97
Chemical Reactions I:
Some Chemical Reactions Involving Oxygen
PRE-LAB ASSIGNMENTS:
To be assigned by your lab instructor.
STUDENT LEARNING OUTCOMES:


Learn how to classify, complete, and balance chemical reactions.
Understand the role of oxygen in many common types of reactions.
EXPERIMENTAL GOALS:
The purpose of this lab is to observe some reactions involving oxygen, and to learn about
different types of chemical reactions.
INTRODUCTION:
In this lab, you will observe a number of chemical reactions, many of which involve
molecular oxygen, O2. The Earth’s atmosphere consists of about 21% oxygen; this oxygen is
involved in many chemical processes that take place on the Earth, from respiration to rusting.
This lab will demonstrate a number of reactions involving molecular oxygen, in both the 20%
oxygen of the atmosphere and in 100% pure oxygen generated in the lab.
In order to understand some of the reactions that we will be looking at, it is important to
understand some of the ways in which chemists classify chemical reactions. One scheme for
classifying chemical reactions is described over the next few pages. In this scheme, reactions are
classified by counting the number and examining the types of reactants and products that are
involved in the reaction, and also by considering the way that groups are exchanged between
reactants and products. This classification scheme does not work well for all reactions — for
instance, many oxidation-reduction reactions do not really fit well under this scheme — but it
works well for many types of common chemical reactions.
[Another classification scheme for chemical reactions, which will be discussed in the
Reactions II experiment, classifies reactions by the driving force of the reaction (precipitation,
neutralization or oxidation-reduction), and provides an additional understanding of chemical
reactions.]
98
Classes of Chemical Reactions:
Classification of Reactions by Numbers of Reactants and Products
and by Exchange of Groups
I. Combination (Synthesis) Reactions — elements or compounds combine to form a new
compound. Two or more reactants form one product.
A + B  C
A. Compound Formation (Element Combination)
element + element  compound
1. metal + non-metal  ionic compound
The formula of the compound is written using the rules for ionic charges.
For example, when aluminum reacts with sulfur, aluminum forms a 3+ charge and
sulfur forms a 2- charge, producing the formula
Al + S  Al3+ + S2- = Al2S3
Balancing the equation gives
2Al + 3S  Al2S3
2Mg + O2  2MgO
2. non-metal + non-metal  molecular compound
The formula of the compound cannot always be predicted, because nonmetals can in
many cases combine in several different ways.
S + O2  SO2 (or SO3)
C + O2  CO2 (or CO)
B. Metal Oxide + Water  Metal Hydroxide (base)
1. The formula for the metal hydroxide is written using the charge rules for metals and
the hydroxide ion.
Al2O3 + 3HOH  2Al(OH)3
MgO + HOH  Mg(OH)2
C. Non-Metal Oxide + Water  Acid
1. The non-metal has the same oxidation state in the oxide as in the acid. To write the
correct formula for the acid, add one molecule of water to the oxide formula, and
simplify if necessary. [For phosphorus oxides, add three molecules of water.]
99
SO2 + H2O  H2SO3
N2O5 + H2O  [H2N2O6]*  2HNO3
* since all of the subscripts are divisible by 2, this can be simplified
P2O3 + 3H2O  [H6P2O6]*  2H3PO3
* since all of the subscripts are divisible by 2, this can be simplified
II. Decomposition — chemical compounds are broken down to simpler substances, either
simpler compounds or elements. One reactant forms many products.
something

 two or more products
single compound do
 current

 2H2(g) + O2(g)
2H2O(l) electric
H2CO3(aq) 
H2O(l) + CO2(g)


A. Decomposition of Binary Compounds. Binary compounds decompose into their
constituent elements. (Check for any diatomic elements that may be produced.)
something

 element + element
binary compound do
HgS 
Hg + S


 current

 2Na + Cl2
2NaCl electric
B. Decomposition of Carbonates. Compounds containing the carbonate ion decompose
when heated to give metal oxides and CO2 gas.
CaCO3(s) 
CaO(s) + CO2(g)


C. Decomposition of Oxoanions. Many oxygen-containing polyatomic ions decompose
when heated to give off oxygen gas. However, the remaining products are not always
easy to predict, since the oxygen may be removed partially or completely.
2KClO3(s) 
2KCl(s) + 3O2(g)


2NaNO3(s) 
2NaNO2(s) + O2(g)


III. Single Displacement — an uncharged metal (that is, a metal in the elemental state) reacts
with an ionic compound (or water, or an acid), displacing the metal (or H) from the ionic
compound, producing a new ionic compound, and a new uncharged metal (or H2 gas). The
activity series determines whether a potential single displacement reaction actually occurs.
Uncharged metal + ionic compound  new metal + new ionic compound
A + BC  B + AC
100
1. Check the activity series to see whether a reaction takes place If the uncharged metal
A is lower on the activity series than the charged metal B, no reaction (NR) takes
place. (There’s a reason why wedding rings are made of gold, and not magnesium.)
2Al + 3CuCl2  2AlCl3 + 3Cu
Cu + AlCl3  NR
2. The formula for the new ionic compound is written using the ionic charge rules.
Check to see if the “new metal” is diatomic H2.
2Na + ZnCl2  2NaCl + Zn
2Na + 2HOH  2NaOH + H2
Mg + 2HBr  MgBr2 + H2
IV. Double Displacement (metathesis, double replacement or partner exchange) — two ionic
compounds react and exchange cation and anion partners, generating two new ionic
compounds (or water from the combination of H+ and OH-). This reaction is common for
ionic compounds in aqueous solutions.
ionic compound + ionic compound  new ionic compound + new ionic compound
AB + CD  AD + CB
1. Formulas for new compounds are written using the ionic charge rules.
2. Do not change charge on any species involved in the reaction.
K2CO3(aq) + Ca(NO3)2(aq)  2KNO3(aq) + CaCO3(s)
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2HOH(l)
V. Combustion — in combustion reactions, organic compounds (which contain carbon,
hydrogen, and perhaps oxygen) react with molecular oxygen, producing carbon dioxide,
water, and energy.
Organic Compound + O2  CO2 + H2O + energy
CH4 + 2O2  CO2 + 2H2O
C6H12O6 + 6O2  6CO2 + 6H2O
VI. Miscellaneous (reactions not falling into the above patterns)
Cu(s) + 4HNO3(aq)  Cu(NO3)2(aq) + 2 NO2(g) + 2 H2O(l)
101
Activity
Series
(Do not memorize)
Li
K
Ba
Ca
Na
Mg
Al
Mn
Zn
Fe
Cr
Cd
Co
Ni
Sn
Pb
H2
Cu
Ag
Hg
Au
Diatomic
Elements
(Memorize)
H2
N2
O2
F2
Cl2
Br2
I2
Balancing Chemical Reactions:
1. Identify all reactants and products.
2. Determine the correct formulas for all reactants and products.
3. Write the unbalanced chemical equation with reactants on left side of the arrow and products
on the right side of the arrow.
H2 + O2  H2O
[unbalanced]
4. Balance equations by adding and/or changing stoichiometric coefficients ONLY.
general, whole-number coefficients are preferred.)
2H2 + O2  2H2O
H2 + ½O2  H2O
(In
[balanced]
[balanced, but has a fractional coefficient]
5. Never change a correct formula to balance an equation (that changes the identity of the
substance).
H2 + O2  H2O
H2 + O2  H2O2
[unbalanced]
[balanced, but we’re making hydrogen peroxide instead of water]
6. Make sure that you recognize polyatomic ions in reactions; in many cases, it is easier to
balance polyatomic ions as a unit than to balance individual atoms. (Consult Appendix V for
a list of polyatomic ions.)
102
PROCEDURE:
NOTE: Fill in the report sheets as you watch the demonstrations or perform the
experiments. Where indicated on the report sheet, write and balance equations for
the observed reactions. Make sure all formulas are correct before you attempt to
balance an equation.
The chemical reactions that we normally call “burning” or “combustion” usually involve
the reaction of an element or compound with molecular oxygen1 (O2). These reactions require
three conditions:



A combustible or flammable material (e.g. fuel).
A supply of oxygen. The source of oxygen is usually air (approximately 21% O2). In
pure oxygen, these reactions are even more vigorous.
Some source of heat or some other initiator to start the process. (Reactions generally go
faster as the temperature is increased.)
A fire can be extinguished by removing any one of these three components. Common firefighting methods seek to either cool the burning object or remove the oxygen supply. Water
both cools and smothers (removes oxygen by displacing it with water vapor), while fire blankets
and chemical fire extinguishers such as CO2 act by smothering the fire. Extinguishing fires by
blasting, such as in oil fields, also works by removal of oxygen. Oxygen itself is not considered
to be flammable, but is said to “support combustion” (it is an oxidizer).
A. Lighting the Match
1. Safety Matches.
KClO3, Sb2S3, glue
paraffin
red phosphorus,
powdered glass, glue
Figure 5. A safety match with its box striker.
1
Other oxidants are possible, but are less common. Halogens, such as Cl2 are good oxidants. Even compounds that
are normally inert can serve as oxidants with very reactive fuels: water can oxidize sodium (a single displacement
reaction), and magnesium will burn in carbon dioxide.
103
The safety match was invented in 1855. When the match is struck, friction converts
some of the red phosphorus on the box into the less stable allotropic form, white
phosphorus, which ignites instantly. The heat thus produced causes the Sb2S3 to react
vigorously with oxygen produced in the decomposition of KClO3, producing Sb2O3, SO2
and SO3, and KCl, and generating even more heat, which ignites the wood or paper match
(which are often coated with paraffin). These matches may sometimes ignite if they are
struck on rough glass surfaces.
2. Strike-Anywhere Matches.
P4S3, , powdered glass, glue
S, KClO3, glue
paraffin
Figure 6. A typical strike-anywhere match.
In a strike-anywhere match, everything needed to start a fire is contained in the match
head. Friction from the strike causes P4S3 to react with oxygen from KClO3, producing
P2O3 and P2O5, SO2 and SO3, and KCl Heat from this brief reaction ignites the match
stick.
B. The Bunsen Burner
1. Reaction (Combustion). Natural gas is made up of about 80-90% methane (CH4) and 515% ethane (C2H6), and traces of other hydrocarbons. When hydrocarbons (organic
compounds containing only C and H) burn completely in oxygen or air (complete
combustion) the only products are carbon dioxide and water. See data sheet.
2. Burner Flame. The most common apparatus for generating heat in the chemistry lab is
the Bunsen burner, invented by Robert Bunsen and Peter Desaga in 1854. Natural gas
enters the base of the burner and mixes with air entering the lower ports on the burner
barrel (see Figure 1). When these ports are closed, not much air mixes with the natural
gas, and the flame is not very hot. The combustion is not complete (i.e., some of the
methane is oxidized to carbon monoxide and elemental carbon instead of carbon
dioxide), resulting in a flame that is yellow. The yellow color comes from light emitted
by particles of hot elemental carbon, which deposits as soot (also known as lampblack or
carbon black) on a piece of glass or porcelain held in the flame. As more air is admitted
through the base of the burner, the natural gas mixes with oxygen in the air, and the flame
becomes hotter, and combustion is complete (i.e., most of the carbon in the methane is
oxidized to carbon dioxide). The blue flame color comes from hot gaseous combustion
104
products. Usually, an inner, brighter blue cone is observed, surrounded by an outer, paler
blue cone. An index card inserted into the flame will illustrate the “hottest” and “coldest”
parts of the flame. A match head left sitting inside the inner cone will illustrate that no
combustion is taking place inside the inner blue cone. See data sheet.
air
air
natural gas
Figure 1. A typical Bunsen Burner.
3. Flame Propagation Speed. A flame will travel along a stream of mixed fuel and air as if
the stream were a fuse. The speed of travel will depend on what the fuel is and whether
air or pure O2 is used. Flames travel 10 to 100 times faster in an O2-H2 or O2-C2H2
(acetylene) mixture than in an air-natural gas mixture. In a torch or burner (see Figure 2),
the mixture of gases moves away from the nozzle, while the flame propagates in the
opposite direction, toward the nozzle:
gas moves
flame moves
fuel
air or O2
Figure 2. A typical torch assembly.
If the gas flow is turned up too high, the flame will sometimes “separate” or “lift off”
105
from the nozzle, and be extinguished. If the gas flow is too low, the flame may travel
back into the body of the torch, causing an explosion. This is especially a problem for
gas mixtures in which the flame travels rapidly. For these mixtures, the fuel and air or O2
must be mixed at the nozzle, and special torches are used. Torches for natural gas/air
mixtures should not be used for O2-H2 or O2-C2H2 mixtures!
C. Decomposition of Potassium Chlorate
1. Reaction. Upon heating, potassium chlorate decomposes to form potassium chloride and
oxygen gas. MnO2 acts as a catalyst, speeding up the reaction without itself being
consumed. See data sheet.
2. Apparatus.
KClO3 + MnO2
O2
O2
H2O
Figure 3. A typical setup for collecting gases by water displacement..
Oxygen gas (O2) enters the upside-down collecting bottle through the rubber tube, and
displaces the water. (Oxygen does not dissolve well in water). Full bottles are stored
upright because O2 is slightly more dense than air, and tends to settle to the bottom of
the bottle. The slight cloudiness comes from water vapor mixed with the oxygen gas.
D. Reactions in Pure Oxygen — Splint, Cloth, Cigarette, Plastic, and Vaseline
1. Reaction. Wood, fabrics, tobacco and some plastics are all made up of compounds that
contain carbon, hydrogen and oxygen, and can be represented by the generic formula
CxHyOz and can burn in oxygen or air readily:
CxHyOz + O2  CO2 + H2O + energy [not balanced]
Notice that the vigor of the reaction shows that smoking (or any other type of flame)
around pure O2, especially in hospitals, should be studiously avoided!
106
E. Iron
1. Iron reacts very slowly with oxygen in the air (rusting), but finely divided iron can be
seen to burn visibly when heated in a Bunsen burner flame in air. Note how much more
dramatic this reaction is when carried out in pure oxygen. Both iron(II) and iron(III)
oxides are formed. Iron(III) oxide is also known as rust. See data sheet.
2. This reaction is similar to one that claimed the lives of three U.S. astronauts when the
inside of their Apollo I command module, which was pressurized with pure oxygen,
caught fire during a routine launch simulation on January 27, 1967. (Since the Apollo
spacecraft were designed to operate in pure oxygen in outer space, albeit at a pressure of
5 psi instead of 14.7 psi, they were originally tested on the ground in a pure oxygen
environment.) In the pure oxygen environment, many things that ordinarily wouldn’t
have been flammable suddenly were flammable — from the Velcro that the astronauts
used to hold things down in zero gravity, to the metal in the spacecraft’s frame.2
F. Magnesium
1. Your instructor will burn a strip of magnesium in the air. DO NOT LOOK DIRECTLY
AT THE INTENSE LIGHT PRODUCED! This reaction is so violent that it is carried
out in air rather than pure O2. Burning magnesium used in flares, and in cutting through
other metals. See data sheet.
2. A reaction occurs when the product of the combustion product of magnesium and oxygen
is mixed with water. Your instructor will use an acid-base indicator to help you
determine what that reaction is.
G. Sulfur
1. Sulfur burns in air with a deep blue flame; this reaction is more intense in a pure oxygen
environment. Two gases, SO2 and SO3, are produced. These same reactions occur when
coal or oil which contains sulfur is burned, producing air pollution. See data sheet.
2. Your instructor will mix the gases produced by burning sulfur with water, and test it
using an acid-base indicator. These reactions can occur in the atmosphere when rain
combines with SO2 and SO3. See data sheet.
H. Red Phosphorus
1. Phosphorus exists in several allotropic forms, two of which are red phosphorus and white
phosphorus. The difference between the two is primarily the way in which the atoms are
2
See www.hq.nasa.gov/office/pao/History/Apollo204/summary.html or
www.hq.nasa.gov/office/pao/History/Apollo204/ for more information.
107
bonded in the solid phase. While each contains only phosphorus atoms, their reactivities
vary considerably: white phosphorus ignites spontaneously upon contact with oxygen in
air, while red phosphorus is more stable. If burning phosphorus touches the skin, it
leaves deep, severe burns. Burning phosphorus is often used for flares, tracer bullets, and
incendiary rounds.
2. When burned in air or oxygen, both P2O3 and P2O5 are formed.
3. Phosphorous oxides also react with water. See data sheet.
I. Commercially Available Oxygen
1. Gas Cylinder Safety. Oxygen and most other laboratory and industrial gases are often
stored under high pressure (2500 psi) in heavy metal cylinders. These cylinders should
always be strapped to a rigid support. If they should fall and the cylinder rupture (the
valves are especially vulnerable), the cylinder will literally behave as a rocket, causing
enormous damage.3 THIS IS TRUE FOR ALL COMPRESSED GAS CYLINDERS!!
(a)
(b)
(c)
(d)
Figure 4. (a) A typical gas cylinder. (b) Gas cylinder with the valve covered by a cap. (c) Gas
cylinder valve exposed. (d) Gas cylinder with a gas regulator connected to the valve.
2. Oxygen in Hospitals. Hospitals maintain large reservoirs of compressed oxygen to be
piped to individual rooms. The oxygen may be given to a patient via nasal tube
(cannula), mask, or tent, depending on physiological factors. Recall the burning of
Vaseline in pure O2; hospital patients should not use petroleum-based cosmetics or
ointments when pure oxygen is being administered.4
Pressurized O2 in gas cylinders contains no moisture; hence if a patient is given
pure O2 directly out of a tank, the dry gas will absorb water from the tissues that it comes
This was demonstrated on Episode 63 of MythBusters, “Air Cylinder Rocker” — see
http://dsc.discovery.com/fansites/mythbusters/mythbusters.html
4
Hospitals don’t like to lose their patients.
3
108
in contact with, causing severe drying of the mucous linings in the respiratory tract.
Before the oxygen is administered to a patient, the gas is usually bubbled through water
or a water-alcohol mixture, so that the gas is already moist before the patient inhales it.
Thus both the moisture content and the flow rate of the gas can be controlled.
J. Candle Trick
1. Wax is a complex mixture of large hydrocarbon molecules. The compound having the
formula C25H52 is typical of the compounds in a paraffin wax candle. Keep in mind that
the wax burns, not just the wick. See data sheet.
2. Your instructor will attempt to relight a candle by holding a match in the vapor trail
emanating from a smoldering candle. Cheap candles produce the best wax vapor trail.
K. Ammonium Dichromate Volcano
1. Ammonium dichromate, (NH4)2Cr2O7, resembles explosives in that it contains both a
source of fuel (NH4+ burns to form N2 and H2O) and an oxidizing agent (the chromium in
the +6 oxidation state in Cr2O72- is reduced to the +3 oxidation state in Cr2O3). It is
unlike explosives in that the reaction is slow so that the release of gases does not produce
a destructive shockwave. The density of Cr2O3(s) is drastically different from that of
ammonium dichromate. This combination results in a dramatic and entertaining
“volcano”-like effect. See data sheet.
L. Nitrogen Triiodide
1. This is another entertaining, but somewhat more dangerous reaction. This compound is
known as a shock-sensitive explosive: one that can be detonated with a slight movement
(e.g. the tickle of a feather). See data sheet.
M. Ether Trough
1. Ether (actually diethyl ether, C4H10O) is a highly volatile, extremely flammable
compound, formerly used as an anesthetic, and still frequently found in chemistry
laboratories and occasionally in hospitals. Fumes from open ether containers can travel
across a room (ether fumes are more dense than air) and act as a fuse from which the
original container can be ignited.
2. Since ether is less dense than water, and does not mix with water, burning ether will float
on any water added to extinguish the fire, causing the fire to spread. Ether fires must be
smothered.
3. If ether is allowed to remain in storage cans for several years, it forms peroxides, which
are explosively unstable at the slightest shock. Old stored cans of ether should be
handled with extreme caution! See data sheet.
109
LAB REPORT
Chemical Reactions I: Some Chemical Reactions Involving Oxygen
Name ________________________________
Date _________
Partner ________________________________
Section _________
Report Grade ______
Special Instructions
Classify the reactions that you write or balance in the data sheet using the following table.
Examples:
3Mg(s) + N2(g)  Mg3N2(s) would be classified as “Combination, Formation from
metal and nonmetal”
2NaCl  2Na + Cl2
would be classified as “Decomposition of a Binary
compound”
I. Combination Reactions
II. Decomposition
A. Formation or Element Combination
A. Binary compound
1. Metal + non-metal
B. Decomposition of Carbonates
2. Non-metal + non-metal
C. Decomposition of Oxoanions
B. Metal Oxide + Water
C. Non-Metal Oxide + Water
III. Single Displacement
IV. Double Displacement (metathesis, double replacement or partner exchange)
V. Combustion
VI. Miscellaneous (reactions not following any of the above patterns)
Part I. Demonstrations of Some Chemical Reactions Involving Oxygen
A. Lighting the Match
1. Safety Matches. Write a balanced equation for the reaction of diantimony trisulfide with
potassium chlorate, assuming that the products are diantimony trioxide, sulfur dioxide,
and potassium chloride.
__ Sb2S3 + __ KClO3  __ Sb2O3 + __ SO2 + __ KCl
Reaction Classification ________________________________________
2. Strike-Anywhere Matches.
Write a balanced equation for the reaction of
tetraphosphorus trisulfide with potassium chlorate, assuming that the products are
diphosphorus trioxide, sulfur dioxide, and potassium chloride.
__ P4S3 + __ KClO3  __ P2O3+ __ SO2 + __ KCl
Reaction Classification ________________________________________
110
B. Bunsen Burner
1. Reaction (Combustion). Write a balanced equation for the complete combustion of
methane and one for the complete combustion of ethane.
Reaction Classification ________________________________________
2.
Burner Flame. According to the diagram to the right,
a. The hottest part of the flame is ______.
b. The coolest part of the flame is ______.
3. Assuming that during this incomplete combustion methane and oxygen react to form only
elemental carbon [C(s)] and water, write a balanced equation in your data sheet showing
the incomplete combustion of methane.
Reaction Classification ________________________________________
C. Decomposition of Potassium Chlorate
1. Reaction. Write a balanced chemical equation for the decomposition of potassium
chlorate. Note that the symbol Δ above (or below) the arrow indicates heating of the
reactant(s), and the MnO2 above the arrow indicates its use as a catalyst.
2 ,Δ
MnO


Reaction Classification ________________________________________
111
E. Iron
1. Write chemical equations for the reactions of iron and oxygen:
formation of iron(II) oxide
Reaction Classification ________________________________________
formation of iron(III) oxide
Reaction Classification ________________________________________
F. Magnesium
1. Write a balanced chemical equation for the combustion of magnesium in air (with
oxygen). (Hint: The product will be an ionic compound.)
Reaction Classification ________________________________________
2. Write a balanced chemical equation for the reaction that occurs when the product of the
combustion of magnesium is mixed with water.
Reaction Classification ________________________________________
G. Sulfur
1. Write two chemical equations for the reaction of sulfur with oxygen, and give the names
of the products.
Reaction Classification ________________________________________
112
2. Write chemical equations for the reaction of water with each of the two sulfur oxides, and
give the names of the products.
Reaction Classification ________________________________________
3. What well known atmospheric pollution problem results from these reactions?
H. Phosphorus
1. Write two chemical equations for the reaction of phosphorus with oxygen, and give the
names of the products.
Reaction Classification ________________________________________
2. Write chemical equations for the reaction of water with each of the two phosphorus
oxides, and give the names of the products.
Reaction Classification ________________________________________
113
J. Candle Trick
1. Write the reaction of oxygen with C25H52.
Reaction Classification ________________________________________
K. Ammonium Dichromate Volcano
1. Write a balanced chemical equation for this reaction.
Reaction Classification ________________________________________
L. Nitrogen Triiodide
1. Write a balanced chemical equation for this reaction.
Reaction Classification ________________________________________
M. Ether Trough
1. Write the reaction of oxygen with diethyl ether (C4H10O).
Reaction Classification ________________________________________
114
Miscellaneous Questions
1. Which allotrope of phosphorus ignites spontaneously in air?
red ___
white ___
blue ___
yellow ___
green ___
2. In the decomposition of KC1O3, the MnO2 acts as a _________________.
3. Bottles full of oxygen should be stored
top up ___
top down ___.
Why?
4. Why is it especially important to avoid smoking, or using any kind of flame, in a hospital?
5. Any torch may be used either for natural gas-air or O2-acetylene.
True ___
False ___
6. Only oxygen cylinders need to be strapped to rigid supports.
True ___
False ___
7. Water should be used to extinguish ether fires.
True ___
False ___
8. Oxygen may be safely administered to patients directly
out of storage tanks with no further treatment.
True ___
False ___
9. Oxygen is a highly flammable gas.
True ___
False ___