8 Exp. 8: The Activity Series of Metals

Pre-Lab: Quiz on the following

Define:

1. Oxidation-reduction reaction
2. Oxidation half reaction
3. Reduction half reaction
4. Do the following half reactions represent oxidation or reduction?
   1. Na(s) à Na+ + e-
   2. Cu2+ + 2e- à Cu(s)
5. Which metal is more reactive if the following oxidation-reduction reaction occurs? Consider hydrogen a metal in this context.

2 HCl(aq) + Pb(s) à H2(g) + PbCl2(aq)

6. How will you avoid injury in this lab?

Introduction:

Over three quarters of the known elements can be classified as metals. In earth’s oxygen rich environment, however, most metals exist as cations in rocks, minerals, or dirt. The red rock features found in the Southwest United States and Australia and the red clay found in the Southeast United States and Eastern Africa contain iron oxide deposits, or iron cations. In their metallic or uncombined form, metals conduct both heat and electricity, and are familiar building, decorative, and conductive materials as they are both ductile and malleable. This reduced form of a metal is typified by a shiny appearance and is most likely what you think of as metallic. Coinage metals, such as gold and silver are readily reduced, and are found uncombined in nature. Alkali metals, such as sodium and potassium are readily oxidized, and so are less familiar in an uncombined form. These reactive examples are also lustrous or shiny when uncombined, however. Metal oxides also differ from porous iron oxides or rust, which allows for continual oxidation to thin dense oxide coatings on aluminum and chromium which protect from further oxidation to the dull grey oxides that forms on calcium or other alkali earth metals.

Metals differ significantly in their ability to lose electrons or oxidize. Reactivity among metals is defined as the ability to lose electrons and to form a cation. You may be familiar with the fact that gold is not very reactive, meaning it does not readily give up electrons, or oxidize, making it a good choice for jewelry, some of which has preserved for millennia. Calcium, on the other hand, is more readily oxidized, and does not last long in an uncombined state. Calcium, therefore, is more reactive, or easily oxidized than gold, and is more likely to be found as a cation combined with another element or elements.

Au(s) à Au3+ + 3e-                                                      Less Reactive                                                                          Reaction 1

Ca(s) à Ca2+ + 2e-                                                       More Reactive                                                                         Reaction 2

This relative reactivity gives rise to the activity series of metals. More reactive, or more easily oxidized metals are found higher on the series. Today you will combine cation solutions of metals as well as reduced (metallic) metal solids to observe their relative reactivity. You will create an activity series based on these observations.

Reactions one and two, listed above, are oxidation half reactions, and show electrons leaving a solid metal sample. The electrons, however, must go somewhere. A more complete oxidation-reduction or redox reaction accounts for the transfer of electrons.

3Ca(s) + 2Au3+ à 3Ca2+ + 2Au(s)                           Reaction          3

In this reaction, the more reactive calcium ends up as a cation in solution, and the less reactive gold ends up as a solid. Calcium, therefore, ranks higher in activity. Six electrons are transferred from solid calcium metal to gold cations. Calcium, therefore, is oxidized, or increases in charge, while gold is reduced, or decreases in charge. Reaction 3 can be divided into oxidation and reduction half reactions.

Oxidation Half Reaction:                               Ca(s) à Ca2+ + 2e-

Reduction Half Reaction:                               Au3+ + 3e- à Au(s)

Balanced Redox Reaction:                                       3Ca(s) + 2Au3+ à 3Ca2+ + 2Au(s)    Reaction 4

Note that the half reactions are multiplied by a coefficient in order to cancel the electrons. The overall redox reaction balances by mass and charge, as do the half reactions.

You probably know that gold is less reactive than calcium. To test the reactivity of these metals you would place a solid piece of calcium metal in a solution of gold cations, which is represented in Reaction 3. If the calcium is ionized into solution and solid gold precipitates, the reactions and assumptions listed are correct. A second testing method is to test the reverse of Reaction 3. If solid gold and magnesium cations in solution are combined, and no reaction occurs, magnesium ranks higher in activity than gold.

3Mg2+ + 2Au(s) à 3Mg(s) + 2Au3+ (No Reaction)                                   Reaction 5

Objective:

Deductively create a brief activity series of metals. Depict the reactions with half reactions and net ionic equations.

Safety:

• 6M HCl is caustic. Wear splash goggles, coats, and gloves
• Some of the reactions occur quickly, and generate heat

Waste/Housekeeping:

• Plan how much reagent you will need, and take a little more
• Dispense solutions into beakers and flasks, not into graduated cylinders at the common lab bench
• Don’t return solutions to the stock containers. It is better to waste than risk contamination.

Procedure:

In this lab you will obtain solutions containing the following cations: Cu2+(as a sulfate), Ag+(as a nitrate), Zn2+ (as a sulfate), and H+ (as a chloride). In addition, you will use solid samples of Cu, Zn, Mg, Ag, and Fe. By combining these solutions as described in the introduction, you will rank Mg(s), Fe(s), Ag(s), Zn(s) Cu(s), and H2(g) based on their relative activities.

You will also write oxidation and reduction half reactions and a balanced redox reaction. See example reaction 4. The relevant half reactions, written as oxidations are as follows:

Cu(s) à Cu2+ + 2e-

Zn(s) à Zn2+ + 2e-

Ag(s) à Ag+ + e-

H2(g) à 2H+ + 2e-

Mg(s) à Mg2+ + 2e-

Fe(s) à Fe2+ + 2e-

1. Dispense approximately 20 mL of the following metal cation solutions into small, labeled Erlenmeyer flasks or beakers: 0.1 M CuSO4, 0.1 M ZnSO4, 0.1 M AgNO3, 6.0 M HCl. Designate a beral pipette for each solution.
2. Design and arrange your test tube rack and test tubes so that each solid metal sample can be submerged in each solution. You don’t need to mix the same solid and cation, such as Cu2+ and Cu(s), for example. You may label test tubes with tape or can draw a key on paper. Design a way to track your reactions. The following example may help.

3. Add one small piece of solid metal (wire, ribbon, etc.) to the test tubes in your grid. Your observations will be qualitative; you will need just enough solid sample to observe a reaction. You may need to scrape or polish the copper or iron solid if an oxidized layer exists. *You cannot obtain a chunk of hydrogen. It is noticeably absent from the metal solids yet must be accounted for in your activity series. How do you evaluate the reactivity of H2(g)?
4. Fill your test tubes with a few mL of solution. You only need enough solution to cover the solid sample and observe what happens.
5. Observe the results. Some combinations react quickly; others may take several minutes. Pay attention to deposition of a solid, evolution of a gas, and temperature changes.

*Some helpful hints as you evaluate the results of your redox reactions:

• An unreacted combination yields valuable information.
• Deposition of a solid signifies the reduction of a metal.
• If iron is not present in the reaction combination, you are not observing rust or iron (II) oxide. If you see something that looks ‘rusty,’ what are you observing?
• H+ reduces to H2(g), what will you observe if hydrogen is less reactive than a metal?

You will find the most information and can corroborate your findings as you evaluate your reactions from an oxidation standpoint as well as a reduction standpoint. Did solid copper, for example, oxidize in Zn2+ solution? If so, did Zn(s) remain reduced or unreactive in the Cu2+ solution? Since you don’t have solid hydrogen, you will assign H2(g) on the activity series based on which metals reduce H+ in solution.

6. Assign your metal samples, not cations, from most reactive to least reactive.
7. Write an oxidation half reaction, reduction half reaction and a balanced redox reaction for each observed reaction. You don’t need to include unreacted combinations.

Post Lab:

1. Based on your lab observations:
   1. Which of the metals used in this lab is most likely to be found in an uncombined or metallic state in nature? Does this metallic state represent an oxidized or reduced state? Find and cite an example of the discovery of this metal in an elemental state.
   2. Which of the metals would be most likely to be found combined with other elements as a cation? Does this represent an oxidized or reduced state? Find and cite an example of the formula and use of this metal in a combined state?
2. Steel, primarily iron metal, is used in the construction of sea going ships and boats. Consider your reactivity data for iron and zinc.
   1. Zn metal is often placed in contact, either directly or through wires, with the steel hull of a ship. Provide a reasoned argument for why this is a widespread practice. Use the concepts that you learned in lab. Be specific in describing the movement of electrons.
   2. Will the zinc need to be replaced periodically? If so, what happens to the missing zinc?
3. Provide a real-world application of the metals activity series.

Your Report:

1. Complete the pre-lab and take the pre-lab quiz in Blackboard before attending lab.
2. Use a lab notebook to take notes and make observations.
3. Create an informal lab document to turn in at the beginning of the next lab. Your lab report should have your name, data arranged in graphs or tables where applicable, conclusive data such as an activity series. Relevant oxidation, reduction, and balanced redox reactions (see Reaction 4). Think about the main point of the lab, or the results that you worked for and be sure to include it or them. Writing the half and redox reactions as outlined (7) is part of your reporting.
4. Numbered responses to any post lab questions. You don’t need to re-write the questions.
5. Submit your document before your next lab appointment under the assignment tab on your laboratory Blackboard shell.

Sources:

Biello, D. The Origin of Oxygen in Earth’s Atmosphere. https://www.scientificamerican.com/article/origin-of-oxygen-in-atmosphere/ (accessed Jul 2, 2020).

(Adapted From) Dieckmann, G., Sibert, J. In An Atoms First Approach to the General Chemistry Laboratory; McGraw Hill: New York, NY, 2014; pp 3–12.