Download Mass/Mole Conversions

Chapter
3
The Atom
The Atom
• Democritus (450 B.C.)
proposed that all matter is
made up of tiny, indivisible
particles. (atomos)
3-2
Laws
• Law of Conservation of
Mass: mass is neither
created nor destroyed during
ordinary chemical rxns or
physical changes.
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Laws
• Law of Definite Proportions:
compounds contain the same
elements in exactly the same
proportions by mass regardless of
the size of the sample or the source
of the compound.
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Laws
• Law of Multiple Proportions: If two or
more different compounds are
composed of the same two elements,
the ratio of the masses of the second
element combined with a certain mass
of the first element is always a ratio of
small whole numbers.
3-5
Daltons Atomic Theory
In 1808 John Dalton, an English
schoolteacher, came up with an
atomic theory to explain these
laws. Many of the tenets of his
theory still hold true today.
3-6
Dalton’s Atomic Theory
~ Each element is composed of atoms.
~ Atoms of a given element are identical, and
different than those of any other element.
~ A given compound forms by combination of
two or more different atoms, always in the
same relative numbers and kinds of atoms.
~ Atoms are neither created nor destroyed in
any chemical reaction, only rearranged.
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The Structure of the Atom
• Atom: the smallest
particle of an element that
retains the chemical
identity of that element.
3-8
The Discovery of the Electron
• Thomson (1897) an English
physicist who discovered
electrons using his famous
cathode ray tube experiment.
He determined the charge to
mass ratio.
3-9
The Discovery of the Electron
The Discovery of the Electron
• Millikan (1909) an American
physicist who determined the
charge of electrons using his
famous oil drop experiment. The
mass of an e- is approximately
1/2000 the mass of an atom.
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The Discovery of the Electron
The Discovery of the Electron
Based on these discoveries, two inferences
were made:
1. Because atoms are electrically neutral, they
must contain a positive charge to balance
the negative electrons.
2. Because e- are so light in mass compared
to atoms, atoms must contain other
particles that account for most of it’s mass.
3-13
Plum Pudding Atomic Model
Discovery of the Nucleus
1911, Ernest Rutherford
conducted experiments
with radioactive materials
that released only
positively charged alpha
particles….
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Discovery of the Nucleus
Discovery of the Nucleus
• Rutherford concluded atoms have a
dense core with a positive charge.
Rutherford’s Atomic Model:
Discovery of the Nucleus
• Rutherford’s student, Neils Bohr, came
up with a way to explain the location of
e- in the atom:
Discovery of the Nucleus
The electron cloud is the current accepted model of the
atom.
The Atom
• Except for H, all atomic nuclei contain
protons and neutrons.
• A proton has a + charge equal in
magnitude to the neg. charged e-.
• Atoms contain equal numbers of p+ and
e-.
• Neutrons are electrically neutral.
• P+ and no have almost identical
masses, electrons weigh 1836 times
less.
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The Atom
• Nuclear Forces: short range p+ to p+,
p+ to no and no to no forces hold the
nuclear particles together.
• Atomic Number (Z): the number of
protons in each atom of a particular
element.
☺The atomic number identifies the element!
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Isotopes
• Isotopes: atoms of the same element that
have different masses due to different
numbers of neutrons.
• Mass Number: the total number of protons
and neutrons that make up the nucleus of an
isotope
~ Isotopes are written with the mass number
written after the element name or symbol with
a hyphen: ex. Uranium-235 or U-235
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Isotopes
Uranium-235 or U-235
Mass number – atomic number = number of neutrons
235 (protons + neutrons) – 92 protons = 143 neutrons
This info could also be portrayed using a
nuclear symbol:
235U
92
Isotopes
~ In nature, elements are almost
always found as a mixture of
isotopes
~ Isotopes have essentially identical
chemical properties
~Isotopes with more neutrons have a
higher mass and are often descibed
as “heavy”
3-24
Isotopes
• Nuclide: a general term for a specific
isotope of an element.
Atomic Mass
• Atomic Mass Unit (AMU): one amu is
exactly 1/12 of the mass of a carbon-12
atom.(p+ = 1amu, no = 1 amu, e- = 0 amu)
• Average Atomic Mass: the weighted
average of the atomic masses of the
naturally occurring isotopes of an
element.
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Calculating the Average
Moles
• Mole: the amount of a substance that
contains as many particles as there are
atoms in exactly 12g of carbon-12.
Moles
• Avogadro’s Number: the number of
particles in exactly one mole of a pure
substance ~ 6.02214179 x 1023 (we’ll
use 6.02 x 1023)
Moles
602,200,000,00
0,000,000,000,
000
If you had 6.022 x 1023 pennies and gave away
1 million a day to every person on earth, it
would take you 3000 years to distribute all
your money!!
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Molar Mass
• Molar Mass: the mass of one mole of a
pure substance.
~ the mass of single atoms or molecules
is measured in amu’s. The mass of a
mole of the same substance is
numerically the same, with the units
g/mol.
Ex. H20 = H x 2 = 1.01 x 2 =
2.02
+ O x 1 = 16.00 x 1 = +16.00
18.02
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Molar Mass
1) What is the molar mass of BaCl2?
2) What is the molar mass of NaI?
1) Ba = 1 x 137.33 g/mol = 137.33
Cl = 2 x 35.45 g/mol = 70.90
208.23 g/mol
2) Na = 1 x 22.99 g/mol = 22.99
I = 1 x 126.90 g/mol = 126.90
149.89 g/mol
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Molar Mass
This photograph
shows one mole of
solid (NaCl
58.44g/mol), liquid
(H2O 18.02 g/mol),
and gas (N2 28.02
g/mol).
3-33
Mass/Mole Conversions
When given the number of
moles, you can find the grams
by:
Moles x _g__ = grams
mole
Ex. 5.0 moles of H2O = X g
5.0 moles x 18.02g = 90. g H2O in 5.0
mol
moles
3-34
Mass/Mole Conversions
Moles x _g_ = moles
mol
Now try these problems:
3) 8.32 moles of barium chloride equals
how many grams?
4) 20.1 moles of sulfur dioxide equals
how many grams?
3-35
Mass/Mole Conversions
3) 8.32 moles of barium chloride equals how
many grams?
8.32 moles BaCl2 x 208.23 g/mol = 1730 grams
BaCl2
4) 20.1 moles of sulfur dioxide equals how
many grams?
20.1 moles SO2 x 64.07 g/mol = 1290 g SO2
3-36
Mass/Mole Conversions
Mass/Mole Conversions
When given the amount in grams, you
can calculate the number of moles by:
g x mol = moles
g
Ex. 11.2 g NaCl = X moles
11.2 g NaCl x 1.0 mol NaCl = 0.191 mols NaCl
58.44g
3-38
Mass/Mole Conversions
g x mol = moles
g
Now try these problems:
5) 50.56 g of sodium chloride
equals how many moles?
6) 329.8 g of ammonia equals how
many moles?
3-39
Mass/Mole Conversions
5) 50.56 g of sodium chloride equals how
many moles?
50.56 g NaCl x mole = 0.8652 moles NaCl
58.44g
6) 329.8 g of ammonia equals how many
moles?
329.8 g NH3 x mole = 19.35 moles NH3
14.03 g
3-40
Particle/Mole Conversions
You can also calculate between
moles and number of particles:
(1.0 moles = 6.02 x 1023 particles)
To enter this number into your
calculator, punch in 6.02 EE button
(one time) 23.
3-41
Particle/Mole Conversions
Ex. 2.59 moles of marble (CaCO3) contains how many
molecules?
2.59 mol CaCO3 x 6.02 x 1023 molecules = 1.56 x1024
1.0 mol CaCO3 molecules
*Particles can be molecules, atoms or formula units
7) How many molecules are in 5.0 moles of carbon dioxide?
5.0 mol CO2 x 6.02 x 1023 molecules =
1.0 mole CO2
3.01 x 1024 molecules CO2
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Particle/Mass Conversions
Ex. What is the mass of 3.25 x1023 molecules of
nitrogen?
3.25 x1023 N2 x 1.0 mol
x
6.02 x 1023
28.02 g = 15.1 g N2
1.0 mol
8) How many molecules are 57.36 g of NaCl?
57.36 g NaCl x 1.0 mol x 6.02 x 1023 = 5.911 x 1023
58.44 g
1.0 mol
molecules
3-43
Ch. 3
The
End!