Activity Series

Names: Andrew Yaksic

Purpose: To become familiar with the relative activities of metals in chemicals.

Equipment: 7 small test tubes, test tube rack

Materials: 0.2 M Ca(NO3)2, 0.2 M Mg(NO3)2, 0.2 M Zn(NO3)2, 0.2 M Fe(NO3)3, 0.2 M FeSO4, 0.2 M SnCl4, 0.2 M Cu(NO3)2, 6 M HCl, 7 small pieces each of calcium, magnesium, zinc, iron, tin, and copper

Introduction: Chemical elements are usually classified by their properties into three groups: metals, nonmetals, and metalloids. Most known elements are metals. Some of the physical properties of metals include high thermal and electrical conductivity, high luster, malleability, and ductility. All common metals are solids at room temperature, except for mercury, which is a liquid.

The periodic table was developed in the 1860s by Dmitri Mendeleev. The periodic table is arranged according to the number of protons in one atom of each element. The columns of the periodic table contain elements with similar properties. The periodic table reveals a great deal of information about the elements based on the placement of elemental symbols.
Figure 15.1 below is a partial periodic table.

This periodic table has shaded elements. To the left of the shaded region are metals, except hydrogen, which is a nonmetal. To the right of the shaded region are nonmetals. In the shaded region are metalloids.

In this experiment, various metals will be placed into solutions of metal ions and into a solution of hydrochloric acid. The reactions will be observed and noted. An activity series will be constructed by ranking the elements according to which elements reacted with the most solutions of metal ions.

Two important terms for this experiment are oxidation and reduction. Oxidation occurs when an element loses electrons, becoming a cation, or positively charged ion. Reduction occurs when an element gains electrons, becoming an anion, or negatively charged ion. Oxidation cannot occur without reduction, so reactions where both processes occur are called oxidation-reduction reactions. This experiment will involve oxidation- reduction reactions and their outcomes.

Procedure:

A. Reactions of Metals with Acid

1. Six test tubes were filled with approximately 1 mL of 6 M HCl. A small piece of calcium, copper, iron, magnesium, tin, and zinc was added to each test tube. The contents of the test tubes were observed. Any changes were noted. The test tubes were emptied and rinsed.

B. Reactions of Metals with Solutions of Metal Ions

1. Seven test tubes were filled with approximately 1 mL of each of the following solutions, one solution per test tube: 0.2 M Ca(NO3)2, 0.2 M Mg(NO3)2, 0.2 M Zn(NO3)2, 0.2 M Fe(NO3)3, 0.2 M FeSO4, 0.2 M SnCl4, 0.2 M Cu(NO3)2.

2. One piece of each of the metals being tested (copper, calcium, iron, magnesium, tin, and zinc) was added to each of the test tubes, one metal at a time.

3. After one piece of one metal had been added to each of the test tubes containing solutions and any changes had been observed and noted, the test tubes were emptied, rinsed, and refilled with solution.

4. Steps 2 and 3 were repeated for each of the six metals.

Observations:

A. Reactions of Metals with Acid

B. Reactions of Metals with Solutions of Metal Ions

Results:

A. Reactions of Metals with Acid

This table outlines the reactions that took place as reported in Observations Part A. Only reactions that occurred are listed below. Net ionic equations remove spectator ions from the chemical equations.

B. Reactions of Metals with Solutions of Metal Ions

This table outlines the reactions that took place as reported in Observations Part B. Only reactions that occurred are listed below. Net ionic equations remove spectator ions from the chemical equations.

C. Relative-Activity Series

This activity series indicates which elements will replace other elements in single replacement (metathesis) reactions. It is ordered from most reactive element to least reactive element. Elements to the left will replace elements to the right.

Ca → Mg → Zn → Fe → Sn → H → Cu

Discussion: There were several possible sources of error in this experiment. Some errors were human errors and were the result of lack of experience. Other errors were chemical. For example, if any of the samples of ionic solutions were contaminated, then the reaction would have been altered by possibly entering another ion into the reaction. The oxide coating on some of the metals, such as calcium, may have delayed or prevented reactions from occurring. This causes discrepancies between what should have happened in a given reaction and what was observed. A major human source of error is ending a reaction too soon. Some reactions take several minutes to take place, and others produce only slight bubbling. Though a reaction is occurring, an untrained eye may disregard the signs of a reaction. Ample time must be allowed for the reaction to occur.

The theory associated with this experiment is the atomic theory, more specifically electron configuration. Different electron configurations, when combined with different ionization energies as required, produce ions of different charges. Valence electrons determine the charge of the ion and the ionization energy quantity. The structure of the atom is responsible for the electron configuration. This experiment verifies that certain electron configurations are more likely to react than others.

Ramifications of this experiment are broad. Laboratory experience was gained. For example, some of the violent reactions demonstrated the necessity of checking labels and being careful while mixing chemicals to prevent dangerous reactions. Some specific ramifications of this experiment include applying the activity series when separating metal from any compound containing that metal. For example, chemical separation of a metal from its oxide is possible by placing it in a solution of a more active ion. Batteries can be created by deferring the flow of electrons from a less reactive ion to a more reactive ion through a wire. Other specific industrial applications of this experiment exist, such as the production of metal from other common metal compounds.

Questions:

1. Calcium should be the most reactive toward oxygen because it is the most reactive of the metals tested.

2. The oxide that would be expected to be thermally unstable and to decompose according to 2MO ∆ 2M + O2 is copper(II) oxide, CuO. Cu is the least reactive of the elements surveyed in this experiment, so heating it would decompose the bonds between copper and oxygen.

3. The following reactions would occur:
2Na + 2H2O → 2NaOH + H2
4Na + O2 → 2Na2O
2Na + 2HCl → 2NaCl + H2
Na + Ca2+ → NR

4. Fe2+ is more reactive than Fe3+. Fe2+ will react with any ion under it in the activity series that has a charge of 2-. Fe3+ will only react with ions under it in the activity series with a charge of 3-. Since many transition metals and all Group 2 elements have ions with charges of 2-, while few elements have ions with charges of 3-, Fe2+ has many more options for reactions than Fe3+.

5. Based on the number of reactions that occurred, aluminum would fall between magnesium and zinc on the activity series.

6. In each of the following oxidation-reduction reactions, substances are oxidized and reduced by oxidizing and reducing agents. They are listed in the following table:

Conclusion: This experiment was successful. Familiarity with relative activities of metals in chemical equations was gained while sharpening laboratory skills. The activity series was also derived correctly.