Chemistry 151
Lab 11: Oxidation-Reduction Reactions
Last Updated Dec. 2012
Introduction
Oxidation-reduction (“redox”) reactions make up a large and diverse part of chemical systems. A few
examples include the following.
Electron transfers between elements and monatomic ions
Perhaps the simplest type of redox reaction is one where electrons are transferred between an element and a
monatomic ion, such as the reaction between nickel (II) ion and magnesium.
Ni2+(aq) +Mg(s)  Ni(s) + Mg2+(aq)
Here, nickel (II) is undergoing reduction because it gains electrons and magnesium is undergoing
oxidation because it loses electrons. Nickel (II) is also referred to as the oxidizing agent, as it is the
chemical removing electron from magnesium. Likewise, magnesium can be referred to as the reducing
agent, since it’s the chemical donating the electrons to nickel (II).
However, combining any ion with an element doesn’t necessarily guarantee a reaction will occur. For
example, the reverse of the above reaction, between magnesium ion and nickel, will not readily occur.
Ni(s) + Mg2+  no reaction
From this, we can conclude that nickel (II) is a stronger oxidizing agent than magnesium ion, and that
magnesium metal is a stronger reducing agent than nickel. Such ideas will be discussed in greater detail in
General Chemistry II.
Reactions between metals and acids
If a metal is a strong enough reducing agent, it will react with an acid to produce hydrogen gas, such as the
reaction between iron and sulfuric acid
Fe(s) + H2SO4(aq)  FeSO4(aq) + H2(g)
The stronger the reducing agent, the more quickly this reaction will occur, which in this case can be
observed by the increased production of gas.
Reactions with water
In this course water is routinely used as the solvent for aqueous reactions, but water can undergo reactions
of its own, many of which are redox reactions. For example, sodium will react vigorously with water to
produce hydrogen gas and sodium hydroxide.
2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Notice that this reaction is similar to the reaction between a metal and an acid (think of water as “HOH”
instead of H2O). Water’s ability to behave as an acid is another concept that will be discussed further in
General II.
Combustion Reactions
Combustion reactions involve the conversion of elemental oxygen, which has an oxidation number of zero,
to either a metal or nonmetal oxide, where oxygen will have an oxidation number of –2. Therefore, all
combustion reactions are also redox reactions. The combustion of methane (CH 4) is shown below.
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Peroxides and bleaches
Two common household oxidizing agents are bleach (sodium hypochlorite, NaClO) and hydrogen peroxide
(H2O2). When stains are removed with bleach, colored compounds are oxidized into colorless ones as
hypochlorite is reduced to chloride ion. H2O2 will also act as an oxidizing agent, as seen in the following
reaction with iron (II).
2Fe2+(aq, colorless) + H2O2(aq) + 2H+(aq)  2Fe3+(aq, yellow) + 2H2O(l)
Indicator solutions
Most reagents that are used as chemical tests or indicators are oxidizing or reducing agents. Benedict’s
reagent, for example, is a common reagent used for testing diabetes. The reagent is a basic solution of
copper (II) that is reduced to copper (I) in the presence of glucose. The copper (I) forms as Cu 2O, a red
solid.
2Cu2+(aq, blue) + C6H12O6(aq) + 5OH-(aq)  Cu2O(s, red) + C6H11O7-(aq) + 3H2O(l)
Early versions of breathalyzer tests used potassium dichromate. Dichromate compounds, which are usually
yellow to orange in color, are reduced to a green chromium (III) compound in the presence of alcohols,
such as ethanol (CH3CH2OH).
3CH3CH2OH(aq) + 2Cr2O72-(aq, orange) + 16H+(aq)  3CH3COOH(aq) + 2Cr3+(aq, green) + 11H2O(l)
Here, ethanol is converted to acetic acid (CH3COOH), the main ingredient in household vinegar.
In this lab you will observe a series of redox reactions, including some of the ones mentioned above. You
will also compare the strengths of different oxidizing and reducing agents.
Procedure
Part I: Reactions Between Elements and Monatomic Ions
A. Metals and Metal Ions
Use a set of small test tubes to prepare the mixtures outlined in the table below. In each case, use
approximately 2 mL of solution and 0.2-0.4 g of metal. If needed, clean the metal with sandpaper. Allow
the solution to set for 5 minutes then record your observations.
Sample
1
2
3
4
5
6
Solution (0.2 M)
CuSO4
ZnSO4
CuSO4
FeSO4
FeSO4
ZnSO4
Metal
Fe (steel wool)
Fe (steel wool)
Zn
Zn
Cu
Cu
B Nonmetals and Nonmetal Ions
1. Transfer 2 mL of 0.1 M potassium iodide to a test tube
2. Add 3-4 drops of bromine solution (the solvent is cyclohexane, which is immiscible with water, so it
will form two layers).
3. Cover the opening of the test tube with a small piece of Parafilm, place your thumb over the opening
and gently shake to mix.
4. Add a few drops of starch (which turns blue in the presence of I 2) to the test tube and mix.
5. Repeat steps 1-3 using 0.1 M potassium bromide and iodine solution.
Part II: Reactions Between Metals and Acids
1. Label three test tubes 1-3. Place a small piece (~0.2 g) of zinc in test tube #1, ~1-2 cm of magnesium
ribbon in #2, and a smaller piece (~0.2 g) of copper in #3. Clean each metal with sandpaper, as
needed.
2. Add 2 mL of 10% hydrochloric acid to each test tube and record your observations.
Part III: Reactions Between Metals and Water
Repeat Part II, substituting deionized water for hydrochloric acid.
Part IV: Combustion Reactions
A. Combustion of Cellulose
1. Attach a ring clamp to a ring stand and set a clay triangle on the ring clamp.
2. Obtain piece of filter paper, approximately 5.5 cm in diameter and crumble it so that it will fit inside a
crucible. Place the crucible on the clay triangle.
3. Heat the crucible with a Bunsen burner and record your observations. Make sure the paper doesn’t
come in direct contact with the flame.
B. Combustion of Magnesium
1. Obtain a small piece of magnesium (~ 1 cm in length) and clean it with sandpaper to give it a shiny
surface.
2. Holding the magnesium with a pair of tongs, use the flame of your Bunsen burner to light the metal.
Record your observations. [Safety Note: Once the magnesium begins to burn, don’t look directly at
the resulting light.]
Part V: The Reaction of Fe2+ with Hydrogen Peroxide
1. Transfer a few crystals of FeSO4 to a medium test tube and dissolve with 2 mL of water.
2. Add 2-3 drops of 3% hydrogen peroxide and record your observations.
Part VI: The Reaction of Cu2+ (Benedict’s Solution) with Glucose
1. Assemble a hot-water bath as follows:
a. Place a Bunsen burner on a ring stand.
b. Attach a metal ring to the ring stand then place a wire gauze on the ring.
c. Half-fill a 250-400 mL beaker with water and set it on the wire gauze. Place a second ring
clamp around the beaker to reduce the chance of an accidental spill.
d. Begin heating the water with the beaker.
2. Place a small amount of glucose in a medium test tube (enough to cover the bottom) and dissolve with
2-4 mL of water.
3. Add 2 mL of Benedict’s solution to the test tube and heat the solution in the hot-water bath. Record
your observations.
Part VII: The Reaction of Dichromate with Ethanol
1. Add a few crystals of potassium dichromate to a medium test tube and dissolve with 2-3 mL of 3.0 M
sulfuric acid.
2. Add 2-4 mL of ethanol and gently swirl to mix. Record your observations.
Waste Disposal
All waste should be disposed of in the waste hood.
Name: _____________________________
Section: ________
Data
Part I: Reactions Between Elements and Monatomic Ions
A. Metals and Metal Ions
Sample
Solution (0.2M)
Metal
1
CuSO4
Fe (steel wool)
2
ZnSO4
Fe (steel wool)
3
CuSO4
Zn
4
FeSO4
Zn
5
FeSO4
Cu
6
ZnSO4
Cu
Results
Based on your results, rank the three cations in order of increasing oxidizing strength
(weakest oxidizing agent) _____ < _____ < _____ (strongest oxidizing agent)
Rank the three metals in order of increasing reducing strength
(weakest reducing agent) _____ < _____ < _____ (strongest reducing agent).
Explain your rankings.
B Nonmetals and Nonmetal Ions
Record your observations in the space below.
From your observations, which is the stronger oxidizing agent, bromine or iodine?
_______
Which is the stronger reducing agent, bromide or iodide?
_______
Part II: Reactions Between Metals and Acids
Sample
Metal
1
Zn
2
Mg
3
Cu
Results
Part III: Reactions Between Metals and Water
Sample
Metal
1
Zn
2
Mg
3
Cu
Results
Rank the three metals in order of increasing reducing strength
(weakest reducing agent) _____ < _____ < _____ (strongest reducing agent).
Part IV: Combustion Reactions
Record your observations in the space below.
A. Combustion of Cellulose
B. Combustion of Magnesium
Part V: The Reaction of Fe2+ with Hydrogen Peroxide
Record your observations in the space below.
Part VI: The Reaction of Cu2+ (Benedict’s Solution) with Glucose
Record your observations in the space below.
Part VII: The Reaction of Dichromate with Ethanol
Record your observations in the space below.
Name: _____________________________
Section: ________
Post-Lab Questions
1. Are the results of Parts I – III consistent with each other? Explain.
2. A chemist accidentally leaves a bottle of iron (II) chloride open overnight. The next morning the once
colorless solution is now light yellow. What happened?
3. You will occasionally hear stories of people suddenly undergoing “spontaneous combustion” (bursting
into flaming without an ignition source). Based you observations in this lab, comment on the likelihood of
such events.
4. Ovens with a self-cleaning feature offer a convenient way to remove grease and other cooking byproducts. Based on your observations today, how do you think they work?
5. A popular item during Independence Day (and other fireworks-based celebrations) are “sparklers” that
emit a bright light when lit. Given today’s observations, what are these sparkler most likely made of?
6. Cooks will usually avoid using aluminum foil or aluminum cookware when prepare acidic foods. Give
two reasons why?
7. In the Introduction section, the equation for the reaction between iron (II) and hydrogen peroxide also
includes hydrogen ion as a reactant, but only solutions of FeSO4 and H2O2 was mixed in a test tube. If an
acid is required for the reaction to work, what was the source of H + for this reaction.
(No, we didn’t add it to the peroxide bottle beforehand; it’s the same stuff you’d buy at a store)
Name: _____________________________
Section: ________
Pre-Lab Questions
1. Based on what was said in the Introduction, how many of the mixtures in Part I should lead to a
reaction?
2. Identify the oxidizing and reducing agent in the following equations. Explain your answers.
a) 2Na(s) + 2H2O(l)  2NaOH(aq) + H2(g)
Oxidizing agent: _______
Reducing agent: _______
b) Al3+(aq) + 3Li(s)  3Li+(aq) + Al(s)
Oxidizing agent: _______
Reducing agent: _______
b) CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
Oxidizing agent: _______
Reducing agent: _______
3. What is the formula of cellulose? What are the predicted products when it undergoes combustion?
4. Write balanced equations for each of the following reactions (assuming one occurs). Phases aren’t
required.
a) Iron (II) ion and copper
b) Magnesium and hydrochloric acid
c) Magnesium and water
d) The combustion of magnesium