Download Unit 3 – History of Atomic Theory

Unit 3 – History of Atomic Theory
Democritus of Abdera (c. 460 – c. 370 BCE)
A greek philosopher that believed all matter was composed of discrete, indivisible
particles called ατομωσ (atomos). The more accomplished philosophers (Aristotle and
Plato) did not accept this theory. They believed all matter was made of four elements:
earth, air, water and fire.
Antoine Lavoisier (1743-1794)
Following extensive studying of combustion and the discovery of carbon, dioxide,
oxygen, nitrogen and hydrogen gases, Lavoisier regarded measurement as the essential
operation of chemistry. Through careful experiments, he weighed the reactants and
products of various reactions and suggested the LAW OF CONSERVATION OF MASS
Joseph Proust (1754-1826)
Proust furthered Lavoisier’s experiments by weighing the ratios of reactants in
specific chemical compounds. He found that for specific compounds the ratio of their
combining was constant. This led to the LAW OF DEFINITE PROPORTION.
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John Dalton (1766-1826)
Dalton developed the first atomic model of the atom that offered explanations for
the simple laws that govern chemistry. solid sphere model
His theory has three main postulates.
1. An element is composed of tiny particles called atoms. All atoms of a given
element show the same chemical properties. *This is not true today because of
the discovery of neutrons and isotopes.
2. Atoms of different elements have different properties. In an ordinary chemical
reaction, no atom of any element disappears or is changed into an atom of
another element.
3. Compounds are formed when atoms of two or more elements combine. In a
given compound, the relative numbers of atoms of each kind are definite and
constant. In general, these relative numbers can be expressed as integers or
simple fractions. This led to the LAW OF CONSTANT COMPOSITION.
Henri Becquerel (1788-1878)
Becqueral accidentally found that a piece of a mineral containing uranium could
produce an image on a piece of photographic paper in the absence of light. He attributed
this phenomenon to a spontaneous emission of radiation by the uranium, which he called
radioactivity. Studies demonstrated three types of radioactive emission: gamma (γ) rays,
beta (β) particles and alpha particles (α) particles. A gamma ray is essentially high energy
“light” (no mass and no charge); a beta particle is a high speed electron (essentially no
mass and a negative charge); and an alpha particle is essentially a helium nucleus (heavy
mass = 4 and 2+ charge). More models of radioactivity have been discovered but the alpha
particles were used in early crucial experiments.
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J J Thomson (1856 – 1940)
Thomson did the first important experiments that led to the understanding of the
composition of the atom. He studied electrical discharges in partially evacuated tubes
called cathode ray tubes.
Thomson found that when high voltage was applied to the tube a “ray” (he called it the
cathode ray) was produced. Due to the fact that the ray emanated from the negative pole of
the electric field the ray was composed of negatively charged particles. This led to the
discovery of electrons. Also from this discovery he calculated the charge-to-mass ratio of
the electron. He reasoned that all atoms contain electrons and that to be electrically
neutral, atoms must also contain positive charges. Thomson postulated that an atom
contained a diffuse cloud of positive charge with negative electrons embedded in it. This is
often called the plum pudding or raisin bun model of the atom because the electrons are
like raisins dispersed in the dough (the positive charge).
plum pudding model
spherical cloud of
positive charge
electrons
s
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Max Planck (1858-1947)
Planck studied the radiation profiles emitted by solid bodies that were heated to
incandescence. From these experiments, he determined that the results could not be
explained by the typical (Newtonian) physics of the time. The loss or gain of energy could
only be obtained by whole number multiples of the quantity, hν, where h is Planck’s
constant (6.626 x 10-34 J·s/particle) and ν (nu) is the frequency of the electromagnetic
radiation. A small “packet” of energy was called a quantum (plural, quanta).
Robert Millikan (1868-1953)
Millikan performed experiments with charged oil drops. These experiments allowed him
to determine the magnitude of the electron charge. Using this value and the charge-tomass ratio determined by Thomson, Millikan was able to determine the mass of an
electron.
Oil drop experiment
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Ernest Rutherford (1871-1937)
Rutherford, who also performed many experiments exploring radioactivity, carried
out an experiment to test Thomson’s model of the atom. The experiment directed alpha
particles at a thin sheet of gold foil. He reasoned that if Thomson’s model were accurate,
the massive alpha particles would crash through the foil with little deflection. Although
most of the particles passed straight through, many of the particles were deflected at large
angles and a very few were reflected back at him. This led to three conclusions:
1. The atom is mostly empty space. Most alpha particles went straight through.
2. The atom has a positively charged center or nucleus. Due to the fact that some
particles came close to the center and the deflection path was more similar to an
electric field of like charges.
3. The nucleus is the concentration of the atom’s mass. The very few that were
deflected were on a collision course with a much more massive center of charge.
Gold Foil Experiment
If Thomson was correct with his model, then a completely different outcome would have
occurred. This Gold-Foil Experiment created the nuclear model of the atom, with the
positively charged nucleus at the center.
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James Chadwick (1891-1974)
Capitalizing on the theories proposed by Rutherford and others, Chadwick produced
experimental evidence of the existence of the atomic particle that had a mass nearly equal
to that of a proton but no charge. This was a result of alpha particle bombardments of
beryllium, lithium and boron. He named this particle the neutron according to the
equation: 9Be + α → 12C + 1n
Louis de Broglie (1892-1987)
Utilizing Albert Einstein’s special theory of relativity (energy has mass) and the
incorporations of Planck’s equation, de Broglie supported the view that light could not only
be quantized but that it also could possess wave properties as all matter can. This is called
the wave-particle duality of light. It depends upon the nature of the quantity you are
investigating. Light, which was previously thought to be purely wavelike, was found to
have particle (masslike properties) and the opposite is also true. The relationship is called
de Broglie’s equation λ = h/mv, where λ is the wavelength (in meters), h is Planck’s
constant, m is the mass of the particle (in kilograms), and v is the velocity of the particle (in
meters/second).
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Niels Bohr (1885-1962)
As Rutherford’s graduate assistant, Bohr held the planetary model of the atom and
incorporated the work of Planck and de Broglie to propose that the electrons must be
balanced by the attraction for the nucleus to resist flying off the atom. However a
constantly accelerating particle (like the electron) should lose energy and then eventually
fall into the nucleus. Due to this apparent contradiction, Bohr applied a theory to the
hydrogen atom that claimed that energy could only occur in specified increments. These
increments were called energy levels and could also have only whole number values.
Electrons in a normal state (ground state) could gain energy and be promoted to a higher
level (excited state) but eventually that energy could not be sustained and would be
emitted as electromagnetic radiation (visible light in some cases). The specific lines of
radiation emitted were consistent with observed emission spectra. This lead to the
quantized planetary model of the (hydrogen) atom and Bohr’s equation: E = -B/n2, where E
is energy of radiation, B is Bohr’s constant, and N is the energy level. Also energy changes
can be calculated for the hydrogen atom: ΔE = Ehi - Elo = -1312 kJ/mol (1/nhi2 - 1/nlo2)
where “hi” represents the excited level and “lo” represents the ground state. Bohr’s
findings were in agreement with 0.01% of observed results. However, when investigating
helium or higher “multi-electron” atoms the percent error grew exponentially.
By the mid – 1920’s it became obvious that Bohr’s model could not be made to work for
these higher atoms. Three physicists, Werner Heisenberg, Louis de Broglie and Erwin
Schrödinger proposed a new branch of physics to help explain the contradictions. This
new branch of physics is called quantum mechanics.
Erwin Schrödinger (1887 - 1961)
Schrödinger gave greater emphasis to the wave properties of the electron and the
probability of finding an electron in a certain volume of the atom. This led to Schrödinger’s
equation that derived certain values for describing the likelihood of locating an electron at
certain radii from the nucleus and describing the electron’s properties in each of these
volumes in 3-D space.
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Schrodinger’s equation for the x-axis:
𝑑2 𝛹
𝑑𝑥 2
+
8𝜋2
ℎ2
(E-V)Ψ
= 0
Solving for all three axes leads to three quantum numbers to explain the location and
properties in multi-electron atoms.
Werner Heisenberg (1901-1976)
Heisenberg stated that due to the physical methods of detection of the electron and
small particles it is impossible to be certain of both its location and momentum according
to the equation: Δx · Δ(mv) ≥ h/4π. This is Heisenberg’s Uncertainty Principle.
Quantum Mechanical Model (current model) of the atom shows the “electron cloud”
as a probability distribution.
Noyau = nucleus
Probabilité de présence de l’électron = probability of the presence of the electro
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Development of Atomic Models
scientist
discovery
Model (name)
Model (picture)
Atom
John Dalton
(atomic theory)
Solid sphere
ee-
e-
e-
e-
JJ Thomson
Electron
e-
Plum pudding
e-
e-
ee-
e-
ee-
e-
ee-
e-
e-
Ernest Rutherford
Nucleus
Nuclear
ee-
eee-
e-
Niels Bohr
Current Model
(Schrodinger,
Heisenberg,
deBroglie)
Quantized energy
levels
QUANTUM
THEORY
-
location of e is a
probability
e-
Planetary
e-
Quantum
Mechanical
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Development of Atomic Theory Studyguide:
Democritus:
 “atomos” (indivisible particles)
 “thought” about it (philosopher – no proof)
Dalton:
 Atomic theory (had experimental evidence)
1. All matter made of small particles called atoms
2. Elements cannot disappear or change in ordinary chemical reactions
3. Elements combine in small whole number ratios to create compounds

solid sphere model
Thomson:
 Cathode ray tube
 Discovered the electron
 plum pudding model
Millikan:
 oil drop experiment
 What two things did he determine? charge and mass of electron
Rutherford:
 gold foil experiment
 What did he think would happen? Alpha particles would pass straight through gold foil.

What really happened? Most alpha particles passed straight through some were deflected
or reflected back.

Discovered the nucleus
(created the nuclear model)
Quantum Theory – extremely small particles traveling at velocities close to the speed of light
Planck: loss or gain of energy in small "packets" called quanta
Bohr:
 Quantized energy levels
 Definite orbits
 planetary model
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deBroglie:
 wave-particle duality
 duality of nature
 light can act like a particle (photon)/electron can act like energy (wave)
 λ = h/mv (matter has wavelength)
Einstein:
 photoelectric effect
 called light particles photons
Current Model: probability distribution (quantum mechanical or electron cloud model)
Heisenberg:
 Heisenberg uncertainty principle: impossible to know location AND momentum of
electron at the same time
Schrödinger
 Wave Equation: described probability of location of electron
Date of Atomic Theory Quiz: _______________________
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Warm-up:
Write the name of each model underneath the picture that represents it:
solid sphere (Dalton), plum pudding (Thomson), nuclear (Rutherford), planetary (Bohr), quantum mechanical
(current)
planetary (Bohr)
nuclear (Rutherford)
solid sphere (Dalton)
(add e-)
plum pudding
(Thomson)
quantum mechanical
Atomic Theory Matching:
1.
e Millikan
2.
b, j Thomson
3.
a, g Dalton
4.
f, k Rutherford
5.
h, l Bohr
6.
d Schrödinger
7.
8.
a.
atomic theory
b.
discovered electron
c.
uncertainty principle
d.
wave equation
e.
oil drop experiment
f.
gold foil experiment
g.
solid sphere model
h.
planetary model
i.
wave-particle duality
j.
plum pudding model
k.
discovered the nucleus
l.
quantized energy levels
c Heisenberg
i deBroglie
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The Atom – Modern Atomic Theory
Subatomic Particles:
symbol
Proton
location
Relative mass
(amu)
1
Actual mass (g)
nucleus
Relative
charge
+1
nucleus
0
1
9.11 x 10-31
e- cloud
-1
0
1.67 x10-27
1.67 x10-27
+
p
Neutron
0
n
Electron
-
e
In a neutral atom, the number of protons is always the same as the number of electrons.
Why does the atom not fall apart? protons (+) attract electrons (-) opposites attract
How does the mass of the electron compare to the mass of the proton?
1/1000
negligible
Heisenberg Uncertainty Principle: cannot simultaneously know the location and momentum of
the electron within an atom
The nucleus makes up most of the mass of an atom. Most of the atom is empty space
The electron cloud makes up most of the volume of an atom.
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Atomic Numbers:
Atomic Number – number of protons in an atom – determines the atom’s identity (located above
element on periodic table.)
Henry Moseley found that each element has a different number of protons in its nucleus.
T or F: The atomic number of an element can change. Why or why not? It would be a different
element with a different number of protons.
An individual atom is electrically neutral, which means the number of protons equals the
number of electrons.
So, an element’s atomic number also indicates the number of electrons in a neutral atom.
Ions:
When an atom gains or loses electrons, it acquires a net electrical charge and is called an ion.
Anion: negatively charged ion (gained electrons - nonmetals)
Cation: positively charged ion (lost electrons - metals)
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Examples of Ions:
1. A magnesium ion loses 2 electrons, what is the charge on a magnesium ion? Mg2+
2. A fluorine atom gains 1 electron, what is the charge on the fluorine ion? F-
3. An iron atom loses 3 electrons, what is the charge of the iron atom? Fe3+
Isotopes:
All atoms of a given element have to have the same number of protons in their nucleus, but not
necessarily the same number of neutrons.
One isotope of each element is more common than another.
In nature, elements are found as a mixture of isotopes.
For example, in a drop of water, 99.9844 % of the hydrogen isotopes are hydrogen – 1.
isotope name
# of protons
# of neutrons
# of electrons
* hydrogen
1
0
1
(hydrogen - 1)
deuterium
1
1
1
(hydrogen – 2)
tritium
1
2
1
(hydrogen – 3)
* indicates the most common isotope of H (look at its average atomic mass)
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The major difference between isotopes of the same element is their mass. The more neutrons
an element has, the higher the mass of that element is.
Mass Number – sum of the protons and neutrons in an atom whole number
Atomic Mass – average mass of all the isotopes of a particular element
Isotope – atom of an element that has a different number of neutrons in the nucleus
Isotopes are named by their mass numbers.
For example: All chlorine nuclei have 17 protons in their nuclei, but some have
18 neutrons
Mass # = 35
and some have
20 neutrons
Mass # = 37
What is the most common isotope of chlorine? chlorine-35
Examples:
1. The chlorine isotope with 17 protons and 18 neutrons is called: chlorine-35
2. The chlorine isotope with 17 protons and 20 neutrons is called: chlorine-37
3. The symbol for chlorine – 35 is:
35
Cl
4. The symbol for chlorine – 37 is: 37Cl
* mass # on top and atomic # on bottom
Because Atomic # = # of protons
and
Mass # = # of protons + # of neutrons
How do you find the # of neutrons in an isotope? Mass number – atomic number
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WS 3-2: Subatomic Particles
Complete the following table.
Element
Symbol
Atomic
Number
Mass
Number
Protons
Neutrons
Electrons
1.
fluorine
F
9
19
9
10
9
2.
iron
Fe
26
56
26
30
26
3.
aluminum
Al
13
27
13
14
13
4.
potassium
K
19
39
19
20
19
5.
Bromine
Br
35
80
35
45
35
6.
Gold
Au
79
197
79
118
79
7.
Silver
Ag
47
108
47
61
47
8.
Silicon
Si
14
29
14
15
14
9.
Titanium
Ti
22
47
22
25
22
10.
Manganese
Mn
25
55
25
30
25
11.
Hydrogen
H
1
2
1
1
1
12.
Chromium
Cr
24
50
24
26
24
13.
Tin
Sn
50
118
50
68
50
14.
Nickel
Ni
28
58
28
30
28
15.
iodine
I
53
127
53
74
53
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3.3 Atomic Particles Worksheet
1. Complete the following table:
symbol
number of
protons in
nucleus
75
number of
neutrons in
nucleus
number of
electrons
net charge
33
42
30
3+
52
76
54
2-
S
16
16
16
0
Tl+
81
123
80
1+
78
117
78
0
As3+
128
Te2-
32
81
195
Pt
2. Give the protons and neutrons in the nucleus of each of the following neutral elements:
a. plutonium-238 94, 144
g. lead-207 82, 125
b. copper-65 29, 36
h. europium-151 (Eu) 63, 88
c. nitrogen-15 7, 8
i. cobalt-60 27, 33
d. hydrogen-3 (also called tritium) 1, 2
j. chromium-54 24, 30
e. chromium-52 24, 28
k. silver-107 47, 60
f. helium-4 2, 2
l. silver-109 47, 62
3. Complete the following table:
symbol
atomic number
238
mass number
number of
neutrons in
nucleus
number of
electrons
U
92
238
146
92
Ca2+
20
40
20
18
40
51
V3+
23
51
28
20
89
Y3+
39
89
50
36
Br1-
35
79
44
35
31
P
15
31
16
18
14
C
6
14
8
6
79
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WS 3-4: Isotopes
Part I: Complete the following table.
(You need to complete the nuclear symbol as well.)
Nuclear Symbol
Hyphen notation
Protons
Neutrons
Electrons
1.
40
Ca
calcium - 40
20
20
20
2.
42
Ca
Calcium-42
20
22
20
3.
43
Ca
Calcium-43
20
23
20
4.
44
Ca
Calcium-44
20
24
20
5.
46
Ca
Calcium-46
20
26
20
6.
16
O
oxygen-16
8
8
8
7.
18
O
oxygen - 18
8
10
8
8.
108
Ag
Silver-108
47
61
47
9.
80
Br
Bromine-80
35
45
35
10.
207
Pb
Lead-207
82
125
82
11.
27
Al
Aluminum-27
13
14
13
12.
32
S
Sulfur-32
16
16
16
Fe
Iron-56
26
30
26
Se
selenium - 78
34
44
34
iron - 55
26
29
26
13.
56
14.
78
15.
55
Fe
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Part II: The isotopes of sulfur are sulfur-32, sulfur-33, sulfur-34, and sulfur-36. Write the nuclear
symbol of each.
32
S, 33S, 34S, 36S
Part III: Answer each question. Show your work.
1. What is the atomic mass of silicon if 92.21% of its atoms have mass 27.977 amu, 4.70% have
mass 28.976 amu, and 3.09% have mass 29.974 amu?
.9221(27.977) + .0470(28.976) + .0309(29.974) = 28.086 amu
2. Determine the average atomic mass for magnesium using the following information.
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Mg
78.99 %
23.985 amu
.7899(23.985) = 18.946
25
Mg
10.00 %
24.986 amu
.1000(24.986) = 2.499
26
Mg
11.01 %
25.983 amu
.1101(25.983) = 2.8607
24.305 amu
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3. Neon has two isotopes. Neon-20 has a mass of 19.992 amu and neon-22 has a mass of
21.991 amu. In an average sample of 100 neon atoms, 90.0% of them are neon-20 and
10.0% of them are neon-22. Calculate the average atomic mass of neon.
.900(19.992) + .100(21.991) = 20.192 amu
4. What is the atomic mass of hafnium if, out of every 100 atoms, 5 atoms have mass 176.0
amu, 19 atoms have mass 177.0 amu, 27 have mass 178.0 amu, 14 have mass 179.0 amu, and
35 have mass 180.0 amu?
.05(176.0) + .19(177.0) + .27(178.0) + .14(179.0) = 178.55 = 178.6 amu
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Radioactivity
As atomic number increases (protons increase), stable nuclei will have more neutrons than
protons. For the larger elements, a point is reached beyond which there are no stable nuclei.
Nuclei with more than 83 protons (after bismuth) are all unstable and will eventually break up
into smaller pieces.
Definition: Radioactvity = nucleus so unstable it “falls apart” we call this radioactive decay
There are three types of radioactive decay:
1. Alpha (α) nucleus emits an alpha particle
An alpha particle is a helium nucleus (two protons and two neutrons)
2.
Beta (β)
a.
Beta-minus decay: a neutron turns into a proton by emitting an electron and
an antineutrino (changes to the next element on the periodic table)
b.
Beta-plus decay: a proton turns into a neutron by emitting an positron and a
neutrino (changes to the previous element on the periodic table)
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3.
Gamma (γ) = nucleus changes from a higher energy state to a lower energy state
The EMR emitted is a gamma ray.
Half-life: (t1/2) is the amount of time required for the amount of something to fall to half its
initial value (radioactive half-life)
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1. What is the half-life of the material in the graphic below? 2 days
2. Can you predict how much material will be left after 6 half-lives?
1.25 (counts per minute)
Unit 3 TEST DATE: _____________________________
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