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Atoms: The Building Block of Matter
3-1 The Atom: From Philosophical
Idea to Scientific Theory
From Philosophical Idea to
Scientific Theory
 The
first idea of matter was simply that all
matter was infinitely divisible.
 Ex.
Folding a piece of paper
 Ex. You could continue to cut a piece of
copper into smaller and smaller pieces forever.
The “Particle Theory” of matter was first
supported by the Greeks and others
scientists (Democritus) around 400 B.C.
From Philosophical Idea to Scientific
He proposed that all matter that makes up the
world is composed of small, indivisible particles.
 Democritus called the building block of matter,
atomos, or the atom.
Interestingly, Aristotle did not agree with Democritus
because there was no evidence to support these claims.
Ex. You cannot continue to cut a piece of
copper into smaller pieces, eventually you
get to copper atoms which cannot be
divided any farther.
Foundations of Atomic Theory
 One
of the biggest speculations at the
beginning of the 1700’s was on whether
elements always combine in the same ratios
when forming compounds.
Reaction – the transformation of one
substance into new substances.
 Chemical
of Conservation of Matter (Mass) –
states that matter is neither created nor
destroyed during a physical or chem. Rxn.
 Law
Foundations of Atomic Theory
of Definite Proportions – a Chem.
cmpd. Contains the same elements in
exactly the same ratio regardless of the
source or sample size.
 Law
– taken from a stream.
 H2O – taken from a paper cup.
 H2O
 NaCl
always contains 39.34% Na by mass;
60.66% Cl by mass.
Foundations of Atomic Theory
 Law
of Multiple Proportions - if 2 or more
different cmpds. are composed of the same
2 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.
In CO2 : 1 g C combines with 2.66 g O
Dalton’s Atomic Theory
I like this guy because he was a schoolteacher.
Dalton proposed an explanation for the laws listed
 He believed that elements are composed of atoms,
and that only whole numbers of atoms can
combine to form cmpds.
 Ex. Water could never a formula H2.124O
Look at Dalton’s Postulates on pg. 66.
Modern Atomic Theory
 Dalton
turned Democritus’s idea into a
scientific theory that could be tested by
 Some parts of Dalton’s theory have actually
been proved NOT to be true. We will
discuss these later.
Atoms: The Building Block of Matter
3-2 The Structure of the Atom
 Although
Dalton thought the atom to be
indivisible, it is actually composed of other
subatomic particles.
 Subatomic Particles  protons, neutrons,
Discovery of the Electron
 Electron
was discovered through
experiments using cathode-ray tubes.
 A stream of charged particles flows from
the cathode to the anode in a cathode ray
tube, causing the fluorescent material inside
the tube to glow.
Discovery of the Electron
 The
negative electrode is the cathode.
 The positive electrode is the anode.
 Cathode rays were deflected by magnetic
 The ray was deflected away from a negative
field and toward a positive field.
 Particles that compose cathode rays are
(-)vely charged.
Discovery of the Electron
Thomson’s Plum
Pudding Model of the
An atom contains a
specific number of
electrons which are in
pool of positive charge.
Ex. Like the raisins in
plum pudding (or the
chocolate chips in a
He knew only that
there was positive
charge, NOT that there
were positive particles.
Discovery of the Atomic Nucleus
Rutherford’s Gold Foil Experiment:
If Thomson’s Model was correct, the alpha
particles should have passed directly through
the foil with only slight deflections. Most of
the particles acted this way, but some were
deflected at wide-angles. The wide angle
deflection could only have been caused if there
was a powerful force in the atom. He reasoned
their must be a small, dense center containing
most of the mass of the atom – nucleus.
The Gold Foil Experiment (figure 3-14)
Most particles passed
through gold without a
1 in 8000 alpha particles
These were sent in ALL
directions including
straight back!
What does this mean?
Most of the atoms positive charge, as well
as the mass is in the middle, called the
Most pass through the empty space but
occasionally one gets close enough to the
positive nucleus to deflect it.
Composition of the Atomic
Nucleus contains positive protons, neutral
 The nucleus has a net positive charge.
 Atoms are electrically neutral because the
positive nucleus is surrounded by a sea of
negative electrons.
 The number of protons in an atom
determines the atom’s identity.
 See Table 3-1 pg. 74
Question: If an atom contains positive
particles, what keeps the atom together?
Don’t like charges repel each other?
Strong Nuclear Force
 Binding Energy
 Size of the atom
Atoms: The Building Block of
3-3 Counting Atoms
The Structure of the Atom
The atom has a positively changed central
Contains Protons and Neutrons
 Protons are positive, equal and opposite to
 Neutrons do not carry a charge and are slightly
more massive
1 proton has the mass of about 2000 electrons
Electrons move in space
around the nucleus
Rutherford visualized it
as a mini solar system.
Atomic Numbers
Henry Moseley found that atoms contain
unique positive charge in their nucleus.
 The number of protons is called the atomic
The atomic number indicates protons.
Chlorine has 17 protons = atomic number
How many protons and electrons are in a
magnesium atom?
What is the name of the element that has
atoms that contain 11 protons.
When an atom gains or loses electrons it
acquires a charge
Fewer electrons means positive charge
 More electrons means negative charge
Charge of ion = # protons - # electrons
Write the chemical symbol for the ion with
9 protons and 10 electrons
What is the symbol of the ion with 13
protons and 10 electrons?
Answer F-
Answer Al3+
7 Protons and 10 electrons?
N 3-
Dalton said all atoms of an element are the
Not quite true, ISOTOPES have a different
number of neutrons
In nature, elements are almost always found
as a mixture of isotopes
 Isotopes are usually in the same
To identify isotopes more specifically
Use the Mass Number
 Mass Number = (# protons) + (# neutrons)
To identify an isotope chemists write the
mass number behind the element symbol for
example Cl-37 indicates that this chlorine
has 20 neutrons, it is written in symbol form
as 3717Cl
 Cl-35 has 18 neutrons and is written as
35 Cl
The Mass of an Atom
Measured by Atomic Mass Units (AMU)
The atomic mass is approximately the same
as the sum of protons and neutrons
This is not very precise so Scientists define
it more precisely.
1 amu is equal to 1/12 the mass of a carbon12 atom.
= 1.66 x 10 –24 grams
 Carbon 12 is the only element with an
AMU equal to protons and neutrons,
because of isotopes
The average mass of an element’s atoms is
called the atomic mass.
AM = (mass isotope x abundance)+(mass
isotope x abundance)+...
Fundamental Subatomic Particles
Mass (g)
+1.602 x 10-19
1.673 x 10-24
1.675 x 10-24
-1.602 x 10-19
9.109 x 10
Relative Mass to Numbers of
Mole – is the amount of a substance that
contains as many particles as there are
atoms in exactly 12 g of carbon-12.
SI unit for amount.
It is a unit which relates atoms and masses.
Relative Mass to Numbers of
Avogadro’s number – is the number of
particles in exactly 1 mole of a substance.
6.02 x 1023
 1 mole of carbon = 6.02 x 1023 atoms
 2 moles of silver = 1.204 x 1024 atoms
 1 mole of water = 6.02 x 1023 water molecules
2 moles of marshmallows = 1.204 x 1024
Relative Mass to Numbers of
Here’s where it could get tricky…
1 mole of water (H2O) = 6.02 x 1023
How many atoms are in 1 mole water?
H2O is composed of 2 H and 1 O atoms = 3
total atoms.
 3 x (6.02 x 1023) = 1.81 x 1024 atoms in
Relative Mass to Numbers of
Molar Mass – the mass of one mole of a
The amount of a substance that contains
Avagadro’s number of particles.
 Usually written with unit g/mol.
 Molar masses are on the periodic table.
Numerically they are the same as atomic mass.
Ex. Molar mass of He = 4.00 g/mol
Molar mass of Al = 26.98 g/mol
Relative Mass to Numbers of
Now is when it gets fun! Gram/mole
Be sure to look at the chart on pg.82. This
is one of the most important ideas which we
cover this year. The only way to understand
this is to practice, practice, practice.
 Practice Problems pgs. 82-85