Download Atoms – Building Blocks of Matter Notes

Atoms – Building Blocks of
Matter Notes - Chapter 3
I. The Atom: From Idea to Theory
A. 400 BC Democritus VS Aristotle
Democritus, an ancient Greek and student of
Aristotle, proposed the 1st atomic theory he said
that the world is composed of 2 things: void
(empty space) and matter. No one supported
him and he had NO experimental evidence to
support his idea.
Greek Model
“To understand the very large,
we must understand the very small.”
Democritus
Greek philosopher
Idea of ‘atomos’
Atomos = ‘indivisible’
‘Atom’ is derived
Democritus’s model of atom
No protons, electrons, or neutrons
Solid and INDESTRUCTABLE
No experiments to support idea
DEMOCRITUS (400 BC) – First Atomic
Hypothesis
Atomos: Greek for “uncuttable”. Chop up a piece of matter until you
reach the atomos.
Properties of atoms:
• indestructible.
• changeable, however, into different forms.
• an infinite number of kinds so there are an infinite number of
elements.
• hard substances have rough, prickly atoms that stick together.
• liquids have round, smooth atoms that slide over one another.
• smell is caused by atoms interacting with the nose – rough
atoms hurt.
• sleep is caused by atoms escaping the brain.
• death – too many escaped or didn’t return.
• the heart is the center of anger.
• the brain is the center of thought.
• the liver is the seat of desire.
“Nothing exists but atoms and space, all else is opinion”.
Aristotle proposed that matter was composed of
one continually flowing substance called hyle.
This idea was widely supported and accepted
until the late 1700’s and he too had NO
experimental evidence to support his idea.
Aristotle - Four Element Theory
FIRE
Thought all matter was
composed of 4 elements:
Earth (cool, heavy)
Water (wet)
Fire (hot)
Air (light)
Ether (close to heaven)
Hot
Dry
‘MATTER’
AIR
Wet
EARTH
Cold
WATER
Relation of the four elements and the four qualities
Blend these “elements” in different proportions to get all substances
B. Late 1700’s Isaac Newton and
Robert Boyle
It was not until the late 1700’s that anyone dared to
question Aristotle’s wisdom. They suggested that
Aristotle was incorrect but did not have their own theory
to submit.
At this time chemist did believe, based on experiments,
that there were different elements and that an element
was a substance that could not be broken down by
chemical means. Chemist knew that some substances
could transform into different or new substances, they
called this a chemical reaction.
C. 1790’s - Basic laws that were
established:
Chemist also knew, via improved balances, that
when a chemical reaction occurred in a closed
space that the mass of the material before the
change equaled the mass of the marital after the
change. Now known as the Law of
Conservation of Mass.
Discovered by Antoine
Laurent Lavoisier (174394) about 1785.
Law of Conservation of Mass
Law of Conservation of Mass
Law of Conservation of Mass
http://www.teachertube.com/viewVideo.php?title=T
esting_Conservation_of_Mass&video_id=85396
Another realization was that substances always
contained their elements in the same proportions
by mass. For example: for any sample of
sodium chloride, the mass of the sample is
always 39.34% Na and 60.66% Cl. Now known
as the Law of Definite Proportions.
It was also known that elements combined to form
more than one compound. Example: carbon
monoxide and carbon dioxide. This is the Law
of Multiple Proportions.
Legos are Similar to Atoms
H
H2
H
H
O
+
H2
H
H
O2
H
O
H 2O
H
O
O
H
H 2O
Lego's can be taken apart and built into many different things.
Atoms can be rearranged into different substances.
D. 1803 John Dalton
British chemist who was the first to have a theory
about matter being composed of atoms and how
atoms might look and behave. Dalton proposed
an explanation for the Law of Conservation of
Mass, Law of Definite Proportions, and Law of
Multiple Proportions. He reasoned that elements
were composed of atoms and that only whole
numbers of atoms can combine to form
compounds. He conceived on the atom as a
solid billiard ball. Here is a summary of his
theory:
1. All matter is composed of atoms.
2. Atoms of the same elements are exactly the
same and atoms of different elements are
different.
3. Atoms cannot be created, destroyed, or
subdivided.
4. Atoms of different elements combine in whole
number ratios to form compounds.
5. In chemical reactions, atoms are combined,
separated, or rearranged.
Dalton’s Symbols
John Dalton
1808
Democritus’s idea, because Dalton was
able to relate atoms to the measurable
property of mass, turned into a scientific
theory!!
The only aspect of Daltons’ Theory that is now
known to be incorrect is the fact that atoms can
be subdivided (into p+, e-, n). And that atoms of
the same element can have deterrent masses
(these are called isotopes).
II. The Structure of Atoms
Atom – smallest particle of an element that retains
the chemical properties of that element. All
atoms consist of 2 regions – the nucleus (p+ & n)
and surrounding the nucleus is the electron
cloud – a region occupied by the negatively
charged particles called electrons. How do we
know this?!
1. Discovery of Electron 1897 (by
J.J. Thomson and Robert Millikan)
1st subatomic particle to be discovered – Thomson was
working with electricity and magnetic fields. He was taking
various gases and sending an electric current through the
gas. When he did this he noticed that a glow was emitted.
(What he was doing, he believed, was separating the
electron from the nucleus of the gas atoms – this caused
the glow!) Thomson went on to prove that the glow was
actually a stream of negatively charged particles – called
electrons. Symbol e-, charge –1, and mass of 0.00055amu
(atomic mass unit,
1amu = 1.66X10 –27 kg)
http://courses.science.fau.edu/%7Erjordan/phy2044/Q_A/ans_14.htm
J. J. Thomson - English physicist. 1897
A Cathode Ray Tube
Source of
Electrical
Potential
Stream of negative
particles (electrons)
Metal Plate
Gas-filled
glass tube
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 58
Metal plate
J.J. Thomson
He proved that atoms of any
element can be made to emit
tiny negative particles.
He knew that atoms did not
have a net negative charge
and so there must be
balancing the negative
charge.
J.J. Thomson
Plum Pudding Model – Thomson proposed that
the atom had negative electrons scattered
throughout a positively charged area (proton
area).
1910 the Plum Pudding
model
Negative electrons
were embedded into a
positively charged
spherical cloud.
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 56
Spherical cloud of
Positive charge
Electrons
2. Protons 1919 (discovered by
Rutherford/J.J. Thomson)
Both Rutherford and Thomson knew that positively
charged particles (protons) must exist (because
an atom is neutral, if there is a negative charged
electron then there has to be a positively
charged proton to make it neutral.) They worked
together to prove they existed. Proton symbol:
+p, charge +1, mass 1.008 amu.
3. Discovery of the Atomic Nucleus
1911
Discovered by Rutherford during his famous gold-foil
experiment and realized that the main part of the atom’s
mass is in the nucleus, and that it is positively charged.
Summary of his experiment:
-Bombarded a thin piece of gold foil with positive alpha
particles
-Most went through as though nothing was there
-Few (1 in 8000) ricochet back toward the source
-Few were deflected off to the side
Rutherford’s Gold Foil Experiment
Rutherford received the 1908 Nobel Prize in Chemistry for his pioneering work in nuclear chemistry.
beam of alpha particles
radioactive
substance
circular ZnS - coated
fluorescent screen
gold foil
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
What he expected…
What he got…
richocheting
alpha particles
Interpreting the
Observed Deflections
.
.
.
.
.
.
beam of
alpha
particles
.
.
.
.
.
undeflected
particles
.
.
.
.
.
gold foil
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 120
.
deflected particle
Rutherford Scattering (cont.)
Rutherford interpreted this result by suggesting that the a
particles interacted with very small and heavy particles
Particle bounces off
of atom?
Case A
Case B
Particle goes through
atom?
Particle attracts
to atom?
Case C
Case D
.
Particle path is altered
as it passes through atom?
Explanation of Alpha-Scattering Results
Alpha particles
Nucleus
+
+
-
-
+
+
-
+
+
-
+
-
+
-
-
Plum-pudding atom
Nuclear atom
Thomson’s model
Rutherford’s model
Rutherford’s Conclusion: the positive alpha particles
had to have hit something else that was positively
charged to cause the ricochet effect. The
“something” was very small and dense because
only a few hit it, therefore the atom must have a
small positively charged nucleus, surrounded by
mostly empty space (because most particles went
through the gold foil.) New model of atom:
Electron 0.00055amu
Proton 1.008 amu.
Neutron 1.008 amu.
4. Neutrons 1932 (Proved by
Chadwick)
New something else existed in an atom because of
the mass of the atom. Neutron is an electrically
neutral particle, symbol n, mass equal that of
protons.
5. Nuclear Forces – the +p and n stay close to
each other due to these short-range forces that
hold the +p and n together.
Current Model of Atom:
Bohr’s Model
Nucleus
Electron
Orbit
Energy Levels
Quantum Mechanical Model
Niels Bohr &
Albert Einstein
Modern atomic theory describes the electronic
structure of the atom as the probability of finding
electrons within certain regions of space (orbitals).
Modern View
The atom is mostly empty space
Two regions
Nucleus

protons and neutrons
Electron cloud

region where you might find an electron
Dalton proposes the
indivisible unit of an
element is the atom.
Review
Models
of the
Atom
Thomson discovers
electrons, believed to
reside within a sphere of
uniform positive charge
(the “plum-pudding model).
Rutherford demonstrates
the existence of a positively
charged nucleus that
contains nearly all the
mass of an atom.
Bohr proposes fixed
circular orbits around
the nucleus for electrons.
In the current model of the atom,
electrons occupy regions of space
(orbitals) around the nucleus
determined by their energies.
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.
Models of the Atom
Dalton’s
model
Greek model
(1803)
(400 B.C.)
1803 John Dalton
pictures atoms as
tiny, indestructible
particles, with no
internal structure.
1800
Thomson’s plum-pudding
model (1897)
Rutherford’s model
(1909)
1897 J.J. Thomson, a British
1911 New Zealander
scientist, discovers the electron,
leading to his "plum-pudding"
model. He pictures electrons
embedded in a sphere of
positive electric charge.
Ernest Rutherford states
that an atom has a dense,
positively charged nucleus.
Electrons move randomly in
the space around the nucleus.
1805 ..................... 1895
1900
1905
1910
1904 Hantaro Nagaoka, a
Japanese physicist, suggests
that an atom has a central
nucleus. Electrons move in
orbits like the rings around Saturn.
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
1915
Bohr’s model
(1913)
1926 Erwin Schrödinger
1913 In Niels Bohr's
model, the electrons move
in spherical orbits at fixed
distances from the nucleus.
1920
1925
Charge-cloud model
(present)
1930
develops mathematical
equations to describe the
motion of electrons in
atoms. His work leads to
the electron cloud model.
1935
1940
1945
1924 Frenchman Louis
1932 James
de Broglie proposes that
moving particles like electrons
have some properties of waves.
Within a few years evidence is
collected to support his idea.
Chadwick, a British
physicist, confirms the
existence of neutrons,
which have no charge.
Atomic nuclei contain
neutrons and positively
charged protons.
Warm-Up
Please complete the table:
Particle
Location
Mass
Charge
Proton
nucleus
1 amu
+
Neutron
nucleus
1 amu
0
Electron
electron
cloud
0
-
III. Counting Atoms
Reading the periodic table
11
atomic number
Na
symbol
Sodium
name
22.990
average atomic mass (in amu’s)
23
mass number (the average atomic mass
rounded to the nearest whole number)
1. Atomic Number
the number of protons in the nucleus. The
atomic number identifies the element!!!!!!!!!
Because atoms are neutral they contain the
same number of electrons as protons.
(Therefore the atomic number is the number of
electrons as well.)
2. Atomic Mass
– average mass of 1 atom of a specific element
(measured in amu’s)
Mass Number
the atomic mass rounded to the nearest whole
number, therefore it is the total number of protons
and neutrons in an atom’s nucleus.
mass # = protons + neutrons
+
+
+
Always a whole number
e+
+
+
NOT on the Periodic Table!
Neutron
Electrons
Nucleus
Proton
e-
e-
Nucleus
e-
ee-
Carbon-12
Neutrons 6
Protons
6
Electrons 6
Nuclear Symbol
Find the
number of protons
=9 +
number of neutrons
= 10
number of electrons
=9
Mass number
= 19
Atomic number = 9
19
9
F
Nuclear Symbol
Find the
number of protons = 11 +
number of neutrons = 12
number of electrons
= 10
Mass number
= 23
Atomic number = 11
23
11
1+
Na
Sodium ion
Nuclear Symbol
Find the
= 11
number of protons
number of neutrons
= 12
number of electrons
= 10
Mass number
= 23
Atomic number = 11
23
11
1+
Na
Sodium ion
Isotopes
Neutron
+
Electrons
Nucleus
+
+
+
+
+
Nucleus
Proton
Proton
Nucleus
Carbon-12
Neutrons 6
Protons
6
Electrons 6
+
+
+
+
Neutron
Electrons
+
+
Carbon-14
Neutrons 8
Protons
6
Electrons 6
Nucleus
Practice: How many protons are in each
of the following? neutrons? electrons?
Symbol
Atomic#
Mass#
#p+
#n
#e-
Be
4
9
4
5
4
Ne
10
20
10
10
10
Na
11
23
11
12
11
4. Ions
– atoms that have lost or gained electrons are
called ions
Positive Ions (when an atom loses electrons)
Example:
23
Na
11 p+ and 11e-
11
Lose 1 electron
11p+ and 10e23
1+
Na
11
4. Ions
– atoms that have lost or gained electrons are
called ions
Negative Ions (when an atom gains electrons)
Example:
19
F
9 p+ and 9 e-
9
Gain 1 electron
9p+ and 10e19
1-
F
9
5. Isotopes
– atoms with the same number of protons (atomic
number is the same) but different numbers of
neutrons (mass number is different). Usually
isotopes are referred to by their name (of symbol)
and their mass number (ex. carbon-12). Every
element on the chart has at least 2 isotopes and
some elements have as many as 25 isotopes.
Example: The isotopes of hydrogen have
separate names rather than being called
hydrogen-1, hydrogen-2, etc. Their names are
protium (H-1), deuterium (H-2), and tritium (H-3).
Name in
Hyphen
Notation
#p+
#e-
#n
(H– 1)
1
1
0
1
(H – 2)
1
1
1
2
(H – 3)
1
1
2
3
Mass #
Nuclear
Symbol
Name in
Hyphen
Notation
(C– 14)
(C – 13)
(C – 12)
#p+
#e-
#n
Mass #
Nuclear
Symbol
Most elements occur naturally as mixtures of
isotopes, as indicated in Table 3-4 of textbook.
The percentage of each isotope in the naturally
occurring element on Earth is nearly always the
same, no matter where the element is found.
The percentage at which each of an element’s
isotopes occurs in nature is taken into account
when calculating the element’s average atomic
mass (which appears on the periodic table).
Isotopic Mass and Natural Abundance
6. Relative atomic masses
a.m.u.- atomic mass unit; One amu is exactly
1/12 the mass of a carbon-12 atom. So the
atomic mass of any nuclide is determined by
comparing it with the mass of the carbon-12
atom. The hydrogen-1 atom has an atomic mass
of about 1/12 that of the carbon-12 atom, or 1
amu.
1 amu = 1.66X10-27kg
7. Average atomic mass
It is the weighted average of the masses of all the isotopes
of that element. A weighted average reflects both the
mass and the abundance of the isotopes as they occur in
nature.
Isotope
Atomic mass abundance (%)
H-1
1.0078amu
99.985%
H-2
2.0141amu
0.015%
H-3
3.0160amu
negligible
The average atomic mass of hydrogen is 1.0079amu. Multiply
each mass number by the percent abundance and add
them up.
Isotope
H-1
H-2
H-3
Atomic mass abundance (%)
1.0078amu
99.985%
2.0141amu
0.015%
3.0160amu
negligible
(1.0078amu)(.99985)
+ (2.0141amu)(0.00015)
1.0079amu
The average atomic mass of hydrogen is 1.0079amu.
Practice: Element Z has 2 natural isotopes. The
isotope with a mass number of 15 has a relative
abundance of 30%. The isotope with a mass
number of 16 has a relative abundance of 70%.
Estimate the average atomic mass for this
element.
The Mole
23
6.02 X 10
STOICHIOMETRY
The study of
the quantitative
aspects of
chemical
reactions.
The Mole
A counting unit
Similar to a dozen, except instead
of 12, it’s 602 billion trillion
602,000,000,000,000,000,000,000
6.02 X 1023 (in scientific notation)
This number is named in honor of
Amedeo _________ (1776 – 1856),
who studied quantities of gases
and discovered that no matter what
the gas was, there were the same
number of molecules present
Just How Big is a Mole?
Enough soft drink cans to cover the
surface of the earth to a depth of
over 200 miles.
If you had Avogadro's number of
unpopped popcorn kernels, and
spread them across the United
States of America, the country would
be covered in popcorn to a depth of
over 9 miles.
If we were able to count atoms at the
rate of 10 million per second, it
would take about 2 billion years to
count the atoms in one mole.
Everybody Has Avogadro’s
Number!
But Where Did it Come From?
It was NOT just picked!
It was MEASURED.
One of the better
methods of measuring
this number was the
Millikan Oil Drop
Experiment
Since then we have
found even better ways
of measuring using xray technology
IV. Relating Mass to Numbers
of Atoms
1. The Mole (can be abbreviated mol, but NOT
m, which is the abbreviation for meter!) - the SI
unit for amount of substance. A mole is the
amount of a substance that contains as many
particles as there are atoms in exactly 12 grams
of carbon-12.
2. Avogadro’s Number
-the number of particles in exactly one mole of a
pure substance. This number was determined
experimentally and its value is 6.02 X 1023,
which means that 12 g of carbon-12 contains
6.02 x 1023 carbon-12 atoms.
3. Using the Mole and Avogadro’s
Number
A mole can be thought of as a counting unit
just like a dozen (12), gross (144), pair (2),
ream (500), mole (6.02X1023 ).
A. How many is a mole? Enough.
If every person living on Earth (6 billion people)
worked to count out one mole of oranges (or anything
else), and if each person counted continually at a rate
of one orange per second, it would take about 4
million years for all the oranges to be counted!
If we had a mole of sand it would cover the earth 7
times over! If you had a mole of dollar bills, you
could spend a million dollars every minute of your life
and never spend it all!
Since the mole is so large, we use it to count
very tiny things – like atoms. Because the mole
is so large, (and we now know that we cannot
count out a mole of anything), how do we know
when we have a mole of anything?
We determine the mass and relate that to the
number of atoms present. (Aluminum cans
example.)
4. Molar Mass
– The mass of one mole of a pure substance.
The pure substance can be an element or a
compound.
The atomic mass is the mass of 1 atom of that
element measured in amu’s.
The atomic mass is also equal to 1 mole of
atoms measured in grams it is called the
molar mass!!!! What a coincidence!!!!
Mass of 1 atom of Pb = 207.2 amu
Mass of 1 mole of Pb atoms = 207.2 g
Mass of 1 atom of N = 14.01 amu
Mass of 1 mole of N atoms = 14.01 g
Mass of 1 atom of Ba = 137.33 amu
Mass of 1 mole of Ba atoms = 137.33 g
Mass of 1 atom of Al = 26.98 amu
Mass of 1 mole of Al atoms = 26.98 g
Let’s prove it: Determine the mass,
in grams, of 6.02X1023 atoms of
aluminum. Use 1amu = 1.66X10-27kg.
With this information we can write some
new conversion ratio’s!!
“!!”…?
1 mole = 6.02X1023 atoms OR molecules OR
formula units
1 mole Al = 26.98 grams
1 atom Al = 26.98 amu
V. Mole Problems – When in
doubt go to the mole.
The MOLE has been defined as 6.02 x l023 atoms of a
pure element or the molar mass of a substance
expressed in grams. It can also be defined as 6.02 x
l023 molecules of a compound or diatomic molecule
(O2, N2, H2, etc)
THE ONLY THING HARD ABOUT UNDERSTANDING
THE DEFINITION OF A MOLE IS THAT YOU
UNDERSTAND THAT THE VALUE OF A MOLE IS
DIFFERENT FOR EVERY DIFFERENT ELEMENT
AND COMPOUND.
1. Gram/Mole conversions-how to convert moles to
grams or grams to moles?
Example: 120 g Ca x 1 mole Ca = 3.0 mole Ca
40 g Ca
Practice:
How many grams of sodium are in 5.00 moles of
sodium?
How many grams of magnesium are in 0.250 moles of
magnesium?
How many moles of lead, Pb, are in 210. g of lead?
How many moles of nitrogen are in 44.0 g of nitrogen?
2. Conversions with Avogadro’s Number
Example: How many atoms of silver, Ag, are in
4.25 moles of Ag?
4.25 moles Ag X 6.02 x 1023 atoms Ag =
1 mole Ag
Practice:
How many atoms of Pb are in 3.80 moles of Pb?
How many moles of Na are in 8.24 x 1024 atoms
of Na?
Two-step conversions:
Ex: How many atoms of sodium, Na, are in 5.25 g of
Na?
5.25 g Na
x 1 mole Na
23 g Na
x 6.02 x 1023 atoms Na =
1 mole Na
Practice:
How many atoms of potassium, K, are in 3.99 g of
K?
How many g of He are in 3.03 x 1021 atoms of He?
3. How many atoms of Li are in 0.755 g of Li?
3. Molar Mass for compounds
How to find the molar mass:
Write a CORRECT formula for the compound
(we’ll do this later)
Look up the atomic mass of each element in the
compound
Multiply the atomic mass by the subscripts, if any.
Add all masses of elements together and use the
unit, g/mol
Example: find the molar mass of NaCl.
Na=22.99 g/mol
Cl=35.45 g/mol
58.44 g/mol
Example: find the molar mass of calcium
phosphate, Ca3(PO4)2.
Ca = 40.08 x 3 = 120.24
P = 30.97 x 2 = 61.94
0 = 16.00 x 8 = 128.00
310.18 g/mol
Practice:
Find the molar mass of amonium sulfate,
(NH4)2SO4
Find the molar mass of Cl2O7
Hydrates - Some compounds trap water
inside their crystal structure and are known
as hydrates. You will not be able to predict
which compounds will form hydrates. All
you have to do is to be able to name them
and find their molar masses (including the
water).
CuSO4 · 5H20 is an example of a hydrate.
This says that one formula unit of cupric
sulfate will trap 5 molecules of water inside
its crystal.
Hydrates are named by naming the ionic compound
by the regular rules and then adding (as a second
word) a prefix indicating the number of water
molecules. You will use the word “hydrate” to
indicate water. The above compound would be
called cupric sulfate pentahydrate.
To find the formula mass of a hydrate ask Medina
(he’s a wise man), simply find the mass of the ionic
compound by itself and then ADHD the mass of
water molecule(s) to that mass.
Practice: What is the formula mass of barium chloride
dehydrate, BaCl2 · 2H2O?
What is the formula mass of aluminum sulfate
octahydrate, Al2(SO4)3 · 8H2O?
VI: Real World Connections – Travels
with Carbon (book page 68)
There is basically is no such thing as “new” air. When a
living thing inhales, molecules are taken in, and when it
exhales, molecules are released back into the
atmosphere to be reused later by some other living
organism. Thus, at least in principle, the molecules of
“air” (nitrogen and oxygen, mostly) that Caesar exhaled
from his last breath have since that time been
redistributed throughout Earth’s entire atmosphere.
When you breathe, it is entirely possible that you will
inhale one or more of these molecules.
www.scifun.ed.ac.uk/card/facts.html
Breathing Everyone's Air
"Breathe out, then wait for a second. Now
breathe in. In that single breath, you have
just inhaled molecules that have been
breathed out by almost every person who
has ever lived. In fact, that breath contains
molecules exhaled by every single organism
that has ever breathed, right back to the first
bacterium.
Mixed Practice Problems:
How many moles of water in 72.0 grams?
How many moles are in 100. grams of nitrogen gas, N2?
How many moles are there in 28.7 grams of lithium nitrate, LiNO3?
How many moles are there in 1.75 tons of magnesium chloride, MgCl2?
How much would 38.0 moles of oxygen gas, O2, weigh in pounds? (454 g
= 1 lb)
How many atoms are there in 75.0 grams of pure iron, Fe?
How many molecules are in 5.00 moles of water?
How many atoms are in 3.50 moles of nitrogen gas, a diatomic molecule?
How much would 7.12 x 1023 molecules of magnesium phosphate,
Mg3(PO4)2 weigh in grams?
10. How many barium atoms are there in 5.89 grams of barium, Ba?
11. What is the formula mass of cupric sulfate petahydrate, CuSO4 ·
5H2O?
12. How boring is this class?
13. How many licks to the center of a tootsie pop? Use Moles.
Learning Check
Suppose we invented a new collection unit
called a rapp. One rapp contains 8 objects.
1. How many paper clips in 1 rapp?
a) 1
b) 4
c) 8
2. How many oranges in 2 rapp?
a) 4
b) 8
c) 16
3. How many rapps contain 40 gummy bears?
a) 5
b) 10
c) 20
The Mole
1 dozen cookies = 12 cookies
1 mole of cookies = 6.02 X 1023 cookies
1 dozen cars = 12 cars
1 mole of cars = 6.02 X 1023 cars
1 dozen Al atoms = 12 Al atoms
1 mole of Al atoms = 6.02 X 1023 atoms
Note that the NUMBER is always the same,
but the MASS is very different!
Mole is abbreviated mol (gee, that’s a lot
quicker to write, huh?)
A Mole of Particles
Contains 6.02 x 1023 particles
1 mole C
= 6.02 x 1023 C atoms
1 mole H2O = 6.02 x 1023 H2O molecules
1 mole NaCl = 6.02 x 1023 NaCl “molecules”
(technically, ionics are compounds not
molecules so they are called formula units)
6.02 x 1023 Na+ ions and
6.02 x 1023 Cl– ions
Avogadro’s Number as
Conversion Factor
6.02 x 1023 particles
1 mole
or
1 mole
6.02 x 1023 particles
Note that a particle could be an atom OR a molecule!
Learning Check
1. Number of atoms in 0.500 mole of Al
a) 500 Al atoms
b) 6.02 x 1023 Al atoms
c) 3.01 x 1023 Al atoms
2.Number of moles of S in 1.8 x 1024 S atoms
a) 1.0 mole S atoms
b) 3.0 mole S atoms
c) 1.1 x 1048 mole S atoms
Molar Mass
The Mass of 1 mole (in grams)
Equal to the numerical value of the average
atomic mass (get from periodic table)
1 mole of C atoms
=
12.0 g
1 mole of Mg atoms
=
24.3 g
1 mole of Cu atoms
=
63.5 g
Other Names Related to Molar Mass
Molecular Mass/Molecular Weight: If you have a single
molecule, mass is measured in amu’s instead of grams. But,
the molecular mass/weight is the same numerical value as 1
mole of molecules. Only the units are different. (This is the
beauty of Avogadro’s Number!)
Formula Mass/Formula Weight: Same goes for
compounds. But again, the numerical value is the same.
Only the units are different.
THE POINT: You may hear all of these terms
which mean the SAME NUMBER… just different units
Learning Check!
Find the molar mass
(usually we round to the tenths place)
A. 1 mole of Br atoms =
B. 1 mole of Sn atoms =
79.9 g/mole
118.7 g/mole
Molar Mass of Molecules and
Compounds
Mass in grams of 1 mole equal numerically to
the sum of the atomic masses
1 mole of CaCl2
= 111.1 g/mol
1 mole Ca x 40.1 g/mol
+ 2 moles Cl x 35.5 g/mol
1 mole of N2O4
= 111.1 g/mol CaCl2
= 92.0 g/mol
Learning Check!
A.
Molar Mass of K2O = ? Grams/mole
B. Molar Mass of antacid Al(OH)3 = ?
Grams/mole
Learning Check
Prozac, C17H18F3NO, is a widely used
antidepressant that inhibits the uptake of
serotonin by the brain. Find its molar
mass.
Calculations with Molar Mass
molar mass
Grams
Moles
Converting Moles and Grams
Aluminum is often used for the structure
of light-weight bicycle frames. How
many grams of Al are in 3.00 moles of
Al?
3.00 moles Al
? g Al
1. Molar mass of Al
1 mole Al = 27.0 g Al
2. Conversion factors for Al
27.0g Al
1 mol Al
or
1 mol Al
27.0 g Al
3. Setup 3.00 moles Al
Answer
x
27.0 g Al
1 mole Al
= 81.0 g Al
Learning Check!
The artificial sweetener aspartame
(Nutra-Sweet) formula C14H18N2O5 is
used to sweeten diet foods, coffee and
soft drinks. How many moles of
aspartame are present in 225 g of
aspartame?
Atoms/Molecules and Grams
Since 6.02 X 1023 particles = 1 mole
AND
1 mole = molar mass (grams)
You can convert atoms/molecules to
moles and then moles to grams! (Two step
process)
You can’t go directly from atoms to
grams!!!! You MUST go thru MOLES.
That’s like asking 2 dozen cookies weigh
how many ounces if 1 cookie weighs 4 oz?
You have to convert to dozen first!
Calculations
molar mass
Grams
Avogadro’s number
Moles
particles
Everything must go through
Moles!!!
Atoms/Molecules and Grams
How many atoms of Cu are
present in 35.4 g of Cu?
35.4 g Cu
1 mol Cu
63.5 g Cu
6.02 X 1023 atoms Cu
1 mol Cu
= 3.4 X 1023 atoms Cu
Learning Check!
How many atoms of K are present in
78.4 g of K?
Learning Check!
What is the mass (in grams) of 1.20 X
1024 molecules of glucose (C6H12O6)?
Learning Check!
How many atoms of O are present in
78.1 g of oxygen?
78.1 g O2 1 mol O2 6.02 X 1023 molecules O2 2 atoms O
32.0 g O2 1 mol O2
1 molecule O2
Percent Composition
What is the percent carbon in C5H8NO4 (the
glutamic acid used to make MSG
monosodium glutamate), a compound used
to flavor foods and tenderize meats?
a) 8.22 %C
b) 24.3 %C
c) 41.1 %C
Chemical Formulas of Compounds
(HONORS only)
Formulas give the relative numbers of atoms or
moles of each element in a formula unit - always a
whole number ratio (the law of definite
proportions).
NO2
2 atoms of O for every 1 atom of N
1 mole of NO2 : 2 moles of O atoms to every 1
mole of N atoms
If we know or can determine the relative number
of moles of each element in a compound, we can
determine a formula for the compound.
Types of Formulas
(HONORS only)
Empirical Formula
The formula of a compound that
expresses the smallest whole number
ratio of the atoms present.
Ionic formula are always empirical formula
Molecular Formula
The formula that states the actual
number of each kind of atom found in one
molecule of the compound.
To obtain an Empirical Formula
(HONORS only)
1. Determine the mass in grams of each
element present, if necessary.
2. Calculate the number of moles of each
element.
3. Divide each by the smallest number of
moles to obtain the simplest whole
number ratio.
4.
If whole numbers are not obtained* in
step 3), multiply through by the smallest
number that will give all whole numbers
* Be
careful! Do not round off numbers prematurely
A sample of a brown gas, a major air pollutant,
is found to contain 2.34 g N and 5.34g O.
Determine a formula for this substance.
require mole ratios so convert grams to moles
moles of N = 2.34g of N = 0.167 moles of N
14.01 g/mole
moles of O = 5.34 g = 0.334 moles of O
16.00 g/mole
N 0.167 O 0.334  NO 2
Formula: N O
0.167
0.334
0.167
0.167
(HONORS only)
Calculation of the Molecular Formula
(HONORS only)
A compound has an empirical formula
of NO2. The colourless liquid, used in
rocket engines has a molar mass of
92.0 g/mole. What is the molecular
formula of this substance?
Empirical Formula from % Composition
(HONORS only)
A substance has the following composition by
mass: 60.80 % Na ; 28.60 % B ; 10.60 % H
What is the empirical formula of the substance?
Consider a sample size of 100 grams
This will contain 28.60 grams of B and
10.60 grams H
Determine the number of moles of each
Determine the simplest whole number ratio