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Atoms: the building blocks of
matter
Chapter 3
Chemistry chapter 3
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The atom
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The atom – smallest piece of matter that has
the properties of an element.
Made of
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Protons
Neutrons
Electrons
Each specimen of a specific subatomic
particle is the same
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 If we split an atom, we no longer have a
specific element
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Early atomic theory - Democritus
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Greek philosopher about 400 B.C.
Gave us the word atom
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Thought
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Atomos - indivisible.
The world was made of empty space and particles
called atoms.
There were different types of atoms for different
types of materials.
Theory was not supported by experimental
evidence.
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Early atomic theory – Aristotle
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Aristotle did not believe in atoms
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thought matter was continuous
He was very influential, so Democritus’s
theory was not accepted for many centuries.
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th
17
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century
People began to express doubts in Aristotle’s
theory.
Isaac Newton and Robert Boyle published
articles stating their belief in the atomic
nature of elements, but they had no proof.
Their theory also had no ability to predict the
unknown.
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Antoine Lavoisier – late 1700s
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Law of conservation of mass
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during a chemical change in a closed system, no
mass is lost
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Joseph Proust – late 1700s to early 1800s
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Law of definite proportions
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specific substances always contain elements in
the same ratio by mass
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Law of multiple proportions
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Some elements form more than one
compound with each other.
If two or more different compounds are
composed of the same two elements, then
the ratio of their masses always contains
small whole numbers
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John Dalton – early 1800s
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Studied experimental observations of
chemical reactions
Proposed explanation of these three laws
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Dalton’s Hypothesis
1.
2.
3.
4.
5.
All matter is composed of very small particles
called atoms.
All atoms of an element are exactly alike; atoms of
different elements are very different.
Atoms cannot be subdivided, created, or
destroyed.
Atoms unite with other atoms in simple ratios to
form compounds
In chemical reactions, atoms are combined,
separated, or rearranged.
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Did Dalton’s theory work?
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Conservation of mass
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the atoms are simply rearranged because they
cannot be created or destroyed
Laws of definite and multiple proportions
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Only whole atoms can combine, giving small
whole numbers in ratios
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Gas research
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J.L. Gay-Lussac
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Under constant temperature and pressure
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Volumes of reacting gases and gaseous products are in
a ratio of small whole numbers.
Amadeo Avogadro explained Gay-Lussac’s
work with Dalton’s theory.
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Equal volumes of gases, under the same
temperature and pressure, have the same number
of molecules.
Helped Dalton’s theory get accepted
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Dalton’s theories
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Atomic theory and law of multiple proportions
have been tested and accepted as correct.
However, there some major exceptions to the
rules.
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Splitting atoms
Different atoms of the same element
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Discussion
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Section review on page 69
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Cathode tubes
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Anode – positive electrode
Cathode – negative electrode
When the tube is on, cathode rays appear
that begin at the cathode and travel to the
anode.
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Cathode rays and electrons
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1897 – J.J Thomson tested cathode rays and
discovered that they were electrons.
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Rays turned a paddlewheel – they had mass
Rays deflected by a magnet just like currentcarrying wire – they were negatively charged
He determined the ratio of the electron’s
charge to its mass.
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Charge on an electron
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Robert Millikan’s famous oil drop experiment.
Tiny oil drops fell through a chamber
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gravitational force offset by applying an opposing
electrical force.
Charge on oil drops determined
This charge was always a whole number
multiple of one small charge
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Charge on an electron
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This small charge was the charge on one
electron.
This is now the standard unit of negative
charge (1-). It can be written e-.
e- can also represent an electron
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Mass of an electron
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Using Thomson’s ratio and Millikan’s charge,
determined to be 9.1 x 10-31 kg
It was found that it’s mass is only 1/1837 the
mass of the lightest atom known – the
hydrogen atom.
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Most of the mass must be somewhere else
Since atoms are neutral, there must be some
positive charge
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Thomson’s plum pudding model
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In this model, the raisins were the electrons
and the pudding was the positive charge.
Sort of like chocolate chip cookie dough.
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The chips are the electrons and the dough is the
positive charge.
Explained the experiments that had been
done so far.
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Testing the plum pudding model
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See page 72
fired alpha particles at a very thin (a few
atoms thick) sheet of gold foil.
They expected the particles to go right
through because the spread out positive
charge in the “pudding” wouldn’t be strong
enough to deflect them.
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What happened
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Most of the particles did go right through
without being deflected at all.
Some were deflected at large angles.
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Ernest Rutherford explained it:
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the positive charge on the atom was concentrated at a
small core – now called the nucleus.
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The atom as we now “know” it
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The nucleus contains all of the positive
charge and most of the mass.
The negatively charged electrons have very
small mass and are located around the
nucleus in the electron cloud.
Most of an atom is empty space.
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Protons
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same charge as an electron; opposite sign.
standard unit of positive charge (1+)
Much larger mass than the electron:
1.67 x 10-27 kg
The number of protons determines the atom’s
identity.
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Neutrons
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Weren’t discovered until the 1930s.
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Neutral – no charge – harder to detect
Slightly more mass than a proton:
1.68 x 10-27 kg
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Nuclear or Strong Force
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The force that holds protons and neutrons
together.
It is effective only for very short distances –
about 10-15 m.
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Dalton’s theory
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Dalton thought that atoms were indivisible
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discovery of electrons, protons, and neutrons did
not fit with his theory.
Led to major revisions in atomic theory
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Isotopes
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Thomson discovered what seemed to be two
kinds of neon atoms.
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Same chemical properties; different masses.
Atoms of the same element that differ in
mass are called isotopes.
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Have the same number of electrons and protons
but different number of neutrons.
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Atomic number
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Number of protons in an atom
Atoms are electrically neutral,  the number
of electrons must equal the number of
protons.
The number of protons determines the
identity of the atom and the number of
neutrons determines the isotope.
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Modification of Dalton’s theory
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All atoms of an element contain the same
number of protons but can contain different
numbers of neutrons.
So we have to use average mass of an atom.
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Nucleons
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Particles in the nucleus – protons and
neutrons
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Mass number
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Total number of nucleons : protons plus
neutrons
Number of neutrons = mass number minus
atomic number
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Designating Isotopes
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Hyphen notation
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Uranium-235
Carbon-14
Carbon-12
The number refers to the mass number
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Nuclide
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General term for any isotope of any element
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Atomic mass units
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There must be a standard for all units of
measurement.
A Carbon-12 atom with 6 protons and 6
neutrons was chosen as the standard
12
6
C
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Atomic mass unit
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Defined as 1/12 the mass of that carbon
atom.
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Average atomic masses
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Many elements have an average atomic
mass close to the number of nucleons in their
nuclei – near whole numbers.
Some don’t – look at Chlorine
The periodic table shows average atomic
masses.
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Weighted averages
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We then use a weighted average to find the
average mass of an atom of a given element.
This is called the average atomic mass or just
atomic mass.
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Finding a weighted average
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A class of 25 students took a test. 10 of
them got 80%. 12 got 90%. 3 got 100%.
What was the average score?
Not 90% - probably less than that.
weighted

10  80%   12  90%   3 100% 
average 
25
 87.2%
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You try
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Neon has two isotopes. Neon-20 has a mass
of 19.992 amu and neon-22 has a mass of
21.991 amu. In any sample of 100 neon
atoms, 90 will be neon-20 and 10 will be
neon-22. Calculate the average atomic mass
of neon.
20.192 amu
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You try
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Compute the average atomic mass of silver,
if 51.83% of the silver atoms occurring in
nature have mass 106.905 amu and 48.17%
of the atoms have mass 108.905 amu.
107.9 amu
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The Mole
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SI unit for amount of substance
Abbreviated mol
A counting unit
6.022 x 1023 particles
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Avogadro’s number
Based on carbon-12, 12 g of C-12 contains
1 mol of atoms
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Molar mass
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The mass of 1 mol of a pure substance
g/mol
Numerically equal to the atomic mass in amu
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On the periodic table the number with a decimal is
the atomic mass in amu AND the molar mass in
g/mol
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conversions
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Grams to moles or moles to grams
Use the molar mass
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Example
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What is the mass in grams of 5.60 mol of
sulfur?
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Example
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How many moles of carbon are in a sample
with a mass of 567 g?
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Example
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How many atoms of lithium are in a sample
with a mass of 76.2 g?
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You try
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How many moles of rubidium are in
3.01 x 1023 atoms of rubidium?
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You try
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How many moles are in 0.255 g of zinc?
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You try
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What is the mass of 1.20 x 1025 atoms of
helium?
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