Transcript Chapter 1: Atoms and Elements

Chapter 1: Atoms and Elements
KTT 111/3 – Inorganic Chemistry I
Dr. Farook Adam
August 2005
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Chapter 1: Atoms and Elements
 Chemistry
is a science that studies the
composition and properties of matter
 Matter is anything that takes up space
and has mass
 Mass is a measure of the amount of
matter in a sample
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
Chemistry holds a unique place among the
sciences because all things are composed
of chemicals
A knowledge of chemistry will be valuable
whatever branch of science you study
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
Chemistry is constantly changing as new
discoveries are made by researchers
 Researchers use a commonsense approach to
the study of natural phenomena called the
scientific method
 A scientific study normally:
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Begins with a question about nature
Involves a search of the work of others
Requires observing the results of experiments
Often results in a conclusion, or a statement based
on what is thought about a series of observations
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 Experiments
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provide empirical facts
Facts are called data
A broad generalization based on the
results of many experiments is called a
(scientific) law
Laws are often expressed as mathematical
equations
Laws summarize the results of
experiments
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
Theoretical models attempt to explain why
substances behave as they do
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A hypothesis is a tentative explanation
A theory is an experimentally tested explanation
of the behavior of nature
The scientific method is dynamic:
observations lead to laws, which
suggest new experiments, which
may lead to or change a hypothesis,
which may produce a theory.
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 Chemical
substances are comprised of
atoms
 Atoms combine to form molecules
which can be represented in a number
of ways, including:
(a) Using chemical symbols and lines for “connections”
(b) A 3-D ball-and-stick model
(c) A 3-D space-filling model
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 Characteristics
or properties of
materials distinguish one type of
substance from another
 Properties can be classified as physical
or chemical
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Physical properties can be observed
without changing the chemical makeup of
the substance
Chemical properties involve a chemical
change and result in different substances
Chemical changes are described by
chemical reactions
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 Properties
can also be described as
intensive or extensive

Intensive properties are independent of
sample size
• Examples: sample color and melting point

Extensive properties depend on sample
size
• Examples: sample volume and mass
 In
general, intensive properties are
more useful in identifying a substance
 Matter is often classified by properties
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
The three common physical states of matter
have different properties:
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Solids have a fixed shape and volume
• Particles are close together and have restricted motion

Liquids have indefinite shape but fixed volume
• Particles are close together but are able to flow
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Gases have indefinite shape and volume
• Particles are separated by lots of empty space
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 Elements
are substances that cannot
be decomposed by chemical means into
simpler substances
 Each
element is assigned a unique
chemical symbol

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Most are one or two letters
First letter is always capitalized
All remaining letters are lowercase
Names and chemical symbols of the
elements are listed on the inside front
cover of the book
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 Compounds
are substances formed from
two or more different elements combined
in a fixed proportion by mass
 The physical and chemical properties of
a compound are, in general, different
than the physical and chemical
properties of the elements of which it is
comprised
 Elements and compounds are examples
of pure substances whose composition
is the same, regardless of source
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 A mixture
consists of varying amounts
of two or more elements or compounds
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Homogeneous mixtures or solutions
have the same properties throughout the
sample
Heterogeneous mixtures consist of two
or more phases
 Matter
can be classified:
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 We
take for granted the existence of
atoms and molecules
 The concept of the atom had limited
scientific usefulness until the discovery
of two important laws: the Law of
conservation of mass and the Law of
Definite Proportions
 These laws summarized the results of
the experimental observations of many
scientists
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 Law
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No detectable gain or loss of mass occurs
in chemical reactions. Mass is conserved.
 Law

of Conservation of Mass:
of Definite Proportions:
In a given chemical compound, the
elements are always combined in the same
proportions by mass.
 In
the sciences mass is measured in
units of grams (symbol, g)
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One pound equals 453.6 g
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The laws of conservation of mass and
definite proportions provided the
experimental foundation for the atomic
theory
Dalton’s Atomic Theory (~1803 AD):
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Matter consists of tiny particles called atoms.
Atoms are indestructible. In chemical reactions, the
atoms rearrange but they do not themselves break
apart.
In any sample of a pure element, all the atoms are
identical in mass and other properties.
The atoms of different elements differ in mass and
other properties.
In a given compound the constituent atoms are always
present in the same fixed numerical ratio.
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Support for Dalton’s Atomic Theory: The Law of Multiple Proportions
Whenever two elements form more than one compound, the different
masses of one element that combine with the same mass of the other
element are in the ratio of small whole numbers.
Each molecule has one sulfur atom, and
therefore the same mass of sulfur. The
oxygen ratio is 3 to 2 by both mass and
atoms:
Sample experimental data:
Mass Mass
Compound
Size
S
O
Sulfur dioxide 2.00 g 1.00 g 1.00 g
Sulfur trioxide 2.50 g 1.00 g 1.50 g
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
It follows from Dalton’s Atomic Theory that
atoms of an element have a constant,
characteristic atomic mass or atomic weight
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For example, for any sample of hydrogen
fluoride:
• F-to-H atom ratio: 1 to 1
• F-to-H mass ratio: 19.0 to 1.00

This is only possible if each fluorine atom is 19.0
times heavier than each hydrogen atom
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 It
turns out that most elements in nature
are uniform mixtures of two or more
kinds of atoms with slightly different
masses
 Atoms of the same element with
different masses are called isotopes
• For example: there are 3 isotopes of hydrogen
and 4 isotopes of iron
 Chemically,
isotopes have virtually
identical chemical properties
 The relative proportions of the different
isotopes are essentially constant
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
A uniform mass scale for atoms requires a
standard
 For atomic mass units (amu, given the
symbol u) the standard is based on carbon:
• 1 atom of carbon-12 = 12 u (exactly)
• 1 u = 1/12 the mass of 1 atom of carbon-12 (exactly)
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This definition results in the assignment of
approximately 1 u for the mass of hydrogen
(the lightest atom)
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Example: Naturally occurring chlorine is a
mixture of two isotopes. In every sample of this
element, 75.77% of the atoms are chlorine-35
and 24.23% are chlorine-37. The measured
mass of chlorine-35 is 34.9689 u and that of
chlorine-37 is 36.9659 u. Calculate the average
atomic mass of chlorine.
Abundance Mass
Isotope
Chlorine-35
Chlorine-37
(%)
75.77
24.23
(u)
34.9689
36.9659
(Rounded) Total
Contribution
0.7577 * 34.9689 = 26.50 u
0.2423 * 36.9659 = 8.957 u
= 35.46 u
The average mass of 1 atom of chlorine in nature is 35.46 u.
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 Experiments
have been performed that
show atoms are comprised of
subatomic particles
 There
are three principal kinds of
subatomic particles:
• Proton – carries a positive charge, found in the
nucleus
• Electron – carries a negative charge, found
outside the nucleus, about 1/1800 the mass of a
proton
• Neutron – carries no charge, found in the
nucleus, a bit heavier than a proton, about 1800
times heavier than an electron
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 An
element can be defined as a
substance whose atoms all contain the
identical number of protons, called the
atomic number (Z)
 Isotopes are distinguished by mass
number (A):
• Atomic number, Z = number of protons
• Mass number, A = (number of protons) +
(number of neutrons)
 For
charge neutrality, the number of
electrons and protons must be equal
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 This
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information can be summarized
Example: For uranium-235
• Number of protons = 92 ( = number of
electrons)
• Number of neutrons = 143
• Atomic number (Z) = 92
• Mass number (A) = 92 + 143 = 235
• Chemical symbol = U
 Summary
for uranium-235:
Mass number, A (protons + neutrons)  235
Chemical Symbol 
U
Atomic number, Z (number of protons)  92
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The Periodic Table summarizes chemical and physical
properties of the elements
 The first Periodic Tables were arranged by increasing
atomic mass
Dmitri Mendeleev
(Rusia)
1869
Julius Lothar Meyer
(German)
Both these researchers drew out the first periodic table
independently of each other.
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Mendeleev have been honored as
the first person to arrange the
elements in the form of a table
because he reported his findings to
the Russian Chemical Society a
few months earlier than Meyer!!!!
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The Modern Periodic table is
arranged by increasing atomic
number:
Elements are arranged in
numbered rows called periods
The vertical columns are called
groups or families (group
labels vary)
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 Modern
Periodic Table with group labels
and chemical families identified
Lanthanides
Note: Placement of elements 58 – 71 and 90 – 103 saves space
Actinides
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Some important classifications:
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A groups = representative elements or main group
elements
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I A = alkali metals
II A = alkaline earth metals
VII A = halogens
VIII = noble gases
B groups = transition elements
Inner transition elements = elements 58 – 71 and
90 – 103
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58 – 71 = lanthanide elements
90 – 103 = actinide elements
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 Classification
as metals, nonmetals, and
metalloids
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 Metals
• Tend to shine (have metallic luster)
• Can be hammered or rolled into thin sheets
(malleable) and can be drawn into wire
(ductile)
• Are solids at room temperature and conduct
electricity
 Nonmetals
• Lack the properties of metals
• React with metals to form (ionic) compounds
 Metalloids
• Have properties between metals and
nonmetals
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The End
Do the exercises at the end of chapter ONE.
For your own practice only!!
You do not have to pass up this assignment.
Any questions?
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