0BCHM 101 GENERAL
CHEMISTRY
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FALL
QUARTER 2008 |
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Section
2 |
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Lecture
Notes – |
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(last revised: 10/6/08) |
2.5
The Modern View of Atomic Structure
- The components of
an atom are protons, neutrons, and electrons.
o Electrons
o The particles discovered by Thompson (Cathode-ray
Emissions).
o Negatively charged.
o Protons
o Same charge as electrons, only the opposite sign
o Approximately 1,800 times the mass of an electron
o Neutrons
o Zero charge (electrically neutral).
o Same mass as the proton.
- Structure of the
Atom:
o A tiny nucleus (diameter ~10-13 cm),
made of protons and neutrons and containing nearly all the mass of the atom
o Surrounded by a sphere (diameter ~ 10-8
cm) containing the electrons.
- How Elements
Differ
o The atoms of a given element (Example: gold) each
contain the same, characteristic number of protons (Example: An atom of gold
always contains 79 protons). This number is called the Atomic Number of the element. (The symbol for atomic number is Z.)
o Since atoms are electrically neutral, an atom has
the same number of electrons surrounding its nucleus as it has protons in its
nucleus.
o The number
and the arrangement of the electrons in the atoms of an element determine the chemistry of that element.
This is because when atoms come together, it is the electrons that “touch” or
interact with each other. To restate this point, the chemistry of a given
element depends critically on the
number of electrons in each of its atoms.
- Isotopes
o For any element, the particular number of protons
in the nucleus of each of its atoms is a characteristic of the element.
However, the number of neutrons can vary without changing the chemistry.
o An example is chlorine. Each chlorine atom contains
17 protons, but some chlorine atoms contain 18 neutrons while others contain
20. If we define the Mass Number of
a particular atom as its number of protons plus its number of neutrons, then we
can compute the mass numbers of these two kinds of chlorine as 35 and 37,
respectively. (The symbol for mass number is A.)
o When the atoms of a given element occur in forms
that have different numbers of neutrons, we call these forms isotopes of that element.
o The isotope of chlorine (atomic symbol: Cl) with
mass number 35 (called chlorine 35) can be symbolized:
o Likewise, chlorine 37 can be symbolized:
o In general, an atom with the chemical symbol, Xx,
the atomic number, Z, and the mass number, A, is written:
o Now we can work Sample Exercise 2.2 (page 52 in the
text). The problem asks us to write the symbol for an atom with an atomic
number of 9 and a mass number of 19. It also asks us to determine the numbers
of electrons and neutrons in the atom.
o Step 1: The data we need to look up are on the two
facing pages of the inside front cover of the text. Looking in the periodic
table, we see that the atomic number (9) is that of the element with the
symbol, F. (If we look up F on the opposite page, we can find that it belongs
to the element, fluorine.). Thus we can write the complete symbol for this atom
(including the atomic number, Z, and the mass number, A, as:
o Step 2: The number of electrons is the same as the
number of protons. This answer is given by the atomic number, 9.
o Step 3: The number of neutrons can be calculated by
subtracting the atomic number from the mass number:
2.6
Molecules and Ions
- Now that we have
studied the nuclear model of the atom and learned that electrons are
responsible for chemistry, we turn to the study of the ways that atoms
interact to form chemical compounds. At this stage we will look only at
some simple ideas of chemical bonding. We will study chemical bonding in
more depth during CHM 102.
- Chemical bonds are what we call the forces that bind atoms
together in compounds. At this point we will consider covalent bonds and
ionic bonds.
- Covalent Bonds: When two or more atoms form a molecule by
sharing electrons, we say that these atoms are covalently bonded. We can
represent these molecules symbolically in several ways:
o The simplest representation is the chemical formula. The atoms are represented
by their chemical symbols. (You will find these in the Table of Atomic Masses
in the inside front cover of your text.) For molecules containing 2 or more of
the same kind of atom, the number is written as a subscript after the chemical
symbol. Thus a water molecule, containing 2 hydrogen (H) atoms and one oxygen
(O) atom has the chemical formula
Another example of a
chemical formula is that of ammonia, a molecule containing 3 hydrogen atoms and
1 nitrogen (N) atom:
Still another example is
methane, a molecule containing 4 hydrogen atoms and 1 carbon (C) atom:
o A topological
formula is a two dimensional representation of a chemical structure. It
shows which atoms are bonded to which. The topological formula for water is:
The topological formula
for ammonia is:
H
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H—N—H
And the topological
formula for methane is:
H
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H—C—H
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H
o A structural
formula not only shows which atoms are bonded to which, but it also
suggests the true shape of the molecule and the actual angles between the
bonds. Here are structural formulas for water, ammonia, and methane:
In these diagrams, the
letters represent the atoms, the light, solid lines represent bonds that lie in
the plane of the screen (or paper, if this is a printout), the dotted lines
represent bonds that project behind the screen, and the dark wedges represent
bonds that poke out of the screen.
Note: The text does not
give a separate definition of topological formula but includes it as a kind of
structural formula without geometric information. (You will not be held
responsible for knowing the difference.)
o A molecular
model gives the truest representation of the structure and composition of a
given molecule. Two commonly used types of models are the space-filling model (Figure 2.17) and the ball-and-stick model (Figure 2.18).
- Ionic Bonds: Not all chemical compounds are covalently bonded. A familiar
example of a chemical compound entirely lacking in covalent bonds is
sodium chloride (NaCl), also known as common table salt. The bonding in
NaCl is ionic, not covalent. Let’s see what leads to ionic bonding in
NaCl.
o Ionization
of Sodium: A sodium atom, with an
atomic number of 11, has 11 electrons. One of the 11 electrons can be dislodged
fairly easily, leaving an entity with 10 electrons that are very tightly held
by a nucleus with a +11 charge. The charge on this entity is +1, and the entity
can be called a sodium positive ion, a sodium cation, or (most commonly) a
sodium ion. This ionization process can be written as a chemical equation:
o Ionization
of Chlorine: A chlorine atom, with
an atomic number of 17, has 17 electrons. These electrons are all tightly held,
and there is a driving force to add one more and produce an entity with 18
electrons in a stable structure about a nucleus with a charge of +17. Thus the
charge on the entity is -1, and the entity can be called a chlorine negative
ion, a chlorine anion, or (most commonly) a chlorine ion. This ionization
process can be written as a chemical equation:
o Ionic
Bonds: Since oppositely charged
ions exert attractive forces on each other, they form bonds without sharing
electrons. These are called ionic bonds. Figure 2.19 illustrates the reaction
of the element, sodium, a soft metal, with chlorine, a pale green gas, to form
the ionically bonded salt, sodium chloride, a colorless, crystalline solid.
o Ionic
Solids: A solid consisting of
oppositely charged ions is an ionic solid or a salt.
o Types of
Ions:
o Simple
Ion: Any ion made from a single
atom. 
and
are good examples.
o Polyatomic
Ion: An ion in which two or more
atoms are covalently bonded together into a molecule with a net charge. A good
example of a compound made from two types of polyatomic ions is ammonium
nitrate,
, which is made of ammonium ions, 
, and nitrate ions, 
. These two ions are illustrated in Figure 2.20:
- Metallic Bonding: (not discussed in Chapter 2, so you don’t
need to know this yet.): Metal atoms tend to cluster into large
aggregates, and the most loosely bound electrons are shared among all the
atoms in an aggregate. This gives rise to metallic bonding, as distinct
from covalent bonding or ionic bonding.
2.7
Introduction to the Periodic Table
- What it is: A Periodic Table is a chart that shows all the known elements
giving lots of information about each and organizing them in a particular
useful manner.
- Where to find it:
o In textbooks: There are at least 2 versions of the
periodic table in your text. (See the inside front cover and page 56.)
o Posted in class: There is one hanging on the
sidewall of D-108 near the door. There is also one posted in the lab (D-103).
o
On the web: If
you do a Google search, you will find many different versions of the periodic
table. Here is one example: HUhttp://www.webelements.com/UH.
- The Periodic Law:
The Periodic Law states that the chemical properties of the elements are a
periodic function of their atomic numbers.
- Arrangement of
the Periodic Table:
o Columns: The elements in a given column of the
periodic table have similar chemical properties. The collection of elements in
a column is called a group or a family.
o Rows: The elements in a given row of the periodic
table are ordered by atomic number, without any gaps in the number sequence.
These rows are also called periods.
o Chemistry of Selected Groups:
o Metals: Most elements are classed as metals. They
are located on the left hand side and in the center of the periodic table, and
as you go down the table to heavier and heavier elements, they occur more and
more toward the right hand side. For example, in Period 6 (elements with atomic
number 55 through 86), fully 30 of the 32 elements are classed as metals. As
elements, metals tend to occur as metallically bound solids that are good
conductors of heat and electricity, they tend to be malleable and ductile so
that they can be hammered into thin sheets or drawn into wires, they often
appear shiny or lustrous, and chemically they tend to react by losing electrons
to form positive ions.
o Group 1 (Alkali Metals, such as sodium and
potassium): These are all very active metals that readily form ions with +1 charges when they react with non-metals.
o Group 2 (Alkaline Earths, such as magnesium and
calcium): These are metals that form ions with +2 charges.
o Groups 3-12 (Transition Metals): There are 10
groups of transition metals, but group to group differences among transition
metals are not as pronounced as those among the main groups (Groups 1 and 2 and
13 through 18). They tend to have complex chemistries. Many of them have two or
more ionic forms, for example: iron (Fe) can lose 2 electrons to form the Fe2+
ion, but it can also lose a third electron to form Fe3+. Many of
them form 2 or more different kinds of oxide, for example: titanium (Ti) forms TiO and TiO2. And many of them are capable of
forming polyatomic anions, for example: chromium (Cr) can form the chromate (an)ion, CrO42-.
o Group 18 (Noble Gases, such as helium and argon):
These elements all occur in nature as monatomic gases. They were thought never
to form chemical bonds until the mid-1960’s when several compounds of xenon and
fluorine were successfully prepared.
o Group 17 (Halogens, the elements fluorine,
chlorine, bromine, and iodine): These non-metals can all form ions with -1 charges when they react with an active metal. They can also
form covalent bonds with themselves and with other non-metals. As elements,
they occur as diatomic molecules, such as Cl2.
o Other Non-Metals: The non-metals other than the
halogens are nearly all found in the upper right corner of the periodic table.
The most important of these are carbon, nitrogen, oxygen, phosphorous, and
sulfur. Like the halogens, they tend to accept electrons and form negative ions
when they react with metals, and they tend to form covalent bonds when they
react with each other.
o What about Hydrogen? The placement of hydrogen in
the periodic table has always been a problem. In some respects its chemistry
resembles the chemistry of the alkali metals, but it also behaves as a
non-metal. As an element, it occurs as the diatomic gas, H2, and it
forms covalent bonds with carbon and other non metals. There are valid
arguments for placing it in Group 17 above fluorine.
- Using the
Periodic Table:
o As a Predictive Tool: Since its earliest days, the
periodic table has been used successfully to predict the properties of unknown,
undiscovered elements. When Mendeleev published the very first version of the
periodic table, he had to leave several gaps in it where there were no known elements
to fit particular combinations of period and group. For example, he could not
fit any known element into Period 4 of Group 14. His answer was to predict that
such an element could be found and that its chemical properties would resemble
those of silicon (Period 3, Group 14) and tin (Period 5, Group 14) and that its
physical properties (atomic mass, density, etc.) would be similar to its
neighboring elements in Period 4. Given this road map of a prediction, other
chemists soon discovered the element, germanium, and found that its actual
properties were a close match for Mendeleev’s predictions.
o As a Data Organizer: The periodic table summarizes
and correlates a wide variety of information in a compact and easy to use
format.
o As a Look-Up Tool: The periodic table charts every
element by atomic number, atomic mass, name, chemical symbol, group, and
period. Thus if you know some of this information about a given element, you
can look up the rest in the periodic table.
o As a Learning Tool: The next section we will cover
in General Chemistry is about Naming Simple Compounds. This is likely to be the
most heavy duty memorization you will encounter all quarter. We will expect you
to know names of elements, their chemical symbols, and the formulas and names of
typical compounds they form when they react with other elements. The good news
in all this is that of the 110 or so known elements, you will be held
responsible for knowing only the 25 or so most important ones. And you can use
the periodic table as a graphical device to help you organize and learn this
information.