WEEK 4
OXYGEN
Oxygen is a nonmetallic element of great importance because it is the most abundant
element on earth. It comprises 21% of the volume of atmospheric air and 89% of the
weight of water. The word oxygen means acid former. An early scientist observed that
all of the acids known to him contained oxygen. The symbol for oxygen is O, but its
formula is O2, since it is a homonuclear diatomic molecule. Its position on the upper
right hand corner of the periodic table indicates that it is highly electronegative. It is a
member of the group VI elements. It, therefore, has six valence electrons and usually an
oxidation number of –2 except when combined with itself.
PHYSICAL PROPERTIES
1.
2.
3.
4.
Oxygen is a colorless, odorless, and tasteless gas.
Oxygen is heavier than air.
It is slightly soluble in water
It can be liquefied and solidified by extreme pressure and low temperature.
CHEMICAL PROPERTIES
1. Reacts with many substances slowly at ordinary temperatures but rapidly at high
temperatures.
2. Reacts with many metallic and nonmetallic elements to form oxides.
3. Supports combustion. This does not mean that oxygen burns. It means that
oxygen helps other things burn. In fact, it is helpful to define burning as the
process of chemically adding oxygen to another substance with the production of
an oxide. Another word for burning is COMBUSTION.
4. Oxygen acts as an oxidizing agent.
Before the last property can be understood, the concepts of oxidation and reduction
must be learned. There are four ways to define OXIDATION:
1.
2.
3.
4.
The addition of oxygen to a substance.
The removal of hydrogen from a substance.
The removal of electrons from a substance.
The increase in oxidation number of a substance.
Burning (combustion) is an example of the first definition. When oxygen is added to
a substance, that substance undergoes oxidation. Since oxygen has caused this
process, it is called an oxidizing agent. When the oxidizing agent has performed its
task, it is reduced. Reduction is the counterpart of oxidation and always
accompanies it. Since reduction is the opposite of oxidation, we obtain definitions of
reduction by reversing the definitions of oxidation. Reduction is:
1.
2.
3.
4.
The removal of oxygen from a substance.
The addition of hydrogen to a substance.
The addition of electrons to a substance.
The decrease in the oxidation number of a substance.
Since oxidation and reduction go hand in hand, the oxidizing agent is reduced and the
reducing agent is oxidized. For example, in burning the substance burned undergoes
oxidation and is the reducing agent. Oxygen simultaneously undergoes reduction and
is the oxidizing agent. The substance burned causes oxygen to be reduced, and
oxygen causes the burned substance to be oxidized. The other definitions of
oxidation and reduction do not necessarily involve oxygen or the formation of oxides.
The second definition will be covered in the chapter on hydrogen.
The last two definitions can be explained by considering an example involving
oxygen and one in which oxygen does not participate.
Consider the formation of carbon dioxide from the elementary substances carbon
and oxygen:
C + O2
CO2
The first definition says that carbon is oxidized, since oxygen is added to it.
Oxygen must at the same time undergo reduction. The same conclusion can be
drawn from the last two definitions of oxidation and reduction. When the
molecule carbon dioxide forms, bonds occur between the one carbon, and the two
oxygens. These bonds are polar covalent. Oxygen is more electronegative than
carbon, so the electrons in the bonds are pulled from the carbon atom and toward
the oxygen atom. Carbon is oxidized because is loses electrons. Oxygen is
reduced because it gains electrons.
Application of the last definition of oxidation and reduction gives the same result.
An oxidation number indicates if electrons have been added or lost by a
substance. Atoms that lose electrons have positive oxidation numbers, and atoms
that gain electrons have negative oxidation numbers. Before the reaction forming
carbon dioxide occurs, carbon and oxygen have oxidation numbers of zero,
because neither has gained or lost electrons in comparison to its normal neutral
state. After the reaction we have:
C0
C+4
and each of the two oxygens undergoes the following:
O0
O-2
The oxidation number of the carbon has increased (from 0 to +4); therefore, it has
been oxidized. The oxidation number of the carbon has decreased (from 0 to –2);
therefore, it has been reduced.
Another example of an oxidation-reduction process is the reaction resulting from
passing electricity through a sodium chloride melt.
2Na + Cl2
2NaCl
In the compound NaCl, sodium is in the +1 oxidation state and Cl is in the –1
oxidation state. The individual oxidation-reduction reactions are:
2Na+ + 2e-
2Na
Cl2 + 2e-
2Cl-
Sodium has gone from a +1 oxidation state in the compound to a zero oxidation
state in the elemental form. Sodium has gained electrons and is reduced. It is the
oxidizing agent.
Chlorine goes from a –1 oxidation state in the compound to a zero oxidation state
in the elemental form. Chlorine has lost electrons and is oxidized. It is the
reducing agent.
A good way to remember the definitions for oxidation and reduction is with the
phrase LEO GERs.
Loss of Electrons is Oxidation
Gain of Electrons is Reduction
FORMS OF OXYGEN
In addition to being a diatomic molecule, oxygen can exist in other forms. When an
element exists in two or more different forms, each with its own physical and chemical
properties, it is called ALLOTROPIC. There are three allotropes of oxygen:
O3
Ozone
(unstable)
O2
O
Molecular Oxygen
(stable)
Nascent Oxygen
(unstable)
Ozone is prepared by passing a high-voltage electrical discharge through oxygen.
electricity
3O2
2O3
The physical properties of ozone are:
1. Pale blue-colored gas.
2. Penetrating odor. The clean smell that you notice after an electrical storm is
caused by ozone.
The chemical properties of ozone are:
1. More active chemically than oxygen, but less stable.
2. Readily decomposes, liberating nascent (newly formed) oxygen.
3. Has great bleaching and deodorizing abilities. This is due to its high oxidizing
capabilities. Has been used to treat drinking and waste water. It is a true
deodorant because it destroys the odor instead of just masking it.
PREPARATION OF OXYGEN
Laboratory method – In the laboratory, we heat a mixture of potassium chlorate
and manganese dioxide.
KClO3 + MnO2

2KCl + 3O2 + MnO2
Manganese dioxide, MnO2, is a catalyst and will be discussed in the next section.
Notice that it does not change during the course of the reaction.
Commercial method – There are two commercial methods for the preparation of
oxygen. One is the evaporation of liquid air. Air is a mixture of about 78%
nitrogen and 21% oxygen. The other method is electrolysis of water:
electricity
2H2O
2H2 + O2
FACTORS AFFECTING REACTION RATES
In the laboratory preparation of oxygen we used manganese dioxide, a catalyst, to
increase the rate of the reaction. Several things can influence how fast our reactions run.
Catalysts – A substance that speeds up a chemical reaction yet appears in
unchanged form among the products of the reaction. All a catalyst does is
increase the rate of the reaction. A catalyst increases the rate of the reaction by
lowering the amount of energy necessary for the reaction to occur. CATALYSTS
DO NOT CAUSE REACTIONS; THEY ALTER THE REACTION RATES.
Catalysts are often written above the arrow in an equation instead of on the
product and reactant side.
Temperature – There is a direct proportion between temperature and the speed of
a chemical reaction. An increase in temperature increases the speed of reactions,
and a decrease in temperature decreases the speed of reactions.
Physical State – This refers to the size of the particles of the reacting substances.
The smaller the size of our reactant, the more surface area we have on which the
reaction can occur. We can dissolve much more finely ground sugar in a glass of
iced tea than course ground sugar. Notice this is a physical change and not a
chemical change.
Concentration – The rate of a chemical reaction is proportional to the molecular
concentration of the reacting substances. The greater the concentration, the faster
the reaction. The lesser the concentration, the slower the reaction.
Light – Many chemical reactions are caused or accelerated by light. These are
called photochemical reactions. Embalmed jaundiced remains may turn green
when exposed to fluorescent lights. Fluorescent lights are missing one component
that ordinary visible light has. This condition may be avoided by the use of
“color-corrected” fluorescent lights that do not lack this component of light.
These lights may actually cause a photochemical reaction that results in a pink
tone to the skin.
Pressure – Pressure has very little effect on chemical reactions other than those
involving gases. Reactions between gaseous reactants are accelerated by an
increase in pressure and slowed by a decrease in pressure.
WEEK 4
HYDROGEN
Oxygen is the most abundant element on earth, but hydrogen is the most abundant
element in the universe. There is very little free hydrogen in the atmosphere. Most
hydrogen occurs in the combined state and may be found in acids, bases, water, plant and
animal tissue, and natural gases. The free form of hydrogen exists in the earth’s crust,
and is released in the gases of volcanic eruptions. Tissue gas, a post-mortem
manifestation of gas gangrene, is largely free hydrogen gas. Chemically hydrogen has
the symbol H and the molecular formula H2. Like oxygen, hydrogen is a homonuclear
diatomic molecule. It has an atomic number of one, an atomic weight of one, and usually
has an oxidation number of +1 except when combined with itself.
PHYSICAL PROPERTIES
1. Hydrogen is a colorless, odorless, and tasteless gas.
2. Hydrogen is lighter than air. It is the lightest gas known.
3. Hydrogen is slightly soluble in water.
Comparing hydrogen with oxygen shows some interesting similarities, as well as
differences. Both are colorless, odorless, and tasteless gases. In order to test for
these gases you must test for one of their chemical properties. Both are slightly
soluble in water. Oxygen is heavier than air, but hydrogen is lighter than air. O2 has
a molecular weight of 32 and H2 has a molecular weight of 2. In addition to being
light, hydrogen is highly diffusible. Remember that diffusibility is a property of
gases that allows them to move through their containers. If we have a balloon filled
with helium (atomic weight of 2), we know that in a day or two the balloon will settle
to the floor. This occurs because the helium has diffused out of the balloon. Imagine
how quickly this would happen if we filled a balloon with hydrogen which only has
an atomic weight of one.
CHEMICAL PROPERTIES
1. Hydrogen burns wit a hot, blue flame, forming water.
2. Hydrogen does not support combustion.
3. Hydrogen acts as a reducing agent.
The description burns with a hot, blue flame forming water, is somewhat
misleading because depending on the amount of oxygen present, this reaction
usually takes the form of an explosion. The most famous exhibition of this property
was the Hindenburg disaster. If helium diffuses through the tight membrane of a
toy balloon, try to imagine an airship approximately two to three football fields in
length containing 6.7 million cubic feet of hydrogen. One small spark of static
electricity was reported to have been the cause of this cataclysmic chemical
reaction.
The second chemical property of hydrogen is that it does not support combustion.
This is an interesting contrast to oxygen. Hydrogen burns but does not support
combustion, whereas oxygen supports combustion but does not burn. This means
that pure hydrogen will extinguish a full flame. However, hydrogen can not be
used in fire extinguishers because it so light that it is very hard to control.
The third property of hydrogen is that it acts as a reducing agent. Remember that
Oxidation is the removal of hydrogen from a substance, and reduction is the
addition of hydrogen to a substance. Hydrogen acts as a reducing agent when it is
added to a substance. Consider the preparation of methyl alcohol:
CO + 2H2
catalyst
CH3OH
Because carbon monoxide gains hydrogen, it is reduced. Hydrogen is the reducing
agent.
We can demonstrate the definition of oxidation as removal of hydrogen from a
substance by a method for the preparation of formaldehyde.
2CH3OH + O2
Methyl alcohol
2HCHO + 2H2O
Formaldehyde
In this reaction, methyl alcohol is oxidized, because it loses hydrogen. Oxygen
undergoes reduction because it gains the hydrogen. Methyl alcohol is the reducing
agent, and oxygen is the oxidizing agent. The loss of hydrogen by reducing agents
occurs in many biological oxidation-reduction reactions. As our bodies break down
ingested foods, these molecules frequently lose hydrogen ions and electrons, which
in turn are accepted by very large molecules that act as oxidizing agents.
PREPARATION OF HYDROGEN
Commercially hydrogen is produced by the electrolysis of water. In the lab, hydrogen is
produced using one of the chemical properties of acids. Common acids used in the
laboratory contain hydrogen like HCl (hydrochloric acid). The property of acids
important in the preparation of hydrogen involves a concept called the
ELECTROCHEMICAL SERIES. This series is a list of the metallic elements in order of
their chemical activity. The more active metals are listed first, the less active last. All
elements above hydrogen in the series displace hydrogen from dilute acids. All elements
below hydrogen in the series do not displace hydrogen from dilute acids. The application
of this principle is the basis for a type of chemical reaction called single replacement.
For example:
ZnCl2 + H2
Zn + 2HCl
Zinc is listed above hydrogen in the electrochemical series, therefore, it displaces the
hydrogen from the acid. This reaction shows us that not all oxidations involve oxygen.
In general, whenever a metal displaces hydrogen from a dilute acid by a singlereplacement reaction, the metal is oxidized and the hydrogen of the acid is reduced. Zinc
is our metal and is oxidized because it has gone from a 0 oxidation state to a +2 oxidation
state. It has lost electrons. Hydrogen is reduced because it has gone from a +1 oxidation
state to a 0 oxidation state. It has gained electrons.
Another feature of the activity series is based on a chemical property of water. As the top
of the series is approached, the metals displace hydrogen not only from dilute acids but
also from water. Very active metals such as potassium, calcium, and sodium, actually
reduce water, liberating free hydrogen. Consider the following:
2K + 2HOH
2KOH + H2
Once again our metal, K, is oxidized and our hydrogen is reduced.
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Tin
Lead
HYDROGEN
Copper
Mercury
Silver
Platinum
Gold
Decreasing
Activity
ACTIVITY SERIES OF METALLIC ELEMENTS