Transcript The Mole - My CCSD

Chapter 3:
The Atom
and the Mole
(with nuclear)
The investigation and understanding of the
atom is what chemistry is all about!
Topics rearranged from your text, pages 62-117.
We come here to be philosophers, and I hope you will always
remember that whenever a result happens, especially if it be new,
you should say, “What is the cause? Why does it occur?” and you
will, in the course of time, find out the reason. -Michael Faraday
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The Mole
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The “mole” represents a number of
things….like a dozen.
How many things is a mole?
6.022137 x 1023… we use 6.02 x1023.
This is Avogadro’s number
– named for a lawyer, Amadoe Avogadro, that
studied molecular gasses as a hobby.

When you have three moles of atoms, you
have (3 x 6.02x1023 =) 1.81x1024 atoms total.
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Recall: Parts of the atom
-1
(subatomic particles)
+1
0
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Proton  charge = +1, mass = 1
Neutron  charge = 0, mass = 1
Electron  charge = -1, mass = 0
In a normal, neutral, unreacted
atom, the number of electrons Ions have more or
equals the number of protons. less electrons than
protons
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They have a charge
Atomic History

Greek philosopher Democritus (400BC)
– coined the term atomon which means “that which
cannot be divided.”

John Dalton (1803) a colorblind chemist.
– Among his interests, Dalton was very interested in a
scientific explanation for his colorblindness the
behavior of gasses.

In his A New System of Chemical Philosophy,
Dalton published five principles of matter.
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Dalton’s Top Five

All matter is made of indestructible and indivisible
atoms.
– (atoms are hard, unbreakable, the smallest thing there is)
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Atoms of a given element have identical physical and
chemical properties.
– (all atoms of X will behave the same anywhere)

Different atoms have different properties.
– (X behaves differently than Y)

Atoms combine in whole-number ratios to form
compounds.
– (two H’s and one O = Water (H2O)
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Atoms cannot be divided, created or destroyed,
– (just rearranged in chemical reactions).
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The Laws:

Constant Composition
– Ratios of atoms in a compound is constant for that
compound.
H 2O hydrogen-oxygen atomic ratio = 2:1
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Conservation of Mass
– Mass is not created or destroyed in a chemical
reaction.

Multiple Proportions:
– Since atoms bond in small, whole number ratios to
form compounds, the ratio of their mass ratios are
small whole numbers.
CO
Oxygen-carbon mass ratio = 1.33
x2
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CO2
Oxygen-carbon mass ratio = 2.66
Conservation of Mass
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Multiple Proportions
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The Cathode Ray Tube

The cathode ray tube
– A new invention suggested the presence of charges
– areas of positive and negative…charge.
– This suggested that atoms must be divisible, and
Dalton’s theory had to be modified.
Electrostatics
 J. J. Thomson (1897)
– English Physicist proposed that the atom is a sphere
of positive charge with small areas of negative
charge.
– This theory become known as the “plum pudding”
model after an English “dessert” of purple bread
and raisins.
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Millikan’s Oil
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Thompson used electrostatics experiments to
determine the electron’s charge-to-mass ratio.
Robert Millikan’s (1909)
– oil-drop experiment allowed the charge of a single
electron to be determined: 1.60 x 10-19 C.

Scientists calculated the mass of an electron to
be 1/2000 of the mass of a proton!
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Ernest Rutherford

Ernest Rutherford (1910)
– New Zealander Physicist, while studying radioactive
elements, found that radioactive alpha particles
deflected when fired at a very thin gold foil.
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The Gold Foil Experiment
– the atom was not a hard sphere but
– was mostly space, with a small concentration of
positively-charged mass (the nucleus).
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Link
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to experiment
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Niels Bohr
The Bohr Model
– A Danish physicist (and student of Rutherford)
rebuilt the model of the atom placing the electrons
in energy levels.

Bohr was one of the founders of quantum
chemistry:
– energy can be taken in and given off in small
packets or quanta of specific size.
– When a specific amount of energy was added to an
atom, an electron could jump into a higher energy
level.
No more…no less!
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Adding the Neutrons

James Chadwick (1932)
– British physicist, proved there was too much mass
in the nucleus
– Suggested the existence of massive, neutral
particles in the nucleus. (neutrons)
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The Modern Model
Dalton’s atom
electron
Thompson’s electrons
neutron
Rutherford’s space and
nucleus
proton
Bohr’s energy levels
(not to scale)
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Chadwick’s neutrons
Elements

112 known elements
– 92 of which are naturally occurring.
– 93 through 112: transuranium.
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Each has an atomic symbol.
Atomic number
8
O
OXYGEN
15.9994
– is number of protons
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Atomic mass
– is the total mass of the protons plus
the neutrons.
Notice that the atomic mass is not a round number,
even though protons and neutrons each have a
mass of 1. This is due to natural abundance.
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Natural Abundance - Isotopes
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Isotopes:
– Each element may have several isotopes
– Isotopes differ in the number of neutrons.
Example:
– the element carbon has 6 protons, but it could have
5, 6, 7, or 8 neutrons, to form Carbon-11, Carbon12, Carbon-13, and Carbon-14. 11C, 12C, 13C, 14C
In nature, there is a mix of different natural
isotopes.
We use this mix to calculate average atomic
mass…
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Calculating Average Atomic Mass
isotope 1 :  relative abundance (%) x isotope mass (amu)
isotope 2 :  relative abundance (%) x isotope mass (amu)
isotope 3 :  relative abundance (%) x isotope mass (amu)
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Sum of the products = average atomic mass
Example:
– The isotopes of element “Bob” are found below:
– Bob-18.0, 25.0%
0.25x18  0.60x19  0.15x20
– Bob-19.0, 60.0%
18
.
9
amu
– Bob-20.0, 15.0%
– What is the average atomic mass of naturally
occurring Bob?
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1 amu = 1.66x10-27kg
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Isotopes
– atoms of the same _______
– different number of _______
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Review …
Ions
– atoms of the same _______
– different number of _______
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Allotropes
– forms of the same _______
– bonded in different _______
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Quanta / Quantum
– Packets of energy of _______ size
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Atomic Mass
– Is the _______ of all _______ found in nature.
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Molar Mass

Molar mass
– expressed in grams per mole (g/mol)
– mass of one mole of a substance.
– link between the atom and the gram.
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No more AMU:
AMU  Molar mass
(we can measure)
The average atomic mass of carbon is 12.01.
What is the mass of a mole of carbon atoms?
What is the molar mass of Copper, Cu?
Cu  63.5 g / mol
What is the molar mass of Nitrogen, N2?
Nx2  14.0 x 2  28.0 g / mol
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Molar Mass Practice
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Determine the Molar Mass of the following
elements and compounds:
Ca  40.41g / mol
Ca
Clx 2  35.5 x 2  71.0 g / mol
Cl2
CaCl2
Ca  Clx 2  40.1  35.5 x 2  111.g / mol
H2O
Hx 2  O  1x 2  16  18 g / mol
Ba(OH)2
Ba  (O  H ) x 2  137  (16  1) x 2  171g / mol
FeSO4
Fe  S  Ox 4  55.8  32.1  16 x 4  152 g / mol
Al2(SO4)3
Alx 2  ( S  Ox 4) x3  27 x 2  (32.1  16 x 4) x3  342 g / mol
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Mole-to-Mass Conversions
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Sodium has an atomic mass of 23 g/mol. How
many moles do you have in 115 grams?
Use a T-chart!
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115 grams  23g / mol  5.0 moles
How many grams are equal to 3.5 moles of
CaCl2? 3.5moles  111g / mol  389 grams
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What is the mass of 0.46 moles of SiO2?
.46mol  60.1g / mol  27.7 grams SiO2
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Mole-Mass Conversion Practice
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Complete the following mole-to-mass conversions:
Mass in grams of 2.25 moles of iron, Fe?
126 grams Fe Use your periodic table to find molar mass
Mass in grams of 0.375 moles of potassium, K?
14.7 grams K
Number of moles in 5.00 grams of calcium, Ca?
0.125 moles Ca
Number of moles in 3.60x10-10 grams of gold,
Au?
1.83x10-12 mol Au
Bires, 2010
End of
Chapter 3
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Mole-Atoms Conversions
Phew!
Mole = 6.02x1023 things, how many atoms are in:
23
24
3
.
0
x
6
.
02
x
10

1
.
8
x
10
atoms
3.0 moles of silver, Ag?
21
0.010 moles of copper, Cu?
6.0 x10 atoms
How many moles do you have in:
24
2
.
4
x
10
24
2.4x10 atoms of helium, He?
 4.0 moles
6.02 x1023
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3.0x1023 atoms of lithium, Li? 0.50 moles
How many moles do you have in 222 grams of
222 grams / 63.55 g / mol  3.49 moles
copper?
How many atoms in 127.1 grams of copper?
127.1grams / 63.55g / mol  2.00 moles  6.02 x1023  1.20 x1024 atoms
Bires, 2010
Isotopes – Nuclides - Radioactivity

Nuclides
mass
– the nucleus of an isotope
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Place the mass above the
charge as seen here.
Nuclides undergo decay:
charge
– transformation into different
nuclides
– Balanced nuclear reactions
– “Radioactive”
– Half Life: time to decay ½
(mass) of a sample
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Images from ChemZone
Alpha Decay

Alpha Decay
mass
– a helium nucleus is released.

Alpha particles:
– move very slowly
– because of their size, can be
blocked with a few pages of
paper or human skin
– cause ionization (damaging!)
– are positively charged
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charge
This is a Nuclear Equation
Alpha Decay occurs
in all elements with
atomic number
above 83.
Images from ChemZone
Beta Decay

Beta Decay
– An electron is ejected
from the nucleus

Beta particles
– move fast
– can penetrate thick lowdensity materials
– but can be blocked with
concrete and metals
Beta Decay occurs when
– are negatively charged
a nucleus has a high
neutron-proton ratio.
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Images from ChemZone
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Gamma Decay
Gamma Decay
No
mass
– High energy photons (gamma rays) are given off.

Gamma rays
– given off as the “spare change” during other
radioactive decays….
– extremely penetrating and powerful. Several
inches of lead is required to slow these particles
down to a stop.
– Don’t get included in nuclear equations.
No
charge
Summary of
three basic
particle decays
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Images from ChemZone
Nuclear Equations Practice
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Sodium-24 undergoes alpha decay
24
11
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Complete the following nuclear equations:
Na  ...
24
11
Na He  ...
4
2
24
11
(help on click)
Na He  F
4
2
20
9
Iodine-131 undergoes beta decay
I  e Xe
131
53
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0
1
131
54
Tungsten-190 undergoes alpha decay
W  He  Hf
190
74
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4
2
186
72
Uranium-238 undergoes alpha decay, then two
beta decays (3 steps)
U  24He  234
90Th
238
92
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Th 10 e 234
91Pa
234
90
234
91
Pa 10 e 234
92 U
Nuclear Fission

Nuclear fission:
– splitting of large, unstable atoms
– releases large amounts of energy

Critical mass (or critical density)
– amount of fissionable fuel needed before
reaction will begin.
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Once fission
begins, it is
Uncontrolled, nuclear fission proceeds to difficult to stop.
completion with great speed.
Nuclear Weapons:
– two half-spheres of fissionable material are
compressed together with conventional
explosives, creating the critical mass.

In order to harness nuclear fission to
create useable electricity, we slow down
the process with control rods…
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Nuclear Fusion
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Nuclear Fusion:
– Joining of smaller nuclei to form
larger nuclei.
– Releases far more energy that
nuclear fission.
– Easier to control than fission.
The sun’s (stars) energy comes
from the fusion of hydrogen atoms
into helium atoms.
The H-bomb is a fusion weapon.
 Fusion:
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– The power supply of the future?
– Why don’t we use it now?
2
1
H  H  He
2
1
4
2
E=mc2

Einstein: Energy and mass are interchangeable.
– E = Energy
– m = mass
– c = speed of light ( 3 x 108 m/s )
– Very small amounts of mass create large amounts
of energy!

We use the formula ΔE= Δmc2 to build new,
artificial elements in supercolliders
(particle accelerators.)
Fermilab cyclotron, Argonne National Laboratory, Chicago
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