Chapter 10. Laboratory: Reduction-Oxidation (Redox) Reactions

A reduction-oxidation reaction or redox reaction is a chemical reaction during which the oxidation state of two or more of the reactants change. Note that the word “oxidation” in this context has nothing to do with oxidation in the sense of combining with oxygen. Most redox reactions do not involve oxygen at all.

The oxidation state of an atom or ion can be thought of as its electrical charge. In simple terms, a neutral atom is in oxidation state 0. An ion with a positive charge has a positive oxidation state overall (although one or more of the component atoms in that ion may have nonpositive oxidation states. For example, if a neutral sodium atom (oxidation state 0, Na⁰) loses an electron, it has been oxidized to oxidation state 1+, and becomes a Na⁺ ion. Conversely, a carbonate ion, CO_3^2–, is made up of one carbon atom in oxidation state 4+ and three oxygen atoms, each in oxidation state 2–, for a net charge (oxidation state) of 2–.

Oxidation in the context of a redox reaction refers to the oxidation state of an atom being increased. Conversely, reduction refers to the oxidation state of an atom being reduced. We saw an example of a redox reaction in the preceding chapter, when we reacted aluminum metal with hydrochloric acid to produce aluminum chloride and hydrogen gas. The balanced equation for that reaction is:

2 Al(s) + 6 HCl(aq) → 2 AlCl₃(aq) + 3 H₂(g)

or, with the oxidation states of each species as a superscript:

2 Al⁰(s) + 6 H⁺(aq) + 6 Cl^–(aq) → 2 Al³⁺(aq) + 6 Cl^–(aq) + 3 H₂⁰(g)

In this reaction, two atoms of aluminum metal (oxidation state 0) each lose three electrons and are oxidized to aluminum ions (oxidation state +3). The six electrons lost by the two aluminum atoms combine with six hydrogen ions (oxidation state +1), reducing the hydrogen ions to form six atoms of hydrogen gas (oxidation state 0), which then combine to form three molecules of diatomic hydrogen gas (still in oxidation state 0). The six chloride ions (oxidation state –1) do not change oxidation states, and so are the same on both sides of the equation.

Many chemical reactions do not involve any change in oxidation state. For example, when we reacted aqueous solutions of copper(II) sulfate and sodium hydroxide to form a precipitate of copper(II) hydroxide and a solution of sodium sulfate, no changes in oxidation state occurred:

CuSO₄(aq) + 2 NaOH(aq) → Cu(OH)₂(s) + Na₂SO₄(aq)

or, showing the oxidation states:

Cu²⁺(aq) + SO₄^2–(aq) + 2 Na⁺(aq) + 2 OH^–(aq) → [Cu²⁺(OH^–)₂](s) + 2 Na⁺(aq) + SO₄^2–(aq)

On both sides of the equation, copper is in the +2 oxidation state, sulfate –2, sodium +1, and hydroxide –1.

Oxidation state 0 is the neutral form of any element, such as (at standard conditions) aluminum in its metallic form or chlorine in its gaseous form. An element in oxidation state 0 possesses the same number of negatively charged electrons in its shell as the number of positively charged protons in its nucleus, leaving a net charge of 0.

If a chlorine atom at oxidation state 0 gains an electron, that chlorine atom becomes a chlorine ion with a charge of –1, and an oxidation state of –1. Conversely, if a chlorine atom at oxidation state 0 loses an electron, that chlorine atom becomes a chlorine ion with a charge of +1, and an oxidation state of +1. With the exception of the noble gases—helium, neon, argon, krypton, xenon, and radon, all in column 18 of the periodic table—all elements are commonly found in at least one oxidation state other than 0. Some elements have only one or two known nonzero oxidation states. Other elements are commonly found in many oxidation states. For example, in its natural form and in its many compounds, carbon can be found in oxidation states –4, –3, –2, –1, 0, +1, +2, +3, and +4.

Oxidation and reduction also apply to atoms that are in oxidation states other than 0. For example, iron is found in three oxidation states. Metallic iron is in oxidation state 0. Metallic iron can be oxidized to the Fe²⁺ ion, also called the ferrous ion or the iron(II) ion. The Fe²⁺ ion can be further oxidized to the Fe³⁺ ion—also called the ferric ion or the iron(III) ion—or it can be reduced to metallic iron. The Fe³⁺ ion can be reduced to the Fe²⁺ ion by gaining one electron, or all the way to Fe⁰ (iron metal) by gaining three electrons.

Ions of a particular element in different oxidation states can have very different chemical and physical properties. For example, solutions that contain vanadium ions in different oxidation states show a rainbow of colors. Solutions of V⁵⁺ ions are yellow; those of V⁴⁺ are blue; V³⁺, green; and V²⁺, violet. Solutions of ions in different oxidations states of chromium, manganese, and other elements show similar color differences.

It’s important to understand that in a redox reaction oxidation and reduction are two sides of the same coin. If one atom is oxidized during the reaction, another must be reduced, and vice versa. Furthermore, the total number of electrons lost by the atom or atoms being oxidized must equal the total number of electrons gained by the atom or atoms that are being reduced. Charge must be conserved. For example, when aluminum reacts with the H⁺ ions from hydrochloric acid to form Al³⁺ ions, the three electrons lost when one aluminum atom is oxidized reduce three H⁺ ions to H⁰ atoms.

The laboratory sessions in this chapter explore various aspects of redox reactions.

Everyday Redox Reactions

Redox reactions, including some that we literally couldn’t live without, are common in everyday life. Here are just a few examples:

• Metals are usually mined in the form of an oxide, carbonate, or other salt of the metal. Refining metals—whether by smelting, electrolysis, or some other method—uses redox reactions; specifically, reduction of the metal in a positive oxidation state to the neutral metallic form.

• Water purification uses redox reactions to oxidize colored, bad-tasting, or otherwise objectionable impurities to forms that are safe and acceptable.

• Bleaches used in laundry, papermaking, and other processes all depend upon redox reactions to oxidize colored stains and impurities to colorless compounds.

• Photographic film works by a redox reaction initiated by light. When photons strike the tiny grains of silver bromide and other silver halides present in the emulsion, they cause silver ions to be reduced to microscopic specks of metallic silver. During development, those tiny specks of silver catalyze reduction of the silver halide grains surrounding them. After washing away the unreduced silver, what remains are the tiny black grains that make the picture.

• Corrosion and rusting are examples of redox reactions that are familiar to everyone. In particular, unprotected iron and steel are prone to react with atmospheric oxygen in the presence of moisture to form iron oxides, familiarly called rust. Many forms of protection against corrosion take advantage of redox reactions as well. For example, plating iron with zinc (called galvanizing) protects the underlying iron, because it will not oxidize while in contact with the more active metal.

What about Fe⁴⁺ and Fe⁺?

It might seem that an Fe³⁺ ion could lose one more electron to produce an Fe⁴⁺ ion. That doesn’t happen, because of the configuration of iron’s electron shell. Only tightly bound electrons remain when iron is in the +3 state. Similarly, it might seem that Fe⁰ could lose one electron (or Fe²⁺ gain one electron) to form Fe⁺. That doesn’t happen, because the electron shell configuration of Fe⁺ is unstable. Any Fe⁺ ions that fleetingly form during an oxidation-reduction reaction immediately shed an electron to form Fe²⁺ ions or grab an electron to form Fe⁰.

Keeping Oxidation and Reduction Straight

Many beginning chemists have trouble remembering whether oxidation means gaining or losing an electron. Generations of chemists have kept this straight using the mnemonic LEO the lion says GER, which you can expand to Lose Electron = Oxidation and Gain Electron = Reduction.

Laboratory 10.1: Reduction of copper ore to copper metal

It may surprise you to learn that at least one redox reaction was discovered and used thousands of years ago. Like many discoveries in chemistry, this one was purely accidental. We can speculate about what happened. One evening long ago, some of our remote ancestors gathered around a fire built in a circle of stones. Those stones happened to be chunks of a copper-bearing ore, probably copper carbonate.

All of the required elements were present. The copper ore provided the copper ions. The charred wood provided the carbon in the form of charcoal. The fire provided the heat needed to initiate and sustain the reaction. Here’s what happened:

2 CuCO₃(s) + C(s) → 2 Cu(s) + 3 CO₂(g)

or, showing the oxidation states:

2 [Cu²⁺[C⁴⁺(O^2–₃)]^2–](s) + C⁰(s) → 2 Cu⁰(s) + 3 [C⁴⁺(O^2–₂)](g)

As the fire burned during the night, some of the copper ore and charcoal disappeared and, as if by magic, were replaced by a shiny puddle of metallic copper. Chances are that this had happened many times before and no one thought much about it, if they even noticed. But this time someone did notice and started wondering. What was this shiny new substance? Could it be used for anything? Can we get more of it just by heating rocks?

Required Equipment and Supplies

	goggles, gloves, and protective clothing
	balance and weighing paper
	ring stand
	support ring
	clay triangle
	crucible and lid
	crucible tongs
	gas burner
	copper carbonate (~5 g)
	activated charcoal (~3 g)

At the time, metallic copper was known, but the only source was small, rare deposits of natural metallic copper. The discovery that copper metal could be obtained by heating copper ore in conjunction with charcoal, a process called smelting, marked the end of the Stone Age.

In this lab, we’ll recreate that first accidental smelting of copper ore to copper metal.

Caution

This experiment uses the hot flame provided by a gas burner. Be careful with the flame, have a fire extinguisher readily available, and use care in handling hot objects (a hot crucible looks exactly like a cold crucible). Wear splash goggles, gloves, and protective clothing.

Procedure

1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.

2. Weigh a clean, dry crucible and lid and record the mass to 0.01 g on line A of Table 10-1.

3. Weigh about 5.0 g of copper carbonate and record the mass to 0.01 g on line B of Table 10-1.

4. Weigh about 3.0 g of activated charcoal and record the mass to 0.01 g on line C of Table 10-1.

5. Mix the copper carbonate and activated charcoal thoroughly, and transfer them to the crucible.

6. Weigh the crucible and contents and record the mass to 0.01 g on line D of Table 10-1.

7. Set up your tripod stand or ring stand with the clay triangle, put the covered crucible in place, and place the gas burner to direct the hottest part of its flame at the bottom of the crucible.

8. Heat the crucible strongly for at least 15 minutes.

9. Allow the crucible to cool to room temperature. You can check the temperature of the crucible by placing your hand near, but not touching, its surface to check for radiated heat.

10. Reweigh the crucible and contents and record the mass to 0.01 g on line E of Table 10-1.

11. Calculate the mass loss (line D minus line E) and enter the value to 0.01 g on line F of Table 10-1.

Figure 10-1. Mixed copper carbonate and charcoal before smelting

Substitutions and Modifications

• You may substitute a test tube and clamp or test tube holder for the crucible, support stand, and clay triangle. If you do use a test tube, choose one that you don’t mind sacrificing, as it will probably not be reusable.

• You may substitute a pair of pliers or similar gripping tool for the crucible tongs. (Be careful. Crucibles are physically fragile.)

• You may substitute a propane torch for the gas burner.

• If you do not have copper carbonate, you can make it by dissolving about 11 g of copper sulfate pentahydrate in about 50 mL of water and adding a solution of about 5 g of anhydrous sodium carbonate in about 20 mL of water. Filter the precipitate of copper carbonate, rinse it two or three times with water, and dry it before use.

• You may substitute an equivalent weight of any form of carbon for the activated charcoal. If you decide to use 3 g of diamonds, send them to me instead, and I’ll send you a whole kilogram of activated charcoal in exchange.

Disposal

The waste material from this lab can be ground to powder and flushed down the drain with plenty of water.

Table 10-1. Reduction of copper ore to copper metal—observed and calculated data

Item	Data
A. Mass of crucible + lid	______.______ g
B. Mass of copper carbonate	______.______ g
C. Mass of carbon (charcoal)	______.______ g
D. Mass of crucible, lid, and reactants (A + B + C)	______.______ g
E. Mass of crucible, lid, and products	______.______ g
F. Mass loss (D – E)	______.______ g

Review Questions

Q:	Q1: Observe and describe the appearance of the products after heating the crucible. (If necessary, use a mortar and pestle to crush the product.) Which products can you identify visually? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________
Q:	Q2: From the balanced equation for this reaction and knowing that the molar masses of carbon and copper(II) carbonate are 12.01 g/mol and 123.56 g/mol respectively, which reactant was in excess, and by how much? Why might we have chosen to use this reactant in excess? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________
Q:	Q3: Assuming that the carbon reacted stoichiometrically with the copper(II) carbonate and that none of the carbon reacted with atmospheric oxygen to form carbon dioxide, calculate the expected mass loss. How does the calculated mass loss compare with the mass loss you observed? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________
Q:	Q4: More than 2,000 years passed from when copper ore was first smelted in about 3500 BC until iron ore was first smelted in about 1200 BC, and another 3,000 years passed before aluminum metal was successfully reduced from aluminum ore in the nineteenth century? Why did it take so long? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________

Laboratory 10.2: Observe the Oxidation States of Manganese

Many elements can exist in several oxidation states, which may differ noticeably in color. In the introduction, we mentioned vanadium as one such element. Manganese is another element whose various oxidation states have different colors, including MnO₄^– (+7, violet), MnO₄^2– (+6, green), MnO₂ (+4, orange), Mn₂O₃ (+3, violet), and Mn²⁺(+2, pale pink).

In this laboratory, we’ll begin with a solution of potassium permanganate, in which manganese is in the +7 oxidation state, and reduce the oxidation state to observe the color changes associated with various oxidation states of manganese. To achieve these changes in oxidation state, we’ll react potassium permanganate with sodium bisulfite in neutral, basic (sodium hydroxide), and acid (sulfuric acid) solutions, which are represented by the equations shown in Table 10-2.

Required Equipment and Supplies

	goggles, gloves, and protective clothing
	test tubes (4)
	test tube rack
	graduated cylinder, 10 mL
	dropper or beral pipette
	stirring rod
	potassium permanganate (KMnO4) solution, 0.01 M (50 mL)
	sodium hydroxide solution, 1.0 M (8 mL)
	sulfuric acid solution, 1.0 M (6 mL)
	sodium bisulfite (NaHSO3) solution, 0.01 M (as required)

Caution

Solutions of sodium hydroxide and sulfuric acid are corrosive. Wear splash goggles, gloves, and protective clothing.

Table 10-2. Oxidation states of manganese

Solution	Oxidation state	Ion	Color	Equation
Neutral	+4	MnO2	orange	2 MnO4–(aq) + 3 HSO3–+ OH–Đ 2 MnO2(s) + 3 SO42–(aq) + 2 H2O(l)
Basic (sodium hydroxide)	+6	MnO42–	green	2 MnO4–(aq) + HSO3–+ 3 OH–Đ 2 MnO42–(aq) + SO42–(aq) + 2 H2O(l)
Acid (sulfuric acid)	+2	Mn2+	pink	2 MnO4–(aq) + 5 HSO3–+ H+Đ 2 Mn2+(aq) + 5 SO42–(aq) + 3 H2O(l)

Substitutions and Modifications

• If you do not have a 0.01 M solution of potassium permanganate, make one up by diluting your potassium permanganate stock solution as required. Alternatively, dissolve 0.08 g of potassium permanganate in about 40 mL of water and make it up to 50 mL. (The exact concentration is not critical for this lab.)

• If you do not have a bench solution of 1.0 M sodium hydroxide, dissolve 4.00 g of sodium hydroxide in 90 mL of water and make it up to 100 mL. (Be careful; this reaction is extremely exothermic.)

• If you do not have a bench solution of 1.0 M sulfuric acid, carefully add 5.6 mL of concentrated (18 M) sulfuric acid to 90 mL of water and make it up to 100 mL. (Be careful; this reaction is extremely exothermic.).

• If you do not have a 0.01 M solution of sodium bisulfite, dissolve 0.10 g of sodium bisulfite in 90 mL of water and make up the solution to 100 mL.

Procedure

1. If you have not already done so, put on your splash goggles, gloves, and protective clothing.

2. Label four test tubes A through D, place them in the test tube rack, and transfer about 10 mL of 0.01 M potassium permanganate solution to each of the tubes.

3. Test tube A will contain only the potassium permanganate solution, and will serve as the control and example of manganese in oxidation state +7. Record the color of this solution on line A of Table 10-3.

4. With stirring, slowly add 0.01 M sodium bisulfite solution to test tube B until no further color change occurs. Record the color and the presence or absence of any precipitate on line B of Table 10-3.

5. Add 8 mL of 1.0 M sodium hydroxide solution to test tube C.

   

   Figure 10-2. Oxidation states of manganese

   Disposal

   You can flush the dilute solutions used in this lab down the drain with plenty of water.

6. With stirring, slowly add 0.01 M sodium bisulfite solution to test tube C until no further color change occurs. Record the color and the presence or absence of any precipitate on line C of Table 10-3.

7. Add 6 mL of 1.0 M sulfuric acid solution to test tube D.

8. With stirring, slowly add 0.01 M sodium bisulfite solution to test tube D until no further color change occurs. Record the color and the presence or absence of any precipitate on line D of Table 10-3.

9. Comparing the appearances of the contents of the four test tubes, decide which oxidation state of manganese is represented by the contents of each test tube and enter that value on the appropriate line of Table 10-3.

Table 10-3. Observe the oxidation states of manganese—observed and calculated data

Test tube	Color	Precipitate?	Oxidation state
A. Potassium permanganate only
B. Potassium permanganate + sodium bisulfite
C. Potassium permanganate + sodium hydroxide + sodium bisulfite
D. Potassium permanganate + sulfuric acid + sodium bisulfite

Optional Activities

If you have extra time, consider performing this optional activity:

• Potassium permanganate has many uses, nearly all of which are based on the fact that it is an extremely strong oxidizer. You can demonstrate this by placing several layers of filter paper or paper towel in the bottom of a Petri dish or large beaker and pouring the excess 0.01 M potassium permanganate solution you prepared for this lab onto the paper. Initially, the paper will have the strong violet color of the permanganate ion (+7). As the permanganate ion oxidizes the cellulose in the paper, it is reduced from the +7 oxidation state to the green +6 oxidation state. If you allow the reaction to continue long enough, the +7 permanganate ion is eventually reduced to the +4 state as orange/brown manganese dioxide (MnO₂).

Review Questions

Q:

Q1: What are the final oxidation states of manganese in test tubes A, B, C, and D? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________

Q:

Q2: What are the oxidation states of each atom in the following species: NH3, H2O, H2O2, K2Cr2O7, NaNO3, FeCl2, and AgNO3? ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________ ____________________________________________________________________