Download Unit 1: Atoms, Molecules, and Ions

Unit 3:
Atoms, Molecules, and Ions
Development of Atomic Theory

It took over 2000 years
from the first proposal of
an “atomic” universe
before actual experimental
evidence had been
accumulated!
Matter according to the Greek Philosophers
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492-375 BC: Empedocles
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supported by Aristotle
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Earth-Air-Fire-Water
400 BC: Zeno (also supported by Aristotle)
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Matter can be infinitely divided
400 BC: Leucippus
470-375 BC: Democritus (student of Leucippus)
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Matter is made of indivisible particles called atoms
Atomic Structure Revolution

1500 Robert Boyle The Skeptical Chemist
He performed detailed experiments with gases and
began the break down of the Greek model of matter.

1700 Antoine Lavoisier
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Law of Conservation of Mass
Lavoisier used an enclosed container to study chemical and physical
changes. After very careful measurements, he concluded that mass
is neither created nor destroyed during these processes.
1790 – 1800 Joseph Proust
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Joseph Proust developed
the Law of Constant
(Definite) Proportions
Compounds are composed
of elements
Proust determined that a
given compound always
contains exactly the same
proportion of elements,
by mass.

For example, water is
composed of a constant
proportion of 1 part
hydrogen to 8 parts
oxygen, by mass.
1800 John Dalton

Proposed the Law of
Multiple Proportions
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If two elements combine
to form different
compounds, the mass
ratios of the two elements
in the compounds can be
expressed as a ratio with
simple whole numbers.
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For example, hydrogen and
oxygen react to make two
different compounds:
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In water, 8 g of oxygen
reacts with 1 g of hydrogen
In hydrogen peroxide, 16 g
of oxygen reacts with 1 g of
hydrogen.
The second mass ratio (16:1)
is exactly double the first
ratio (8:1)
Since water is H2O,
Hydrogen peroxide must be
H2O2
2
Law of Multiple Proportions, cont’d

Copper forms two different oxides. An analysis of
the two compounds shows that in the first compound
63.5 g of copper combines with 16.0 g of oxygen,
and in the second compound, 10.0 g of copper
combines with 5.04 g of oxygen.
Show that the two substances obey the law of
multiple proportions.
Dalton’s Atomic Theory

John Dalton proposed explanations for the laws
of mass conservation, definite & multiple
proportions:
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An element is composed of tiny particles called
atoms.
All atoms of a given element are identical and
atoms of different elements are different.
Atoms of different elements combine in ratios of
small whole numbers when forming compounds.
Chemical reactions only rearrange the way atoms
are combined; the atoms themselves are not
changed.
1897 J.J. Thomson

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J.J. Thomson is studied
mysterious “cathode rays”
in a cathode ray tube
(CRT).
He discovered that these
cathode rays carried a
negative charge, and that
they had mass
His conclusion: cathode
rays are really streams of
negatively charged
particles, “electrons”
Cathode Ray Tubes
Thomson’s Conclusions
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Thomson saw that the
cathode ray beam could be
deflected using either an
external magnetic or
electric field
Based on the direction of
the beam deflection, he
concluded it carried
negative charge
J.J. Thomson

He determined the “charge
to mass ratio” for an
electron:
e/m = 1.758819 x 108 C/g

This is important because if
either charge or mass can
be determined then the
other can be calculated.
Thomson’s Atomic Model
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Thomson’s discovery
showed that there was
matter even smaller than
atoms.
After discovering the
electron, J.J. Thomson
proposed a “plum
pudding” or “raisin bun”
model of the atom
He envisioned a positively
charged “dough” with
negatively charged
electrons scattered
throughout.
1909 Robert Millikan
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Determined the charge
on a electron using an
oil drop experiment.
After he charge the oil
drop then he balanced it
between a negative and
a positive plate.
e=1.602177 x 10-19 C
m=9.101390 x 10-28 g
1911 Earnest Rutherford
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Rutherford studied the “new”
phenomenon of radioactivity,
discovered by Henri
Bequerel.
In 1898, Rutherford
discovered 3 different types
of radioactivity: a, b, g
He later used the alpha
particles in a famous
experiment that ended with
the discovery of the “nuclear
atom”
Next: Shockwave Animation
Rutherford’s Gold Foil Experiment
The Gold Foil Experiment
Thomson’s Model?
Nuclear Model?
"It was as if you fired a 15-inch shell at a sheet of tissue paper
and it came back to hit you."
- E. Rutherford
Gold Foil
Rutherford’s Nuclear Model
If an atom was as big as a football
stadium, the nucleus would be the
size of an ant on the 50-yard line!
1932 James Chadwick
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The last of the three main
subatomic particles was
discovered in 1932 by
James Chadwick
Worked with Rutherford
looking for an uncharged
subatomic particle, but
failed
In 1932, he revisited some
earlier experiments and
was successful – he
discovered the neutron!
Thanks,
Mr. Chadwick!
Three Important “Subatomic” Particles
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The electron was discovered first by Thomson – it is negatively charged
and the lightest of these three subatomic particles.
The proton was deduced by Rutherford. It is positively charged and has
a mass of approximately 1 amu.
The neutron was the last particle to be discovered (by James Chadwick)
in 1932. It has approximately the same mass as a proton, but is neutral.
Important Vocabulary…

Atomic Number (Z)
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The number of protons in
the nucleus of an atom
Also equal to the number of
electrons around the nucleus
if the atom is neutral
Determines the identity of
an element
Mass Number (A)

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The number of protons &
neutrons (i.e. “nucleons”)
within the nucleus of an
atom
Isotopes
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Atoms of the same element
(with the same # of protons),
with different numbers of
neutrons (i.e. different mass
numbers)
Carbon-14 is an isotope that
has 14 protons and neutrons
Atomic Mass

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The mass of an atom.
May be measured in
“grams”
More conveniently
expressed in “atomic mass
units”, amu
Three Isotopes of Hydrogen
Two Isotopes of Sodium
Understanding Isotopes

How many protons,
electrons and neutrons are
there in a neutral atom of
phosphorus-32?
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15 Protons (Z)
17 Neutrons (32 – 15)
15 Electrons (neutral)

How many protons,
electrons, and neutrons are
in a neutral atom of
potassium-39?
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19 protons (Z)
20 neutrons (39 – 19)
19 electrons (neutral)
The Mass Spectrometer
An instrument that separates and analyzes particles based on their masses.
Pictured below is the separation of neon gas into its three isotopes.
The Mass Spectrum of Neon

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A computer displays a
graph of Abundance vs
Mass Number for neon gas
The three peaks suggest
three different isotopes:
20Ne, 21Ne and 22Ne.
The area under each peak
represents the “abundance”
or the “fraction” of the
neon gas contributed by
each isotope
Average Atomic Mass & Natural Abundances

Chromium exists as four
stable isotopes. Use the
information below to
calculate the average atomic
mass of chromium.

Mass
49.946 amu
51.941 amu
52.941 amu
53.939 amu

Abundance
4.35%
83.79%
9.50 %
2.36%

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Naturally occurring lithium
consists of two isotopes
with atomic masses of
6.015 and 7.016 amu.
The average atomic mass
of lithium is 6.941 amu.
Calculate the natural
abundance of each isotope.
Ans. 7.42% & 92.58%
Atomic Mass Units

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The atomic mass unit is
DEFINED as one-twelfth
the mass of the nucleus of
an atom of carbon-12.
A proton or neutron has a
mass of approximately 1
amu
1 amu = 1.66054 x 10-27 kg
Originally defined with
respect to hydrogen, and
then oxygen.
1922 Niels Bohr
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Electrons closer to the nucleus have lower
energy and are more stable due to strong
electrostatic attraction with the nucleus.
The letter “n” is used to designate energy levels
(orbits), with lower whole number values of n
representing lower energy orbits closer to the
nucleus.
The key to Bohr’s Quantum Model was that
electrons are restricted to certain “allowed”
energy levels in atoms.
Whereas Rutherford suggested electrons orbit
the nucleus, Bohr’s model required that they
occupy only certain orbits!
The Bohr Atom
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The first electron orbit
holds a maximum of 2
electrons
The second orbit holds a
maximum of 8 electrons
The third orbit may hold up
to 18 electrons
The nth orbit can hold up to
2n2 electrons!
Draw a Bohr diagram for Carbon-14
6p
8n
Draw a Bohr diagram for Oxygen-15
8p
7n
Evolution of the Atomic Model


Bohr’s model of the atom is very useful for understanding
elementary concepts in bonding and chemical reactions
However, it is fundamentally flawed and he soon helped
replace it with the “Quantum Mechanical Model” of the atom
which no longer views the electron as simply a “particle”, but
acknowledges and depends on the wave-nature of the electron
also.
Energy Level Diagrams


Very similar to Bohr diagrams, without the “orbits”
The energy level diagram for neutral fluorine is…
7__e2__e-
9p
F
The electrons in the highest
energy level are called the
VALENCE ELECTRONS …
so a neutral fluorine atom
has 7 valence electrons!
Ions

Remember…
ELECTRONS are
gained and lost …
nothing happens
to the protons!
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During many chemical reactions,
atoms gain or lose electrons.
Since the # of electrons no longer
equals the # of protons, these atoms
must be charged. A charged atom is
called an ion.
Cations are positively charged ions –
they have lost electrons (and now
have more protons than electrons)
Anions are negatively charged ions –
they have gained electrons (and now
have more electrons than protons)
Predicting Ion Charges
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Atoms are most stable when their valence
orbit (the outermost orbit) is FULL of
electrons.
The Noble Gas family has full valence
energy levels and these elements are
highly un-reactive.
The other elements would like to achieve
a Noble Gas electron configuration.
For example, a fluorine atom gains 1
electron to become a fluoride anion, FA sodium atom loses 1 electron to
become a sodium cation, Na+
F
Na
Predicting Ion Charges
Count the valence electrons, add or subtract electrons
(whichever is the smaller number) to achieve a full valence level.
Note that metals tend to lose electrons (become cations)
1+
n/a
While nonmentals gain electrons (become anions)!
2+
3+ ?? 3- 2- 1-
1869 Dmitri Mendeleev
In 1869 Mendeleev and Lothar
Meyer (Germany) published
nearly identical classification
schemes for elements known to
date. The periodic table is
based on the similarity of
properties and reactivities
exhibited by certain elements.
Later, Moseley (1912)
established that each elements
has a unique atomic number,
which is how the modern
periodic table is organized.
Periodic Table: Metals & Nonmetals
What are the properties of metals and nonmetals?
Where are the “semimetals” (metalloids)?
1
IA
1
2
IIA
13
IIIA
14
IVA
15
VA
18
VIIIA
16
VIA
17
VIIA
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
Metals
10
11
IB
12
IIB
Nonmetals
Metals & Nonmetals

Properties of Metals…
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Lustrous (shiny)
Malleable
Ductile
Conductors of heat &
electricity
Tend to lose electrons
(become cations)
Note: only Iron, Nickel,
Cobalt are strongly magnetic
metals

Properties of Nonmetals…
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Dull
Brittle
Insulators
Tend to gain electrons
(become anions)
Note: Semimetals
(metalloids) are elements
with properties of BOTH
metals and nonmetals.
Representative, Transition, & Rare Earth Elements
Periods: Across the Periodic Table
1
IA
1
18
VIIIA
2
IIA
13
IIIA
2nd Period
2
3
3
IIIB
4
IVB
5
VB
6
VIB
4
5
6th Period
6
7
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA
Groups: Down the Periodic Table
1
IA
1
18
VIIIA
Alkali Family:
1 e- in the valence shell
2
IIA
13
IIIA
14
IVA
15
VA
16
VIA
2
3
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
4
5
6
7
Halogen Family:
7 e- in the valence shell
17
VIIA
Chemical Families of the Periodic Table
Halogen
Noble Gas
Chalcogens
Alkali
Alkaline
(earth)
1
IA
1
18
VIIIA
2
IIA
13
IIIA
Transition Metals
2
3
4
5
6
7
3
IIIB
4
IVB
5
VB
6
VIB
7
VIIB
8
9
VIIIB
10
11
IB
12
IIB
14
IVA
15
VA
16
VIA
17
VIIA