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Atoms: The Building
Blocks of Matter
Chapter 3
 Matter
 Anything that occupies
space
and has mass
 Atoms
 Tiny particles too small
to see
 Molecules
 Atoms bonded together
 States of Matter
 Solid
 Fixed Volume
 Rigid Shape
 Liquid
 Fixed volume
 Assume the shape of the
container
 Gas
 Atoms separated by large distances
 Assume the volume and shape of the
container
 Composition
 Pure substance
 One type of atom
or molecule
 Mixture
 Two or more types
of atoms or
molecules
 Mixtures
 Heterogeneous Mixture
 Two or more regions with
different composition
 Homogeneous Mixture
 Uniform composition
 Elements
 Cannot be broken into simpler
substances
 Compounds
 Composed of two or more
elements
 Antoine Lavoisier
 Law of Conservation of
Mass
 “Matter is neither created
nor destroyed”
 The transformation of a substance or substances
into one or more new substances is known as a
chemical reaction.
 Law of conservation of mass: mass is neither
created nor destroyed during ordinary chemical
reactions or physical changes
 Total amount of matter remains constant in a
chemical reaction
 58 grams of butane burns in 208 grams of oxygen to
form 176 grams of carbon dioxide and 90 grams of
water.
58 grams + 208 grams = 176 grams + 90 grams
266 grams
=
266 grams
 Energy
 Law of Conservation
of Energy
 “Energy is neither
created nor
destroyed”
 Law of definite proportions: a chemical
compound contains the same elements in exactly
the same proportions by mass regardless of the size
of the sample or source of the compound
 Law of multiple proportions: if two or more
different compounds are composed of the same two
elements, then 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
1.
2.
3.
4.
5.
All matter is composed of extremely small
particles called atoms.
Atoms of a given element are identical in
size, mass, and other properties; atoms of
different elements differ in size, mass, and
other properties.
Atoms cannot be subdivided, created, or
destroyed.
Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds.
In chemical reactions, atoms are combined,
separated, or rearranged.
 Not all aspects of Dalton’s atomic theory
have proven to be correct. We now know
that:
• Atoms are divisible into even smaller
particles.
• A given element can have atoms with
different masses.
• Atoms of any one element differ in
properties from atoms of another
element.
• All matter is composed of atoms.
Think about it….
 Even though the two shapes look different, the
characteristics of the various parts that compose
them are the same.
 The same is true with the atom.
 Though atoms of different elements display
different properties, isolated subatomic particles
have the same properties.
 An atom is the smallest particle of an element that retains the




chemical properties of that element.
The nucleus is a very small region located at the center of an
atom.
The nucleus is made up of at least one positively charged
particle called a proton and usually one or more neutral
particles called neutrons.
Surrounding the nucleus is a region occupied by negatively
charged particles called electrons.
Protons, neutrons, and electrons are often referred to as
subatomic particles.
Cathode Rays and Electrons:
• Experiments in the late 1800s showed
that cathode rays were composed of
negatively charged particles.
• These particles were named electrons.
 Charge and Mass of the Electron
• Joseph John Thomson’s cathoderay tube experiments measured the
charge-to-mass ratio of an electron.
Plum-pudding model of an atom.
• Robert A. Millikan’s oil drop experiment
measured the charge of an electron.
• With this information, scientists were able to
determine the mass of an electron.
 More detail of the atom’s structure was provided in
1911 by Ernest Rutherford and his associates Hans
Geiger and Ernest Marsden.
 The results of their gold foil experiment led to the
discovery of a very densely packed bundle of matter
with a positive electric charge.
 Rutherford called this positive bundle of matter the
nucleus.
 Except for the nucleus of the simplest type of
hydrogen atom, all atomic nuclei are made of protons
and neutrons.
 A proton has a positive charge equal in magnitude to
the negative charge of an electron.
 Atoms are electrically neutral because they contain
equal numbers of protons and electrons.
 A neutron is electrically neutral.
Properties of Subatomic
Particles
 The nuclei of atoms of different elements differ in
their number of protons and therefore in the amount
of positive charge they possess.
 Thus, the number of protons determines that
atom’s identity.
Forces in the Nucleus:
• When two protons are extremely close to
each other, there is a strong attraction
between them.
• A similar attraction exists when neutrons are
very close to each other or when protons and
neutrons are very close together.
• The short-range proton-neutron, protonproton, and neutron-neutron forces that hold
the nuclear particles together are referred to
as nuclear forces.
 The radius of an atom is the distance from the
center of the nucleus to the outer portion of its
electron cloud.
 Because atomic radii are so small, they are
expressed using a unit that is more convenient for
the sizes of atoms.
 This unit is the picometer, pm.
 Atoms of different elements have different numbers
of protons.
 Atoms of the same element all have the same
number of protons.
 The atomic number (Z) of an element is the number
of protons of each atom of that element.
•The mass number is the total number of protons and
neutrons that make up the nucleus of an isotope.
 Isotopes are atoms of the same element that have
different masses.
 The isotopes of a particular element all have the same
number of protons and electrons but different numbers
of neutrons.
 Most of the elements consist of mixtures of isotopes.
 Average atomic mass is the weighted average of the
atomic masses of the naturally occurring isotopes of an
element.
 Calculating Weighted Average Atomic Mass
 The average atomic mass of an element depends on both
the mass and the relative abundance of each of the
element’s isotopes.
 Imagine that your semester grade depends 60% on
exam scores and 40% on laboratory explorations.
 Your exam scores would count more heavily toward
your final grade.
 The atomic mass of an element is a weighted
average of the masses of the naturally occurring
isotopes of that element.
 If you had three sumo wrestlers each weighing 500
pounds and one ballerina weighing 100 pounds, what
is the average weight?
 A weighted average would be
 (3 X 500 + 1 X 100)/4 = 400
 (small sumo wrestlers or large ballerinas)
 The chlorine atom has two isotopes. One of
the isotopes weighs 35 and other isotope
weighs 37. Since they are both chlorine atoms,
they each have 17 protons (the atomic number
for chlorine). One of the isotopes has 18
neutrons and the other isotope has 20
neutrons. The average weight of
chlorine
is 35.5. Why?
17
Cl
Chlorine
35.5
 The isotope that weighs 35 comprises 75% of the atoms
of chlorine and the isotope that weighs 37 comprises
the other 25%. Convert the percentages to decimal
fractions and proceed:
 0.75 X 35 + 0.25 X 37 = 35.5
 That is the weighted average of atomic weights.
Gallium has two naturally occurring isotopes: Ga-69 with mass
68.9256 amu and a natural abundance of 60.11% and Ga-71 with
mass 70.9247 amu and a natural abundance of 39.89%. Calculate
the atomic mass of gallium.
Solution:
Convert the percent natural abundance into decimal form.
Ga-69  0.6011
Ga-71  0.3989
Atomic Mass
Atomic Mass = 0.6011 (68.9256 amu) + 0.3989 (70.9247 amu)
= 69.72 amu
 Hyphen notation: The mass number is written with a
hyphen after the name of the element.
 uranium-235
 Nuclear symbol: The superscript indicates the mass
number and the subscript indicates the atomic
number.
235
92
U



The number of neutrons is found by subtracting
the atomic number from the mass number.
mass number − atomic number = number of
neutrons
 235 (protons + neutrons) − 92 protons = 143
neutrons
Nuclide is a general term for a specific isotope of an
element.
Sample Problem
How many protons, electrons, and neutrons are there in
an atom of chlorine-37?
Solution:
atomic number = number of protons = number of
electrons
mass number = number of neutrons + number of protons
 The standard used by scientists to compare units of
atomic mass is the carbon-12 atom, which has been
arbitrarily assigned a mass of exactly 12 atomic mass
units, or 12 amu.
 One atomic mass unit, or 1 amu, is exactly 1/12 the
mass of a carbon-12 atom.
 The atomic mass of any atom is determined by
comparing it with the mass of the carbon-12 atom.
 Stores sell nails by the pound
 Easier than selling
individually
 Timely
 How many nails in a pound?
 Need a certain number for a
project
 Analogy
 How many atoms in a given
mass of an element?
 Atoms are too small to count
A hardware store customer buys 2.60 pounds
of nails. A dozen of the nails has a mass of
0.150 pounds. How many nails did the
customer buy?
Solution map:
2.60 lb. x 1 doz
x
0.150 lbs
12 nails = 208 nails
1 doz.
The customer bought 2.60 lbs of nails and received
208 nails. He counted the nails by taking a mass.
If he purchased a different size nails, the mass of 1
dozen of them would change.
The mass of the same idea works for atoms of
different elements. Different elements have
different masses.
 Mole
 6.022 x 1023
 Avogadro’s Number
 Convenient number for dealing
with atoms
 Mole vs. Dozen
 Certain number of objects
 Atoms
Amedeo
Avogadro
1776-1856
There are other numbers similar to the mole
that we use all of the time:
 pair
= two
 dozen
= 12
 score
= twenty
 Gross
= twelve dozen or 144
and least but not last:
 MOLE
= 6.02 X 1023
This is called Avogadro’s number after the man
who discovered this number.
 There are 6.02 X 1023 atoms in one mole of atoms
 There are 6.02 X 1023 molecules in one mole of
molecules
 There are 6.02 X 1023 elephants in one mole of
elephants
 There are 6.02 X 1023 doughnuts in one mole of
doughnuts
 There are 6.02 X 1023 “thingys” in one mole of “thingys”
!
 1 mole of atoms = 6.022 X 1023 atoms
 Convert 3.5 moles of He to He atoms
3.5 moles X 6.022 X 1023 atoms = 2.1 X 1024 atoms
1 mole of atoms
 Nail Analogy
 Mass of nails
 Number of nails
 Molar Mass
 Converts number of atoms to mass
 Molar mass in grams per mole
 Copper
 63.55 amu
 63.55 grams per mole
 Sulfur
 32.06 amu
 32.06 grams per mole (g/mol)
 6.022 X 1023 atoms
 Carbon
 12.01 amu
 12.01 grams per mole (g/mol)
 6.022 X 1023 atoms
 Gram/Mole Conversions
• Chemists use molar mass as a conversion factor
in chemical calculations.
• For example, the molar mass of helium is 4.00 g
He/mol He.
• To find how many grams of helium there are in
two moles of helium, multiply by the molar mass.
4.00 g He
2.00 mol He 
= 8.00 g He
1 mol He
Sample Problem:
What is the mass in grams of 3.50 mol of the element
copper, Cu?
Sample Problem:
A chemist produced 11.9 g of aluminum, Al. How many
moles of aluminum were produced?
Sample Problem:
How many moles of silver, Ag, are in 3.01  1023 atoms of
silver?
Carbon and Diamond:
 How many carbon atoms in 0.58 grams of diamond?
 (Use conversion factors to solve.)
grams  moles  atoms
0.58 g X 1 mol X 6.022 X 1023 atoms = 2.9 X 1022 atoms
12.01 g
1 mol
Sample Problem
What is the mass in grams of 1.20  108 atoms of copper,
Cu?