Atoms: The

Building Blocks
of Matter

1
The Atom: From Philosophical Idea to
Scientific Theory
WHEN YOU CRUSH A LUMP OF SUGAR, YOU
CAN SEE THAT IT IS MADE UP OF MANY
SMALLER PIECES OF SUGAR. YOU MAY GRIND
THESE PARTICLES INTO A VERY FINE POWDER,
BUT EACH TINY PIECE IS STILL SUGAR.

2
NOW SUPPOSE YOU
DISSOLVE THE SUGAR IN
WATER. THE TINY
PARTICLES SEEM TO
DISAPPEAR COMPLETELY.

3
EVEN IF YOU LOOK AT THE SUGAR-WATER
SOLUTION THROUGH A POWERFUL MICROSCOPE
YOU CANNOT SEE ANY SUGAR PARTICLES. YET
IF YOU WERE TO TASTE THE SOLUTION, YOU’D
KNOW THAT THE SUGAR IS STILL THERE.

4
OBSERVATIONS LIKE THESE LED EARLY
PHILOSOPHERS TO PONDER THE FUNDAMENTAL
NATURE OF MATTER. IS IT CONTINUOUS AND
INFINITELY DIVISIBLE, OR IS IT DIVISIBLE ONLY
UNTIL A BASIC, INVISIBLE PARTICLE THAT
CANNOT BE DIVIDED FURTHER IS REACHED?

5
 Foundations of
Atomic Theory
Nearly all chemists in late 1700s
accepted the definition of an element
as a substance that cannot be broken
down further
6
They knew about chemical reactions
but there was great disagreement
as to whether elements always
combine in the same ratio when
forming a specific compound
7
 Law of
Conservation of
Mass
With the help of improved balances,
investigators could accurately measure
the masses of the elements and
compounds they were studying 8
This lead to discovery of several
basic laws
Law of conservation of mass -
states that mass is neither
destroyed nor created during
ordinary chemical reactions or physical
changes 9
Law of Definite
Proportions
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 10
EACH OF THE SALT
CRYSTALS SHOWN
HERE CONTAINS
EXACTLY
39.34% SODIUM AND
60.66% CHLORINE BY
MASS.

11
 Law of Multiple
Proportions
Two elements sometimes combine to form more
than one compound
For example, the elements carbon and oxygen
form two compounds, carbon dioxide and carbon
monoxide
Consider samples of each of these compounds, each
containing 1.0 g of carbon 12
In carbon dioxide, 2.66 g of oxygen
combine with 1.0 g of carbon
In carbon monoxide, 1.33 g of oxygen
combine with 1.0 g of carbon
The ratio of the masses of oxygen in
these two compounds is exactly 2.66 to
1.33, or 2 to 1
13
 1808 John Dalton
Proposed an explanation for the law of
conservation of mass, the law of definite
proportions, and the law of multiple proportions
He reasoned that elements were composed of
atoms and that only whole numbers of atoms can
combine to form compounds
14
 Dalton’s Atomic
Theory
1. All matter is composed of extremely
small particles called atoms.
2. Atoms of a given element are identical
in size, mass, and other properties; atoms
of different elements differ in size, mass,
and other properties.
15
3. Atoms cannot be divided, created or
destroyed.
4. Atoms of different elements combine
in simple whole-number ratios to form
chemical compounds.
5. In chemical reactions, atoms are
combined, separated, or rearranged.
16
 Modern Atomic
Theory
Dalton turned Democritus’s idea into a
scientific theory which was testable
Not all parts of his theory have been
proven correct

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Ex. We know atoms are divisible into
even smaller particles - subatomic
particles

18
We know an element can have atoms
with different masses - isotopes

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 Isotone
different atoms with the same number of
neutrons

20
 how to remember?
isotoPe - same number of Protons
isotoNe - same number of Neutrons

21
 isobar
different elements, same atomic mass

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ALTHOUGH JOHN DALTON THOUGHT ATOMS WERE
INDIVISIBLE, INVESTIGATORS IN THE LATE 1800S
PROVED OTHERWISE. AS SCIENTIFIC ADVANCES
ALLOWED A DEEPER EXPLORATION OF MATTER,
IT BECAME CLEAR THAT ATOMS ARE ACTUALLY
COMPOSED OF SEVERAL BASIC TYPES OF
SMALLER PARTICLES AND THAT THE NUMBER AND
ARRANGEMENT OF THESE PARTICLES WITHIN AN
ATOM DETERMINE THAT ATOM’S CHEMICAL
PROPERTIES.
The Structure of the Atom
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TODAY WE DEFINE AN ATOM AS
THE SMALLEST PARTICLE OF AN
ELEMENT THAT RETAINS THE
CHEMICAL PROPERTIES OF THAT
ELEMENT.
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 All atoms consist of
two regions

Nucleus - very small region located near the center

of an atom
In the nucleus there is at least one positively
charged particle called the proton
25
 Usually at least one neutral
particle called the neutron
Surrounding the nucleus is a
region occupied by negatively
charged particles called electrons
26
 Discovery of the
Electron
Resulted from investigations into the relationship
between electricity and matter
Late 1800s, many experiments were performed:
electric current was passed through different gases
at low pressures

27
These experiments were carried out in
glass tubes known as cathode-ray tubes

28
 Cathode Rays and
Electrons
Investigators noticed that when current was passed
through a cathode-ray tube, the opposite end of
the tube glowed

29
Hypothesized that the glow was caused by a
stream of particles, which they called a cathode ray
The ray traveled from the cathode to the anode
when current was passed through the tube

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 Observations
1. cathode rays deflected by magnetic field in same
way as wire carrying electric current (known to
have negative charge)
2. rays deflected away from negatively charged
object

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32
Observations led to the hypothesis that the
particles that compose cathode rays are negatively
charged
Strongly supported by a series of experiments
carried out in 1897 by the English physicist Joseph
John Thomson

33
He was able to measure the ratio of the charge
of cathode-ray particles to their mass
He found that this ratio was always the same,
regardless of the metal used to make the
cathode or the nature of the gas inside the
cathode-ray tube

34
Thomson concluded that all
cathode rays are composed of
identical negatively charged
particles, which were later
named electrons
35
 Charge and Mass of
the Electron
Confirmed that the electron carries a
negative electric charge
Because cathode rays have identical properties
regardless of the element used to produce
them, it was concluded that electrons are
present in atoms of all elements
36
Cathode-ray experiments provided
evidence that atoms are divisible and
that one of the atom’s basic
constituents is the negatively charged
electron

37
Thomson’s experiment revealed that the
electron has a very large charge for its tiny
mass
Mass of the electron is about one two-
thousandth the mass of the simplest type of
hydrogen atom (the smallest atom known)
Since then found that the electron has a
mass of 9.109 × 10 − 31 kg, or 1/1837 the mass
of the hydrogen atom 38
Based on information about electrons, two
inferences made about atomic structure
b/c atoms are neutral they must have
positive charge to balance negative
electrons
b/c electrons have very little mass, atoms
must have some other particles that
make up most of the mass 39
 Thomson’s Atom
Plum pudding model (based on English dessert)

Negative electrons spread evenly through positive

charge of the rest of the atom
Like seeds in a watermelon 40
 Discovery of Atomic
Nucleus
1911 by New Zealander Ernest Rutherford and his
associates Hans Geiger and Ernest Marsden
Bombarded a thin, gold foil with fast-moving alpha
particles (positively charged particles with about four
times the mass of a hydrogen atom)

41
Assume mass and charge were
uniformly distributed throughout
atoms of gold foil (from Thomson’s
model of the atom)
Expected alpha particles to pass
through with only slight deflection
42
 What Really
Happened…
Most particles passed with only slight
deflection
However, 1/8,000 were found to
have a wide deflection

43
Rutherford explained later it was “as
if you have fired a 15-inch artillery
shell at a piece of tissue paper and
it came back and hit you.”

44
45
 Explanation
After 2 years, Rutherford finally came up
with an explanation
The rebounded alpha particles must have
experienced some powerful force within the
atom
46
The source of this force must occupy a very
small amount of space because so few of
the total number of alpha particles had been
affected by it
The force must be caused by a very densely
packed bundle of matter with a positive
electric charge
Rutherford called this positive bundle of
matter the nucleus 47
Rutherford had discovered that the
volume of a nucleus was very small
compared with the total volume of an
atom

48
If the nucleus were the size of a
marble, then the size of the atom
would be about the size of a football
field 49
But where were the electrons?
Rutherford suggested that the electrons
surrounded the positively charged nucleus
like planets around the sun
He could not explain, however, what
kept the electrons in motion around
the nucleus
50
Rutherford’s Atom

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 Composition of
Atomic Nucleus
Except hydrogen, all atomic nuclei
made of two kinds of particles
Protons
Neutrons
52
Protons = positive
Neutrons = neutral
Electrons = negative

Atoms are electrically neutral, so number of

protons and electrons IS ALWAYS THE SAME
53
The nuclei of atoms of different elements
differ in the number of protons they
contain and therefore in the amount of
positive charge they possess

So the number of protons in an atom’s

nucleus determines that atom’s identity
54
 Forces in the
Nucleus
Usually, particles that have the
same electric charge repel one
another
Would expect a nucleus with more
than one proton to be unstable
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When two protons are extremely close to each
other, there is a strong attraction between them
Nuclear forces - short-range proton-neutron,
proton-proton, and neutron-neutron forces that
hold the nuclear particles together

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 The Sizes of atoms
Area occupied by electrons is electron cloud – cloud
of negative charge
Radius of atom is distance from center of nucleus
to outer portion of cloud

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Unit – picometer (10 -12 m)
Atomic radii range from 40-270 pm
Very high densities – 2 x 108 tons/cm 3

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 Counting Atoms
Consider Neon, Ne, the gas used in many
illuminated signs. Neon is a minor part
of the atmosphere. In fact, dry air
contains only about 0.002% Ne. And yet
17
there are about 5 x 10 atoms of neon
present in each breath you inhale.
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In most experiments, atoms are
too small to be measured
individually. Chemists can analyze
atoms quantitatively, however, by
knowing fundamental properties of
the atoms of each element.
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In this section, you will be
introduced to some of the basic
properties of atoms. You will then
discover how to use this
information to count the number
of atoms of an element in a
sample with a known mass. 61
You will also become familiar the
the mole, a special unit used by
chemists to express amounts of
particles, such as atoms and
molecules.
62
 Atomic Number
Atoms of different elements have
different numbers of protons
Atomic number (Z) - number of
protons in the nucleus of each atom
of that element
63
Shown on
periodic table
Atomic number
identifies an
element
64
65
 Mass Number
Mass number -
total number of
protons and
neutrons in the
nucleus of an
isotope 66
Shows the composition of a nucleus as the
isotope’s nuclear symbol
235
Uranium-235 is written as 92U

Nuclide – general term for specific isotope of

an element

67
The superscript
indicates the mass
number

68
the subscript
indicates the
atomic number

69
70
 Relating Mass to
Numbers of Atoms
The relative atomic mass scale
makes it possible to know how
many atoms of an element are
present in a sample of the
element with a measurable mass
71
Three very important concepts provide
the basis for relating masses in grams
to numbers of atoms
1. The mole
2. Avogadro’s number
3. Molar mass
72
 The Mole
SI unit for an amount of substance
(like 1 dozen = 12)
Mole (mol) - amount of a substance
that contains as many particles as
there are atoms in exactly 12g of C-12
73
 Avogadro’s Number
Avogadro’s number - the number of
particles in exactly one mole of a pure
substance

6.022 x 10 23

How big is that? 74

If 5 billion people worked to count
the atoms in one mole of an
element, and if each person counted
continuously at a rate of one atom
per second, it would take about 4
million years for all the atoms to be
counted
75
 Molar Mass
Molar mass - the mass of one mole of a pure
substance
Written in unit g/mol

Found on periodic table (atomic mass)

Ex. Molar mass of H = 1.008 g/mol
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 Example

How many grams of helium are there in 2 moles

of helium?
2.00 mol He x = ? g He

2.00 mol He x 4.00 g He = 8.00 g He

1 mole He
77
 Practice Problems
What is the mass in grams of 3.50 mol of the
element copper, Cu?
222g Cu

78
What is the mass in grams of 2.25 mol of the
element iron, Fe?
126g Fe
What is the mass in grams of 0.375 mol of the
element potassium, K?
14.7g K

79
What is the mass in grams of 0.0135 mol of the
element sodium, Na?
0.310g Na
What is the mass in grams of 16.3 mol of the
element nickel, Ni?
957g Ni

80
1. A chemist produced 11.9 g of aluminum, Al. How
many moles of aluminum were produced?
0.441 mol Al
2. How many moles of calcium, Ca, are in 5.00 g
of calcium?
0.125 mol Ca
3. How many moles of gold,Au, are in 3.60 ×
− 10
10 g of gold?
1.83 × 10− 12 mol Au
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 Conversions with
Avogadro’s Number
How many moles of silver, Ag, are in 3.01 x 1023

atoms of silver?
Given: 3.01x10 23 atoms of Ag
Unknown: amount of Ag in moles

Ag atoms × = moles Ag
82
 23
3.01x10 atomsAg x 1 mol Ag =0.500 mol Ag
23
6.02x10 atomsAg

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 Practice problems
1. How many moles of lead, Pb, are in 1.50×10 12

atoms of lead?
2.49×10− 12 mol Pb

84
2. How many moles of tin, Sn, are in 2500 atoms
of tin?
4.2×10− 21 mol Sn

85
3. How many atoms of aluminum, Al, are in 2.75
mol of aluminum?
1.66×1024 atoms Al

86
1. What is the mass in grams of 7.5 × 1015

atoms of nickel, Ni?

−
7.3×10 g7 Ni
2. How many atoms of sulfur, S, are in 4.00 g of
sulfur?
22
7.51×10 atoms S
3. What mass of gold,Au, contains the same
number of atoms as 9.0 g of aluminum,Al?
66g Au
87