Concepts to explore: Observe an oxidation reduction reaction Use the properties of a reaction product to verify its identity Rank the reactivity of certain metals in a weak acid, and compare it to their order in the Activity Series of Metals Introduction Have you ever wondered why some rings will turn a finger green, while others won t? The ring you just bought a couple of weeks ago is already turning your finger green. Another ring that you have had for years still looks almost new. Why is this? Knowing how reactive different metals are is extremely important. This property helps us to decide which metal to use for a particular application. Some metals are so nonreactive that they will not react with even the strongest acids. Many will only react with certain acids, while others are so reactive that they will react with water. This is why jewelry is often made out of gold and not ordinary iron. Iron is more reactive, and will gradually rust when exposed to the oxygen and moisture in the air. Gold, on the other hand, will not react easily with anything. You know that gold is less reactive than iron, but what about other metals? How do you know which ones are more reactive? The activity series of metals places elements in order by their reactivity. Metals higher on the list give up their valence electrons more easily than the metals below them, meaning that any metal on this list will displace a metal below it in a reaction. This can easily be observed if the less reactive metal is in an aqueous solution during the reaction, as it will precipitate out of the solution when replaced. A partial activity series list is shown in Figure 2. As a part of this lab, you will observe the reaction of an acid and zinc metal and compare it to the reaction of the same acid with iron metal. The reaction between the zinc and the acid can be written as follows: Figure 1: The Statue of Liberty in New York Harbor is made of copper, and was originally the color of a penny. The statue has gradually acquired its green color due to the natural oxidation of copper when exposed to air and water. 2 H + ( aq) + Zn (s) Zn 2+ (aq) + H 2(g) In this reaction, the zinc transfers its valence electrons to the hydrogen in an acid solution. (Remember a superscript plus sign indicates that the compound is missing a valence electron a minus sign indicates that it has an extra valence electron.) This causes the hydrogen in the acid solution to separate as hydrogen gas while the zinc metal forms zinc ions dissolved in the solution. The zinc is said to be more electropositive than the hydrogen it displaces. 169
Reactions like this that involve the transfer of electrons are called oxidation reduction reactions, or redox for short. Oxidation generally describes the loss of electrons by a molecule, while reduction describes the gain of electrons. In this reaction, since zinc is losing electrons, it is being oxidized. Zn (s) Zn 2+ (aq) + 2e - Hydrogen is gaining electrons in this reaction, so it is being reduced: 2 H + (aq) + 2e - H 2 (g) This is often remembered by the phrase LEO the Lion goes GER. LEO stands for Losing Electrons Oxidation, and GER stands for Gain Electrons Reduction. In Lab 15 we called this same reaction a single replacement reaction. How can this be? It has to do with how the reactions are sorted. Just like laundry, there are several different ways to categorize and sort reactions depending on the application. For example, when doing laundry, you may need to separate red clothes from white ones. Other times you separate the light colored clothes from the dark colored clothes. Sometimes reactions are categorized as they were in Lab 15, and sometimes they are categorized as redox and non redox reactions. But how do we know just by looking at this reaction equation that it is an redox reaction and that electrons are being transferred? The oxidation number must first be assigned to each of the atoms involved on both sides of the reaction equation. Let s look at a reaction between aluminum and hydrochloric acid. 6 HCl (aq) + 2 Al (s) 2 AlCl 3 (aq) + 3 H 2(g) There are several rules to help you determine oxidation numbers in a reaction: The oxidation number of an element by itself is zero. This means that the aluminum metal (Al) on the right side of the above equation has an oxidation number of 0. The hydrogen gas (H 2 ) on the right side of the equation also has an oxidation number of 0. When an atom exists as a simple ion in a substance, the oxidation number is the same as its charge in the compound. The chloride (Cl) ion has a negative charge, so its oxidation number in is 1 in each instance it appears. For a neutral compound, the sum of the oxidation numbers is always zero. A polyatomic ion s oxidation number equals the charge on the ion. Since HCl is neutral, the sum of the oxidation numbers should equal zero. We know from above that the Cl ion has an oxidation number of 1; For the sum of the oxidation numbers in HCl to equal 0, the H must have an oxidation number of +1. Similarly, for the oxidation of AlCl 3 to be zero, the Al must have an oxidation number of +3 to balance out the three Cl atoms (each with oxidation numbers of 1). 170
Figure 2: Activity Series of Metals (Partial List) Lithium Potassium Barium Calcium Sodium Magnesium Aluminum Zinc Chromium Iron Release hydrogen from cold water, steam, and acids Release hydrogen from steam and acids Loses electrons easily (more easily oxidized) Cobalt Nickel Tin Lead Hydrogen Copper Mercury Silver Platinum Gold Release hydrogen from acids Do not release hydrogen from acids Do not lose electrons easily (not easily oxidized) Sometimes it is helpful to write the oxidation numbers over the atoms as is shown below: +1 1 0 +3 3( 1) 0 6 HCl (aq) + 2 Al (s) 2 AlCl 3 (aq) + 3 H 2(g) The activity series (Figure 2) of metals also indicates how easily a metal will cause a release of hydrogen in redox reactions. Metals at the top portion of the list release hydrogen merely by being placed in cold water. As you go down the list more harsh conditions are required to release hydrogen. Metals towards the bottom do not release hydrogen even when placed in acids. Sometimes a metal forms a thin oxide coating that protects it from reacting any further with its surroundings. Aluminum is fairly reactive and does this quickly. On the other hand, when iron is exposed to oxygen it corrodes entirely into iron oxide (rust) over time. For this reason, iron is commonly coated in zinc to prevent oxidation. In this laboratory exercise, you will compare the reactivity of the redox reactions of zinc and iron with a weak acid solution: saturated citric acid. Citric acid is a different type of acid than HCl, but works in a similar way. 171
Pre lab Questions 1. What is the oxidation state for each atom in the following reaction: 4Fe + 3O 2 2Fe 2 O 3 a. Elemental iron (Fe) b. Elemental oxygen (O 2 ) c. One iron atom in Fe 2 O 3 d. One oxygen atom in Fe 2 O 3 2. Which element was oxidized and which element was reduced in the above reaction equation? a. Element oxidized b. Element reduced 3. From the Activity Series of Metals, determine the order of reactivity of the following metals: Ni, Au, Fe, Ca, Zn, and Al. Most Reactive Least Reactive 172
Experiment: Metal Reactivity In this lab you will use what you know about chemical reactions and oxidation to examine how two metals (zinc and iron) react differently with a citric acid solution. You will then draw conclusions about the reactivity of the metals based on what you observe. Materials Safety Equipment: Safety goggles, gloves 10 ml graduated cylinder Iron nail (uncoated) 250 ml beaker Saturated citric acid solution 2 test tubes Sandpaper Test tube rack Stopwatch Galvanized nail Baking soda Procedure 1. Label two test tubes Zn and Fe with the permanent marker, and place them in a test tube rack. 2. Tip the test tube labeled Zn slightly and let a galvanized nail gently slide into it with the top of the nail going down first. 3. If necessary, lightly sand the iron nail that is not galvanized to remove any rust and wipe it clean. Gently place it in the test tube labeled Fe the same way you did previously. 4. Add approximately 3 to 5 ml of saturated citric acid solution to each test tube. 5. Make initial observations and continue recording observations after one minute, three minutes, and five minutes. HINT: The most notable observations are how quickly bubbling occurs and how violently the bubbling of each continues. 6. After observing the reactions for five minutes, rank the two metals in order of their reactivity. Compare your results with the actual reactivity series. 7. To clean up, separate the acid solution from the metals by pouring them into a 250 ml beaker while leaving the metals in their test tubes. This is called decanting. Rinse the test tube containing the metals several times with water and add the rinses to the beaker. To neutralize the acid, add small amounts of baking soda to the acid solution in the beaker and stir. Continue this until no more gas forms. Pour the liquid down the drain, and throw the metals in the trash. 173
Data Table 1: Observations of Reactions Time (minutes) Observations Initial 1 3 5 Summarize your observations and list the order of reactivity of the metals that you observed: 174
Post lab Questions 1. Based on what you observed, what is one of the products formed in Part 1? How do you know? 2. Did the order of reactivity you determined in Part 2 match the order given in the Activity Series of Metals? Explain. 2. Do you think the following reaction would occur? Explain your answer. FeCl 2 (aq) + Cu (s) Fe (s) + CuCl 2 (aq) 4. How do you think acid rain might affect the rate of rusting of metal? 175