Download - Catalyst

10/2/2011
Chapter 2
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Fundamental Chemical Laws (2.2)
Dalton’s Atomic Theory (2.3)
Defining the Atom (2.5)
Atomic Structure (2.6)
Molecules and Ions (2.7)
The Periodic Table (2.8)
Nomenclature (2.9)
This is the outline for the content we will cover in
lecture. Please read the entire chapter.
Law of Conservation of Mass
The total mass of substances does not change during a
chemical reaction – mass is neither created nor destroyed.
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Law of Definite Proportions
No matter what its source, a particular chemical
compound is composed of the same elements in
the same parts (fractions) by mass.
WATER  H2O
No matter what the source
water is ALWAYS
1 part hydrogen to 8 parts oxygen
(by mass: one molecule is 2 g H and 16 g O)
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Law of Definite Proportions
Chemical analysis of a 9.07 g sample of calcium phosphate
shows that it contains 3.52 g of Ca. How much Ca could be
obtained from a 1.000 kg sample?
Mass fraction of Ca =
3.52 g Ca = 0.388 * 100 = 38.8%
9.07 g sample
(i.e., ANY sample of calcium phosphate is 38.8% Ca by mass )
Mass of Ca in 1.000 kg of sample =
1.000 kg sample x 38.8 kg Ca
= 0.388 kg Ca
100 kg sample
or 388 g Ca
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Law of Multiple Proportions
In a nutshell, two (or more) compounds can contain
different relative amounts of the same elements:
If elements A and B react to form two compounds, the
different masses of B that combine with a fixed mass of A
can be expressed as a ratio of small whole numbers.
(Evidence of the existence tiny individual particles…)
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Law of Multiple Proportions
Mass of oxygen that combines with 1.00 g of carbon:
Compound #1
1.33 g O per g C
Compound #2
2.66 g O per g C
mass of O in compound #2 = 2.66 g = 2
mass of O in compound #1
1.33 g
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EXACT 2:1 RATIO
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Dalton’s Atomic Theory
1.
All matter consists of tiny particles
called atoms.
2.
Atoms of an element are identical in
mass and other properties and are
different from atoms of any other
element.
3.
Compounds result from the chemical
combination of a specific ratio of
atoms of different elements.
4.
Chemical reactions involve
reorganization of the atoms – changes
in the way they are bonded. Atoms of
one element do not change and
cannot be converted into atoms of
another element during chemical
reactions.
John Dalton
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Thomson and Cathode Rays
• Thomson used partially
evacuated glass tubes to
discover the existence of
negatively charged particles
called electrons.
• Concluded that atoms must
also possess positively charged
particles, which led to the plum
pudding model.
Plum Pudding Model
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Millikan and Oil Drops
• Millikan used oil droplets and
X-rays to determine the
magnitude of charge on an
electron.
• With Thomson’s cathode ray
experiment, determined the
mass of an electron:
9.11x10-31 kg
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Rutherford Experiment
Aim:
• To study the internal
structure of the atom.
• Investigate the mass
distribution in the atom.
Procedure:
• Use a radioactive source to
bombard a thin piece of gold
foil with a particles.
• If Plum Pudding is correct, a
particles (high energy and
positively charged) would go
through the foil.
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Nuclear Atom Model
Original Theory:
– Plum pudding model
Revised Theory:
– Nucleus (dense positive
charge) at the center of the
atom.
– Large amount of space
between nucleus and
electrons.
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Atomic Structure
The Angstrom
(Å) is a unit of
atomic distance.
1. Every atom contains small, dense nucleus.
2. All of the positive charge and most of the mass
are concentrated in the nucleus.
3. The nucleus is surrounded by a large volume of
nearly empty space that makes up the rest of the
atom.
4. The rest of the atom is thinly populated by
electrons, the total charge of which exactly
balances the positive charge of the nucleus.
This is a
femtometer (fm).
If an atom was the
size of a baseball
stadium, the nucleus
would be the size of a
fly on home plate.
1 amu = 1.66 x 10-24 g
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What roles do the different particles play?
# Protons = chemical identity of the atom (which element is it?)
– in an electrically-neutral atom, the number of protons in the
nucleus is exactly balanced by the number of electrons around it
# Electrons = ionic character of the atom
An ion has either more or fewer electrons than the electricallyneutral atom.
– anion = more electrons, so ion is negatively-charged
– cation = fewer electrons, so ion is positively-charged
# Neutrons = isotopic character of the atom
– isotopes have the same number of protons, but different
number of neutrons (chemically indistinguishable)
– an atom of an element usually comes in at least 2 or 3 different
isotopes (sometimes more)
– usually there will be one isotope that is far more abundant than
the others
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Atomic Structure Definitions
• Atomic number (Z): the number of protons in the nucleus
of an atom
• Mass number (A): the sum of the numbers of protons and
neutrons in the nucleus of an atom
• Atomic Mass: an average of the atomic masses of the most
common isotopes
Nuclear
Symbol
A
X
Z
In the periodic
table...
Atomic
symbol (X)
8
O
16.00
For example:
16
O
8
or 16O
Atomic number (Z)
Atomic symbol (X)
Atomic Mass
(related to A)
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Isotopes of sodium
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Available on the course website under “Exam Info” and “Lecture Notes”
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Let’s count some particles
Cobalt-60
Uranium-238
60
# protons
# electrons
27
27
60 – 27 = 33
U
92
92
238 – 92 = 146
Co
238
# neutrons
Chlorine-37
anion
37
Cl–
17
18
37 – 17 = 20
Copper-63
cation
63
Cu2+
29
27
63 – 29 = 34
Copper-65
cation
65
Cu2+
29
27
65 – 29 = 36
Isotopes and Atomic Mass
• Atoms of the same element that differ in mass
(e.g. 12C, 13C, 14C)
– isotopes are the same element
– isotopes have the same number of protons
– isotopes differ in the number of neutrons, and
therefore they differ in mass (more on amu in Ch 3).
• Many isotopes occur in nature. Most natural isotopes are
not radioactive, nor are they necessarily harmful.
• A sample of an element will contain some percentage of all
its isotopes.
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Modern Reassessment of the Atomic Theory
1. All matter is composed of atoms. Although atoms are composed
of smaller particles (electrons, protons, and neutrons), the atom
is the smallest body that retains the unique identity of the element.
2. Atoms of one element cannot be converted into atoms of another
element in a chemical reaction. Elements can only be converted into
other elements in nuclear reactions in which protons are changed.
3. All atoms of an element have the same number of protons and
electrons, which determine the chemical behavior of the element.
Isotopes of an element differ in the number of neutrons, and thus
in mass number, but not in chemical behavior (much). A sample of
the element is treated as though its atoms have an average mass.
4. Compounds are formed by the chemical combination of two or more
elements in specific ratios, as originally stated by Dalton.
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