Download 3-3 Molar Mass

Chapter 3
Atoms: The Building Blocks of
Matter
3-1 An Ancient Idea
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Water: atoms smooth
and round so it flowed
with no permanent
shape
Fire: atoms
thorny,making burns
painful
Earth: atoms rough
and jagged so they
held together to make
hard stable substances
3-1 Opposing View
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Did not believe in
atoms
Believed matter
was continuous
Very influential
Neither
Democritus nor
Aristotle supported
their ideas with
experimentation.
3-1 Law of Conservation
of Mass
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Antoine Laurent Lavoisier
(1743-1794) – the father of
modern chemistry
Mass is neither created nor
destroyed during ordinary
chemical reactions or
physical changes
Arrested by French
Revolutionary Tribunal for
his membership in the
Ferme Generale and
executed
3-1 Law of Definite Proportions
(Constant Composition)
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A chemical compound
contains the same
proportions by mass
regardless of the size
of the sample or
source of the
compound
Sodium chloride (NaCl)
is always 39.34%
sodium and 60.66%
chlorine
3-1 John Dalton
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English
schoolteacher –
1808
Proposed
explanation for
these laws
3-1 Dalton’s Atomic
Theory
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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.
3-1 Dalton and Conservation of
Mass
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In chemical reactions, atoms are rearranged – bonds are broken
and new bonds are formed.
Same number of atoms of each type exist before and after the
reaction.
No atoms are created or destroyed.
3-1 Dalton and Definite
Proportions
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Each kind of atom has its own mass
and atoms in a compound always exist
in fixed whole number ratios.
Sodium chloride always 1:1
sodium:chlorine
Na: 23, Cl: 35.5
Na: 23/58.5 *100 = 39.34%
Cl: 35.5/58.5 *100 = 60.66%
3-1 Law of Multiple
Proportions
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If two or more different compounds are composed of the same
elements, then the ratio of the masses of the second element
combined with a certain mass of the first element is always a whole
number ratio.
3-1 Dalton and Multiple
Proportions
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Each kind of atom has a unique mass and
atoms in compounds are combined in fixed
whole number ratios.
3-1 Impact of Dalton’s
Atomic Theory
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Related Democritus’s idea of atoms to
the MEASURABLE property of mass –
atomic theory could then be TESTED
by EXPERIMENT
Dalton’s theory has been modified
over the years as new information has
become available, but the fundamental
principles hold true today
3-1 Dalton’s Atomic
Model
3-2 Structure of the Atom
ATOM - The
smallest
particle of an
element that
retains the
chemical
properties of
that element.
3-2 Discovery of the
Electron
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1897
J. J. Thomson,
English physicist
Did experiments
with cathode ray
tubes – glass tubes
containing gases at
low pressure
3-2 Cathode Ray Tubes
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Cathode – negative electrode, anode –
positive electrode
When a current is passed through the tube,
the end opposite the cathode glows
Glow caused by stream of particles called
cathode ray because is originated at
cathode – ray traveled from cathode to
anode
3-2 CRT Experiments
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Observation – a paddle wheel placed on rails
between electrodes rolled along rails from cathode
to anode
Conclusion – existence of cathode ray supported,
cathode ray has MASS
3-2 CRT Experiments
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Observations – cathode rays deflected by a
magnetic field in the same way as a wire carrying
an electric current (cathode ray acts NEGATIVE);
deflected away from negatively charged objects
Conclusion – cathode ray is made of negatively
charged particles
3-2 Charge and Mass of
the Electron
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J. J. Thomson called these tiny
negatively charged particles
“electrons”
He calculated the mass to charge
ratio of the electron – it is always the
same no matter what metal is used for
the electrodes
3-2 Charge and Mass of
the Electron
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Robert Millikan,
American physicist
1909
Showed that mass
of electron is
1/2000 the mass of
the simplest known
atom (hydrogen)
9.109 x 10-31 kg
3-2 Important Points
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Atoms contain tiny negatively charged
particles called electrons.
Electrons are present in atoms of all
elements.
Atoms are divisible and one of parts is
negatively charged.
Because atoms are neutral, there must also
be a positive component.
Because electrons have such small mass,
atoms must contain other parts that make
up most of their mass.
3-2 The Plum Pudding
Model of the Atom
3-2 Discovery of the
Atomic Nucleus
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1911
Ernest Rutherford,
New Zealand (with
Hans Geiger and
Ernest Marsden)
Important
experiment
providing more
detail into atom’s
structure.
3-2 The Gold Foil
Experiment
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Thin, gold foil
bombarded with
alpha particles
(positively
charged
particles with 4x
mass of
hydrogen atom)
3-2 The Gold Foil
Experiment - Results
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Most particles went right through gold
foil
A few were slightly deflected
A few bounced off the gold foil!
3-2 The Gold Foil
Experiment - Conclusions
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Particles that pass through foil – hit nothing
Particles slightly deflected – come close to a
positive charge
Particles that bounce back – hitting a positive
charge
3-2 Rutherford’s
Explanation
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The atom has a densely packed
bundle of matter with a positive
charge (he called it the nucleus).
The nucleus contains all of the positive
charge and most of the mass.
The nucleus has very little volume –
the atom is mostly empty space.
3-2 Rutherford’s
Atomic Model
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Small nucleus in the
center
Electrons orbit
nucleus like planets
around the Sun.
Sometimes called
planetary atomic
model.
3-2 Composition of the
Atomic Nucleus
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Nuclei contain two kinds of particles – protons and neutrons
Electrons are outside the nucleus, in the electron cloud
Table 3-1 on p. 74
Particle
Symbol
electron
e-
proton
p+
neutron
n0
Relative Mass
Relative Actual
Charge Number Mass
Mass
(amu)
(kg)
-1
+1
0
0
1
1
0.0005486 9.109x10-31
1.007276
1.673x10-27
1.008665
1.675x10-27
3-2 Forces in the Nucleus
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Like charges generally repel each
other BUT at very close range there is
an attraction between them
Nuclear Forces – short range protonneutron, proton-proton, and neutronneutron forces that hold the nuclear
particle together
3-2 The Sizes of Atoms
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Expressed as atomic radius – distance
from center of nucleus to outer edge
of electron cloud
Measured in picometers (10-12 m)
Generally range from 40 pm to 270 pm
Nuclear radius is ~0.001 pm (like a
dime in a football field)
3-3 Counting Atoms
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Atoms of the same
element all have the
same number of
protons – this is the
atomic number
Atomic number is
found on the periodic
table – elements
arranged in order of
increasing atomic
number
3-3 Isotopes
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Atoms of the same element that have
different masses (different numbers of
neutrons) – same chemical behavior
3-3 Isotopes
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Most of the elements
consist of mixtures of
isotopes.
To identify an isotope,
must know atomic
number AND mass
number (protons +
neutrons)
Isotopes usually
identified by specifying
mass number
3-3 Isotopes
Nuclide – general term for any isotope of any element
Hyphen Notation
neon-20, uranium-235
Nuclear Symbol
3-3 Relative Atomic Mass
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One atom has been chosen as a
standard, masses of all other atoms
expressed in relation to this standard
Atomic mass unit – exactly 1/12 the
mass of a carbon-12 atom
Atomic mass of any other atom is
determined by comparing it with the
mass of carbon-12
3-3 Average Atomic Mass
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Most elements occur naturally as mixtures
of isotopes
The percentage of each isotope in an
element is constant – same no matter
where sample comes from
Average atomic mass is the weighted
average of the atomic masses of the
naturally occurring isotopes of an element
3-3 Calculating Average
Atomic Mass
Copper has 2 isotopes – copper-63 and
copper-65
Copper-63: 69.17%, 62.929599 amu
Copper-65: 30.83%, 64.927793 amu
(.6917)(62.929599 amu) + (.3083)(64.927793 amu) = 63.55 amu
3-3 Average Atomic Mass
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Atomic masses that appear on the periodic table are AVERAGE atomic masses
3-3 Relating Mass to
Number of Atoms
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THE MOLE!
The amount of a
substance that
contains as many
particles as there are
atoms in exactly 12 g
of carbon-12
The mole is a counting
unit, like a dozen or a
ream
1 mole of something is
6.022 x 1023
3-3 Avogadro’s Number
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12 g of carbon-12
contains exactly
6.0221367 x 1023
atoms.
Named Avogadro’s
number after
Amedeo Avogadro,
Italian scientist
3-3 Molar Mass
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A mole is the amount of a substance that contains
Avogadro’s number of particles
The mass of one mole of a pure substance is called
the molar mass of that substance (g/mol)
3-3 Molar Mass
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To find the molar
mass of an atom,
find the atomic
mass in amu on the
periodic table and
change the unit to
g/mol
The molar mass is
numerically equal
to the atomic mass
element
Atomic
mass
(amu)
Molar
mass
(g/mol)
carbon
12.011
12.011
copper
63.55
63.55
iron
55.85
55.85
neon
20.18
20.18
3-3 Gram/Mole
Conversions
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Molar mass can be used as a
conversion factor
Can convert between mass and moles
for any substance
Moles to Grams Multiply (by molar
mass)
Grams to Moles Divide (by molar
mass)
3-3 Conversions with
Avogadro’s Number
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If amount in moles is known, can
calculate number of particles (and vice
versa)
Moles to Particles Multiply (by
Avogadro’s number)
Particles to Moles Divide (by
Avogadro’s number)
Mole Conversions