Unit 2 - Organization of
Matter
Chapter 3 - Atoms: The Building
Blocks of Matter
Atoms: The Building Blocks of
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The Law of Conservation of
Mass

The transformation of substances into one
or more new substances is known as a
chemical reaction.
 The law of conservation of mass states
that matter cannot be created or destroyed
in ordinary chemical reactions
 In simpler terms, the mass of the
reactants = the mass of the products
after a reactions has taken place.
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Anatomy of Chemical
Formulas

A chemical formula indicates the
relative number and type of atoms in a
chemical compound
Al2(SO4)3
2 Al’s, 3 (SO4) ions
Subscripts - below the chemical symbol
Superscripts - above the chemical symbol
(like an exponent)
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Anatomy of Chemical
Equations

A chemical equation represents, with
symbols and formulas, the identities and
relative amounts of the reactants and
products in a chemical reaction.
Example:
(NH4)2Cr2O7 (s)
N2(g) + Cr2O3(s) + 4H2O(g)
Reactant
Products
Ammonium Dichromate(solid) yields or produces N(gas),
chromium(III) oxide(solid) and gaseous water
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The Law of Definite
Proportions
The Law of Definite Proportions
states that a chemical compound
contains the same elements in the
exactly the same proportions by mass
regardless of the size of the sample or
source of the compound.
 Example : NaCl - Table salt always
consists of 39.34% Na and 60.66% Cl

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The Law of Multiple
Proportions

If 2 or more compounds (combined
elements) are composed of the same 2
elements, then the ratio of the masses is
always a small whole number.
 For example: CO and CO2 - Lets say that
each of these compounds contain 1 gram of
C, a 1:1 ratio. In CO, 1.33g of O combines
with 1 g of C. In CO2, 2.66 grams of O
combines with 1 g of C. The ratio of O in the
two compounds is 2.66 to 1.33, or 2:1.
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Dalton’s Atomic Theory

He reasoned that only whole numbers of
atoms can combine to form compounds

1. All matter is composed of extremely small
particles called atoms.
 2. Atoms of the same element are identical in
size, mass and other properties.
 3. Atoms cannot be subdivided, created, or
destroyed.
 4. Atoms of different elements combine in simple
whole number ratios to form compounds.
 5. In chemical reactions, atoms are combined,
separated, or rearranged.
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Modern Atomic Theory

Two differences from Dalton’s Theory
 We
know that atoms are divisible into
even smaller particles.
 A given element can have atoms with
different masses (Isotopes - more about
these later)
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3.2 -The Structure of the Atom

An atom is the smallest particle of an
element that retains the chemical properties
of the element.
 Subatomic particles:

Protons (+ charge) and neutrons (no charge)
located in the nucleus
 Electrons surrounding the nucleus with a
negative charge.
 Atoms in their natural state are neutral. (the # of
E’s = the # of P’s)
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Discovery of the Electron
Electricity was passed through certain
(noble) gases at low pressures in a
cathode ray tube.
 The surface directly opposite the
cathode glowed. This was caused by a
stream of particles, which they called a
cathode ray.

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Discovery of an Electron ContThe cathode ray deflected away from a
negatively charged object (Like
charges repel)
 These observations led to the
hypothesis that the particles that
compose cathode rays are negatively
charged.
 These were later renamed electrons.

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Charge and Mass of the
Electron

J.J. Thompson revealed that the electron
has a very large charge for its tiny mass.
He found that the ratio was always the same
no matter what kind of metal or gas was
used in the tube.
 Robert Milikan showed that the mass of an
electron is in fact about one two-thousandth
(1/1837) the mass of the simplest type
hydrogen atom, which is the smallest atom
known.
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Based on what we learned
about electrons,

2 other inferences were made about
atomic structure. (About 1909)
 Because
atoms are electrically neutral,
they must contain a + charge to balance
the negative electrons.
 Because electrons have so much less
mass than atoms, atoms must contain
other particles that account for most of
their mass.
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Discovery of the Atomic
Nucleus

In 1911, Ernest Rutherford (and associates
Gieger and Marsden) bombarded a thin,
gold foil with fast moving alpha particles,
which are positively charged particles
 They expected the particles to pass through
with only slight deflection.
 For the vast majority of the particles this
was the case.
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Rutherford’s Gold Foil
Experiment
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Discovery of the Nucleus cont

They were shocked to find that 1 in 8000
deflected back toward the source.
 Rutherford reasoned that the rebounded
particles must have experienced some
powerful repulsive force within the atom.
 He conclude that the force must be caused
by a very densely packed bundle of matter
with a + electric charge.
 He called this bundle of matter the nucleus.
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Discovery of the Nucleus cont
Rutherford had discovered that the volume
of the nucleus was small compared to the
total volume of the atom.
 If the nucleus was the size of a marble, the
atoms was the size of a football field.
 What about the electrons? Rutherford
suggested they orbit the nucleus like planets
around the sun.
 Although, he couldn’t explain what kept the
electrons in motion.
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Composition of the Atomic Nucleus Properties of Subatomic Particles
Particle
Symbols
Relative
electric
charge
Mass
Number
Relative
mass
(amu)
Electron
E-
-1
0
0.000548 9.109 X
Proton
P+
+1
1
1.007276 1.673 X
10-27
Neutron
n0
0
1
1.008665 1.675 X
10-27
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Actual
mass
(kg)
10-31
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Section 3.3 - Counting Atoms

The atomic number (Z) of an element is
the number of protons in the nucleus of
each atom of that element. (the whole
number)
 Mass Number (A) is the average atomic
mass (number with a decimal point)
rounded off.
 The most identifiable characteristic of an
atom is it’s atomic number.
 Example: Helium’s (He) atomic # is Z=2.
There is not another atom of any element
that has a atomic number of 2. It’s mass
number is A=4.
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Determining Number of Protons, Neutrons and
Electrons in a NEUTRAL Atom Using Your
Periodic Table




Whole number = # of
protons and electrons
# with decimal is the
average atomic mass
Average atomic mass
rounded off is the mass
number.
Mass # = protons +
neutrons
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Isotopes
Isotopes are atoms of the same element
(same number of protons) that have
different masses because of a difference in
number of neutrons. (Deviation from
Dalton’s Atomic theory! Postulate #2)
Example: All atoms of hydrogen have 1 proton

Hydrogen has three main isotopes: Protium,
Deuterium, and Tritium.
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Mass Numbers of H Isotopes
Atomic #
# of
neutrons
Mass #
Protium
1
0
1+0=1
Deuterium
1
1
1+1=2
Tritium
1
2
1+2=3
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Mass Number & Atomic
Number
Identifying an isotope requires knowing
both the atomic number and the mass
number of the isotope.
 Although isotopes have different
masses, they do not differ significantly
in their chemical behavior because
they are after all the same atom!!!

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Designating Isotopes
Isotopes are usually identified by
specifying their mass number.
1. The mass number is written with a
hyphen after the name of the element.
This is called hyphen notation.
Hydrogen-3
Uraniam-235
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Designating Isotopes - cont
2. Using the nuclear
symbol
Nuclide is the
general term for
any isotope of any
element.
The A and Z number can be on
the either side of the symbol.
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Determining Number of Neutrons, Protons, and
Electrons in Isotopes
P+ = 54
E- = 54
N0 = 82
Xenon - 136
Mass number
239U
Mass number
Can appear on
left or right of
symbol
P+ = 92
E- = 92
N0 = 147
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Average Atomic Masses of
Elements
Most elements occur naturally as
mixtures of isotopes.
 Average atomic mass is the weighted
average of the atomic masses of all
the naturally occurring isotopes of an
element.

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Mass Numbers of H Isotopes
Atomic #
# of
neutrons
Mass #
Protium
(99.98%)
1
0
1+0=1
Deuterium
(0.0156%)
1
1
1+1=2
Tritium
(4 x10-15%)
1
2
1+2=3
The weighted average of all three of these isotopes of
Hydrogen is it’s average atomic mass, which is
1.00797 g/mol.
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Calculating Average Atomic
Masses
The average atomic mass depends on
both the mass and the relative
abundance of each of the elements’
isotopes.
Example: Cu consists of 69.17% copper63, which has a mass of 62.929598 amu,
and 30.83% copper-65, which has a
mass of 64.927793 amu.
atomic mass units - amu = mass of the isotope
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Calculating - cont

The average atomic mass can be
calculated by multiplying the atomic
mass of each isotope by its’ relative
abundance (expressed in decimal
form) and adding the results.
(0.6917 x 62.929599 amu) + (0.3083 x 64.927793 amu) = 63.55 amu
(copper-63)+(copper-65) = the weighted average atomic mass for
Cu
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IMPORTANT! COPY INTO YOUR NOTES!!!!!
VERY HELPFUL CHART!!!
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Molar Mass

The mass of one mole of a pure substance
(element or compound/molecule) is called
the molar mass.
 Expressed in g/mol (grams per mole)
 It is numerically = to the average atomic
mass of an element or masses added up in
a compound. So, Molar Mass is the average
atomic mass or masses with the units of
g/mol.
 Example:
Li (element) has a molar mass of 6.94 g/mol
(grams (mass) per mole)
 6.94 g is the mass of 1 mole of Lithium.
 Or, for every 1 mole of Li it has a mass of 6.94g
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Conversions - Moles to Gram - Gram to Moles
Using Molar Mass as a Conversion Factor
How many grams are there in 3.6 mol of Ca?
Conversion factor - choice of two
1 mole Ca or
40.08g Ca
Write what you are
Given with unit and
40.08g Ca
1 mol Ca
substance!
3.6 mol Ca
X 40.08g Ca
1 mol Ca
= 144.29 g Ca
How many moles are there in 7.8g of Si?
7.8g Si
X
1mol Si
28.09 g Si
= 0.28` mol Si
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Molar Mass of a Compound
What is the mass of one mole of NaCl?
Na’s average atomic mass is 22.990 g/mol
 Cl’s is 35.453 g/mol
 35.453+22.990 =58.443 g/mol (grams per
mole)
 AS YOU ADD UP THE MASSES, MAKE SURE
THAT YOU USE ALL THE NUMBERS YOU
SEE IN THE PERIODIC TABLE. THEN, YOU
CAN ROUND AND REPORT TO THE
HUNDRETH! (ACCURACY AND PRECISION)
 1 mole of NaCl has a mass of 58.44 g/mol
 A molar mass of 58.44 g/mol.

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Calculation of Molar Mass of a Compound
What is the molar mass of Ba(NO3)2?
Ba = 137.327 g/mol
(NO3) = ( N = 14.0067 + 3(15.9994)) = 62.0049 g/mol
2 (62.0049) = 124.0098 g/mol
Ba
+
(NO3)2
137.327 + 124.0098 = 261.34 g/mol
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Conversions - Moles to Gram - Gram to Moles
Using Molar Mass as a Conversion Factor
COMPOUNDS/MOLECULES
How many grams are there in 3.6 mol of Ca(OH)2?
Conversion factor - choice of two
1 mole Ca(OH)2
or
74.09g Ca(OH)2
Write what you are
Given with unit and
74.09g Ca(OH)2
1 mol Ca(OH)2
substance!
3.6 mol Ca(OH)2 X 74.09g Ca(OH)2 = 266.73 g Ca(OH)2
1 mol Ca(OH)2
How many moles are there in 7.8g of SiO2?
7.8g SiO2 X
1mol SiO2
= 0.13 mol SiO2
60.08 g SiO2
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What Is A Mole?
The term mole is classified as a counting number, a
number used to specify a certain number of objects.
Pair and dozen are other examples of counting numbers.
Seldom does anyone purchase computer paper by the sheet.
Instead you buy a ream, or package of 500 sheets. At the
grocery store you buy eggs by the dozen. Many other objects
are identified in packages of this size—rolls, ears of corn, file
folders. When you buy a package of a dozen, you know you
will get twelve objects.
A mole equals 6.022 x 1023 objects!
Atoms, molecules, ions or electrons are counted in moles since they
are so small!
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Avogadro’s Number (Constant)

1 mole of any element or compound (pure
substance) contains 6.022 X 1023 particles,
molecules, atoms or ions, depending if we are
talking about a element, molecule/compound, or
bonded together atoms. This number is
Avogadro’s Number or Constant.
 1 mole of carbon (element) equals 12 grams, so in
those 12 grams there are 6.022 X 1023 particles,
 1 mole of CO2 (molecule) has a mass of 44.00g, so
in those 44.00 grams there are 6.022 x 1023
molecules.
 1 mole of Ba has 6.022 x 1023 atoms

Let's say that real clearly: one mole of
ANYTHING contains 6.022 x 1023
entities!!!
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Conversions - Moles to Atoms/Particles - Vise Versa
Memorize : 1 mole of any substance contains 6.022 x 1023
atoms/particles/molecules/ions
How many atoms are contained in 3.2 mol of F?
Possible Conversion Factors (1 mole = 6.022 x 1023 atoms)
1 mol of substance or
6.022 x 1023 atoms
6.022 x 1023 atoms
1 mol of substance
3.2 mol F
X 6.022 x 1023 atoms = 1.9 x 1024 atoms
1 mol F
of F
How many moles are in 4.2 x 1024 atoms of C?
4.2 x 1024 atoms C X 1 mol of C
6.022 x 1023 atoms of C
= 6.97 mol
of C
Note: A molar mass conversion factor is not used!!!
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Conversions - Grams to Atoms/Particles - Vise Versa
How many atoms are contained in 3.2 g of F?
3.2 g F
X 1mol F X 6.022 x 1023 atoms
19.00g F
1 mol F
= 1.0 x 1023 atoms
of F
How many grams are in 4.2 x 1024 atoms of C?
4.2 x 1024 atoms C X 1 mol of C
X 12.01 g C
6.022 x 1023 atoms of C
1 mol C
= 83.8 g
of C
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Conversions - Grams to Atoms/Particles - Vise Versa
How many atoms are contained in 5.23 g of KNO3?
5.23 g KNO3 X 1mol KNO3 X 6.022 x 1023 atoms KNO3
101.10g KNO3
1 mol KNO3
= 3.1 x 1022 atoms
of KNO3
How many grams are in 4.2 x 1024 atoms of CO?
4.2 x 1024 atoms CO X 1 mol of CO
X 28.01 g CO
6.022 x 1023 atoms of CO 1 mol CO
= 195.4 g
of CO
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Relating Mass to Number of
Atoms



The mole is the SI Unit of amount of substance.
Scientifically, A mole is the amount of pure
substance (element or compound/molecule) that
contains as many particles as there are atoms in
an isotope of Carbon -12.
The modern atomic weight scale (masses
reported in your PTs) is based on C-12.
 Example: the relative mass of a hydrogen
atom compared to a carbon atom is 1.0074.
Therefore….
 1 mole of H has a mass of 1.0074 g
 1 mole of oxygen has a mass of 15.9994 g
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