
Nuclear Fusion = the process of binding smaller atomic nuclei into a single larger and more stable
nucleus, requiring a huge amount of energy to initiate, followed by a large release of energy
1. Creation of natural elements
a. Hydrogen, other light elements - from the Big Bang
b. Elements #2-92 (except Fr, Pr, Te, At)
Nuclear fusion occurs in stars (naturally)
Occurs in hydrogen bomb (artificially) > 2 x 107oC
The sun converts 3 x 1014 g of H into He every second.
4
1
4
H  He + energy
1
2
Mass is not conserved.2Mass is converted into energy via E = mc2
Other fusion reactions occur in the sun:
4He + 4He  8 Be + 
4
2
2
4He + 8Be 122C + 
4
6
2
2
-8-
(gamma ray)
H/Chemistry
Atomic Theory
2. Synthetic Elements
a. Nuclear bullets
i. Bombard nuclei of elements with small particles such as p+, n , 4 He ( particles)
2
& e- (0-1 particles)
ii.
Used to create Elements # 93-100
17
14
iii. 1919 first experiment: N + 4 He  O + 1 H
8
7
2
1
2
(Rutherford)
2
b. Crashing nuclei
i.
Accelerators hurl nuclei into each other at very high speeds.
12
244
254
1
e.g. 6 C + 96 Cm 102 No + 2 0 n
carbon
2curium nobelium neutron
ii. Elements beyond # 100
iii. These elements are very unstable:
e.g. Element 109 existed for only 3.4 x 10-3 sec (3 atoms)
c. Superheavy elements (“transuranium” elements)
Stability of nucleus of atom depends on filling "shells" within nucleus with alternating p+ and
n. The more filled shells, the more stable it would be.
e.g. Element 114
244 Pu
94
+
48 Ca
20

289
114Fl
+ 3 10n (1999, Russia)
Production of the Transuranium Elements
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H/Chemistry
Atomic Theory
1. Does the diagram illustrate a natural transmutation reaction or an induced transmutation reaction?
______________________________________________________________________________
2. What is the name and nuclear symbol of the isotope produced in the reaction?
____________________________________________________________
3. Write nuclear equations to show how dubnium-263, lawrencium-262, and seaborgium-266 can be
produced from a nuclear reaction of neon-22 and americium-244.
_____________________________________________________________________________
_____________________________________________________________________________
_____________________________________________________________________________
5. Each of the radioisotopes in the table decays within 20 seconds to 10 hours.
Write a nuclear equation for each decay.
_______________________________________________________________________________
_______________________________________________________________________________
________________________________________________________________________________
________________________________________________________________________________
6. Which, if any, of the four isotopes listed in the table would you expect to find at Earth’s surface? Why?
__________________________________________________________________________
Nuclear equations
Complete the following equations :
 42He + _____
214
83Bi
239
93Np
 23994Pu + ______
Write a balanced nuclear equation for the alpha decay of americium-241.
Write a balanced nuclear equation for the beta decay of bromine-84.
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H/Chemistry
Atomic Theory
Next Steps in Discovering the Structure of Atoms – Properties of Electrons
• Wave nature of light – EMR (James Maxwell, 1864)
• Particle nature of light – quantum (Max Planck, late 1800s)
• Emission of light and other EMR from heated elements  emission spectra
V. Electromagnetic Radiation (EMR)
A.
B.
Definition: energy that exhibits wave-like (or oscillating) behavior as it travels through
space
1.
James Maxwell in 1864 – unified the electric and magnetic forces 
electromagnetic force
2.
Speed of EMR is always the same: c = 3.00 x 108 m/s
3.
e.g. light, microwaves, TV, radio, X-rays
 MEMORIZE
Remember:
Waves have wavelength,  (lambda) in nm and frequency,  (nu) or f
# waves/sec = Hz (hertz) = cycle/sec or s-1
** c = 
 MEMORIZE
note: and  are inversely proportional.
Electromagnetic Radiation Practice Problems from Textbook (pp. 121, 124)
c = 
Microwaves are used to transmit information. What is the wavelength of a microwave having a
frequency of 3.44 x 109 Hz? (8.72 x 10-2 m)
1.
What is the frequency of green light which has a wavelength of 4.90 x 10-7 m?
(6.12 x 1014 s-1)
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H/Chemistry
Atomic Theory
2.
An X-ray has a wavelength of 1.14 x 10-10 m. What is its frequency?
(2.63 x 1018 s-1)
3.
What is the speed of an electromagnetic wave that has a frequency of 7.8 x 10 6 Hz?
(3.00 x 108 m/s)
4.
A popular radio station broadcasts with a frequency of 94.7 MHz.
What is the wavelength of the broadcast? (1 MHz = 106 Hz) (3.17 m)
The Idea of the Quantum
 Quantum = the smallest discrete amount of energy that can exist independently, esp. as EMR
 1 quantum of light = 1 photon
E = h
h = Planck’s constant = 6.626 x 10-34 J.s
(Energy emitted = Planck’s constant x frequency of light hitting it)
 there is a direct relationship between E and 
 The amount of energy in EMR increases with increasing frequency:
E = hPlanck’s constant = h = 6.626 x 10-34 J·s
Example: Tiny water drops in the air disperse the white light of the sun into a rainbow. What is the energy
of a photon from the violet portion of the rainbow if it has a frequency of 7.23 x 1014 s-1? (4.79 x 10-19 J)
5.
What is the energy of each of the following types of radiation?
a. 6.32 x 1020 s-1 (4.19 x 10-13 J)
b. 9.50 x 1013 Hz (6.29 x 10-20 J)
c. 1.05 x 1016 s-1 (6.96 x 10-18 J)
6.
Name the types of radiation in each part of #5.
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H/Chemistry
Atomic Theory
Electromagnetic Radiation Practice Problems
1.
What is the frequency of EMR with a wavelength of 235 pm? What type of EMR is this?
2.
What is the frequency of EMR with a wavelength of 0.614 cm? What type of EMR is this?
3.
What is the wavelength of EMR with a frequency of 8,512 Hz? What type of EMR is this?
4.
What is the wavelength of EMR with a frequency of 625 x 10 17 Hz? What type of EMR is this?
5.
If the speed of light is 3.00 x 108 m/s, calculate the wavelength of the electromagnetic radiation whose
frequency is 7.500 x 1012 Hz.
6.
Determine the frequency of light with a wavelength of 4.257 x 10-7 cm.
7.
For the following sources, calculate the missing member of the wavelength/frequency pair.
a)
FM radio waves with a frequency of 94.7 Hz.
b)
A laser with a wavelength of 1064 nm.
c)
An X-ray source, emitting X-rays with a wavelength of 175.4 pm.
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H/Chemistry
VI.
Atomic Theory
Bohr’s model
Neils Bohr (Danish physicist)
– studied with both Thomson and Rutherford
– refined Rutherford’s model and published his own model in 1913
– incorporated Planck’s idea of quanta of energy
– received the Nobel Prize in 1922
– provided an explanation for the spectral lines of hydrogen
POGIL – Electron Energy and Light
1.
Electrons are arranged in circular paths, “orbits” around the nucleus.
2.
Electrons have fixed energy levels which prevent them from falling into the nucleus.
Electrons closest to the nucleus have lowest Etotal = KE + PE (most stable)
3.
To move from one energy level to another, electrons must gain or lose energy.
EMR is emitted from the atom when electrons fall down to a lower energy level.
4.
The distance between energy levels is not equal.
A “quantum” is the amount of energy needed to make the leap.
5.
Bohr’s model did not explain the spectra of atoms with > 1 electron.
Definitions related to Spectra
Spectrum = whole range of related qualities [Latin: appearance, from specere – to view]
Electromagnetic spectrum = all EMR arranged according to 
Emission = any radiation of energy by EM waves
[Latin: emitto – to send out, to utter, to hurl, to set free]
Emission spectrum = the spectrum into which light or other EMR from any source can be
separated
Continuous spectrum = a spectrum whose source emits light of every  in a continuous band
Bright-line spectrum = pattern of bright lines on a dark background. Source = glowing gas that
radiates in special ’s characteristic of the chemical composition of the gas
Line spectrum
_______E2________
Energy given off  each photon (quantum of light)
emitted has a characteristic  which contributes a
line to the spectrum.
_______E1________
Each substance (element) has a set of ’s as distinct
as a fingerprint.
Visible
Balmer Series
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e- drop to n=2
o
n
t
r
i
H/Chemistry
Atomic Theory
VII.
Quantum Theory of the Atom (current model)
A.
Quantum Mechanics
Louis de Broglie: discovered the wave nature of electrons
Erwin Schrödinger, 1926: complex mathematical equation which treats electrons as waves
 foundation of quantum theory
B.
1.
Uses probabilities (90%)
2.
Heisenberg Uncertainty Principle: we cannot simultaneously measure an e -‘s
position and its velocity
Orbitals and Quantum #’s
1.
All probable locations of an electron = orbital in the shape of an “electron cloud”
2.
Each e- in an atom has its own set of four quantum #’s: n, l, m, s
3.
Principal energy levels (n) correspond to Bohr’s orbits.
Total # e- in one principal energy level = 2n2, where n = principal energy level
4.
Sublevels (l), “magnetic” (position in space: m) and spin (s) are additions and take
into account the observed energies of the e-s.
Orbitals
1. What is the shape of an s orbital? ___________________________________________________
2. What is the relationship between the size of an s orbital and the principal energy level in which it is
found?
______________________________________________________________________________
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H/Chemistry
Atomic Theory
3. What is the shape of a p orbital? How many p orbitals are there in a sublevel?
______________________________________________________________________________
4. How many electrons can each orbital hold? _____________________
5. Look at the diagrams of the p orbitals. What do x, y, and z refer to?
______________________________________________________________________________
6. How many d orbitals are there in a given sublevel? How many total electrons can the d orbitals in a
sublevel hold?
______________________________________________________________________________
7. Which d orbitals have the same shape? _____________________________________________
8. What point in each diagram represents an atom’s nucleus? ______________________________
9. How likely is it that an electron occupying a p or a d orb ital would be found very near an atom’s
nucleus? What part of the diagram supports your conclusion?
_____________________________________________________________________________
_____________________________________________________________________________
Principal Energy
Level (n)
(or “Shell”)
3-D (m)
(position in
space)
x, y, z
only 1
orientation
3 orientations
for p: px, py, pz
etc.
Spin (s)
+ ½
- ½
Total # e-
1
Sublevel (l)
(or “Subshell”)
shape of probability
cloud
s
2
s, p
3
s, p, d
4
s, p, d, f
2(4)2 = 32
5
s, p, d, f
2(5)2 = 50
6
s, p, d, (f)
2(6)2 = 72
1 e- + ½
1 e- - ½
Sublevels
e-
(# orbitals in each sublevel)
x 1 orbital
= 2 e- total
A.
s:
holds 2
B.
p:
holds 2 e-
x
3 orbitals
= 6 e- total
C.
d:
holds 2 e-
x
5 orbitals
= 10 e- total
D.
f:
holds 2 e-
x
7 orbitals
= 14 e- total
?
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2n2
2(1)2 = 2
2(2)2 = 8
2(3)2 = 18
H/Chemistry
VIII.
Atomic Theory
Electron Configuration – how electrons are distributed within an atom
Three Rules for Filling Orbitals
1.
Aufbau: Fill in order of increasing energy levels.
2.
Hund’s Rule: Fill all orbitals at same energy level with at least 1 e -, before adding the
second e-.
3.
Pauli Exclusion Principle: only 2 e- per orbital, of opposite spin
1. What does each small box in the diagram represent? _____________________
2. How many electrons can each orbital hold? _________________
3. How many electrons can the d sublevel hold? __________________
4. Which is associated with more energy: a 2s or a 2p orbital? ______________
5. Which is associated with more energy: a 2s or a 3s orbital? ______________
6. According to the Aufbau Principle, which orbital should fill first, a 4s or a 3d orbital? _________
7. Which orbital has the least amount of energy? __________________
8. What is the likelihood that an atom contains a 1s orbital? ___________________________
9. Sequence the following orbitals in the order that they should fill up according to the Aufbau
Principle: 4d, 4p, 4f, 5s, 6s, 3d, 4s.
_____________________________________________________________________________
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