GENERAL CHEMISTRY SUBCOURSE MD0803 EDITION 100

GENERAL CHEMISTRY  SUBCOURSE MD0803 EDITION 100
U.S. ARMY MEDICAL DEPARTMENT CENTER AND SCHOOL
FORT SAM HOUSTON, TEXAS 78234-6100
GENERAL CHEMISTRY
SUBCOURSE MD0803
EDITION 100
DEVELOPMENT
This subcourse is approved for resident and correspondence course instruction. It
reflects the current thought of the Academy of Health Sciences and conforms to printed
Department of the Army doctrine as closely as currently possible. Development and
progress render such doctrine continuously subject to change.
ADMINISTRATION
For comments or questions regarding enrollment, student records, or shipments,
contact the Nonresident Instruction Branch at DSN 471-5877, commercial (210) 2215877, toll-free 1-800-344-2380; fax: 210-221-4012 or DSN 471-4012, e-mail
[email protected], or write to:
COMMANDER
AMEDDC&S
ATTN MCCS HSN
2105 11TH STREET SUITE 4192
FORT SAM HOUSTON TX 78234-5064
Approved students whose enrollments remain in good standing may apply to the
Nonresident Instruction Branch for subsequent courses by telephone, letter, or e-mail.
Be sure your social security number is on all correspondence sent to the Academy of
Health Sciences.
CLARIFICATION OF TRAINING LITERATURE TERMINOLOGY
When used in this publication, words such as "he," "him," "his," and "men" are intended
to include both the masculine and feminine genders, unless specifically stated otherwise
or when obvious in context.
.
TABLE OF CONTENTS
Lesson
Paragraph
INTRODUCTION
1
ELEMENTS OF CHEMICAL STRUCTURE AND INORGANIC
NOMENCLATURE
Section l.
Section II.
Exercises
2
Elements of Chemical Structure
Rules Of Inorganic Nomenclature
ELEMENTS OF CHEMICAL CHANGE
1-1—1-7
1-8—1-18
2-1 --2-13
Exercises
3
ELEMENTS OF ORGANIC CHEMISTRY
Exercises
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3-1 --3-18
SUBCOURSE MD0803
GENERAL CHEMISTRY
INTRODUCTION
In the process of achieving and maintaining proficiency in your military
occupational specialty (MOS), you will be learning concepts and performing tasks that
are based on important chemical principles. As you become more proficient with these
principles, you may reach the point where you will not need to give them much
conscious thought. Meanwhile, however, you should study this subcourse to gain a
working knowledge of the fundamental principles of chemistry.
Subcourse Components:
This subcourse consists of 3 lessons. The lessons are:
Lesson 1, Elements of Chemical Structure and Inorganic Nomenclature.
Lesson 2, Elements of Chemical Change.
Lesson 3, Elements of Organic Chemistry.
Credit Awarded:
To receive credit hours, you must be officially enrolled and complete an
examination furnished by the Nonresident Instruction Branch at Fort Sam Houston,
Texas. Upon successful completion of the examination for this subcourse, you will be
awarded 14 credit hours.
You can enroll by going to the web site http://atrrs.army.mil and enrolling under
"Self Development" (School Code 555).
A listing of correspondence courses and subcourses available through the
Nonresident Instruction Section is found in Chapter 4 of DA Pamphlet 350-59, Army
Correspondence Course Program Catalog. The DA PAM is available at the following
website: http://www.usapa.army.mil/pdffiles/p350-59.pdf.
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LESSON ASSIGNMENT
LESSON 1
Elements of Chemical Structure and Inorganic
Nomenclature.
LESSON ASSIGNMENT
Paragraphs 1-1 through 1-18 and exercises.
LESSON OBJECTIVES
SUGGESTION
MD0803
After completing this lesson, you should be able to:
1-1.
Define: atom, molecule, element, compound
matter, energy, atomic number, atomic weight,
electron configuration, isotope, valence
octet rule, ion, cation, anion, radical.
1-2.
List the three states of matter and the
characteristics of each.
1-3.
List the three basic particles in an atom and the
charge and mass of each.
1-4.
State the maximum number of electrons a given
electron shell may contain.
1-5.
Given a block for an element from the periodic
table, write the name of each piece of
information which may be obtained about the
element.
1-6.
Given the name of an element or radical
commonly encountered in medicine, state the
symbol or formula and common valence(s) for
that element or radical.
1-7.
List the three types of chemical bonds and state
whether the electrons are shared or transferred.
1-8.
Given the name of an inorganic compound
commonly encountered in medicine, write the
chemical formula for the compound.
1-9.
Given a chemical formula of an inorganic
compound commonly encountered in pharmacy,
state the name for that compound.
After completing the assignment, complete the
exercises at the end of this lesson. These exercises
will help you to achieve the lesson objectives.
1-1
LESSON 1
ELEMENTS OF CHEMICAL STRUCTURE AND INORGANIC NOMENCLATURE
Section I. ELEMENTS OF CHEMICAL STRUCTURE
1-1.
INTRODUCTION
Chemistry is the science that studies the composition and changes in
composition of the substances around us. Man's natural curiosity about the things and
transformations that he observed was the original impetus for the development of this
science, but its true beginning was in the work of the alchemists of the Middle Ages.
These men searched for a way to change the base metals such as lead into gold. In
the large span of time since then, chemistry has developed into a true science and we
have amassed a tremendous volume of knowledge. To facilitate the study of chemistry,
we can divide it into two divisions: Inorganic chemistry, which deals with the elements
and mineral materials, and organic chemistry, which deals with compounds containing
carbon. More divisions of chemistry exist, but we will be primarily concerned with these
two.
1-2.
IMPORTANCE OF CHEMISTRY
Why do we study chemistry? The answer to this question will be obvious when
you consider the various classes of compounds we encounter in medicine and in our
daily lives. For example, we are concerned with compounds such as drugs and the
changes they undergo. Here are some things chemistry will tell us about drugs.
a. Actions. Chemistry may tell us about the actions of drugs on the body. Drug
effects are determined by the chemical structure of a drug; changes in structure may
alter the actions of the drug.
b. Safety and Storage Procedures. Special safety or storage precautions may
be necessary for particular drugs. These can be identified by the chemical structure.
c. Incompatibilities. Sometimes, two or more drugs cannot be mixed because
of undesirable consequences. There are three types of incompatibilities:
(1) Chemical. Alterations of chemical properties may occur when two or
more drugs are mixed.
(2) Physical. Physical properties of ingredients may produce a mixture
unacceptable in appearance or accuracy of dosage.
(3) Therapeutic. When two or more drugs are given to a patient, they may
interact in some way to change the effects of one of the drugs.
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1-3.
MATTER
Matter is anything that occupies space and has weight. If you look around you,
you will see matter. The table, books, walls, and your body are all composed of matter.
Obviously, the matter around you is not all the same.
a. Physical States of Matter. In general, we can group all matter into three
groups called states of matter.
(1) Solids. Solids have a definite shape and volume. Examples of solids
are books, rocks, pieces of steel, and sand.
(2) Liquids. Liquids have a definite volume but indefinite shape. That is,
they take the shape of their container. Water, mercury, alcohol, and oils are liquids.
(3) Gases. Gases have neither a definite shape nor a definite volume.
They assume not only the shape of their container, but also the volume of their
container. Gases may be expanded or compressed to fit the container in which they are
being placed. Therefore, the air in an automobile tire would, if released, expand to fill a
large weather balloon.
b. Properties of Matter. Matter possesses two types of properties, physical
and chemical. Characteristics such as smell, color, shape, freezing point, boiling point,
and solubility are said to be physical properties of matter. Energy content, reactions
with other substances, and chemical reactions due to light, heat, and electricity are said
to be chemical properties of matter. From the physical and chemical properties
exhibited by a substance, it is possible to isolate, identify, and classify the particular
substance.
c. Classification of Pure Matter. Matter that cannot be separated into two or
more types of matter by physical means is called pure matter. Pure matter consists of
two types, elements and compounds.
(1) Elements. Elements are substances that cannot be separated into two
or more types of matter by physical or chemical methods. Another way to say this is
that elements consist of only one type of atom. An atom is a chemical building block
and can be defined as the smallest part of an element that remains unchanged during
any chemical reaction and exhibits or displays the chemical properties of that element.
Examples of common elements are oxygen, gold, iron, mercury, hydrogen, and carbon.
Table 1-1 lists the elements with their symbols, atomic numbers, and atomic weights.
(2) Compounds. Compounds are composed of two or more elements
chemically combined. Compounds are substances that have been purified by physical
means, but not by chemical methods. They can be separated into two or more types of
matter by chemical methods because their basic unit, the molecule, is a combination of
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two or more types of atoms. A molecule is composed of two or more atoms and is the
smallest part of a compound that can exist and still retain the chemical properties of that
compound. Illustrated in Table 1-1 are the relationships of these building blocks and
classifications of matter.
ELEMENT
Atoms
*
*
*
*
*
*
*
ELEMENT
COMPOUND
Atoms
Molecules
ELEMENT
SYMBOL
ATOMIC
NUMBER
Actinium
Aluminum
Americium
Antimony
Argon
Arsenic
Astatine
Barium
Berkelium
Beryllium
Bismuth
Boron
Bromine
Cadmium
Calcium
Californium
Ac
Al
Am
Sb
Ar
As
At
Ba
Bk
Be
Bi
B
Br
Cd
Ca
Cf
89
13
95
51
18
33
85
56
97
4
83
5
35`
48
20
98
ATOMIC
WEIGHT
227
26.9815
243
121.75
39.948
74.9216
210
137.34
247
9.0122
208.980
10.811
79.909
112.40
40.08
249
* Denotes elements most common to medicine.
Table 1-1. Elements, symbols, atomic numbers, and atomic
Weights in alphabetical order (continued).
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ELEMENT
SYMBOL
ATOMIC
NUMBER
* Carbon
Cerium
Cesium
* Chlorine
Chromium
* Cobalt
* Copper
Curium
Dysprosium
Einsteinium
Erbium
Europium
Fermium
* Fluorine
Francium
Gadolinium
Gallium
Germanium
* Gold
Hafnium
Helium
Holmium
* Hydrogen
Indium
* Iodine
Iridium
* Iron
Krypton
Kurchatovium
Lanthanum
Lawrencium
* Lead
* Lithium
Lutetium
C
Ce
Cs
Cl
Cr
Co
Cu
Cm
Dy
Es
Er
Eu
Fm
F
Fr
Gd
Ga
Ge
Au
Hf
He
Ho
H
In
I
Ir
Fe
Kr
Ku
La
Lw
Pb
Li
Lu
6
58
55
17
24
27
29
96
66
99
68
63
100
9
87
64
31
32
79
72
2
67
1
49
53
77
26
36
104
57
103
82
3
71
ATOMIC
WEIGHT
12.01115
140.12
132.905
35.453
51.996
58.9332
63.54
247
162.50
254
167.26
151.96
253
18.9984
223
157.25
69.72
72.59
196.967
178.49
4.006
164.930
1.00797
114.82
126.9044
192.2
55.847
83.80
257
138.91
257
207.19
6.939
174.97
* Denotes elements most common to medicine.
Table 1-1. Elements, symbols, atomic numbers, and atomic
Weights in alphabetical order (continued).
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ELEMENT
ATOMIC
NUMBER
SYMBOL
* Magnesium
* Manganese
Mendelevium
* Mercury
Molybdenum
Neodymium
Neon
Neptunium
Nickel
Niobium
* Nitrogen
Nobelium
Osmium
* Oxygen
Palladium
* Phosphorus
Platinum
Plutonium
Polonium
* Potassium
Praseodymium
Promethium
Protactinium
* Radium
Radon
Rhenium
Rhodium
Rubidium
Ruthenium
Samarium
Scandium
* Selenium
* Silicon
* Silver
Mg
Mn
Md, Mv
Hg
Mo
Nd
Ne
Np
Ni
Nb, Cb
N
No
Os
O
Pd
P
Pt
Pu
Po
K
Pr
Pm
Pa
Ra
Rn
Re
Rh
Rb
Ru
Sm
Sc
Se
Si
Ag
12
25
101
80
42
60
10
93
28
41
7
102
76
8
46
15
78
94
84
19
59
61
91
88
86
75
45
37
44
62
21
34
14
47
ATOMIC
WEIGHT
24.312
54.9380
256
200.59
95.94
144.24
20.183
237
58.71
92.906
14.0067
254
190.2
15.9994
106.4
30.9738
195.09
242
210
39.102
140.907
147
231
226
222
186.2
102.905
85.47
101.07
150.35
44.956
78.96
28.086
107.870
* Denotes elements most common to medicine.
Table 1-1. Elements, symbols, atomic numbers, and atomic
Weights in alphabetical order (continued).
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1-6
ELEMENT
ATOMIC
NUMBER
SYMBOL
* Sodium
* Strontium
* Sulfur
Tantalum
Technetium
Tellurium
Terbium
Thallium
Thorium
Thulium
Tin
Titanium
Tungsten
Uranium
Vanadium
Xenon
Ytterbium
Yttrium
* Zinc
Zirconium
Na
Sr
S
Ta
Tc
Te
Tb
Tl
Th
Tm
Sn
Ti
W
U
V
Xe
Yb
Y
Zn
Zr
11
38
16
73
43
52
65
81
90
69
50
22
74
92
23
54
70
39
30
40
ATOMIC
WEIGHT
22.9898
87.62
32.064
180.948
99
127.60
158.924
204.37
232.038
168.934
118.69
47.90
183.85
238.03
50.942
131.30
173.04
88.905
65.37
91.22
* Denotes elements most common to medicine.
Table 1-1. Elements, symbols, atomic numbers, and atomic
Weights in alphabetical order (concluded).
d. Classification of Mixed Matter. Matter that can be separated by physical
means is called mixed matter and may be homogeneous or heterogeneous.
(1) Homogeneous mixtures. Mixtures that are uniform throughout are called
homogeneous. An example of a homogeneous mixture is a solution of sugar in water.
Any small part of this solution would exhibit the same properties as any other small part;
therefore, it would be uniform throughout the mixture.
(2) Heterogeneous mixtures. Mixtures that are not uniform are called
heterogeneous. An example of a heterogeneous mixture is a mixture of water and oil.
If a small sample is taken, it may not be the same as another small sample taken from
elsewhere in the mixture. This is because oil and water do not mix well--they give a
nonuniform mixture.
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1-4.
ENERGY
There are many things in our surroundings that we know exist, yet are not matter.
They are forms of energy. Heat, light, and electricity are examples of energy. Energy
may be simply defined as the ability to do work or overcome resistance.
1-5.
ATOMIC STRUCTURE
Early scientists felt that all matter must be built from some basic unit, just as a
wall may be constructed from a basic unit, the brick. In trying to find this basic unit, they
separated matter by all the methods (chemical and physical) available to them until they
could not separate it any further. They felt this separation must result in the building
block of matter, which they called the atom (from the Greek word for indivisible). They
also observed that the basic units or atoms for various elements differed in their
properties, as iron was certainly different from carbon. This led them to try to find the
structure of the atom. The difficulty of this problem can be seen when you consider that
one cubic centimeter of gold contains as many as 59,000,000,000,000,000,000,000
atoms. The atom is so small that it defies conception. Through ingenious methods,
particularly in the last 100 years, we have discovered many facts about this tiny particle,
which enables us to understand many of the changes that occur around us.
a. Atomic Model. In order for us to picture what an atom looks like, we can use
a description with which most people are familiar--the solar system model. In this
model, the atom is thought of as a tiny solar system in which there is a central core (like
the sun) with other particles traveling in circular paths or orbits (like the planets). While
more complex and exact models have been developed, this is the best approximation
for general use.
b. The Nucleus. The central core from the solar system model is called the
nucleus (which is derived from the Latin word nucis meaning nut or kernel). The
nucleus contains two types of particles, the proton and the neutron.
(1) The proton. The proton is a particle that has a mass (or weight) of one
amu (atomic mass unit) and a positive one (+1) electrical charge. The symbol for the
proton is p, p+ or H+.
(2) The neutron. The neutron has a mass of one amu (atomic mass unit)
but has no electrical charge; that is, it is a neutral particle. In an atom that has more
than one proton, the positive charges tend to repel each other. The neutrons serve to
bind the protons so that this electrical repulsion does not cause them to fly off into
space. The symbol for the neutron is n.
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(3) Atomic number and atomic weight. Two important figures commonly
used when discussing an atom are its atomic number and its atomic weight.
(a) Atomic number. The atomic number of an atom is equal to the
number of protons in the nucleus of the atom. For example, a carbon atom has six
protons in its nucleus; therefore, the atomic number of carbon is six.
(b) Atomic weight. The atomic weight of an atom is equal to the
number of protons in the nucleus of the atom (one amu each) plus the number of
neutrons in the nucleus of the atom (one amu each). Therefore, a carbon atom with six
protons and six neutrons has an atomic weight of 12.
c. The Outer Structure. The particles that orbit the nucleus (as the planets
orbit the sun) are called electrons. These particles have an electrical charge of negative
one (-1), but their mass is so small that it is considered to be zero. Actually, the mass of
the electron is 1/1837 of the mass of a proton, but the mass, which contributes to the
atom is so small that it is not important. The symbol for the electrons is e or .
(1) Electron configuration. Since we may have many electrons going
around the nucleus, It might appear that there could be a collision of electrons.
Collisions do not occur because the electrons are located in orbits, which are different
distances from the nucleus and because of the repulsion between like charges. The
number of electrons and their locations are called the electron configuration. This
electron configuration is different for each element.
(2) Electron shell. The term electron shell (or energy level) describes where
electrons are located (i.e., a specific region around the nucleus). Since electrons can
be forced to leave their atoms, the term energy level indicated the amount of energy
required to remove the electrons from the various levels or shells. A nucleus can have
seven shells, but more chemicals of medicinal importance contain electrons in the first
four, which are labeled the K, L, N, and N shells. The K shell is the closest to the
nucleus and the N shell is the farthest from the nucleus (figure 1-1). These shells
contain different numbers of electrons. The maximum number each shell can hold is
equal to 2N2, where N is the number of the shell (K=1, L=2, M=3, and so forth.). Thus
the maximum number of electrons that each of the first four shells can hold Is:
K=
L=
N=
N=
2(1)2
2(2)2
2(3)2
2(4)2
= 2
= 8
= 18
= 32
Since, for example, the M shell can contain as many as 18 electrons, the possibility for
collision might still appear to exist. The reason collisions do not occur is that a shell is
subdivided into smaller energy levels, called subshells and orbitals, which we will not
need to consider.
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(3) Number of electrons. What determines the number of electrons an atom
will contain? For an atom to exist freely in nature, it must be electrically neutral (without
a charge). There are two particles in an atom that have charges--the proton, which is
positive, and the electron, which is negative. For electrical neutrality, the sum of the
charges must be zero. In other words, the number of electrons (negative charges) must
equal the number of protons (positive charges).
Figure 1-1. First four electron shells.
d. Atomic Structure of Elements. As previously stated, each element consists
of a single type of atom. Since all atoms consist of the three basic particles we have
just discussed (except hydrogen, which usually has no neutrons), the only ways in
which elements can differ are atomic number (the number of protons) and atomic
weight, (the number of protons and neutrons). There are over 106 different elements
which scientists know to have a different atomic number and atomic weight. These
elements have a large assortment of properties. Two elements are liquids at room
temperature, eleven are gases, and all others are solids.
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e. Periodic Law. While investigating the properties of the elements, scientists
discovered an interesting fact that is now called the periodic law. This law states that
the properties of the elements are periodic functions of the atomic number. As the
atomic number increases, the properties of the elements repeat themselves at regular
Intervals.
f. Periodic Table. The periodic law allowed the scientists to group together the
elements that had similar properties and form a systematic table of the elements. This
table is the periodic table (Table 1-2). The vertical columns are called groups, and the
horizontal rows are called periods. This table contains a lot of information that we will
not generally use; however, we are concerned with the basic information we can obtain
about the elements. Figure 1-2 includes four blocks for elements from the periodic table
showing the information, which can be obtained from it. You should note that the
number of neutrons is not given in the periodic table. This can be determined by
subtracting the atomic number from the atomic weight.
Figure 1-2. Identifying the components of the periodic table.
g. Isotopes. All the atoms of a particular element are not identical. Slight
variations in the number of neutrons are found to occur naturally. Variations can also
be produced in reactors. Atoms that have the same number of protons but a different
number of neutrons (same atomic number, but different atomic weights) are called
2
3
239
isotopes. Sometimes isotopes are referred to by their mass numbers, H , H , U , and
so forth. All of the isotopes of a particular element have identical electronic
configurations; and since electronic configurations determine chemical properties,
isotopes of an element
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Table 1-2. Periodic table of the elements.
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exhibit identical chemical behavior. Induced nuclear reactions can produce both stable
and radioactive nuclei. If the nucleus of the atom is unbalanced during the
bombardment reaction the atom is called a radioisotope. Radioisotopes, such as
cobalt66 for treatment of cancer and iodine131 for diagnosing of thyroid tumors, are of
vital importance in the medical field. The presence of isotopes helps to explain why
many atomic weights in the periodic table are not whole numbers since all of the
isotopes must be considered when computing the average atomic weight of the
element.
1-6.
VALENCE AND CHEMICAL BONDING
We have now developed the concept that matter was built from a basic unit
called the atom, and we have discussed the nature of the atom. We know, however,
that very little matter exists as free elements. Most of the things around us are
combinations of elements. Logically, the next step is to consider how things combine.
a. Valence. The valence of an element can be defined as a measure of its
combining power or the number of electrons an atom must gain, lose, or share to have
a full or stable outer electron shell. The reason atoms combine is contained in this
definition. There are certain electron configurations in nature that are unusually stable
(unreactive). The elements that have these configurations are in Group VIII A of the
periodic table. They are sometimes referred to as the inert or noble gasses because
they are found in very few combinations in nature. Other elements, by gaining, losing,
or sharing electrons, can try to make their outer electron shells resemble the shells of
the noble gases and hence become very stable. We can see how this works by
considering the two simplest elements, hydrogen, and helium. Hydrogen has one
electron in the K shell since it has only one proton. Therefore, hydrogen is a very
reactive element, occurring naturally in many compounds. Helium, a noble gas, has two
electrons in the K shell since it has two protons. Helium is very unreactive. Note that
helium, by having two electrons, has a completed outer shell, since the K shell can hold
only two electrons. Hydrogen would like to be as stable as helium and could be if it
could gain or share one more electron to give it a completed outer shell. Hydrogen
seeks this electron in nature by combining with other elements.
b. Octet Rule. If you examine the noble (inert) gases (like helium), you will see
that not all have a completed (full) electron shell. Except for helium, the noble gases
have eight electrons in their outer shell, yet they are still very stable. Chemists have
observed that other elements sometimes gain, lose, or share electrons in order to have
eight electrons in their outer shell. This observation led to the development of the octet
rule, which states that outer electron shells prefer to have eight electrons even though
the shell may not be full. (Octet means a group of eight.) On the next page are some
examples of the electron configurations for various elements which indicate to us how
many electrons they can gain, lose, or share to fit the octet rule or have a completed
outer shell.
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c. Positive Valence. An atom that must give up electrons to become stable will
have more protons than electrons in its stable configuration and will not be electrically
neutral. It will be positively charged since there are more positive charges than
negative charges. This is indicated by a + sign. The number of electrons it gives up is
the numerical value of its valence. Consider, for example, the element sodium, which
has 11 protons and 11 electrons in its free state. It has one electron in the M shell,
which it loses easily to become stable. After it loses the electron (that is, gives up a
negative charge), it will have a positive one charge and its valence will be +1.
d. Negative Valence. An atom that must gain electrons to become stable will
have more electrons than protons in its stable configuration and will not be electrically
neutral. It will be negatively charged since there are more negative than positive
charges. This is indicated by a "-" sign. The number of electrons it gains is the
numerical value of its valence. Consider, for example, the element chlorine, which has
17 protons and 17 electrons in its free state. It is one electron short of fitting the octet
rule in the M shell as that shell contains 7 electrons. After it gains the electron, it will
have a negative one charge and its valence will be -1.
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e. Important Symbols and Valences. Since it is very tedious to continually
write complete names for elements, chemists developed the symbols for the elements
which you observed on the periodic table. It will not be necessary for you to know all
the symbols for your work but a number of them appear frequently enough that they
should be memorized. Table 1-3 lists important elements with their symbols and
valences. These should be committed to memory. (Please note that most, but not all,
valences conform either to the completed shell or octet rules.)
f. Ions. Any atom that gains or loses electrons becomes charged (electrical
charge) and is called an ion. An ion can be defined as any charged atom or group of
atoms. If the ion is positively charged, it is called a cation. If it is negatively charged, it
is called an anion. A group of atoms that has a charge and goes through a reaction
unchanged is called a radical. Whenever we write the symbol for an element and wish
to indicate it is an ion, we write the charge as a superscript to the symbol, for example,
+
Cl 1 or Na 1.
g. Chemical Bonding. When elements combine to form chemical compounds,
the electrons in the outer shell may be transferred from one atom to another or there
may be a mutual sharing of the electrons. In either case, a chemical bond is produced.
This means the two atoms do not travel or react independently of one another but are
held together by the exchange or sharing of the electrons. Both atoms involved in the
reaction attain a completed outer orbit, and stability results. There are three types of
chemical bonds--electrovalent, covalent, and coordinate covalent.
(1) Electrovalent (ionic) bonding. A transfer of one electron from one atom
to another resulting in opposite charges on the two atoms that holds them together by
electrostatic (opposite charges attract) attraction is called an electrovalent or ionic bond.
A good example of this is the bond formed between a Na (sodium) and a Cl (chlorine)
atom.
-
1 e in M shell
MD0803
-
7 e in M Shell
1-15
Sodium has a complete outer shell
and a charge of +1. Chlorine has
met the octet rule in the M shell and
has a charge of –1.
(2) Covalent bond. If two atoms each donate an electron that is shared with
the other atom, the bond is a covalent bond. An example of this is the bond between
two H (hydrogen) atoms. Double and triple covalent bonds are also possible.
Both atoms have 1 e- in the K shell.
By sharing their electrons each
hydrogen has 2 e in the K shell and
both are stable because of the
completed shell.
(3) Coordinate covalent bond. If one atom donates two electrons for sharing
with another atom (which donates no electrons), it is called a coordinate covalent bond.
An example of this type of bond between N (nitrogen) in ammonia and a hydrogen ion
(proton).
Ammonia -N
has a complete
outer shell
MD0803
+
-
H has an
empty K
shell
By sharing the two e ,
H+ has a completed K shell
and N still has a completed
outer shell.
1-16
NAME
SYMBOL
VALENCE
Acetate
Aluminum
Ammonium
Antimony
Arsenic
Barium
Bicarbonate
Bismuth
Bromine
Calcium
Carbon
Carbonate
Chlorine
Copper
Fluorine
Gold
Hydrogen
Hydroxide (Hydroxyl)
Iodine
Iron
Lead
Lithium
Magnesium
Manganese
Mercury
Nickel
Nitrate
Nitrogen
Oxygen
Permanganate
Phosphate
Phosphorus
Potassium
Silver
Sodium
Strontium
Sulfate
Sulfur
Zinc
C2H3O2
Al
NH4
Sb
As
Ba
HCO3
Bi
Br
Ca
C
CO3
Cl
Cu
F
Au
H
OH
I
Fe
Pb
Li
Mg
Mn
Hg
Ni
NO3
N
0
MnO4
PO4
P
K
Ag
Na
Sr
SO4
S
Zn
-1
+3
+1
-3,
-3,
+2
-1
+3,
-1,
+2
+2,
-2
-1,
+1,
-1
+1,
+1
-1
-1,
+2,
+2,
+1
+2
+2,
+1,
+2,
-1
+1,
-2
-1
-3
-3,
+1
+1
+1
+2
-2
-2,
+2
+3, +5
+3, +5
+5
+1, +3, +5, +7
+4, -4
+1, +3, +5, +7
+2
+3
+1, +3, +5, +7
+3
+4
+3, +4, +6, +7
+2
+3
-3, +3, +5
+3, +5
+2, +4, +6
NOTE: The most common valences are underlined where there may be more than one
valence.
Table 1-3. Valences.
MD0803
1-17
1-7.
FORMULA WRITING
a. Formulas. Formulas are combinations of symbols that represent a
compound. A formula indicates which elements are involved and the number of atoms
of each element contained in the compound. In writing formulas, we use subscripts,
coefficients, and parentheses in addition to the symbols of the elements. Subscripts
indicate the number of atoms of an element, as in H2 where two is the subscript
meaning two hydrogen atoms. If there is no subscript with a symbol, it is assumed
there is only one atom of that element. Coefficients, numbers in front of the formula,
indicate the number of molecules of compound, as in 4HCl where four is the coefficient
indicating four molecules of HCl. Parentheses are used to separate a radical from the
rest of the formula when it would be confusing not to do so. In HNO3, it is not necessary
to include parentheses for the NO3 - radical since there is little chance for confusion.
However, we use parentheses for the same radical if it appears NO3 in a compound
such as Hg(NO3)2 where the 2 indicates that we have two NO3 - radicals.
b. Steps in Formula Writing. In writing formulas for compounds, there are four
steps that should be followed.
(1)
Determine the symbols for the elements in a compound.
(2)
Determine the valence of each of the atoms or radicals.
(3)
Write the positive element's symbol first, followed by that of the negative
(4)
Make the compound electrically neutral by using subscripts.
element.
c. Example. Write the formula for calcium chloride.
(1)
Calcium = Ca, Chloride = Cl.
(2)
Ca valence is +2, Cl valence is -1.
-1
(3) Ca+2Cl . If we add the charges, we find that this compound is not
neutral (+2 - 1 = +1). Therefore, we must proceed to step (4).
(4) To have two negative charges to balance the two positive charges, we
-1
must have two Cl ions (-1 x 2 = -2). Thus, the formula would be CaCl2.
MD0803
1-18
d. Rule of Crossing Valences. A convenient rule for determining what
subscripts are necessary in writing formulas is the rule of crossing valences. This rule
states that one can take the valence of the element at the left and make it the subscript
of the element at the right, and in like manner take the valence of the element at the
right and make it the subscript of the element at the left. For example:
Fe+3
SO4 –2
becomes Fe2(SO4)3
Section II. RULES OF INORGANIC NOMENCLATURE
1-8.
INTRODUCTION
a. This section discusses how to name a compound from its formula. The
interrelationship of names and formulas is very important to you. You will be required to
recognize both, in interpreting, preparing, and using these chemicals.
b. This section is in the format of programmed instruction. Each frame
presents some material, and then asks some questions in which you apply the material
presented. The correct answers follow so that you can check your answers for
accuracy. It is important that you use a piece of paper to cover the answers as you
work the program. You should fill in the answers as you work each frame and then
check your answers. If you answered any questions incorrectly, go back and review the
frame so that you understand the correct answer.
1-9.
GENERAL TERMS
There are several general terms we use that give us information about inorganic
compounds. To describe the number of different elements in a compound we use the
terms binary, ternary, and quaternary. A binary compound contains two different
elements, such as NaCl. A ternary compound contains three different elements, such
as H2SO4. A quaternary compound contains four different elements such as NaHCO3.
a. Questions.
(1) CO2 is a ______________________ compound because it contains
________ different elements.
(2) Al(OH)2Cl is a ______________________ compound because it
contains ________ different elements.
(3) KNO3 is a ______________________ compound because it contains
________ different elements.
MD0803
1-19
b. Answers.
(1)
Binary, two (C,O).
(2)
Quaternary, four (Al,O,H,Cl).
(3)
Ternary, three (K,N,O).
1-10. NUMBER PREFIXES
We often use prefixes to denote the number of atoms of an element in a
compound. For example, CO contains one oxygen atom and is named carbon
monoxide. Mon or mono indicates one atom. Here is a list of the commonly used
number prefixes.
Examples
Mono, mon
Di
Tri
Tetra
Penta
Hexa
Hepta
Octa
Nona
Deca
= one
= two
= three
= four
= five
=six
= seven
= eight
= nine
= ten
CO
CO2
SO3
Carbon monoxide
Carbon dioxide
Sulfur trioxide
a. Questions.
(1)
NCl3 is named nitrogen
(2)
SO2 is named sulfur ______________________ oxide.
(3)
CF4 is names carbon ______________________ flouride.
chloride.
b. Answers.
MD0803
(1)
Tri.
(2)
Di.
(3)
Tetra.
1-20
1-11. NAMING METALLIC CATIONS
Many metallic elements have only one possible valence. The names for the
cations formed by these metals are given the name of the element. For example, Na+1
is called sodium ion; Ca+2 is called calcium ion. Other metallic elements, however, may
have more than one valence. Since valence is a measure of combining power, these
elements may form more than one compound with the same anion. Therefore, we must
have some way to differentiate between the varying valences when we name them.
There are two common methods for doing this.
a. The first method uses a root word from the name of the element (or the Latin
name for the element) with a suffix to indicate the valence state. The suffix --ous
indicates the lower valence; the suffix --ic indicates the higher valence. For example,
Hg+1 is called mercurous ion, but Hg+2 is called mercuric ion.
(1)
Questions. (You may wish to refer to table 1-3.)
(a) Al+3 is called ______________________ ion.
(b) Fe+2 is called ferr-- ______________________ ion.
(c)
Fe+3 is called ferr-- ______________________ ion.
(d) K+1 is called ______________________ ion.
(e) Cu+1 is called cupr- ______________________ ion.
(f)
Cu+2 is called cupr- ______________________ ion.
(g) Ba+2 is called ______________________ ion.
(2)
Answers.
(a)
Aluminum.
(b)
--ous.
(c)
--ic.
(d) Potassium.
(e)
--ous.
(f)
--ic.
(g) Barium.
MD0803
1-21
b. The second method for naming metallic cations uses the name of the element
followed by a roman numeral in parentheses to indicate the valence. For example, Cu+1
is written as copper (I) and Cu+2 is written as copper (II). Remember, these methods for
specifying valence need be used only when there is more than one valence possible.
(1)
Questions.
(a) Fe+2 is written
ion.
(b) Fe+3 is written
ion.
Mg+2 is wrjtten
ion.
(d) Hg+1 is written
ion.
(e) Ag+1 is written
ion.
Pb+4 is written
ion.
(c)
(f)
(2)
Answers.
(a)
Iron (II) (ferrous).
(b)
Iron (III) (ferric).
(c) Magnesium.
(d)
Mercury (I) (mercurous).
(e) Silver.
(f)
Lead (IV) (plumbic).
1-12. NAMING ANIONS
There are generally two types of anions. Many anions are elemental; that is
they are made of only one atom of one element. Others are composed of groups of
atoms of one or more elements that pass through a reaction unchanged in most cases.
This latter group of anions is called radicals. We will concern ourselves first with the
naming of elemental or monatomic anions.
MD0803
1-22
a. The names of the elemental anions are made by adding the --ide suffix to the
root of the element's name. Thus anions formed by chlorine (Cl-1) are called chloride
ion; anions formed by oxygen (O-2) are called oxide ion.
(1)
Questions.
(a) Br -1 is called ______________________ion.
(b) S -2 is called ______________________ ion.
(c)
H -1 is called ion.
(d) N -3 is called ion.
(2)
Answers.
(a) Bromide.
(b) Sulfide.
(c)
Hydride.
(d) Nitride.
b. The most common type of anionic radicals consists of a central atom
covalently bonded to a number of atoms of oxygen. Monovalent anionic radicals
(Valence = -1) normally contain three oxygen atoms; radicals with negative valences
greater than one normally contain four oxygen atoms. The names for these normal
types of radicals are formed from the root for the name of the central atom plus the
suffix -ate. Thus, ClO3 -1 is named chlorate and SO4 -2 is named sulfate. It is important
to note that these generalizations have exceptions. The best way to remember the
names and formulas for the radicals is to memorize the common ones. Most of these
are listed in this subcourse.
(1) Sometimes a central atom may be bonded to a different number of
oxygen atoms than normal; in other words, a series of radicals may be formed with the
same central atom. Different suffixes and prefixes are used to name these different
radicals. When there is one less oxygen atom than normal, the suffix -ite is used. The
name for ClO2 -1 is chlorite; SO3 -2 is called sulfite.
(2) Occasionally, there are other radicals in a series. This is especially true
of the halides (fluoride, chloride, bromide, and iodide ions). If there are two less oxygen
atoms than usual, the -ite suffix is used with the prefix hypo-. For example, ClO-1 is
called hypochlorite. If there is one more oxygen atom than normal, the -ate suffix is
used in combination with the prefix per-, so ClO4 -1 is named perchlorate.
MD0803
1-23
(3) A chart summarizing the use of the prefixes and the series of radicals
formed by chlorine as examples follows:
PREFIX
SUFFIX
NAME OF ION
RADICAL
hypo-
--ite
hypochlorite
ClO
--ite
chlorite
ClO2
--ate
chlorate
ClO3
--ate
perchlorate
ClO4
per-
(a) Questions.
1 IO3 -1 is called __________________ ion.
2 IO2 -1 is called __________________ ion.
3 IO4 -1 is called __________________ ion.
4 PO4 -3 is called _________________ ion.
5 PO3 -3 is called _________________ ion.
6 NO3 -1 is called _________________ ion.
7 CO3 -2 is called _________________ ion.
(b) Answers.
1 lodate.
2 lodite.
3 Periodate.
4 Phosphate.
5 Phosphite.
6. Nitrate.
7 Carbonate. (Be sure to learn the exceptions!)
MD0803
1-24
-1
-1
-1
-1
c. There are several significant exceptions to the rules for the naming of anionic
radicals. The most important is the previously mentioned carbonate radical (CO3 -2).
Several others bear mentioning because you are likely to see them in medicine.
(1) Certain radicals are derived when hydrogen is removed from an acid to
form a charge group of atoms (radical). If one hydrogen is removed, the radical gets the
prefix “bi.” This indicates that one hydrogen is missing.
(a) Example 1: When carbonic acid (H2CO3) gives up one hydrogen
ion, it loses a positively charged hydrogen atom. It becomes a radical HCO3 -1 and is
assigned the name bicarbonate. “Bi” indicates one hydrogen was removed.
(b) Example 2: H2PO4 -1 is called the biphosphate radical because it
was derived from phosphoric acid (H3PO4) by removing one hydrogen atom.
(2) Several radicals do not follow any of the above rules. Their names and
formulas must be memorized. Some of the most common are hydroxide (OH -1),
peroxide (O2 -2), and thiosulfate (S2O3 -2).
(3) Occasionally, metals with valences higher than +1 will form salts that
contain oxide or hydroxide ion. When these occur in the middle of the formula, they are
referred to as either oxy- or hydroxy-, respectively. Number prefixes are used to denote
the number of them.
1-13. NAMING SALTS
A salt is an ionic compound containing some cation other than hydrogen and
some anion other than hydroxide and oxide. Since the compound must be electrically
neutral, the total positive valence (from all of the cations) must equal the total negative
valence (from all the anions). This gives us a method for determining the valence of
any particular ion in the formula. The names for salts are made by writing the name of
the cation followed by the name of the anion. For example, CaCl2 has calcium as the
cation and chloride as the anion, so the compound is called calcium chloride. FeSO4
has sulfate as the anion, but we need to know whether the ion is ferrous ion or ferric ion.
This is easy for us to do: since we know the total negative valence (from sulfate) is -2,
the total positive valence (for iron) must be +2; therefore, it is ferrous ion. The
compound is ferrous sulfate.
a. Questions.
MD0803
(1)
KBr is
(2)
Mg(NO3)2 is
(3)
BaSO4 is
.
.
.
1-25
(4)
BiOCl is
.
(5)
HgCl2 is
.
(6)
CuSO4 is
.
(7)
Al(OH)2Cl is
.
(8)
NaHCO3 is
.
(9)
PbSO4 is
.
(10) KBrO3 is
.
b. Answers.
(1)
Potassium bromide.
(2)
Magnesium nitrate.
(3)
Barium sulfate.
(4)
Bismuth oxychloride.
(5)
Mercuric chloride (mercury (II) chloride).
(6)
Cupric sulfate (copper (II) sulfate).
(7)
Aluminum dihydroxychloride.
(8)
Sodium bicarbonate (sodium hydrogen carbonate).
(9)
Plumbous sulfate (lead (II) sulfate).
(10) Potassium bromate.
1-14. NAMING BINARY ACIDS
All acids have hydrogen as the only cation. Binary acids are those acids that are
composed of only two elements; that is, they consist of hydrogen in combination with
some elemental anion. Usually the anion is a halide (F, Cl, Br, I), but binary acids with
other anions also occur.
MD0803
1-26
a. The names for the binary acids are formed by using the prefix hydro-, the root
name for the anion, and the suffix -ic, followed by the word “acjd.” For example, HCl is
called hydrochloric acid.
b. An exception to this rule is hydrocyanic acid which has the formula HCN.
Although this is a ternary acid, the cyanide radical (CN -1) is usually treated like a halide
ion when naming its compounds.
c. The binary acids are really covalent compounds which act as acids only when
they are in solution, especially in water. When you know that one of the binary acids is
by itself, you can properly name it in a similar manner to the salts; thus, HCl as a pure
gas would be called hydrogen chloride.
(1)
Questions.
(a) HBr is called _________________________________________ .
(b)
HI is called ___________________________________________ .
(c)
H2S is called _________________________________________ .
(d) HF gas is called _______________________________________ .
(2)
Answers.
(a) Hydrobromic acid.
(b) Hydriodic acid.
(c)
Hydrosulfuric acid.
(d) Hydrogen fluoride.
1-15. NAMING TERNARY ACIDS
a. The ternary acids generally are made of hydrogen ion combined with one of
the radicals that contain oxygen. For this reason, they are often referred to as
“oxyacids.”
b. When naming the ternary acids, the suffixes on the names of the radicals are
changed and followed by the word "acid" to show the presence of the hydrogen.
Radicals ending in -ate change their suffix to -ic; radicals ending in -ite change their
suffix to -ous. The prefixes, if there are any, are not changed. Occasionally, an extra
syllable is added in the middle of the name for pronunciation purposes--these do not
follow any pattern and must be learned. Here are some examples of naming ternary
acids from the radicals:
MD0803
1-27
RADICAL
NAME OF RADICAL
ACID
NAME OF ACID
SO4-2
Sulfate
H2SO4
Sulfuric acid
SO3-2
Sulfite
H2SO3
Sulfurous acid
ClO -1
Hypochlorite
HClO
Hypochlorous acid
(1)
Questions.
(a) HNO3 is called ____________________________________ .
(b) HNO2 is called ____________________________________ .
(c)
HClO4 is called ____________________________________ .
(d) H2CO3 is called ____________________________________ .
(e) H3PO3 is called ____________________________________ .
(f)
(2)
H3PO4 is called ____________________________________ .
Answers.
(a) Nitric acid.
(b)
Nitrous acid.
(c)
Perchloric acid.
(d) Carbonic acid.
(e) Phosphorous acid.
(f)
Phosphoric acid.
1-16. NAMING BASES
a. The most common bases are those included by the Classical Theory of Acids
and Bases; that is, they are hydroxyl ion (OH -1) donors. Thus most of the bases are
composed of the hydroxyl radical combined with a metallic cation.
MD0803
1-28
b. The names for these bases are made by writing the name of the cation
followed by "hydroxide.” It is not normally necessary to use number prefixes because
the valence of the cation tells us the number of hydroxyl radicals in each molecule. You
can see that this method of naming bases is very similar to the method used for naming
salts, except that the anion is always hydroxide. For example, NaOH is called sodium
hydroxide and Ca(OH)2 is called calcium hydroxide.
(1)
Questions.
(a) KOH is called ____________________________________ .
(b) Mg(OH)2 is called ____________________________________ .
(c)
Fe(OH)2 is called ____________________________________ .
(d) Al(OH)3 is called ____________________________________ .
(e) Fe(OH)3 is called ____________________________________ .
(2)
Answers.
(a) Potassium hydroxide.
(b) Magnesium hydroxide.
(c)
Ferrous hydroxide.
(d) Aluminum hydroxide.
(e) Ferric hydroxide.
1-17. NAMING COVALENT INORGANIC COMPOUNDS
There are a number of inorganic compounds that are bonded into molecules by
covalent bonds. Most of these are the oxides, sulfides, and halides of the nonmetallic
elements.
a. Generally, these compounds are named by writing the name of the central
atom (usually the first one in the formula) followed by the name of the anion formed by
the other element. Number prefixes are used when necessary to avoid confusion
between different compounds formed by the same elements. Here are some examples:
COMPOUND
H2S gas (see para 1-14a(3))
CO
CO2
MD0803
1-29
NAME OF COMPOUND
Hydrogen sulfide
Carbon monoxide
Carbon dioxide
b. There are two very important exceptions to this which you have probably
already seen. These are water (H2O) and ammonia (NH3). Both of these have common
names, which are firmly established in the nomenclature, property of these two
compounds which makes them different from almost all others is the ability to readily
accept a proton (coordinate covalent bond with a hydrogen cation) to form cations.
Thus water becomes hydronium ion (H3O+1); ammonia becomes ammonium ion
(NH4+1) very easily in the right conditions.
(1)
Questions.
(a) SO2 is called ____________________________________ .
(b) SO3 is called ____________________________________ .
(c)
CCl4 is called ____________________________________ .
(d) NI3 is called ____________________________________ .
(e) CS2 is called ____________________________________ .
(f)
NH3 is called ____________________________________ .
(g) NH4+1 is called ____________________________________ .
(h) NH4Cl is called
(2)
.
Answers.
(a) Sulfur dioxide.
(b) Sulfur trioxide.
(c)
Carbon tetrachloride.
(d) Nitrogen triiodide.
(e) Carbon disulfide.
(f)
Ammonia.
(9)
Ammonium.
(h) Ammonium chloride.
MD0803
1-30
1-18. WATERS OF HYDRATION
Many times when a substance crystallizes into a solid, molecules of water are
included in the crystal. These molecules of water combine with the substance in a fixed
ratio, similar to the fixed ratios between the atoms in a molecule. Whenever weighing or
doing calculations based on compounds that have waters of hydration, the amount of
water in the crystals must be taken into consideration.
a. When writing formulas for these compounds, the waters of hydration are
shown by placing a dot (or dash) after the formula for the compound, followed by the
formula for water with a coefficient to indicate the number of waters of hydration. For
example, cupric sulfate forms crystals that contain five molecules of water for each
molecule of cupric sulfate--its formula is written CuSO4.5H2O.
b. Compounds that contain waters of hydration are called hydrates. (If all the
water has been removed by drying, they are called anhydrous.) When writing the
names for these compounds, the number of waters of hydration is indicated by using
number prefixes. Thus, the name for CuSO4.5H2O is cupric sulfate pentahydrate.
Another number prefix seen occasionally in the names of hydrates is hemi-, which
means one-half (1/2).
(1)
Questions.
(a) AlCl3.6H2O is called ____________________________________ .
(b)
Mg3(PO4)2.5H2O is called ________________________________ .
(c)
Na2HPO4.7H2O is called ______________________________ .
(d) FeSO4.7H2O is called ___________________________________ .
(e) Na2CO3.1OH2O is called ________________________________ .
(f)
(2)
CaSO4.1/2H2O is called _________________________________ .
Answers.
(a) Aluminum chloride hexahydrate.
MD0803
(b)
Magnesium phosphate pentahydrate.
(c)
Disodium hydrogen phosphate heptahydrate.
1-31
(d)
Ferrous sulfate heptahydrate.
(e) Sodium carbonate decahydrate.
(f) Calcium sulfate hemihydrate (two molecules of calcium sulfate for
each molecule of water).
Continue with Exercises
MD0803
1-32
EXERCISES, LESSON 1
INSTRUCTIONS. Write the word, words, symbols, or numbers that properly completes
the statement in the space provided or mark the correct word/phrase from those given.
After you complete the exercises, turn to Solutions to Exercises and check your
answers. Reread the material referenced for each exercise answered incorrectly.
1.
An atom is a c__________________ building block. An atom is the
__________________est part of an (element) (compound) that remains
unchanged during any __________________ reaction and exhibits the
(chemical) (physical) properties of that (element) (compound).
2.
A compound is a combination of two or more types of __________________ .
3.
An element is a substance that (can)(cannot) be separated into two or more
types of matter by __________________ or __________________ methods.
4.
A compound is a substance that has been __________________ by physical
methods, but not by __________________ methods.
5.
Matter is anything which occupies ______________ and has ______________ .
6.
Energy is the ability to do __________________ . Examples of energy are
e__________________y, h ________________ , and l ________________ .
7.
The three physical states of matter are s _______________ , l _____________ ,
and g _______________ . A s __________________ has a definite shape and
a/an __________________ volume. A liquid has a/an __________________
shape and a/an __________________ volume. A gas has a/an
__________________ shape and volume.
8.
Examples of physical properties are s _______________ l, c _______________ ,
s _______________ e, f _______________ p _______________ ,
b _______________ p _______________ , and s _______________ y.
MD0803
1-33
9.
An example of a chemical property is _______________ content. To observe a
chemical property, it is necessary that a c ______________ r _______________
occur. Such reactions may be due to l _______________ , h _______________,
or e _______________ .
10.
The three basic parts of an atom are the p _______________ ,
e _______________ , and n _______________ . A p _______________ has a
mass of _______________ and a charge of (-1) (0) (+1). A n _______________
has a mass of _______________ and a charge of (-1) (0) (+1). An
e _______________ has a mass, rounded to the nearest unit of
_______________ and a charge of _______________ .
11.
The atomic number of an element is the number of _______________ in each
atom of the element. In an electrically neutral atom, it is also the number of
_______________ in that atom.
12.
The atomic weight of an element is the number of _______________ and
_______________ in each atom of the element. An atom containing six protons
and six neutrons has an atomic weight of _______________ .
13.
The term electron configuration refers to the _______________ and
_______________ of electrons in the atoms of an element.
14.
Isotopes of an element have in their atoms the same number of
_______________ but different numbers of _______________ .
15.
The maximum number of electrons in the K shell is _______________ , in the L
shell is _______________ , in the M shell is _______________ , and in the N
shell is _______________ .
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1-34
16.
Below is a block from the periodic table.
IV
A
6
2
4
C
12.01115
The number 12.01115 is the atomic _________________ of the element. The
2
numbers 4 represent the e _______________ c _______________ of the
element. There are (two) (four) electrons in the K shell and (two) (four) electrons
in the L shell. The number 6 in the block is the atomic ____________ of the
element. Each carbon atom has (six)(twelve) protons. The letter C is the
____________ for the element carbon. IV A is the ____________ .
17.
If an element's atomic number is 18 and its atomic weight is 40, the number of
neutrons in each atom is _______________ .
18.
The valence of an element is a measure of its c _______________ power.
Valence is the number of e _______________ that an atom must
g _______________ , l
, or sh _______________ to have a
full or stable outer electron _______________ .
19.
An ion is any _______________ atom or group of _______________ . It has
g _______________ or l _______________ at least one electron.
20.
A cation is an ion with a (positive) (negative) charge.
21.
An anion is an ion with a (positive) (negative) charge.
22.
A radical is a charged _______________ of atoms that goes through many
reactions without being _______________ .
23.
According to the octet rule, the outer electron shell of an atom “prefers” to have
_______________ electrons.
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1-35
24.
What are the symbols the following?
a. Barium:
b. Iron:
c.
Sulfate:
,+
____ e, + ____ , + _____
___ ___4, ___ ___
d. Phosphorus:____, +____, ____5
e. Hydrogen: _____, +_____
f.
Potassium: _____, _____1
g. Oxygen: _____ _____
h. Copper: _____ u, +1, + _____
i.
Bromine: ____ ____, ____ 1
j.
Mercury: ____ g, +____, ____2
k.
Iodine: ____, ____ ____
l.
Sulfur: ____, ____ ____
m. Silver: ____ ____, +____
n. Calcium: ____ a, ____ 2
o. Nitrate: N ____ , - ____
p. Aluminum: A ____ , + ____
q. Chlorine: ____ ____, ____ ____
r.
Zinc:
____ ____, ____ ____
25.
The three types of chemical bonds are e________ (i ___________),
c___________, and c________ c___________ .
26.
In _______________ (I ___________) bonding, the electrons are not shared.
They are donated by (one element) (both elements).
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1-36
27.
In c___________ bonding, the electrons (are) (are not) shared. They are
donated by (one element) (both elements).
28.
In c_________ c________ bonding, the electrons (are) (are not) shared. They
are donated by (one element) (both elements).
29.
Listed below are chemical symbols for parts of 12 different molecules. First, label
each part with its valence. Then, write the formula of the molecule with the proper
subscripts to make it electrically neutral.
PARTS
a. H, Cl
FORMULA
_____________________
b. H, SO4 _____________________
c. Na, Br
_____________________
d. H, NO3 _____________________
e. Ca, Cl _____________________
f.
Na, Cl
_____________________
g. Mg, CO3 _____________________
h. Ca, NO3 _____________________
i.
NH4, SO4 _____________________
j.
K, PO4
_____________________
k. Al, SO4 _____________________
l.
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Zn, PO4 _____________________
1-37
30.
Listed below are the names of 10 compounds. First, write down the symbols for
each part of the compound. Then, label each part with its valence. Add subscripts
to make each molecule electrically neutral. The result is the formula for the
compound.
NAME
PARTS
FORMULA
a. Calcium bromide
________
________
b. Sodium carbonate
________
________
c.
________
________
d. Calcium hydroxide
________
________
e. Barium hydroxide
________
________
f.
________
________
________
________
h. Magnesium phosphate ________
________
i.
Sulfuric acid
________
________
j.
Sodium sulfide
________
________
Aluminum hydroxide
Potassium bromide
g. Silver chloride
31. The number of atoms of oxygen in H2SO4 is. ______. The number of atoms of
chlorine in NaCl is ___________. The number of atoms of iron in Fe2(SO4)3 is
______________. The number of atoms of sulfur in K2S is ________. The
number of atoms of oxygen in Al2(SO4)3 is ____________. The number of atoms
of hydrogen in (NH4)2SO4 is ____________.
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1-38
32. Listed below are the names of 10 compounds. First, write down the symbols for
each part of the compound. Then, label each part with its valence. Add subscripts
to make each molecule electrically neutral. The result is the formula for the
compound.
NAME
PARTS
FORMULA
a. Sulfur dioxide
_____________ _____________
b. Mercurous oxide
_____________ _____________
c. Hydrobromic acid
_____________ _____________
d. Mercury (1) chloride
_____________ _____________
e. Potassium bicarbonate ____________ _____________
f. Ammonium iodide
_____________ _____________
g. Aluminum oxychloride _____________ _____________
h. Nitrous acid
_____________ _____________
i.
Potassium permanganate __________ _____________
j.
Magnesium nitrate
hexahydrate
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___________ _____________
1-39
33. Listed below are the formulas for eight different compounds. Write the name for
each one.
FORMULA
a. Fe(HCO3)3
NAME
fer ____ ___ carbon ___
b. MgCl2.6H2O
________ ium chlor ___
_______ hydrate
c. HI
___dri _____acid
d. KOH
___ t ____ ium ___ dro ______
e. K2HPO4
___ tassium _____ogen phos______
f. FeCO3
fer _____ _____ bon _____
g. Ca(OH)NO3
cal ____ ____ dro _____ nit _____
h. NaOH
sod _____ hydro ______
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1-40
34. Listed below are the formulas for 12 different compounds. Write the name for each
one.
FORMULA
NAME
a. HNO3
n __________ a __________
b. FePO4
f ___________ p __________
c. Al(OH)2Cl
a ___________ d __________
d. (NH4)2SO3
a ___________ s __________
e. Hg3PO4
____________ p __________
f. NaHCO3
s ____________ b __________
g. NCl
h ____________ a __________
h. MgO
m ____________ o __________
i.
Ba(OH)2
b ____________ h __________
j.
CaHPO3
c _______ h _______ p _______
k. CaCO3
c __________ c __________
l.
p __________ c __________
KCl
Check Your Answers on Next Page
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1-41
SOLUTIONS TO EXERCISES, LESSON 1
1.
chemical,
small(est), element, chemical, chemical, element (para 1-3c(1))
2.
elements (para 1-3c(2))
3.
cannot, physical, chemical (para 1-3c(1))
4.
purified, chemical (para 1-3c(2))
5.
space, weight (para 1-3)
6.
work, electricity, heat, light (para 1-4)
7.
solid, liquid, gas
solid, definite
indefinite; definite
indefinite, (para 1-3a)
8.
smell, color; shape; freezing point; boiling point; solubility (para 1-3b)
9.
energy
chemical reaction
light; heat; electricity (para 1-3b)
10.
proton; electron, neutron
proton; one; +1
neutron; one; 0
electron; zero; -1 (para 1-5b,c)
11.
protons, (para 1-5b(3)(a))
electrons (para 1-5c(3))
12.
protons, neutrons
12 (para 1-5b(3)(b))
13.
Number, locations (para 1-5c(1))
14.
Protons; neutrons (para 1-5g)
15.
2, 8, 18, 32 (para 1-5c(2))
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1-42
16.
weight
electron configuration
two, four,
number
six
symbol
group (para 1-5; fig 1-2)
17.
22 (para 1-5f)
18.
combining
electrons; gain, loose, share, shell (para 1-6a)
19.
charged atoms
gained, lost (para 1-6f)
20.
positive (para 1-6f)
21.
negative (para 1-6f)
22.
group, changed (para 1-6f)
23.
eight (para 1-6b)
24.
a
b
c
d
e
f
g
h
i
j
k
l
m
n
o
p
q
r
25.
electrovalent, ionic; covalent, coordinate covalent (para 1-6g)
MD0803
Ba: +2
Fe; +2, +3
SO4: -2
P +3 +5
H +1
K +1
O; -2
Cu +1 +2
Br -1
Hg +1 +2
I; -1
S: -2
Ag; +1
Ca; +2
NO 3; -1
Al; +3
Cl; -1
Zn; +2 (para 1-6e; Table 1-3)
1-43
26.
electrovalent (ionic); one element (para 1-6g(1))
27.
covalent; are, both elements (para 1-6g(2))
28.
coordinate covalent, are; one element (para 1-6g(3))
29.
a
b
c
d
e
f
g
h
i
j
k
l
30
Formula
a CaBr2
b Na2CO3
c Al (OH)3
d Ca(OH)2
e Ba(OH)2
f
KBr
g AgCl
h Mg3(PO4)2
i
H2SO4
j
Na2S (paras 1-7, 1-12, 1-13, 1-15, Table 1-3)
H + , Cl - ; HCl
H + , SO4 -2; H2SO4
Na +, Br - ; NaBr
H +, NO3 - ; HNO3
Ca +2 , Cl - ; CaCl2
Na + , Cl - ; NaCl
Mg +2 , CO3 -2 ; MgCO3
Ca +2 , NO3 -1 ; Ca(NO3)2
NH4 + , SO4 -2; (NH4)2 SO4
K + , PO4 -3 , K3PO4
Al +3 , SO4 -2 ; Al2(SO4)3
Zn +2 , PO4 -3; ZN3(PO4)2 (para 1-7; Table 1-3)
31. 4
1
2
1
12
8 (para 1-7a)
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1-44
32.
Formula
a SO2 (Table 1-3, paras 1-10, 1-17)
b Hg2O (Table 1-3; para 1-11a)
c HBr (Table 1-3; para 1-14)
d HgCl (Table 1-3; para 1-11b)
e KHCO3 (Table 1-3; para 1-12c(1))
f
NH4I (Table 1-3; Para 1-17b)
g AlOCl (Table 1-3; para 1-12c(3))
h HNO2 (Table 1-3; para 1-15)
i
KMnO4 (Table 1-3; para 1-12b)
j
Mg(NO3)2.6H2O (Table 1-3; paras 1-12b, 1-18)
33.
Name
a ferric bicarbonate (Table 1-3; paras 1-11, 1-13)
b magnesium chloride, hexahydrate (Table 1-3; paras 1-10, 1-18)
c hydriodic acid (para 1-14a(2))
d potassium hydroxide (para 1-16a(1))
e potassium monohydrogen phosphate or dipotassium phosphate
(paras 1-12b(1), 1-13)
f
ferrous carbonate (para 1-13)
g calcium hydroxynitrate (paras 1-12c(3), 1-13)
h sodium hydroxide (para 1-16)
34.
Name
a nitric acid (para 1-15b(1)(a)
b ferric phosphate (paras 1-12, 1-13)
c aluminum dihydroxy-chloride (para 1-12c(3))
d ammonium sulfite (paras 1-12b, 1-17)
e mercurous phosphate (paras 1-11, 1-12, 1-13)
f
sodium bicarbonate (para 1-13a(8))
g hydrochloric acid (para 1-14)
h magnesium oxide (para 1-17)
i
barium hydroxide (para 1-16)
j
calcium monohydrogen phosphite (paras 1-12, 1-13)
k calcium carbonate (para 1-13)
l
potassium chloride (para 1-13)
End of Lesson 1
MD0803
1-45
LESSON ASSIGNMENT
LESSON 2
Elements of Chemical Change.
LESSON ASSIGNMENT
Paragraphs 2-1 through 2-13 and exercises.
LESSON OBJECTIVES
After completing this lesson, you should be able to:
2-1.
Given a chemical equation, describe the
chemical events occurring in the reaction, to
include names of reactants and products and any
special conditions indicated.
2-2.
Given a description of a chemical reaction, write
and balance the equation for the reaction.
2-3.
Define equilibrium exothermic, endothermic,
milligram molecular weight, and milliequivalent
weight.
2-4.
Given the name of an inorganic compound
commonly encountered in medicine, calculate
the milligram formula weight and milliequivalent
weight of that compound.
2-5.
Define oxidation and reduction.
2-6.
Given an inorganic chemical reaction, identify the
oxidizing and reducing agents.
2-7
Define an acid and a base according to the
classical theory and according to the BronstedLowry theory.
2-8.
Given a chemical formula, indicate whether it is
an acid or a base.
2-9.
List three properties of bases and five properties
of acids.
2-10. State the antidotes for external or internal contact
with a strong acid or a strong base.
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2-1
SUGGESTION
MD0803
2-11.
Correctly define the term salt.
2-12.
Given a chemical formula, indicate whether it is
a salt, acid, or base.
2-13.
Given a pH value, indicate whether it is acidic,
basic, or neutral.
2-14.
List three methods for measuring pH.
2-15.
State the function and two general components
of a buffer system.
2-16.
From a list of pairs of compounds, select which
represents a buffer system.
2-17.
State four important properties of water and two
major methods of water purification.
2-18.
Define solute, solvent, solubility, dissociation,
and electrolyte.
After completing the assignment, complete the
exercises at the end of this lesson. These exercises
will help you to achieve the lesson objectives.
2-2
LESSON 2
ELEMENTS OF CHEMICAL CHANGE
2-1.
CHEMICAL REACTIONS
As a provider of health care, you will not be required in most cases, to write and
balance chemical equations. You will, however, be using and/or seeing the effects of
chemical reactions on a daily basis. Chemical reactions are frequently used to explain
various concepts of pharmacology and physiology. Consider drugs. All drugs are
chemicals and any pharmacological reference you consult will refer to the chemical
changes drugs undergo in the body. Consequently, it is essential that you have a basic
knowledge of what a chemical reaction involves and how that chemical reaction can be
expressed as a chemical equation.
a. Definite Composition. When atoms combine, they do so in definite ratios of
intact atoms to produce compounds with definite composition. Note that this
combination is by number of atoms, not by weights of atoms. What the individual atoms
happen to weigh is not important. Atoms do not know what they weigh. When they do
interact and combine, it is always as whole particles, and the particle-to-particle or
atom-to-atom ratio can always be expressed in simple, whole numbers. Chemical
changes do not split atoms into fractional pieces. This is the reason we are able to write
a formula such as HCl for the compound hydrochloric acid. Hydrochloric acid is always
formed from one atom of hydrogen and one atom of chlorine. Since a chemical reaction
is merely a change in matter, and matter consists of atoms or molecules, we can
discuss chemical reactions by talking about interactions of individual molecules or
atoms.
b. Chemical Equations. In discussing a chemical reaction, it would be very
cumbersome to write it out in the same manner as we state it verbally. To get around
this problem, chemists have developed chemical equations. Chemical equations are
abbreviated ways of writing chemical reactions. They save much writing and effort and
give at least as much information as a verbally stated reaction. Chemical equations
show:
MD0803
(1)
The kinds of atoms or molecules reacting.
(2)
The products formed.
(3)
The number of atoms entering the reaction.
(4)
The number of molecules formed in the product.
(5)
The proportion in which the substances react to give definite products.
2-3
c.
Chemical Symbols. In writing chemical equations, we use a number of
symbols. The most common symbols are shown below with their meanings.
SYMBOL
∆
→
↑
↓
MEANING
Heat (a form of energy)
“yields,” indicates direction of reaction
given off as a gas
given off as a precipitate
As we illustrate several types of reactions, the uses of these symbols will become
apparent.
d. Types of Reactions. There are four types of chemical reactions, which are
possible: Combination reactions, decomposition reactions, single replacement
reactions, and double replacement reactions.
(1) Combination reactions. A combination reaction can be represented by
the chemical equation A + B --> AB (one atom of A plus one atom of B yield one
molecule of AB). A specific example of this type of reaction is the combination of a
metal with oxygen to yield a metallic oxide.
2 Mg + O2 --> 2 MgO
This equation tells us that two atoms of magnesium and one molecule of oxygen react
to form two molecules of magnesium oxide.
(2) Decomposition reactions. The general equation representing
decomposition reactions is AB → A + B. Here is a good example:
CaCO3 → CaO + CO2
∆
↑
This equation tells us that calcium carbonate will yield calcium oxide and carbon
dioxide. The ∆ also tells us that this reaction occurs when heat is applied to calcium
carbonate. The ↑ indicates that the carbon dioxide is given off as a gas.
MD0803
2-4
(3) Single replacement reactions. The general equation for a single
replacement reaction is A + BC → AC + B. An example is:
Zn + CuSO4 → ZnSO4 + Cu
This equation tells us that one atom of zinc and one molecule of cupric sulfate yield one
molecule of zinc sulfate and one atom of copper.
(4) Double replacement reactions. The most commonly occurring reaction
is the double replacement reaction. The general equation for this reaction is AB + CO
→ AD + CB. Double replacement reactions can be further subdivided into several
classes. The most common of these classes are the precipitation reaction, the acidbase reaction, and the oxidation-reduction reaction. An example of the precipitation
reaction is:
BaCl2 + Na2SO4 → 2 NaCl + BaSO4
This equation tells us that one molecule of barium chloride reacts with one molecule of
sodium sulfate to yield two molecules of sodium chloride and one molecule of barium
sulfate as a precipitate. Acid-base and oxidation-reduction reactions will be covered
later.
2-2.
WRITING CHEMICAL EQUATIONS
At this point, you have seen several examples of chemical equations and should
be familiar with the symbols used in an equation. We will now examine the process of
writing an equation when we are given a verbal description of the reaction. One general
rule that must be kept in mind is that there will always be the same number and kinds of
atoms in the products of a reaction as in the reactants. This is because matter can
neither be created nor destroyed in a chemical reaction and atoms always combine in
certain proportions. When given a written verbal description of a chemical reaction, the
following steps are used to write the equation for the reaction.
a. Write the symbols for all elements involved.
b. Write the correct formulas for any compounds and check for diatomic
molecules. (Some elements never exist as single atoms but only as diatomic
molecules. These elements can be identified from their names, which end in -gen or
-ine. The common diatomic molecules are hydrogen (H2), nitrogen (N2), oxygen (O2),
chlorine (Cl2), fluorine (F2), and bromine (Br2).)
c. Balance the equation by placing coefficients where appropriate.
Remember that there must be equal numbers of atoms of each kind on both sides of the
equation. In this step, the subscripts that were used in writing the correct formulas
cannot be changed.
MD0803
2-5
2-3.
EXAMPLE
For application of these steps, consider this description of a reaction. Calcium
metal and water react to yield calcium hydroxide and hydrogen gas.
a. Write the symbols for all elements involved.
Ca, O, H
b. Write the correct formulas for any compounds and check for diatomic
molecules.
Ca + H2O → Ca(OH)2 + H2
↑
c. Balance the equation by placing coefficients where appropriate. Look at the
number of atoms of each element in the products and reactants.
REACTANTS
PRODUCTS
1 Ca
1 O
2 H
1 Ca
2 O
4 H
It is apparent here that there are twice as many oxygen and hydrogen atoms in the
products as reactants. How can this equation be balanced to give equal numbers of
atoms on both sides? Fill in the coefficients of the molecules in the equation below.
Ca +
H2O → _____
Ca(OH)2 + ________ H2 ↑
Since there are twice as many hydrogen and oxygen atoms on the right as on the left, if
we could double the numbers of these atoms on the left, we would have a balanced
equation. This can be done by placing a two in front of H2O. All the other coefficients
would be one (if there is no coefficient, we assume it is one, so there is no need to write
it in front of each molecule).
2-4.
EQUILIBRIUM REACTIONS
We have implied that all reactions only go in the direction of the products, but this
is not always the case. Sometimes as products are formed, they react with one another
or decompose to form the reactants. Thus, the reaction is going in both directions at the
same time, and if allowed to continue indefinitely, would result in a constant amount of
products and reactants. Reactions that go in both directions are called equilibrium
reactions, and when the rate of formation of product is the same as the rate of formation
MD0803
2-6
of reactant, they are said to be in equilibrium. In writing an equation, we indicate
equilibrium by drawing arrows pointing in opposite directions
( ------> )
.
As an example of an equilibrium reaction, consider the dissociation of a compound into
ions:
( <------- )
Na2CO3
------->
<------
2Na+ + CO3 -2
Sodium carbonate in solution dissociates into sodium ions and carbonate ions. Some of
the ions come back together to form sodium carbonate. Thus, an equilibrium is
established.
2-5.
EXTERNAL CONDITIONS AFFECTING CHEMICAL REACTIONS
External conditions that affect reactions are usually types of energy that are put
into a reaction, such as heat or light. Chemical reactions are always accompanied by
an energy change. Either energy is released or it is acquired. When the amount of
energy is changed, so is the amount of matter. This is called the Law of Conservation of
Matter and Energy. However, ordinary chemical reactions involve such small matter
changes that they go undetected and may be ignored.
a. Heat. Generally, heat is the form of energy we are most concerns us most. It
may affect a reaction in one of two ways.
(1) Exothermic reactions. If a reaction gives off heat, it is called an
exothermic reaction. External heat, if supplied to this type of reaction, will slow down
the rate of reaction.
(2) Endothermic reactions. If a reaction takes in heat, it is an endothermic
reaction. lf heat is added to an endothermic reaction, the rate of reaction will increase.
This may be of value in the preparation of medicinal products.
b. Light. Light is a form of energy that may cause many chemicals to
decompose. For this reason, it is necessary to protect some drugs from contact with
light by placing them in dark-colored or opaque containers. These containers prevent
most or all of the outside light from coming into contact with the drug.
2-6.
REACTING QUANTITIES
It has already been emphasized that all reactions occur on an atom-to-atom
level. This presents a small problem to us, since we cannot hold an atom in our hand,
or count out a specific number of atoms to put into a reaction. How then do we
measure amounts of material that will react together? Chemists long ago solved this
MD0803
2-7
problem by learning how to count particles indirectly. They did this by measuring
samples of the chemicals in particular ratios by their weights. To understand the means
of doing this, we need to expand our concept of atomic weight to compounds in the form
of the formula (or molecular) weight.
a. Milligram Formula (Milligram Molecular) Weight. When atoms combine to
form compounds, the atomic nuclei are not affected. There is no net loss of weight.
Regardless of whether the particle formed is a molecule or an ion group, it will have a
formula and a formula weight. The formula weight of a compound is the sum of the
atomic weights of all the atoms that appear in its chemical formula. Consider, for
example, carbon dioxide:
Atoms: C + O + O = CO2 (molecule)
Atomic weights: 12 + 16 + 16 = 44 (formula weight)
While we have arrived at a formula weight which is in terms of atomic mass units, it is
much more useful to express it in terms of milligrams. This is known as the milligram
formula weight. For the example above, CO2, the milligram formula weight is 44 mg.
This is a quantity that we can measure and see, and thus can easily work with. It also
represents a reacting unit of the compound.
b. Molarity. A molar solution, or a one molar (1M) solution, consists of onegram molecular weight (GMW) of solute dissolved in enough water to make 1 liter of
finished solution. Molarity, then, is the number of GMWs dissolved in enough water to
make a finished solution of 1000 ml. Molar solutions may have as a solute a solid, a
liquid, or a gas. Later in this subcourse, we will use the concept of molarity to explain
the measurement of acidity, called the pH.
(1) Calculating the gram molecular weight. One-gram molecular weight of a
substance is its molecular weight expressed in grams. Thus, a GMW of NaOH would
be 40 grams, where the atomic weights are as follows: Na = 23, O = 16, and H = 1.
Thus, .5 GMW of NaOH would be 20 grams, and so forth. A mole is one-gram
molecular weight of a substance. Thus, a mole of NaOH is 40 grams of NaOH; a halfmole (.5 mole) is 20 grams; two moles of NaOH are 80 grams, and so on.
(2) Calculating the molarity of a solution. To find the molarity of a solution,
we divide the number of gram molecular weights of solute by the number of liters of total
solution. The formula may be written:
Molarity =
MD0803
no. of GMWs of solute
no. of liters of solution
2-8
Since many problems are stated in terms of the weight of solute and require you to
determine the number of gram molecular weights (moles), the following formula will be
of benefit:
weight of solute
GMW
No. of GMWs =
(3) Example. What is the molarity of a solution containing 29.25 grams of
sodium chloride in 500 ml. of total solution?
Step 1. Find the number of GMWs.
GMW of NaCl = 58.4 grams
No. of GMWs =
weight of solute
GMW
29.25
58.4
No. of GMWs =
Step 2.
= 0.5
Find the molarity.
Molarity =
no. of GMWs of solute
no. of liters of solution
500 ml = 0.5 liter
Molarity =
0.5 =
0.5
1 molar or 1M
c. Milligram Equivalent Weight (Milliequivalent Weight). Sometimes we are
interested in more than just the weight ratios of reacting compounds. Since the valence
of an element is a measure of that element's combining power, the valences in a
compound should be indicative of their reactivity. Therefore, chemists have modified the
milligram formula weight to include the positive or negative valence of a compound.
This value is called the milligram equivalent weight and is defined as the milligram
molecular weight divided by the total positive or negative valence. Consider, for
example, sodium hydroxide:
Milligram molecular weight = 40 mg
Total positive valence = 1
Milligram equivalent weight = 40 mg = 40 mg
1
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Another example is potassium phosphate (K3PO4):
Milligram molecular weight = 212 mg
Total positive valence = 3
Milligram equivalent weight = 212 mg = 70.7 mg
3
In a reaction, one milliequivalent (mEq) weight of one compound will react with one
milliequivalent weight of another. If we are reacting two compounds, then, we can
determine how much of each compound should be used to obtain a desired amount of
product.
2-7.
OXIDATION-REDUCTION REACTIONS
Previously, we have examined the processes involved in writing, balancing, and
interpreting reactions and looked at examples of several types of reactions. One type of
reaction we did not examine closely was the oxidation-reduction reaction (sometimes
called redox reaction). Even though this type of reaction is very important in the
chemistry of drug molecules, it is beyond the scope of our instruction to study them in
detail. However, a basic understanding of this process will be valuable to you in
understanding many of the incompatibilities, storage problems, and some disease
states that you will encounter later.
a. Review of Valence. Before these reactions are studied, valence should be
reviewed briefly. The following two valence concepts are especially important in
oxidation-reduction reactions:
(1) All elements in their free and uncombined state are considered to have a
valence of zero. This holds even for those elements that are diatomic molecules in their
free state.
(2) All atoms can exist in a number of valence states. The common
valences which you learned previously are the preferred and most stable valences
under normal conditions, but other valences can and do occur.
(3) These two concepts are important because oxidation-reduction reactions
always involve a change in the valence numbers of some of the elements involved in
the reaction.
b. Oxidation. Oxidation, in inorganic chemistry, is defined as the loss of
electrons or an increase in the valence of an element. Consider, for example, the
oxidation of elemental iron:
FeO-2e
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-
2-10
---> Fe+2
Iron in its free state has a valence of zero and is very reactive since its common valence
state is +2 or +3. It loses two electrons to become the ferrous ion. The valence has
gone from 0 to +2, thus iron has been oxidized. It can undergo further oxidation to the
+3 valence state:
Fe+2 -le
-
---> Fe+3
Here the ferrous ion has lost another electron to become a ferric ion.
c. Reduction. In inorganic chemistry, reduction is defined as the gain of
electrons or a decrease in the valence of an element. Consider the reduction of
elemental oxygen:
O2 + 4e - -----> 2 O -2
Observe that oxygen is a diatomic molecule in its free elemental form and has a valence
of zero. Since the most common valence state of oxygen is -2, oxygen accepts
electrons readily to become the oxygen anion. The valence of each oxygen atom has
gone from 0 to -2, thus oxygen had been reduced. If the valence is made smaller
(reduced), reduction has occurred.
d. Oxidizing and Reducing Agents. For all practical purposes, it is impossible
to simply add or subtract electrons from an element except in an electrolytic cell. In
fact, the oxidation of one element and the reduction of another always occur
simultaneously. One element loses the electrons; the other element gains the electrons
that are lost by the first. Consider these two reactions when they are combined:
2Fe - 4e - -----> 2Fe+2
O2 + 4e - ------> 2O-2
2Fe + O2 ------> 2FeO
This is an oxidation-reduction reaction that is very common in our industrialized society.
The oxidation of iron by atmospheric oxygen gives us iron oxide, commonly known as
rust. In this reaction, oxygen was reduced, going from a zero to a -2 state by receiving
electrons from iron. Because it accepted the electrons from iron and allowed the iron to
oxidize, oxygen is called an oxidizing agent. Iron, which gave up electrons, is called the
reducing agent. General characteristics of reducing and oxidizing are shown in the
following table.
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2-8.
REDUCING AGENT
OXIDIZING AGENT
(1)
(2)
(3)
(1)
(2)
(3)
Gives up electrons
Oxidized during reaction
Unusually low valence
state compared to most
common state
Gains electrons
Reduced during reaction
Unusually high valence
state compared to most
common state
ACIDS AND BASES
The two most important classifications of compounds in inorganic chemistry are
acids and bases. The following discussion forms the groundwork for understanding
some of the most important chemical changes which you will encounter.
a. Classical Acid-Base Theory. Svante Arrhenius, in 1887, published the first
satisfactory explanation of the acid-base phenomena that had been observed by
chemists.
(1) Acids. Arrhenius defined an acid as a compound that donates protons
(H+) in solution. Examples would be any of the compounds you learned to name as
acids earlier in this subcourse.
HOH
HCl -----> H+ + Cl H2SO4
HOH
-----> H+ + HSO4 -
NOTE: The HOH (H2O), which indicates that water is the solvent in these reactions.
Both HCl and H 2SO4 contribute protons in solution.
(2) Bases. Arrhenius defined a base as any compound that donates
hydroxyl (OH-) ions in solution. Again, you should be familiar with several examples
from your nomenclature studies.
HOH
NaOH -----> Na+ + OH HOH
KOH -----> K+ + OH (3) Discussion. These classical definitions are based on the
dissociation of the compounds into ions in solution. This implies that all acids and
bases must contain exchangeable hydrogen and hydroxyl ions, respectively, in their
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formulas. This theory did explain the majority of the compounds known at the time, but
there were some exceptions. Chemists knew, for example, that metal oxides (MgO,
CaO, etc.) dissolved in water exhibited base-like properties. Also, ammonia (NH3) in
solution exhibited the properties of a base. The attempts to explain these exceptions
led to new definitions of acids and bases.
b. Modern Acid-Base Theory. In 1923, Bronsted and Lowry, two chemists in
different countries, independently derived new definitions of acids and bases to explain
the exceptions to the classical theory. The new theory they developed was
named, appropriately, the Bronsted-Lowry theory. This theory differs from the classical
theory in that the dissociation of water is considered as well as the dissociation of the
compound.
(1) Dissociation of water. Even though we often think of water as merely
being an inert solvent, it does dissociate into ions.
-------->
H2O <-------- H+ + OH This is an equilibrium type reaction as indicated by the double arrow. Actually, very few
ions exist at any time since they rapidly recombine to form molecular water. If we put
numbers in this reaction, there are 500 million molecules of water for each hydrogen or
hydroxyl ion.
(2) Bronsted-Lowry acid. By the Bronsted-Lowry theory, an acid
is any compound (charged or uncharged) capable of donating a proton. This is
essentially the same as the classical definition.
(3) Bronsted-Lowry base. The real value of the Bronsted-Lowry theory is in
the definition of a base. A base is defined as a charged or uncharged substance
capable of accepting a proton. Generally, the proton a base accepts comes from the
dissociation of water.
(a) Consider, for example, ammonia dissolved in water:
-------->
-------->
NH3 + H2O <------- NH3 + H+ + OH - <------- NH4 + + OH By accepting a proton from water, ammonia has effectively increased the concentration
of hydroxyl ions in the solution. This would account for the properties like those of a
classical base.
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(b) A second example would be magnesium oxide dissolved in water.
--------->
MgO + H2O <---------
MgOH+ + OH -
By accepting a proton from water, magnesium oxide has likewise increased the
concentration of hydroxyl ions in the solution.
NOTE:
The two theories explain all the properties of acids and bases that will be
utilized in medicine. It deserves mention that there are other theories of acids
and bases that explain more complex phenomena. If these are of interest to
you, a college chemistry text should have a discussion of some of them.
c. Properties of Acids. We have defined all acids based on one common
property, the ability to donate hydrogen ions in solution. Therefore, you should expect
them all to exhibit a set of common properties, which they do. The properties we are
concerned with are as follows:
(1) Acids change blue litmus paper to red. Litmus paper, which contains
dyes sensitive to hydrogen ion concentration, turns red when there is a high
concentration, blue when there is a low concentration.
(2) Acids have a sour taste. This property is familiar to you if you have ever
tasted a lemon. Lemons contain citric acid, which gives them their sour taste.
(3)
Acids react with metals to release hydrogen gas. For example:
Zn + 2H+ ------> Zn ++ + H2
You will notice that this reaction is an oxidation-reduction reaction. For practice, pick
out the oxidizing and reducing agents.
(4) Acids react with carbonates and bicarbonates to form carbon dioxide.
For example:
CaCO3 + 2HCl ------> CaCl2 + H2O + CO2 ↑
(5) Acids react with bases to form salts and water (neutralization reaction).
For example:
HCl + NaOH -------> NaCl + H2O
(salt)
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d. Properties of Bases. In the same manner that all acids had certain
properties in common, all bases have related properties. The ones that are important to
the medical personnel are as follows:
(1) Bases change red litmus paper to blue. This is just the opposite of the
change which acids cause in litmus paper.
(2)
Bases possess a bitter taste and feel soapy when in contact with the
skin.
(3) Bases react with acids to form salts and water (neutralization reaction).
This is the same type reaction as previously discussed under acids.
e. Classification of Acids and Bases. Even though all acids possess certain
properties in common, as do bases, not all possess them to the same degree. Some
acids, for example, will completely neutralize sodium hydroxide with equal
concentrations while others will only partially neutralize this base. As you might
suspect, the differences in the strengths of acids results from differing abilities to donate
hydrogen ions and the differences in bases from differing abilities to donate hydroxyl
ions or accept hydrogen ions.
(1) Some acids and bases dissociate more readily than others when placed
in solution. Those that dissociate at a rate greater than 50 percent are considered to be
strong acids or bases. Weak acids and bases dissociate at a rate that is less than 50
percent. Examples:
(a) When hydrochloric acid (HCl) is placed in solution, most of the
molecules will dissociate to form free H+ ions and Cl- ions. Hydrochloric acid is
therefore considered a strong acid.
(b) When carbonic acid (H2CO3) is placed in solution, less than 50%
will ionize into free H+ ions and HCO3 - ions. Most of the molecules will remain in
molecular form.
HOH
+
H2CO3 -----> H + HCO3 -----> H2CO3
(2) This means one mole (gram molecular weight) of HCl will produce more
hydrogen ion in solution than will one mole of H2CO3 and will consequently exhibit acidic
properties to a greater degree than will carbonic acid. A simpler way to say this is that
HCL is a stronger acid than H2CO3.
(3) The same rationale holds for bases as well as acids. Therefore, we can
divide or classify acids or bases into groups based on their dissociation--strong acids or
bases (those that dissociate completely) and weak acids or bases (those that dissociate
to a small degree).
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f. Acids and Bases of Medicinal Importance. One may come in contact with
a number of important acids and bases. You must be able to identify them as acids or
bases and know their relative strengths. Table 2-1 shows these acids and bases.
There is not an easy way to differentiate between strong and weak acids, but strong and
weak bases can be differentiated based on valence. Strong bases have a positive
valence of one; weak bases have a positive valence greater than one.
g. Safety and Antidotes. Acids and bases should be handled with care to
avoid spilling on skin. They should not be taken internally unless intended for that
purpose. If the skin is exposed to these compounds or is ingested, the following
antidotes are recommended for first aid treatment.
(1)
Acids.
(a) External. Use large amounts of water to wash acids off the skin.
Exception: If phenol (an organic acid) is spilled on the skin, wash off with alcohol.
(b) Internal. Give an antacid, other than a carbonate or bicarbonate,
such as milk of magnesia or magnesium oxide. DO NOT give an emetic or induce
vomiting.
(2)
Bases.
(a) External. Wash the area with large amounts of water.
(b) Internal. Give a weak acid such as vinegar or fruit juice. Weak
acids (or weak bases) are only effective if administered within 10-15 minutes of
ingestion of strong base (or strong acid). DO NOT give an emetic or induce vomiting.
2-9.
SALTS
Previously, it has been stated that one of the properties associated with acids
and bases is the neutralization reaction. This reaction involves the production of a salt
and water from the reaction of an acid and a base. We will now examine various types
of salts produced in neutralization reactions. Salts are the third major classification of
inorganic compounds (acids and bases being the first two). They are important in the
physiology of the body and are often used as therapeutic agents.
a. Definition. We have already given one definition of a salt in our discussion,
that is, the product of a reaction between an acid and a base. A more specific
definition, however, would be an ionic compound formed by the replacement of part or
all of the acid hydrogen of an acid by a metal or a radical acting like a metal. It is an
ionic compound that contains a positive ion other than hydrogen and a negative ion
other than hydroxyl (OH-) or "O-2," as in MgO.
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b. Types of Salts. There are four types of reactions possible between acids
and bases as we classified them (strong or weak) earlier. These are as follows:
Relative strength of common acids and bases
STRONG ACIDS
HCl
H2SO4
H3PO4
WEAK ACIDS
Hydrochloric acid
Sulfuric acid
Phosphoric acid
HC2H3O2 (HAC) Acetic acid
H2CO3
Carbonic acid
H3BO3
Boric acid
STRONG BASES
KOH
NaOH
WEAK BASES
Potassium hydroxide
Sodium hydroxide
Fe(OH) 2
Al(OH)3
NH3
Ferrous hydroxide
Aluminum
hydroxide
Ammonia
_____________________________________
* Ca(OH) 2
Calcium hydroxide
Mg(OH) 2
Magnesium hydroxide
MgO
Magnesium oxide___________
* Notice that Ca(OH) 2, Mg(OH) 2, and MgO which forms Mg(OH) 2 in water are
chemically classified as strong bases because of their high degree of dissociation.
Because they are only slightly soluble in water, they produce low concentrations of
the hydroxide (OH-) ion in solution. Since calcium hydroxide and magnesium
hydroxide do not produce tissue damage, they can be safely used as therapeutic
agents (e.g., antacids).
Table 2-1. Relative strength of common acids and bases.
(1)
Strong acid and strong base.
HCl+ NaOH ---> NaCl + H2O
(2)
Weak acid and weak base.
2H2CO3 + Fe(OH)2 ---> Fe(HCO3)2 + 2H2O
(3)
Strong acid and weak base.
2HCl + Fe(OH)2 -----> FeCl2 + 2H2O
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(4)
Weak acid and strong base.
H2CO3 + NaOH -----> NaHCO3 + H2O
NOTE: These four reactions result in three types of salts. Reactions (1) and (2) result
in neutral salts (that is, in terms of pH), which means a solution of the salt
in water will be a neutral solution. Reactions such as (3) result in acidic salts,
which produce acidic solutions. Reaction (4) results in basic salts, which produce
basic solutions.
c. Determination of Salt Type. To determine the type of salt from a chemical
formula, we employ the following steps:
(1) The first element comes from a base. Determine which base and
whether it is weak or strong.
(2) The remainder of the formula comes from the acid. Determine which
acid and whether it is weak or strong.
(3) By knowing the strengths of the acid and base that formed the salt,
the salt type can be assigned. Table 2-2 is a summary of salt types resulting from
various acid-base combinations.
d. Example. Al2(SO4)3.
(1) The first element, aluminum, comes from the base Al(OH)3. Since it has
a valence of +3, it is a weak base.
(2)
The sulfate radical comes from H2SO4, sulfuric acid, which is a strong
acid.
(3) This compound is an acidic salt since it is the product of a reaction
between a strong acid and a weak base.
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e. Example. FeBO3.
(1) The first element, iron, comes from the base Fe(OH)3, and since its
valence is +3, it is a weak base.
(2)
The borate radical comes from boric acid, which is a weak acid.
(3) Thus, this is a neutral salt, since it is the product of a reaction between a
weak acid and a weak base.
f. Importance of Type of Salt. The type of salt is very important when a salt is
used medicinally, since the body maintains a specific acidity in the tissues and fluids.
The type of salt is also important in the prediction and understanding of
incompatibilities. It is important for you to identify the type of salt from its formula. The
importance and use of the type will become clear to you as you progress through the
course.
Table 2-2. Salt types resulting from various acid-base combinations.
2-10. pH AND ACIDITY
In discussing acids, bases, and salts, we often refer to a solution or compound
being acidic, neutral, or basic in a qualitative manner. This concept is useful to us in a
general sense, but would be of much greater value if we could speak in quantitative
terms. It would be valuable if we could answer the question of how acidic one solution
is in relation to another solution.
a. pH. The solution to this problem is not as difficult as it may seem. Acids
donate protons (hydrogen ions, H +) in solution. Thus, the acidity of a solution must be
related to this property.
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(1) In fact, the acidity of a solution is the concentration of hydrogen ions in
that solution. Since we can calculate the hydrogen ion concentration, as you will learn
later, we can now determine a numerical value of the acidity of a solution. The
concentrations of hydrogen ions in both acidic and basic solutions are generally very
small. A strong solution of HCl, for example, may contain only 0.01 mole of hydrogen
ions per liter of solution. A solution of NaOH may have as little as 0.00000000001 mole
of hydrogen ion per liter of solution.
(2) To simplify the expression of such terms, chemists have transformed the
concentration values into numbers, called pH numbers, which are easier to utilize. This
is done according to the following equation:
pH = -log[H+]
The abbreviation log stands for logarithm. (For example, log 1 = 0, log 0.1 = -1, log 0.01
= -2, log 0.001 = -3, log 0.0001 = -4.) The expression [H+] here is the concentration of
hydrogen ions in moles per liter. If we consider the two previous examples, you can see
how this transformation aids us. The pH of the HCl solution would be -(-2.0) = 2.0; the
pH of the NaOH solution would be -(-11.0) = 11.0. These numbers, 2 and 11, are
certainly easier to work with than 0.01 and 0.00000000001.
b. pH Scale. This transformation results in a range of pH numbers from 0 to 14,
which is called the pH scale.
(1) The limits of the scale are related to the dissociation; how they are
arrived at is beyond our scope. Further information on this relationship can be found in
an inorganic chemistry textbook.
(2) While you will not need to calculate a pH value, you will need to interpret
what a pH value means at times. To learn this function, examine the following pH scale:
(3) A pH value less than 7.0 means the solution is acidic; the lower the
number, the more acidic. A solution with a pH of 2.0 is more acidic than one with a pH
of 4.0. Any pH value greater than 7.0 means the solution is basic with larger numbers
indicating solutions that are more basic. The only value on the scale that indicates a
neutral solution is 7.0. The pH values for some common pharmaceutical products are
given below.
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PRODUCT
pH
Cherry Syrup
Benylin® Expectorant
Glycyrrhiza® Syrup
3.5 - 4
5.0 - 5.5
6.0 - 6.5
Iso-Alcoholic Elixir
Orange Syrup
Terpin Hydrate Elixir with Codeine
5.0
2.5 - 3.0
8.0
c. Measurement of pH. There are three common methods for measuring pH,
which you may encounter in medicine.
(1) Litmus paper. Litmus paper is a paper coated with a dye, which is red in
an acid pH or blue in a basic pH. It will only indicate whether a solution is acidic or
basic; it will not give an actual pH value.
(2) pH paper. pH paper works on the same principle as litmus paper but
uses several different dyes. By comparing paper color with a chart, the pH of a solution
can be determined within one pH unit. If a closer measurement is needed, special
narrow-range papers can be used to determine the pH within 0.1-pH unit.
(3) pH meter. The most accurate tool for pH measurement is the pH meter.
This makes use of an electrode dipped into solution and is accurate to about 0.01-pH
unit, depending on the particular machine.
2-11. BUFFERS
Many drugs are stable in solution only at certain pHs or in narrow pH ranges. If a
solution of one these drugs is desired, the manufacturer must find a way to maintain this
certain pH over a period of time. This is accomplished by the use of buffer systems. A
buffer is a solution of a weak acid and the salt of that weak acid (weak bases could also
be used, but usually are not practical). The function of the buffer is to resist changes in
pH by reacting with any hydrogen or hydroxyl ions that are added to the solution. Two
of the most common buffer systems are:
a. Acetic Acid/Sodium Acetate. This is a common buffer found in many drug
solutions.
b. Carbonic Acid/Sodium Bicarbonate. This is the buffer system that is most
common in the fluids and tissues of the body and is used to keep the pH of the blood
and body fluids constant.
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2-12. WATER
Water is the most important liquid to all living organisms. It comprises about 57
percent of your body weight. It is the solvent for or is contained in most of the nutrients
your body requires for growth or maintenance. It is also the primary vehicle for almost
all liquid pharmaceutical preparations. Because of the inherent importance of water in
the practice of medicine, it is essential to acquire some knowledge of the properties of
water and aqueous (water-based) solutions.
a. Properties of Water. All of us are familiar with some properties of water.
We know that generally water is a bland-tasting, colorless liquid. Other specific
properties of water are of importance in medicine.
(1)
Its boiling point is 100ºC (212ºF).
(2)
Its freezing point is 0ºC (320F).
(3)
It is a polar solvent (dissolves ionic compounds).
(4)
Generally, it is chemically inert (unreactive) in biological or drug
systems.
b. Importance of Properties. The properties above are the specific reasons
that water is so valuable to living systems and to pharmaceutical preparations. The
wide difference between the freezing point (water as ice) and the boiling point (water as
steam or vapor) indicates that water will be a liquid at most of the temperatures
encountered under normal conditions. An example should help emphasize the
importance of these properties. If we wanted to prepare a liquid drug solution for a
patient who could not swallow capsules, we used a liquid vehicle with a freezing point of
25o C (77o F) and a boiling point of 30o C (86o F), we would be giving the patient a
worthless product. As the patient left home, the drug solution would boil if it were a
normal summer day (temperature = 86o F), and when the patient entered his airconditioned home, the remaining solution might become a solid which could not be
poured from the bottle. We also want our vehicle to be as unreactive as possible so
that only the drug is exerting a pharmacological effect.
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c. Structure of Water. The properties of water may best be explained by
examining the structure of the water molecule. The water molecule consists of two
hydrogen atoms bonded covalently to one oxygen atom. The three atoms are bound
together as shown below.
O
H
H
This arrangement leads to an electron-rich atom, oxygen, on one end and two electronpoor atoms, hydrogen, on the other end. This results in a molecule that resembles a
bar magnet in that it has a negative pole and a positive pole, as shown below.
_
O
H
+
H
+
Actually, there are not distinct electrical charges on the molecule, only partial charges,
referred to as δ+ and δ- (the Greek letter delta, δ, meaning partial). While these charges
are only partial, they are still strong enough for water to be referred to as a polar
molecule, meaning that it has a positive and negative end.
d. Hydrogen Bond. The polarity of the water molecule gives rise to an unused
type of bond between water molecules, the hydrogen bond. This bond is the electrical
attraction between the partially negative oxygen atom of one molecule and the partially
positive hydrogen atom of another molecule.
δO
H
H
δ+ - - - - δ-
H
H
O
O
δ+
H
H
δH
O
H
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The hydrogen bond is a very weak attraction about 1/10 to 1/20 the strength of the
hydrogen-oxygen covalent bond. The hydrogen bond explains why water has such a
high boiling point in relation to other compounds of similar molecular weight. For
example, methane (CH4, molecular weight = 16) boils at a temperature below 0ºC, while
water (molecular weight = 18) boils at 100ºC. Methane does not exhibit hydrogen
bonding.
e. Water Purification. We are all familiar with some of the ecological problems
facing the world today. Water is subject to mineral and biological contamination. Since
we will often be using water in the preparation of our products, we must be concerned
with its purity and the methods utilized for its purification. There are two common
methods of water purification used at Army medical treatment facilities--distillation and
ion exchange.
(1) Distillation. Distillation is the process of boiling water, collecting the
vapor, and then condensing the vapor back into water. Minerals and some of the
bacterial contamination will remain in the boiling vessel as a residue. Very pure water
may be prepared by repeating the distillation process several times. If sterile water is
desired, the water must be sterilized, because the process of distillation does not
necessarily sterilize water.
(2) Ion exchange (deionization). Less common than distillation because it is
less efficient, ion exchange involves passing water through a column containing a
charged resin. Ions in the water are held by electrical attraction and are thus removed
from the water.
2-13. SOLUTIONS
We are seldom concerned with just water in the hospital. We are generally more
concerned with substances dissolved in water. These are solutions. When we speak of
a solution, there are several terms, which are important to understand.
a. Solute. A solute is the substance, which is dissolved in a solution.
b. Solvent. The solvent is, the substance, which dissolves the solute. It is
usually water in pharmaceutical solutions, but not always.
c. Solubility. The maximum amount of a compound, which will dissolve in a
given amount of solvent at a given temperature is the solubility of that compound.
d. Dissociation. (Ionization). In general, two things can happen to a solute in a
solution. It can dissolve and exist in solution as molecules or it can dissociate and exist
entirely or partially as ions. The process of splitting a molecule into ions is known as
dissociation.
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e. Electrolyte. When a substance dissociates to a fair extent in water, it will
produce enough ions to support an electric current. We can use this property to
differentiate between substances that are molecular and substances that are ionic in
solution. An electrolyte is a substance, which dissociates sufficiently in solution to carry
an electric current. It is therefore ionic in nature (figure 2-1).
f. Non-Electrolyte. A substance that does not dissociate or carry an electric
current in solution is called a non-electrolyte.
g. Hydrolysis. Some compounds form ions in solution by reacting with water.
This reaction with water is called hydrolysis. Hydrolysis is the dissociation of a
compound through the splitting and incorporation of water. Hydrolysis occurs when
acidic, basic, or neutral (weak acid/weak base) salts are dissolved in water. Consider,
for example, the basic salt, sodium bicarbonate.
NaHCO3 + H2O ---> Na+ + HCO3 - + H2O ---> Na+ + OH- + H2CO3
The hydroxyl ion from water is associated with the sodium ion from the salt. The
hydrogen ion from water is associated with the bicarbonate radical, and these two exist
primarily as undissociated carbonic acid. The net result of this hydrolysis reaction is a
basic solution containing sodium and hydroxyl ions and undissociated carbonic acid.
Figure 2-1. Flow of electric current through electrolyte solution.
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h. Electrolyte Strength. It should be apparent, recalling the discussion of acids
and bases that not all electrolytes will dissociate to the same extent in solution.
(1) Those that dissociate and exist entirely as ions in solution are called
strong electrolytes. Strong electrolytes include strong acids, strong bases, and their
neutral salts.
(2) Compounds that dissociate to a small extent and exist only partially as
ions in solution are called weak electrolytes. Weak electrolytes include weak acids,
weak bases, and salts of weak acids and/or weak bases.
(3) To identify whether a compound is a strong or weak electrolyte, it is first
necessary to identify what type of compound it is. For example, consider NaCl. This
compound is a salt formed from a strong base (NaOH) and a strong acid (HCl) and is,
therefore, a neutral salt. Since it is a neutral salt of a strong acid and a strong base, it is
a strong electrolyte as defined above. Other salts can be determined in a like manner.
Continue with Exercises
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EXERCISES, LESSON 2
INSTRUCTIONS. Write the word, words, symbols, or numbers that properly completes
the statement in the space provided or mark the correct word/phrase from those given.
After you complete the exercises, turn to Solutions to Exercises and check your
answers. Reread the material referenced for each exercise answered incorrectly.
1.
We are given the following chemical equation:
NaOH + HCl ---> NaCl + H2O
It means that the base s ________ h_________ reacts with the acid
h___________ a__________ to yield the salt s ____________ c____________
and the compound ___________ . To form one molecule of sodium chloride and
one molecule of water, we need (one) (two) molecule(s) of sodium hydroxide and
(one) (two) molecule(s) of hydrochloric acid.
2.
We are given the following chemical reaction:
AgNO3 + KCl ---> KNO3 + AgCl ↓
This means that s___________ n__________ and p____________
c__________ react to yield p____________ n__________ and s____________
c____________ . The arrow next to AgCl means that AgCl p_______________ s.
3. Calcium hydroxide and nitric acid react to yield calcium nitrate and water. The
formula for calcium hydroxide is ______. The formula for nitric acid is ______.
The formula for calcium nitrate is ________. The formula for water is ________.
Before balancing, the equation for this reaction is _________________________,
In the columns below are listed atoms and radicals of the reactants and products.
Indicate the number in the formula of each compound. Consider water to be HOH.
REACTANTS
MD0803
PRODUCTS
_____ Ca
_____ Ca
_____ OH
_____ OH
_____ H
_____ H
_____NO3
_____NO3
2-27
Thus, on the left side of the equation, we need twice as much
_________________. On the right side of the equation, we need twice as much
_____________. In order to satisfy this requirement, the equation becomes
__________________.
On each side of the equation, we now have (one) (two) calcium atoms(s), (one)
(two) hydroxide radical(s), (one) (two) hydrogen atoms, and (one) (two) (three)
nitrate radicals. Since we now have an equal number of each type of atom and
radical on both sides of the equation, we can say that the equation is
__________________.
4.
Sulfuric acid and ferric hydroxide react to produce ferric sulfate and water. The
equation without balancing is:
_________+________ ------> _________+________
Now, list the reactants and the products as in the previous exercise.
REACTANTS
PRODUCTS
_______ H
_______ H
_______ SO4
_______SO4
_______ Fe
_______ Fe
_______ OH
_______ OH
Now, balance the equation. (Before looking at the answer, be sure you have the
same number of each type of atom or radical on both sides of the equation.)
__H2SO4 + __ Fe (OH)3 ----> __Fe2(SO4)3 + __ HOH
5.
Iron metal and sulfur react to yield ferrous sulfide. The balanced equation for this
reaction is ______________________________.
6. Sodium carbonate and hydrochloric acid react to yield sodium chloride, water, and
carbon dioxide gas. Use a piece of scratch paper to write the balanced equation.
The balanced equation is ________________________.
Did you forget the arrow after CO2? The arrow means that CO2 is given off as a
_______________.
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2-28
7.
An exothermic reaction (gives off) (takes in) heat.
8.
An endothermic reaction ______________heat.
9. A reaction is in equilibrium if the reactants are being formed from the
___________ at the same rate as the products are being formed from the
_______________. The symbol used to indicate that a reaction may go in both
directions is ________________________.
10.
The formula weight of a compound is also known as its ________________
weight. If the molecular weight is expressed in milligrams, it is known as the
_____________ molecular weight. If it is expressed in grams, it is known as the
_______________ molecular weight. The formula weight of carbon dioxide is
12 + 16 + 16 = 44. The milligram formula weight of carbon dioxide is ______.
The gram formula weight of carbon dioxide is __________.
11. The milligram molecular weight of a compound is the sum of the a________
w_______s of all the atoms that appear in its chemical formula, with the weights
expressed in ________________.
12. Using Table 1-1 in this subcourse and rounding the atomic weights to the nearest
whole number, what is the milligram molecular weight of sodium bicarbonate,
NaHCO3?
Atoms:
Na + H + C + O + O + O
Atomic weights: __+ __+ __+ __+
+ __
Adding the atomic weights, we get a formula weight of __________.
Thus, the milligram molecular weight is _____________.
The gram milecular weight is _____________________.
13. A mole of NaHCO3 weighs _______ grams. A liter of a 1M solution of NaHCO3
contains ________grams of NaHCO3.
14. The milligram equivalent weight of a compound is its milligram molecular weight
divided by the total ____________ or _______________v ___________.
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2-29
15.
The bicarbonate radical has a total negative valence of______. Sodium
bicarbonate has a total positive valence of ________. Since the milligram
molecular weight of NaHCO3 is 84 mg, its milligram equivalent weight is ______
______.
16. Sulfuric acid, H2SO4, has a milligram molecular weight of _________ ________ .
Its total positive valence is ___________. Its milligram equivalent weight is
______ _______.
17. If an element is oxidized, it (gains) (loses) electrons and its valence (increases)
(decreases). If an element is reduced, it (gains) (loses) electrons and its valence
(increases) (decreases). If an element is oxidized, there is an increase in its
_______________. If it is reduced, there is a decrease in its ____________.
When elemental iron reacts with diatomic oxygen, the valence of Fe goes
from______________ to _____________. The valence of O goes from
___________ to _____________. Iron (gains) (loses) electrons. Oxygen (gains)
(loses) electrons. Therefore, iron is (oxidized) (reduced) and oxygen is (oxidized)
(reduced). The element that gains electrons and loses valence is said to be
___________________. The element that loses electrons and gains valence is
said to be _________________________.
18. Oxidation may be defined as a ________________ of electrons or a
______________of valence. Reduction may be defined as a _____________ of
electrons or a _____________ of valence.
19. In an oxidation-reduction reaction, the oxidizing agent is (oxidized) (reduced) and
the reducing agent is ____________________.
20. When elemental magnesium Mg reacts with diatomic iodine, we have the following
reaction:
Mg + I2 ---> MgI2
In this reaction, the valence of magnesium goes from _______ to ______. The
valence of iodine (I) goes from _______ to ____________. Since the valence of
magnesium increases, magnesium is said to be ______________. Since the
valence of iodine decreases, iodine is said to be __________. The reducing agent
in this reaction is _______. The oxidizing agent is _____________
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2-30
21.
According to the classical theory, an acid is a compound that donates
___________________ and a base is a compound which donates__________
____________. The symbol for a proton is _________. The symbol for a
hydroxyl (hydroxide) ion is _____. According to the classical theory, an acid
donates ___________________ ions and a base donates ___________________
ions.
22. According to the Bronsted-Lowry theory, an acid is a compound which
donates_______________ and a base is a compound which
__________________s protons.
23.
According to the Bronsted-Lowry theory, which substance acts as an acid in the
following reaction?
H2CO3 + OH - ---> HCO3 - + H2O
Which substance acts as a base in the following reaction?
H+ + HCO3 - ---> H2CO3
24.
Below is a list of nine compounds. For each, indicate whether it is best described
as an acid, a base, or salt.
a. HCl ____________
b. NaOH
____________
c.
___________
HNO3
d. Ca(OH)2
e. KOH
____________
f.
____________
H3PO4
g. Fe(OH)3
_____________
h. MgO
___________
i.
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NaCl
2-31
25.
26.
Below are listed five properties of acids with some words missing. Complete the
list.
a.
Acids change the color of litmus paper from __________ to ___________.
b.
Acids taste _______________.
c.
Acids react with metals to release ________gas.
d.
Acids react with carbonates and bicarbonates to form ____________
___________ gas.
e.
Acids react with bases to form __________ and water.
Below are listed four properties of bases some words missing. Complete the list.
a.
Bases change the color of litmus paper from ___________ to ___________.
b.
Bases taste _________________________.
c.
Bases feel ____________________.
d.
Bases react with acids to form ______________ and ______________.
27.
If a person spills a strong acid (except phenol) or a strong base on his skin, he
should wash the acid or base off with large amounts of ___________________ .
28.
If a person swallows a strong acid, you should give him an antacid such as
____________ of ____________ or _____________________
__________________. The antacid should NOT be a __________________ or
______________________. Vomiting (should) (should not) be induced.
29. If a person swallows a large amount of a strong base, such as sodium hydroxide,
you should give him a _________ such as v__________ or f ______________
j__________. Vomiting (should) (should note) be induced.
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2-32
30.
A salt is a compound formed by the replacement of _______________ions in
an___________ with a __________ or a radical acting like a metal. For example,
the salt NaCl is formed by the replacement of the hydrogen in______ by the metal
________. The same salt NaCl results when NaOH reacts with _______. In
general, we can say that an ionic compound is a salt if it contains a positive ion
other than _________________ and a negative ion other than
____________________ or __________________ .
31.
For each of the eight compounds listed below; indicate whether it is an acid, a
base, or a salt.
a.
NaNO3 _______________
b.
Al(HCO3) 3 _________________
c.
K2CO3 _______________________
d.
FePO4 ______________________
e.
Ca(OH) 2 ______________________
f.
Na(C2H3O2) ____________________
g.
NH4Cl _______________________
h.
H2SO4 ___________________
32. If the pH of a solution is 6.9, it is slightly ____________. If the pH is 7.5, it is
slightly _____________. If the pH is 8.8, the solution is __________ .
33. According to the subcourse, the devices used to measure pH are
____________________ paper, and a ___________ ____________
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2-33
34. A mixture used to keep the pH of a solution nearly constant is called a
__________________ system. Such a system usually contains a weak
__________and a ___________of that weak _______________. The function of
a buffer system is to keep the __________ at a nearly _____________ level.
Which of the following represents a buffer system?
HCl/NaHCO3
HBr/KBr
NaOH/NaCl
H2CO3/NaHCO3
35. The boiling point of water is ____________. The freezing point of water
is___________. Because water dissolves ______________ compounds, it is
called a __________ solvent. In biological and drug systems, water is generally
(inert) (reactive).
36. Two major methods of water purification for medical use are di
______________________ and I ___________ e _____________ (de
______________).
37. A solute is a substance that is _______________ in a solution. A solvent is the
substance in which the ____________ is dissolved. The solubility of a compound
is the maximum amount that will ____________ in a given amount of
________________ at a given ____________. The substance in which the solute
is dissolved is called the _________________. The amount of a substance that
will dissolve under specified conditions is called its _________________.
38. The dissociation of a molecule means that it is ______________ into ions. When
a compound is dissolved and its molecules split into ions, we call this process
______________. When an electrolyte is dissolved, many of its molecules
_______________ and form _________. If a compound forms enough ions in
solution to make the solution capable of carrying an electric current, then we say
that compound is an ____________________ .
Check Your Answers on Next Page
MD0803
2-34
SOLUTIONS TO EXERCISES, LESSON 2
1.
sodium hydroxide; hydrochloric acid; sodium chloride; water
one, one (para 2-1)
2.
silver nitrate; potassium chloride; potassium nitrate; silver chloride
precipitates (para 2-1)
3.
Ca(OH)2
HNO3
Ca(NO3)2; H2O
Ca(OH)2 + HNO3
Ca(NO3)2 + H2O
1
2
1
1
---->
Ca; 1 Ca
OH; 1 OH
H; 1 H
NO3; 2 NO3
NO3
OH
Ca(OH)2 + 2HNO3 ---->Ca(NO3)2 + 2 HOH
one, two; two, two
balanced (paras 2-2, 2-3)
4.
H2SO4 + Fe(OH)3 --->Fe2(SO4)3 + HOH
2
1
1
3
H; 1 H
SO4; 3 SO4
Fe; 2 Fe
OH; 1 OH
3H2SO4
+
2Fe(OH)3 ---->Fe2(SO4)3 + 6HOH (paras 2-2, 2-3)
5.
Fe + S ----> FeS (paras 2-2, 2-3)
6.
Na2CO3 + HCl ----> 2 NaCl + H2O + CO2 ↑
gas (paras 2-2, 2-3)
7.
gives off (para 2-5a(1))
8.
takes in (para 2-5a(2))
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2-35
9.
products, reactants,
------>
<------- )
(para 2-4)
10.
molecular
milligram
gram
44 mg
44 grams (para 2-6)
11.
atomic weights, milligrams (para 2-6a)
12.
23 + 1 + 12 + 16 + 16 + 16
84
84 mg
84 grams (para 2-6, table 1-1)
13.
84
84 (para 2-6b)
14.
positive, negative valence (para 2-6b)
15.
-1
+1
84 mg (para 2-6c)
16.
98 mg
+2
49 mg (para 2-6, Table 1-1)
17.
loses, increases
gains; decreases
valence
valence
0; +2
0; -2
loses
gains
oxidized, reduced
reduced
oxidized (para 2-7)
18.
loss, gain
gain, loss (para 2-7b, c)
MD0803
2-36
19.
reduced; oxidized (para 2-7d)
20.
0; +2
0; -1
oxidized
reduced
Mg
I2 (para 2-7)
21.
protons
hydroxyl ions
H+
OHH+ (hydrogen), OH- (hydroxyl) (para 2-8a)
22.
protons accepts (para 2-8b)
23.
H2CO3
HCO3 - (para 2-8)
24.
a acid
b base
c acid
d base
e base
f acid
g base
h base (MgO + H
2O ----> MgOH+ + (OH -)
i salt (para 2-8a, b)
25. a
b
c
d
e
blue to red
sour
hydrogen
carbon dioxide
salts (para 2-8c)
26. a
b
c
d
red to blue
bitter
soapy
salts, water (para 2-8d)
27. water (para 2-8d)
MD0803
2-37
28.
milk of magnesia; magnesium oxide
carbonate, bicarbonate;
should not (para 2-8g(1)(b))
29.
weak acid, vinegar; fruit juice
should not (para 2-8g(2))
30.
hydrogen; acid; metal
HCl, Na
HCl
hydrogen (H +); hydroxyl (OH -); oxide (O-2) (para 2-9a)
31. a
b
c
d
e
f
g
h
salt
salt
salt
salt
base
salt
salt
acid (para 2-9a)
32.
acidic
basic
basic (para 2-10b)
33.
litmus; pH; pH meter (para 2-10c)
34.
buffer
acid, salt; acid
pH; constant
H 2CO3/NaHCO3 (HCl and HBr are strong acids; NaOH is a strong
base) (para 2-11)
35.
100 º C (212 ºF)
0 º C (32 ºF)
ionic, polar
inert (para 2-12e)
36.
distillation; ion exchange (deionization) (para 2-12a)
37.
dissolved
solute
dissolve; solvent temperature
solvent
solubility (para 2-13c-c)
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2-38
38.
split
dissociation
dissociate; ions
electrolyte (para 2-13d, e)
End of Lessson 2
MD0803
2-39
LESSON ASSIGNMENT
LESSON 3
Elements of Organic Chemistry.
LESSON ASSIGNMENT
Paragraphs 3-1 through 3-18 and exercises.
LESSON OBJECTIVES
After completing this lesson, you should be able to:
SUGGESTION
MD0803
3-1.
State the type of bond most prevalent in
inorganic chemistry and the type of bond most
prevalent in organic chemistry.
3-2.
State the reason for the importance of
structural formulas for organic compounds.
3-3.
List the three types of carbon-carbon bonds
and state the class of organic compounds
representative of each type.
3-4.
Define the terms hydrocarbon, aliphatic,
saturated, unsaturated, and aromatic.
3-5.
Given a list of structural formulas for organic
compounds, match each to the name of the
class of compounds to which it belongs.
3-6.
Given a list of structural formulas for organic
compounds, select those whose groups are
acidic or basic.
3-7.
Given a list of classes of organic compounds,
select those that have the ability to form
hydrogen bonds between themselves.
3-8.
Given a list of classes of organic compounds,
select the special reactions which each will
undergo, to include oxidation, reduction, salt
formation, esterification, amide formation, and
hydrolysis.
3-9.
Given a list of definitions for chemical terms,
select the definitions for oxidation and
reduction as they apply to organic chemistry.
After completing the assignment, complete the
exercises at the end of this lesson. These exercises
will help you to achieve the lesson objectives.
3-1
LESSON 3
ELEMENTS OF ORGANIC CHEMISTRY
3-1.
INTRODUCTION
Carbon is one of the most abundant elements in our world. It is part of the
molecular structure of all living organisms. It is the basis for our fuel and energy
production and it plays a large role in the chemistry of many of the synthetic fabrics and
plastics that have become so important to our lifestyle. Carbon compounds also
account for a vast majority of today's drugs. It is very important that you, as a health
care provider, have a basic understanding of the chemistry of carbon compounds and
organic chemistry due to the roles that carbon plays.
3-2.
CONTRAST WITH INORGANIC CHEMISTRY
There are several general differences between the chemistries of carbon
compounds and inorganic compounds, which will help give you an overall view of
organic chemistry (Table 3-1).
NORGANIC
CHEMISTRY
ORGANIC
CHEMISTRY
TYPE OF BONDING
Ionic
Covalent
MOLECULAR SIZE
Small
Large
WATER SOLUBILITY
Soluble
Insoluble
SOLUBILITY IN
ORGANIC SOLVENTS
Insoluble
Soluble
CLASSES OF COMPOUNDS
Acid, base, or salt
groups)
Many (functional
STRUCTURAL FORMULAS
Unimportant
Very important
Table 3-1. Comparison of organic and inorganic chemistry.
MD0803
3-2
3-3.
STRUCTURAL FORMULAS
A structural formula is a chemical formula that shows how atoms are bonded to
each other. For example, we might write AlOHCl2 as
Cl1
OH
Al
Cl1
to show the bonds. However, in inorganic chemistry, the compounds are such that
there is generally only one possible way to combine the atoms. This is not the case in
organic chemistry, where very often there are many possible combinations for the
atoms in the compound. Consider, for example, the formula C4H10. This formula could
represent either of the following compounds.
CH3-CH-CH3
or
CH3-CH2-CH2-CH3
CH3
These compounds have slightly different properties. As the formulas become more
complex, the differences are even greater. For this reason, it is often better to use a
structural formula in organic chemistry rather than the simple chemical formula.
3-4.
CARBON
Before we examine carbon compounds, we first need to examine the structure
and mention some properties of the carbon atom. Carbon has an atomic number of six,
meaning it has six protons, and consequently has six electrons. These electrons are
distributed with two in the K shell and four in the L shell. In forming compounds, carbon
would appear to gain or lose the four electrons in its outer shell. Thus, we have the +4,
-4 valences you learned for carbon earlier in this subcourse.
In fact, carbon does not usually exchange electrons with other elements but prefers to
share four electrons to complete its L shell. This is the reason that covalent bonding is
predominant in organic chemistry.
MD0803
3-3
3-5.
CARBON-CARBON BONDING
Carbon atoms have the unique ability to bond to other carbon atoms and form
chains which may also have branches.
C-C-C-C-C-C-C
or
C-C-C-C-C-C-C-C-C-C-C
I
I
C
C
I
I
C-C
C
This is the reason that the molecular size is so great in organic chemistry. Molecular
weights in the thousands are not uncommon. Three types of bonds are formed
between carbon atoms.
a. Single Bonds. A single bond is a covalent bond formed by two carbon
atoms sharing two electrons. Compounds that contain only single bonds between
carbon atoms are called alkanes.
b. Double Bonds. A double bond consists of two covalent bonds formed by two
carbon atoms sharing four electrons as shown below.
Compounds that contain at least one carbon-carbon double bond are referred to as
alkenes.
c. Triple Bonds. A triple bond consists of three covalent bonds formed by two
carbon atoms sharing six electrons as shown below.
Compounds that contain at least one triple bond between carbon atoms are called
alkynes.
3-6.
HYDROCARBONS
The simplest organic compounds are the hydrocarbons, which are composed
solely of carbon and hydrogen. Since there are only two elements involved, one might
expect there would be only a few different compounds. However, carbon does bond to
itself and form long chains. So there are many, many different hydrocarbons. They can
be classed in two general groups, aliphatic and aromatic. These compounds are the
starting point for all organic compounds.
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3-4
a. Aliphatic Hydrocarbons. Aliphatic hydrocarbons consist of straight or
branched chains of carbon atoms with the other valence electrons involved in bonds
with hydrogen. Examples are:
CH3-CH2-CH2-CH2-CH3
CH2 = CH2
CH3
CH2 -CH3
CH--CH
CH3
CH3
We can subdivide aliphatic hydrocarbons into two groups based on the types of carboncarbon bonds the compounds contain.
(1) Saturated aliphatic hydrocarbons. Saturated aliphatic hydrocarbons are
hydrocarbons in which all of the carbon-carbon bonds are single bonds. These
compounds are also referred to as alkanes, as mentioned for single bonds earlier. We
often refer to the alkanes as the methane series. Methane is the simplest hydrocarbon
with the formula CH4. All other alkanes are formed by adding CH2s to the formula (table
3-2). In this series, the names from C5 to C10 all begin with the Greek prefix for the
number (e.g., penta- for five) and end in -ane from “alkane.” The two low-molecularweight alkanes are gases. Alkanes are not very reactive chemically and are insoluble in
water. About the most important reaction they undergo is that they burn to form carbon
dioxide and water (combustion reaction). Some typical saturated compounds you might
encounter are:
FORMULA
BOILING POINT (oC)
Methane
CH4
-161.5
Ethane
C2H6
- 88.3
Propane
C3H8
- 44.5
Butane
C4HlO
- .5
Pentane
C5H12
+ 36.2
Hexane
C6H14
Heptane
C7H16
Octane
C8H18
Nonane
C9H20
Decane
C10OH22
NAME
+125.8
+174.0
Table 3-2. Common alkanes.
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3-5
(a)
Liquid petrolatum (mineral oil). This liquid is used as a solvent and
(b)
Petrolatum (petroleum jelly). This semisolid is used as an ointment
(c)
Paraffin (wax). This solid is used in pharmacy as a stiffening agent.
as a laxative.
base.
(2) Unsaturated aliphatic hydrocarbons. The second type of aliphatic
hydrocarbon is unsaturated hydrocarbons. These are hydrocarbons, which contain at
least one double or triple bond (that is, they are alkenes or alkynes). An example of an
alkene is ethene, the simplest alkene, which consists of two double-bonded carbon
atoms and four hydrogen atoms.
CH2 = CH2
Note that the name is similar to the saturated compound ethane. The -ene ending
comes from the word alkene and denotes that it contains a double bond. Similarly, if
there were a triple bond between the two carbon atoms, the name would be ethyne with
the -yne ending denoting the triple bond (from alkyne). The physical properties of
alkenes and alkynes are similar to the properties of alkanes of similar molecular
weights. Chemically, the word unsaturated implies that these compounds can form
additional bonds. This is the case, for alkenes and alkynes are much more reactive and
undergo many reactions not possible with alkanes.
b. Aromatic Hydrocarbons. The second major group of the hydrocarbons is
the aromatic hydrocarbons, which are hydrocarbons that contain a benzene ring as part
of their structure. Benzene has the formula C6H6 and consists of six carbon atoms in a
ring with three alternating double bonds.
H
H
C
C
H-C
C-H
C
C
H
H
The benzene ring is also represented with the following symbols:
or
MD0803
or
3-6
Ø
(1) Benzene is completely insoluble in water, it is a volatile liquid at room
temperature, and it is fairly unreactive. The properties of other aromatics are reflective
of benzene but vary according to the substituents added to the ring in place of one of
the hydrogen atoms.
(2) The term “aromatic” has its origin in the fact that certain aromatic
substances (for example: oil of bitter almonds, vanilla, and oil of wintergreen) contain
the benzene ring. The possession of an odor is not characteristic, however, of all
aromatic substances.
(3) Aromatic hydrocarbons are the starting point for many medicinally
important compounds, as the following examples indicate.
Morphine
Narcotic Analgesic–For
relief of severe pain.
Non-Narcotic Analgesic–
For relief of moderate pain.
Aspirin
Epinephrine
Adrenergic Drug–Used to
treat shock.
Stimulant–For diet control,
narcolepsy, and
hyperkinesis.
Dextroamphetamine
MD0803
3-7
(4) You will notice in the compounds above that they are not pure aromatic
hydrocarbons because they contain elements other than carbon and hydrogen. These
additional elements are the basis for the classification of substituted organic compounds
and are called functional groups. The important functional groups will be considered in
the following paragraphs.
3-7.
INTRODUCTION TO FUNCTIONAL GROUPS
There are millions of organic compounds known to exist, and most are more
complex than the simple hydrocarbons we have discussed. To facilitate the study of
their reactions and properties, they are conveniently classed according to the functional
groups they contain. A functional group is a group of atoms or a single atom that is
substituted for a hydrogen or a hydrocarbon. These groups generally determine the
types of reactions and properties of these more complex compounds. (A summary of
the properties and reactions of the compounds contained in paragraphs 3-8 through
3-16 is tabulated in Table 3-3, pages 3-14 and 3-15.)
3-8.
ALCOHOLS
Alcohols are hydroxyl (-OH) derivatives of hydrocarbons formed by replacing a
hydrogen with the hydroxyl radical and are of the general form R-OH where R
represents the hydrocarbon. There are three classes of alcohols: primary, secondary,
and tertiary. A primary alcohol is one in which the hydroxyl group is attached to a
carbon atom which, in turn, is attached to not more than one other carbon atom. A
secondary alcohol is one in which the hydroxyl group is attached to a carbon atom
which in turn, is connected to two carbon atoms. A tertiary alcohol is one in which the
hydroxyl group is attached to a carbon atom which in turn, is attached to three other
carbon atoms.
CH3-CH2-OH
Primary Alcohol
Ethanol
(ethyl alcohol)
CH3-CH-CH3
CH3
OH
CH3—C—CH3
Secondary Alcohol
2-propanol
(isopropyl alcohol)
OH
Tertiary Alcohol
Alcohols that contain two or more hydroxyl groups are referred to as polyhydroxy
alcohols. An example you will encounter frequently in this course and on the job is
glycerin, which is:
CH2—CH—CH2
OH
MD0803
OH OH
3-8
a. Properties of Alcohols. The low-molecular-eight alcohols are volatile
liquids, and the high-molecular-weight alcohols (more than 13 carbons) are solids. The
first three alcohols (C1 to C3) are completely miscible (mix in any proportion) with water.
The water solubility decreases as the number of carbons increases, and the largemolecular-eight alcohols are insoluble in water. Alcohols have higher boiling points and
melting points than alkanes with the same or similar molecular weights (MW). For
example:
MW
MELTING
POINT
BOILING
POINT
CH3-CH2-CH2-CH2-OH
74
-9O0C
1180C
CH3-CH2-CH2-CH2-CH3
72
-1300C
360C
The water solubility and the high melting and boiling points of alcohols result from their
ability to form hydrogen bonds with water and to form hydrogen bonds intermolecularly
(between themselves).
b. Reactions of Alcohols. Chemically, the alcohols can be considered to be
neutral (in terms of acids and bases) even though they can act as very weak acids or
bases as water does. They undergo several kinds of chemical reactions, the most
important of which is oxidation. Oxidation in organic chemistry is defined as the
elimination of hydrogen from or the addition of oxygen to a compound.
(1)
The oxidation of a primary alcohol can be expressed by the following
example:
KMnO4
CH3-OH -------->
H
NOTE:
O
ll
HCH
O2
O
---------> H—C—OH
∆
O
C=O and -C-OH are two more functional groups, indicating aldehydes
and carboxylic acids, respectively.
The first step in this oxidation is the removal of two hydrogen atoms from the alcohol to
form an aldehyde, and the second step is the addition of one oxygen atom to the
aldehyde to form a carboxylic acid.
(2) Secondary alcohols undergo only the first step. For example, a threecarbon alcohol is oxidized to form CH3 –C –CH3 , which is an example of a new class of
compounds called ketones.
ll
O
(3)
MD0803
Tertiary alcohols are not oxidized.
3-9
(4) One reason the oxidation reaction is important is that it is the means the
body uses to eliminate the popular liquid, ethyl alcohol, or ethanol (CH3-CH2-OH).
c. Uses of Alcohols. Alcohols are most commonly used as solvents in the
pharmacy. They are also used as disinfectants and antiseptics.
3-9.
PHENOLS
Phenols are hydroxyl derivatives of hydrocarbons formed by replacing a
hydrogen on the benzene ring of an aromatic hydrocarbon with the hydroxyl radical.
Phenols have the general formula Ar-OH, where Ar represents a substituted or
nonsubstituted aromatic hydrocarbon. Thus, phenols are really just a special class of
alcohols. However, they have enough unique properties that they deserve to be
considered as a separate class of compounds. Below are some examples of typical
phenols.
a. Properties of Phenols. All phenols are white solids with moderately high
melting points and are soluble in water. They also have the property of being able to
form eutectics with camphor, menthol, or thymol, which are solid alcohols. (A eutectic is
a uniform mixture formed from two compounds that melt at a temperature lower than the
melting point of either of the two compounds.) Thus, phenol (a solid) and camphor (a
solid) form a liquid mixture at room temperature which is called a eutectic.
b. Reactions of Phenols. Chemically, phenols are weakly acidic compounds.
The hydrogen dissociates to a small degree from the hydroxyl radical to act as an acid
as shown below.
Since phenols are weak acids, they will form salts with inorganic bases. Phenols with
two hydroxyl groups also undergo oxidation reactions.
c. Uses of Phenols. Medicinally, phenolic compounds have three uses: as
keratolytics (compounds that remove hornified or scaling outer layers of skin),
antipruritics (relieve itching), and disinfectants. These uses arise from the fact that
phenols are very caustic to animal tissues. Precautions must therefore be taken when
you are using phenols in preparations. These properties, possessed to different
degrees by various phenols, depend on which other functional groups are present and
the number of hydroxyl groups.
MD0803
3-10
3-10. ETHERS
An ether can be thought of as a hydrocarbon derivative of water where the two
hydrogens of water are replaced by hydrocarbon groups. Thus, ethers have the general
structural formula R-O-R' where R and R' represent any two hydrocarbons, which may
be alike or different. Some examples of ethers are:
CH3-CH2-O-CH2-CH3
Ø-O-CH3
CH2=CH-O-CH3
a. Properties of Ethers. Ether molecules are slightly polar, but cannot form
hydrogen bonds with each other since they do not have a hydrogen atom attached
directly to an oxygen atom. Therefore, they have about the same boiling points and
melting points as alkanes of similar molecular weights.
M.W.
Boiling Point
CH3-CH2-CH2-CH2-CH2-CH2-CH3
100
980C
CH3-O-CH2-CH2-CH2-CH2-CH3
102
1000C
b. Reactions of Ethers. Since ether molecules are slightly polar and have an
oxygen atom in their structure, they can form hydrogen bonds with water. This property
accounts for the fact that ethers are slightly soluble in water. Chemically, ethers are
inert except for the oxidation reaction. Ethers are oxidized in the presence of oxygen to
form peroxides, which are explosive when concentrated.
c. Uses of Ethers. Medicinally, ethers are used as general anesthetics. They
are also used as solvents. Many of you are involved with ordering and storing ethers.
3-11. AMINES
Amines result from the replacement of one or more of the hydrogen atoms of
ammonia with hydrocarbons and have the general formula R-NH2. There are four
classifications of amines: primary, secondary, tertiary, and quaternary. Primary amines
result from replacing one of the hydrogens of ammonia by a hydrocarbon, as in CH3NH2; secondary amines result from the replacement of two hydrogens of ammonia by
two hydrocarbons, as in CH3-NH-CH3; and tertiary amines result from the replacement
of all three hydrogens of ammonia by hydrocarbon groups. The fourth classification of
amines is sometimes encountered in drug structures. This classification is the
quaternary amine that is formed by replacing the four hydrogens of the ammonium ion
(NH4 +) by hydrocarbon groups. Whenever one of the hydrocarbon groups connected
to the nitrogen atom contains a benzene ring, the compound is referred to as an
aromatic hydrocarbon.
a. Properties of Amines. The low-molecular-weight amines are all volatile
liquids, and those having up to five carbons are soluble in water. The element nitrogen
is in the same period of the periodic table as oxygen and has some similar properties--
MD0803
3-11
the most significant being the ability to form hydrogen bonds. The formation of
hydrogen bonds between amines, and between amines and water, accounts for their
higher boiling points (than alkanes) and their water solubility.
b. Reactions of Amines. Since amines are derivatives of ammonia, they are
bases as defined by the Bronsted-Lowry theory. The nitrogen of the amine can accept
a proton to form a substituted ammonium ion.
CH3 –CH2 –NH2 + H+ ---> CH3 –CH2 –NH3 +
Amines will thus react with inorganic acids to form salts. (Amines react with organic
acids to form amides, a class of organic compounds discussed later in this subcourse.)
CH3 –NH2 + HCl ---> CH3 –NH3 +ClThe reaction in the example above results in a hydrochloride salt of the amine and is a
very important reaction in pharmacy. Many drugs contain an amine functional group,
and if they contain many carbon atoms, they are not very soluble in water. The salts
formed from amines, however, are very soluble in water. Therefore, if we wish to use a
water solution of an amine drug that is insoluble, we can make it soluble by forming the
salt of the amine.
c. Use of Amines. As already stated, the amine functional group is contained
in many different drugs that have quite different actions in the body. Generally, these
drugs are very complex and you would never be expected to draw or know the structure
for these drugs. You should, however, recognize the -NH2 group of an amine and be
cognizant of its basic properties.
3-12. CARBOXYLIC ACIDS
Carboxylic acids are formed by the two-step oxidation of alcohols as stated
previously and have the general structural formula
O
or R –COOH.
Some examples of carboxylic acids are:
II
R–C –OH
O
II
CH3 –C –OH
Ethanoic Acid
(Acetic Acid)
O
O –C –OH
Benozoic Acid
O
II
CH3 –CH –C –OH
CH3
2-Methylpropionic Acid
MD0803
3-12
a. Properties of Carboxylic Acids. Carboxylic acids are very polar
compounds due to the two oxygen atoms and can form two hydrogen bonds between
themselves as shown below.
They have the highest melting points of any of the classes of compounds in table 3-3; a
carboxylic acid has a higher melting point than a different type of organic compound
with a similar molecular weight. Consequently, they are all solids under normal
conditions. The compounds with four carbons or less are miscible with water; those
with five carbons are slightly soluble, and those with more than five carbons are
generally insoluble in water.
b. Reactions of Carboxylic Acids. As their name implies, carboxylic acids are
the most acidic of all organic compounds but are still weak acids when compared to
inorganic acids.
Carboxylic acids will form salts with inorganic bases, and as with the basic amines, this
property is often used to make insoluble organic acids soluble in water as their salt.
This pair, ethanoic acid (acetic acid) and its salt sodium ethanoate (sodium acetate), is
used as a buffer system.
Carboxylic acids undergo three other important chemical reactions: reduction, ester
formation, and amide formation.
(1) Reduction in organic chemistry is the opposite of oxidation and is the
addition of hydrogen to or the elimination of oxygen from a compound. In the case of
carboxylic acids, the removal of oxygen first results in an aldehyde, which may be
reduced further by the addition of hydrogen to form an alcohol.
(2) Ester formation, as illustrated by the reaction below, is the reaction of a
carboxylic acid with an alcohol to yield a new class of compound called an ester.
O
ll
CH3 –COOH + CH3 –CH2 –OH ------> CH3 –C –O –CH2 –CH3 + H2O
Acid
Alcohol
Ester
MD0803
3-13
CLASS
OF
COMPOUNDS
ALCOHOL
PHENOL
ETHER
AMINE
CARBOXYLIC
ACID
ALDEHYDE
KETONE
ESTER
AMIDE
O
║
R-C-H
O
║
R-C-R
O
O
R-C-O-R
R-C-NH2
GENERAL
STRUCTURE
PRODUCT
OR
COMPOUND
R-OH
Ar-OH
R-O-R
R-NH2
O
║
R-C-OH
ROH
ArOH
ROR
RNH2
RCOOH
RCOH
RCOR
RCOOR
RCONH2
O
║
CH3COH
O
║
CH3CH
0
║
CH3CCH3
0
║
CH3COCH3
O
║
CH3CNH2
Ethanic
Acid
Ethanol
Methyl
EthanoatE
Ethanamide
(Acetic
Acid)
(Acetaldehyde)
2Propanone
(Methyl
Ketone)
-
-
Ø-OH
C2H5OC2H5
CH3CH2NH2
(Ethyl
Alcohol)
Phenol
Ethyl
Ether
Ethylamin
e
OTHER
COMMON
NAMES
Grain
Alcohol
Carbolic
Acid
Diethyl
Ether
-
pH
Neutral
Slightly
Acidic
Neutral
Basic
Acidic
Yes
Yes
No
Yes
Higher
Higher
Same
NAME IN
COMMON
SYSTEM
HYDROGEN
BONDING
BETWEEN
THEMSELVES
COMPARISON
OF BOILING
POINT TO
CORRESPONDING
ALKANE
OXIDIZED
TO
CH3CH2O
H
Ethanol
(Acetamide)
Dimethyl
Detone
Dimethyl
Ester
-
Neutral
Neutral
Neutral
Neutral
Yes
2–H
Bonds
No
No
No
Yes
Highest
Same
Same
Same
Same
High
Acid
Acid
(Very
Difficult)
-
-
2 Alcohol
Alcohol
+
Alcohol
-
Alcohol
+
Acid
Acid
+
Amine
Acetone
O
1
Aldehyde
And/Or
Acid
o
2
Ketone
REDUCED
TO
HYDROLYSIS
(Methyl
Acetate)
-
-
-
-
CO2
+
H20
-
-
-
Alcohol
Alcohol
-
-
-
-
-
O
Table 3-3. Summary of properties for functional groups
MD0803
3-14
(3) Amide formation, as illustrated by the reaction below, is the reaction of
a carboxylic acid with an amine to yield a new class called an amide.
CH3 –COOH
Acid
+ CH3 –CH2 –NH2
Amine
O
ll
------> CH3 –C–NH–CH2 –CH3 + H2O
Amide
c. Uses of Carboxylic Acids. Many acids, such as acetic, salicylic, and lactic,
are used topically to treat local conditions. Others are used systemically. Still others,
like citric acid, which is found naturally in lemons, are used to flavor syrups for
administration of other drugs. They are also used in many analytical procedures in the
clinical laboratory.
3-13. ALDEHYDES
Aldehydes result from the first oxidation of alcohols and have the general
structural formula R-C-H. Since aldehydes cannot form hydrogen bonds between
ll
O
themselves, they have lower boiling points than corresponding alcohols or acids.
Again, as with the other classes or organic compounds in table 3-3, the lowermolecular-weight aldehydes (up to five carbons) are soluble in water. Aldehydes are
neutral in pH and undergo both oxidation and reduction reactions. They are easily
oxidized to acids and reduced to alcohols. Some aldehydes, such as vanillin and
benzaldehyde, are frequently used in the pharmacy as flavoring agents.
O
II
Ø-C-H
Benza ldehyde
O
Il
HCH
Methanal
(Formaldehyde)
Others, such as formaldehyde, are often used as disinfectants.
3-14. KETONES
Ketones result from the oxidation of a secondary alcohol and have the general
structural formula
O where R and R' can be the same or different
hydrocarbon groups.
║
R-C-R'
a. Ketones are similar to aldehydes in their boiling points, which are lower than
those of corresponding alcohols and carboxylic acids.
b. Ketones are neutral compounds, being neither acids nor bases. They
undergo the process of reduction, by which they are converted to secondary alcohols.
MD0803
3-15
c. The ketone functional group appears in the structure of many complex
drugs, such as steroid compounds and vitamins. Simple ketones, with the exception
of acetone, are seldom used. Acetone is used as a solvent and cleaning fluid.
O
║
CH3–C–CH3
2-propanone
(Acetone)
3-15. ESTERS
Esters, as previously mentioned, are formed from the reaction of a carboxylic
acid with an alcohol and have the general structural formula RCOOR' or O where R
and R' can be the same or different hydrocarbon groups.
II
R-COR'
a. Properties of Esters. The simplest esters are liquids and have fragrant
odors. An example is ethyl acetate, CH3-CH2-OOC-CH3, which has the odor of
pineapple. Esters cannot form hydrogen bonds between themselves; consequently,
they have boiling points similar to alkanes of similar molecular weight. They can form
hydrogen bonds with water. Therefore, esters that contain less than five carbon atoms
are soluble.
b. Reactions of Esters. Esters are neutral in pH and undergo two important
chemical reactions, hydrolysis, and reduction. Hydrolysis is the splitting of an ester
with the incorporation of water to form a carboxylic acid and an alcohol.
O
O
II
H2O
ll
R-C-O-R ---> R-C-OH + R'-OH
After hydrolysis, the acid product can undergo reduction to form a second alcohol as
described previously.
c. Uses of Esters. The ester functional group is found in many complex
molecules which you will be studying in the pharmacology subcourses if you take
them.
Acetylsalicylic Acid (Aspirin)--an analgesic
O
ǁ
C –OH
Ο
O–C–CH3
ll
O
MD0803
3-16
Nitroglycerin--a cardiac drug
CH2-O-NO2
l
CH2-O-NO2
l
CH2-O-NO2
3-16. AMIDES
Amides are formed from the reaction of a carboxylic acid with an amine or
ammonia and have the general formula
O
where R and R' can be the same
or different hydrocarbon groups.
R-C-NH-R'
a. Properties of Amides. Amides, because of the hydrogen attached to the
nitrogen atom, can form hydrogen bonds between themselves. They have higher
boiling and melting points than corresponding alkanes. Since they can also form
hydrogen bonds with water, amides containing up to five carbon atoms are soluble in
water.
b. Reactions of Amides. Amides are neutral in pH and undergo the
hydrolysis reaction. For amides, hydrolysis is the splitting of the compound with the
incorporation of water to form a carboxylic acid and an amine.
O
O
ll
ll
R-C-O-NHR' ------> R-C-OH + R'-NH2
H2O
c. Uses of Amides. Some examples of drug molecules containing the amide
functional group are shown below1
CH3
Lidocaine (Xylocaine) = local
anesthetic
О
O
ll
NH-C-CH2-N
CH3
C2H5
C2H5
Dibucaine (Nupercaine) = local
anesthetic
Ο
O H
ll ll
C-N-CH2-CH2-N
N
O-C4H9
MD0803
3-17
C2H5
C2H5
3-17. HALOGENATED HYDROCARBONS
Halogenated hydrocarbons are compounds with the general formula R-X where
R is any hydrocarbon group and X is a halogen (Cl, Br, F or I). The most significant
property of halogenated hydrocarbons is that as you increase the number of halogens
on the compound, the flammability of the compound decreases. This property has
been used to produce ethers that are nonflammable to be used as general anesthetics
such as:
F Cl
H
l l
l
Methoxyflurane (Penthrane)
H–C–C–O–C–H
l l
l
F Cl
H
3-18. SUMMARY
Functional groups, when attached to various hydrocarbons, increase the
reactivity and water solubility of the hydrocarbon. Carboxylic acids and phenols are
the only organic acids; they are weak acids. Amines are the only significant organic
bases. All functional groups that contain a hydrogen connected to a nitrogen or
oxygen atom have the ability to form hydrogen bonds between themselves. All
functional groups that contain a nitrogen or oxygen atom can form hydrogen bonds
with water, which increase their solubility. In general, organic compounds of low
molecular weight (less than five carbons) which contain functional groups, are soluble
in water. Table 3-3 summarizes the properties and some reactions of the organic
compounds we have studied.
Continue with Exercises
MD0803
3-18
EXERCISES, LESSON 3
INSTRUCTIONS. Write the word, words, symbols, or numbers that properly
completes the statement in the space provided or mark the correct word/phrase from
those given. After you complete the exercises, turn to Solutions to Exercises and
check your answers. Reread the material referenced for each exercise answered
incorrectly.
1. The most common type of bonding in inorganic chemistry is __________
bonding. The most common type of bonding in organic chemistry is
______________ bonding. This means that in organic compounds, the electrons
involved in bonds are usually ____________ by the atoms.
2. In inorganic chemistry, a compound is often completely identified by a simple
formula such as H3PO4, NaCl, or Fe(OH)3. However, in organic chemistry, we
frequently need formulas that indicate the molecule's ____________. In organic
chemistry, several compounds may have the same makeup in terms of
e__________ and number of each _________, but their chemical and physical
properties may be quite _____________.
3. In water, organic compounds are generally (soluble) (insoluble) and inorganic
compounds tend to be _____________. In organic solvents, organic compounds
are generally _________ and inorganic compounds are generally __________ .
4. A hydrocarbon is a compound that contains only two elements, _________ and
_________. If a hydrocarbon has only open chains (straight or branched), it is
called an _______________ hydrocarbon.
5. In an organic compound, two carbon atoms may share only one pair of electrons.
This is called a ______ bond. If two carbon atoms share two pairs of electrons, it
is called a _________ bond. If two carbon atoms share three pairs of electrons,
it is called a __________ bond.
6. An aliphatic hydrocarbon, which has only single bonds between carbon atoms is
called an ____________. An aliphatic hydrocarbon with at least one double
bond between carbon atoms is called an ________. If an aliphatic hydrocarbon
has at least one triple bond between two carbon atoms, the compound is called
an __________ .
MD0803
3-19
7.
In an aliphatic hydrocarbon, the chains of carbon atoms are open, that is, either
_______________ or ____________. If an aliphatic hydrocarbon has only
single carbon-to-carbon bonds, the compound is said to be s_____________.
A saturated compound has only ___________ bonds. Saturated aliphatic
hydrocarbons are also called ____________.
8. An unsaturated aliphatic hydrocarbon contains at least one bond, which is
__________ or _________. If it contains at least one double bond, it is called an
______. If it contains as least one triple bond, it is called an .
9. As a part of their structure, aromatic compounds include a __________ _______.
The number of carbon atoms in a benzene ring is ________
Ο
is a comon symbol for the _____ _____.
Compounds including such a structure are called ________________
compounds.
MD0803
3-20
10.
Below is a list of organic compounds.
Amide Carboxylic Acid
Ether Phenol
Halogenated Hydrocarbon Amine
Alcohol Aldehyde
Ester Ketone
Using this list, decide which class is most appropriate for each of the chemical
formulas below. Write the name of the class beside the chemical formula.
a. CH3-CH2-NH2
b. CH3-CH2-OH
c.
CH3-CH2-O-CH2-CH3
O
║
d. CH3-C-O-CH3
O
║
e. CH3-C-H
f.
О
OH
O
║
g. CH3-C-CH3
h. CH3-CH2-CH2-Cl
i.
MD0803
O
║
CH3-C-NH-CH3
3-21
11.
When a hydrogen atom in a hydrocarbon is replaced with a hydroxyl group
(OH-), the resulting compound is an __________. The general form of an alcohol
is represented by the symbol __________. CH3OH is an _________.
H3C
CH-CH2-OH is a (primary) (secondary) alcohol.
H3C
H3C
CH2-OH is a (primary) (secondary) alcohol.
H3C
12. If we compare an alcohol with an alkane of similar molecular weight, we find that
the boiling point and melting point of the alcohol are (lower) (higher). The
solubility of the alcohol is (lower) (higher). These properties are due to the fact
that alcohols can form ___________ bonds, both with w_________ molecules
and with other _______________ molecules.
13. If an organic compound has a hydrogen atom bonded to an oxygen or nitrogen
atom, then that compound can form ___________ bonds between its own
molecules.
14. If an organic compound includes a nitrogen or oxygen atom, then the compound
can form hydrogen bonds with __________. This makes the compound more
__________ in water if its molecular weight is (low) (high). By low molecular
weight, we mean that the number of carbons is generally less than ________.
Since all of the classes of organic compounds in Table 3-3 (alcohols, phenols,
ethers, amines, carboxylic acids, aldehydes, ketones, ester, and amides) have at
least one oxygen or nitrogen atom, lower molecular-weight in these classes tend
to be _____________ in water.
MD0803
3-22
15. Write yes or no by each of the following classes or compounds to indicate
whether compounds of that class generally form hydrogen bonds between their
own molecules.
a.
Alcohol (ROH)
b.
Aldehyde (RCOH)
c.
Amide •(RCONH2)
d.
Amine (RNH2)
e.
Carboxylic acid (RCOOH)
f.
Ester (RCOOR)
g.
Ether (ROR)
h.
Ketone (RCOR)
i.
Phenol (ArOH)
16. If compounds of a single class can form hydrogen bonds between their own
molecules, then the boiling point of that compound is (lower) (higher) than the
alkane of similar molecular weight. Compounds with one particular functional
group are able to form two hydrogen bonds between a pair of its molecules; this
functional group is _______________. Thus, carboxylic acids tend to have a
very (high) (low) boiling point, compared to alkanes of similar molecular weight.
As a result, carboxylic acids are usually (solid) (liquid).
17. The boiling points of ethers, aldehydes, ketones, and esters are (about the same
as) (more than) the boiling points of alkanes of similar molecular weight.
18. The boiling points of alcohols, phenols, amines, amides, and carboxylic acids are
(lower) (higher) than the boiling points of alkanes of similar molecular weight.
MD0803
3-23
19.
Label each of the following compounds as to whether it is acidic, basic, or
neutral.
a.
OH
b.
CH3-CH2-OH
c.
CH3-O-CH3
d.
CH3-CH2-NH2
e.
O
║
CH3-C-OH
20.
Of the compounds in the exercise above, salts would be mostly easily formed
with the ones lettered ___, ___, and ___. The two classes of organic
compounds which will form salts when reacted with inorganic bases are
_________ and ___________ __________. The class of organic compounds
which will form salts when reacted with inorganic acids is _________. This
property of carboxylic acids and amines is most useful in pharmacy because it
helps to make organic compounds that are very ________________ in water.
Such salts tend to be quite soluble even when the molecular weight is high.
They can thus be more easily utilized by the body.
21.
In organic chemistry, a molecule is oxidized if it (gains) (loses) oxygen atoms or
if it (gains) (loses) hydrogen atoms. Reduction is the opposite of oxidation. In
organic chemistry, a molecule is reduced if it (gains) (loses) oxygen atoms or
(gains) (loses) hydrogen atoms.
O
O
[O]
║
[O]
║
22. CH3-CH2-OH ----> CH3-C-H -----> CH3-C-OH
ethyl alcohol
acetaldehyde
acetic acid
As illustrated above, the first step in the oxidation of a primary alcohol is the
formation of an
. The second step results in the
formation of an
_____ ___________.
MD0803
3-24
OH
O
ll
[O]
ll
23. CH3-CH-CH3 --------> CH3-C-CH3
isopropyl
acetone (dimethyl)
alcohol
ketone)
As illustrated above, the oxidation of a secondary alcohol results in the formation
of a ___________ .
24. The first step in reduction of a carboxylic acid results in the formation of an
___________ . The second step results in the formation of an ____________ .
25. The reduction of ketones result in the formation of a ____________ __________
.
26.
O
H+
O
ll
ll
CH3-C-O-CH2-CH3 + H2O -----> CH3-C-OH
ethyl acetate
acetic acid
+ CH3CH2OH
ethyl alcohol
The above reaction is an example of (oxidation) (reduction) (hydrolysis). The
general form of an ester is ____________.
The ester in the above reaction is __________ _________. Hydrolysis of an
ester results in the formation of an ___________ and a __________
__________. We know that a carboxylic acid can be reduced to form another
__________. Therefore, the combined processes of hydrolysis and reduction of
an ester results in the formation of two ___________.
Hydrolysis is the splitting of a compound with the incorporation of __________ to
form (two) (three) new compounds.
MD0803
3-25
O
ll
CH3-C-OH + CH3CH2OH
acetic acid
ethyl alcohol
27.
H+
----->
O
ll
CH3-C-O-CH2CH3 + H2O
ethyl acetate
Ethyl acetate is an example of the class of compounds called ___________.
Therefore, the reaction above can be called ester _______________ or
__________________. It is almost exactly the opposite of a reaction between
an ester and water called ______________. Esterification is an important
property of both alcohols and carboxylic acids.
O
ll
28. CH3-C-NH2 + H2O
acetamide
------->
O
ll
CH3-C-OH
acetic acid
+ H-NH2
ammonia
The above reaction is an example of ____________. Acetamide is an example
of the class of compounds called _________. Ammonia may be considered the
simplest example of the class of compounds called __________. Hydrolysis of
an amide results in the formation of a ___________ _______ and an
___________.
O
O
ll
ll
29. CH3-C-OH + CH3-NH2 -----> CH3-C-NH-CH3 + H2O
acetic acid methylamine
methylacetamide
Methylacetamide is an example of the class of compounds called
___________. Therefore, the reaction above can be called ____________
_____________. Whereas a carboxylic acid and an alcohol may react to form
an _____________, a carboxylic acid and an amine may react to form an
_______. This reaction is the opposite of a reaction between an amide and
water called _____________. Amide formation is an important property of both
________________ _________ and _______________.
30. As hydrogen atoms in a hydrocarbon molecule are replaced by halogen atoms,
the flammability of the molecule (decreases) (increases).
Check Your Answers on Next Page
MD0803
3-26
SOLUTIONS TO EXERCISES, LESSON 3
1.
ionic
covalent
shared (para 3-4, Table 3-1)
2.
structure
elements, atom, different (para 3-3)
3.
insoluble; soluble
soluble, insoluble (Table 3-1)
4.
hydrogen; carbon
aliphatic (para 3-6)
5.
single
double
triple (para 3-5)
6.
alkane
alkene
alkyne (para 3-5)
7.
straight, branched
saturated;
single
alkanes (para 3-6a)
8.
double, triple
alkene
alkyne (para 3-6a(2))
9.
benzene ring
six
benzene ring
aromatic (para 3-6b)
10.
a Amine (para 3-11)
b Alcohol (para 3-8)
c Ether (para 3-10)
d Ester (para 3-15)
e Aldehyde (para 3-13)
f. Phenol (para 3-9)
g Ketone (para 3-14)
h Halogenated hydrocarbon (para 3-17)
i Amide (para 3-16)
MD0803
3-27
11.
alcohol
ROH
alcohol
primary
secondary (para 3-8)
12.
higher;
higher
hydrogen, water, alcohol (para 3-8a)
13.
hydrogen (para 3-18)
14.
water
soluble low
five
soluble (para 3-18)
15
a
b
c
d
e
f
g
h
i
16.
higher
O
║
-C-OH
(carboxyl)
high;
solid (paras 3-8a, 3-12a;Table 3-3)
17.
about the same as (Table 3-3)
18.
higher (Table 3-3)
19
a
b
c
d
e
MD0803
yes (para 3-8a; Table 3-3)
no (para 3-13; Table 3-3)
yes (para 3-16a; Table 3-3)
yes (para~3-11a; Table 3-3)
yes (para 3-12a; Table 3-3)
no (para 3-isa; Table 3-3)
no (para 3-l0a; Table 3-3)
no (Table 3-3)
yes (Table 3-3)
acidic
neutral
neutral
basic
acidic (Table 3-3)
3-28
20.
a, d. ephenols; carboxylic acids,
amines
soluble (paras 3-9b, 3-11b, 3-12)
21.
gains, loses
loses; gains (paras 3-8b; 3-12b(1))
22.
aldehyde
carboxylic acid (para 3-8b(1))
23.
ketone (para 3-8b(2))
24.
aldehyde
alcohol (para 3-12b(1))
25.
secondary alcohol (para 3-14b)
26.
hydrolysis
O
ll
R-C-O-R
ehyl acetate
alcohol; carboxylic acid
alcohol
alcohols
water; two (para 3-15)
27.
esters
formation; esterification
hydrolysis (paras 3-12b(2); 3-15)
28.
hydrolysis
amides
amines
carboxylic acid, amine (para 3-16)
29.
amides
amide formation
ester, amide
hydrolysis
carboxylic acids, amines (para 3-12b(3); 3-16)
30.
decreases (para 3-17)
End of Lesson 3
MD0803
3-29
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