Oxidation / Reduction

Experiment 9 Oxidation / Reduction

Electron Transfer Reactions Expt 9 Redox.wpd

INTENT In this experiment you will focus on the atomic model interpretation of the most easily recognized chemical reactions of common experience, namely the oxidation/reduction reactions.

Key Terms and Concepts (familiarize yourself with them): C conjugate oxidizing agent & reducing agent relationship C oxidation & reduction (oxidation states) C oxidizing agent & reducing agent C metal & nonmetal activity series

DISCUSSION Most people have no trouble recognizing the combustion of fuels, the rusting of iron, or the digestion of foods by an organism as examples of chemical reactions. These reactions are all examples of redox reactions. Reduction-oxidation (redox) reactions create new compounds by exchanging electrons and thereby fundamentally changing the reacting particles. These kinds of reactions will be examined in this experiment.

Before we look at any specific reactions, let us define some of the basic terms of redox reactions.

Originally the term oxidation described what happens to any substance that reacted with oxygen, O2, and in so doing lost electrons. The term has now been generalized so that:

Oxidation is any process which increases the oxidation state of an atom due to the loss of electrons. Reduction is any process which reduces the oxidation state of an atom due to the gain of electrons.

The actual species involved in these reactions are the oxidizing agents and the reducing agents.

By taking on electrons the oxidizing agent causes something else to be oxidized (lose electrons). C The oxidizing agent is the electron acceptor and it is reduced.

By giving up electrons, the reducing agent causes something else to be reduced (gain electrons). C The reducing agent is the electron donor and it is oxidized.

These processes may seem very confusing at first. To avoid confusion we treat redox reactions as those in which electrons are transferred from the reducing agent to the oxidizing agent.

121 In order to understand the stoichiometry of redox reactions we will focus on the number of electrons gained by the oxidizing agent and lost by the reducing agent. We will need to identify how many moles of electrons are given up by each mole of reducing agent and how many moles of electrons are accepted by each mole of oxidizing agent. Since mass and energy are neither created nor destroyed in a chemical reaction, the number of electrons given up by the reducing agent must be the same as the number accepted by the oxidizing agent. That is, they must exchange an equal number of electrons. Thus their coefficients in the chemical equation must represent equivalent quantities of electrons. For example, consider the reaction of Zn and Cu2+ ion in solution:

! 2+ & oxidation: Zn (s) Zn (aq) + 2e reduction: Cu2+ + 2e& ! Cu (aq) (s) 2+ ! 2+ net reaction: Zn (s) + Cu (aq) Cu (s) + Zn (aq)

In the example above the electrons lost in the oxidation equation equal the electrons gained in the reduction equation. Sometimes, in order for the number of electrons accepted to be equal to the number of electrons given up, the reactions must be multiplied by a whole number. For example, consider the reaction of F2 and H2O:

& ! & oxidation: 2e + F2 2 F ! + & reduction: 6 H2O O2 + 4 H3O + 4e

In order for the transfer to be equal, the first reaction must be multiplied by two before adding the two half-reactions to get the net equation.

& ! & 2 (2 e + F2 2 F ) 6 H O ! O + 4 H O+ + 4e& 2 2 3 ! & + Net reaction: 2 F2 + 6 H2O 4 F + O2 + 4 H3O Note: 4 e&s absorbed = 4 e&s ejected

Not every species has the same attraction for electrons. This is why electrons can be gained or lost. Even though we may sometimes refer to certain species as "wanting to give up electrons," substances do not give up electrons willingly. A species loses or gains electrons to achieve a lower energy state.

Note these generalizations: The greater the tendency for a substance to take on electrons, the less its tendency, once having taken on the electrons, to give them back. That is, the stronger the oxidizing agent is, the weaker its conjugate reducing agent will be (and vice versa).

The stronger oxidizing agent and stronger reducing agent will always react (eventually) to produce a weaker oxidizing agent and reducing agent. The weaker oxidizing agent and the weaker reducing agent will not react.

& For the F2 and H2O reaction, F2 and O2 are the oxidizing agents; F and H2O are the reducing agents.

For the reaction to proceed as written, F2 must be a stronger oxidizing agent than O2.

122 Since metals are good reducing agents, few metals are found pure in nature. So far you have learned that metals tend to lose electrons. The result of this is that very few metals are truly stable in our environment. What in the environment will take on these electrons? Some metals react with H2O to produce H2 gas. Those same metals and many additional ones may react with O2 and moisture in the air.

Note that O2 in the presence of acid is a better oxidizing agent than just O2 and water. Consequently, it is easy to understand the concerns about acid emissions from combustion reactions that result in "acid rain."

Rust formation on the surface of iron is actually the product of a reaction of iron with oxygen and moisture from the air. For this reason you would not want to make water pipes out of iron. You would want to make water pipes out of something which is less reactive under these conditions, like copper. The following equations illustrate the difference:

! 2 Fe (s) + 3 O2 (g) 2 Fe2O3 (s) But, ! Cu (s) + O2 (g) no appreciable reaction

Therefore, iron is a more reactive metal, or is "more active," than copper. In other words, iron is a better reducing agent than copper.

The reason that many metals even exist in our environment has to do with the rate at which they react. Many electron-transfer reactions are fast, but those which involve covalent bond making or breaking may be slow. (More about this in General Chemistry II.) The rate may also be affected by any oxide coating on the metal surface, which prevents further reaction.

We will be looking at the relative abilities of substances to attract electrons. If we consider substances two at a time, we can establish a relative order of reactivity for them. From this approach, we can develop a complete order of reactivity, known as the "activity series." Those species which are higher on the list of reducing agent are "more active."

How can we experimentally determine order of reactivity? In this experiment you will use observations to determine an order of activity, first for the metals, then the nonmetals. Some metals react with water or acid in water to produce H2. The evolution of hydrogen gas will prove which metal is more reactive. Before you start, give some thought to experimental procedure. Should you try to keep the size and shape of all the metals pieces approximately the same? Why? How will you distinguish between air bubbles on the metal surface and H2 gas being produced?

Let' s inspect the reaction of K metal in water. The reaction is very exothermic. The H2 actually ignites when we toss a small piece of K in the water.

! + & 2 K (s) + 2 H2O (R) 2 K (aq) + 2 OH (aq) + H2 (g) reducing agent oxidizing agent conjugate oxidizing agent conjugate reducing agent

Since both K and H2 are reducing agents and the reaction occurs as written (and not the other way around), we can state that K is a stronger reducing agent than H2. Potassium would be placed above hydrogen gas in an activity series. Note that if we try the reverse reaction, nothing will happen since hydrogen gas is less reactive. This conjugate relationship is very useful because it allows us to gain the

123 same information from the experiment no matter which reaction direction of the reaction we choose to investigate. Note that the reducing agents on both sides are being compared, not the reducing/oxidizing pair.

As you saw in the example above, if one adds a very reactive metal to pure water, the metal reduces the & +n water to OH and H2 as the metal is oxidized to M .

! +n & M (s) + H2O M + OH (aq) + H2 (g)

What happens to nonmetals, such as the halogens? If one adds Cl2, a very strong oxidizer, to pure water, + & the Cl2 oxidizes the water to H3O and O2 as it is reduced to Cl . So pure water can commonly function as either an oxidizer or a reducer. We will study several reductions of water. However, when we use chlorine water, the expected oxidation of water by Cl2 is slow (Cl-Cl bonds and O-H bonds are broken while O-O bonds are formed). The oxidation of water is not a factor in the fast reactions between & halogens (X2) and halide ions (X ).

In this experiment you will discover the reactivity of halogens by reacting aqueous halogen solutions with halide solutions. For example, chlorine water will be mixed with bromide ions from a sodium bromide solution. Does the reaction written below occur?

& ! & Cl2 (aq) + 2 Br (aq) 2 Cl (aq) + Br2 (aq)

How will you know? You will again use visual clues. Unfortunately there is no bubbling in this reaction to help you out. So you will use the color of the halogens dissolved in hexane to tell you which halogen is present. In other words, the hexane layer, formed after adding hexane to the aqueous solution, will tell you whether chlorine or bromine is present. If the color of the hexane layer indicates the presence of

Br2, then the reaction has occurred as written. If the color of the hexane layer indicates the presence of

Cl2, no reaction has occurred.

124 Data / Report Name: Experiment 9 Partner: CHM 1025 Section: Date: Oxidation / Reduction Electron Transfer Reactions

Pre-Laboratory Exercises

Please do the following exercises before your lab session and get an OK signature from your instructor.

1. Show the change in oxidation number in the following (in other words, give the number of electrons gained or lost, per atom).

&2 ! & ! + ! a) S S8 b) NO3 NO c) NH4 NH3

2. Consider the following reaction:

! +3 & 2 Al (s) + 3 Cl2 (g) 2 Al (aq) + 6 Cl (aq)

a) Which species is oxidized?

b) Which species is reduced?

c) How many electrons were transferred in this process?

d) Which species are reducing agents?

e) Which species are oxidizing agents?

f) If the reaction occurs as written, which is the strongest reducing agent?

3. Which of the following reactions are redox? Label them R for redox and X for non-redox.

& & ! & a) 2 MnO4 (aq) + 6 Cl (aq) + 4 H2O (R) 2 MnO2 (s) + 3 Cl2 (aq) + 8 OH (aq)

+ & ! b) H3O (aq) + OH (aq) 2 H2O (R)

+2 ! +2 c) Zn (s) + Cu (aq) Zn (aq) + Cu (s)

& +2 ! d) 2 I (aq) + Pb (aq) PbI2 (s)

& &2 ! &2 & e) NO3 (aq) + SO3 (aq) SO4 (aq) + NO2 (aq) OK:

125 EXPERIMENTAL You will perform a series of simple tests in this experiment to determine the relative reactivity of several metals (as reducing agent) and halogens (as oxidizing agents). First you will investigate the reactivity of some metals with water and with acid solution. A thin wire or flat piece of foil would work best, but most important is that comparisons are made of samples similar in size and shape. In this way, any observed differences in the reactions can be linked directly to the identity of the metals, and not be confused by any other factor. Any metal that is higher than hydrogen in the activity series will displace hydrogen from water, creating a metal ion/hydroxide ion solution.

PART I Reactions of metals with water / 1 M HCl Your instructor will demonstrate the reaction of potassium (K) metal and sodium (Na) metal with water. The reaction generates a large amount of energy as well as hydrogen gas. This reaction can be explosive and must be performed under controlled conditions. Record your observations during the demonstration in the chart below.

Label four small test tubes with the formulas of your metals and obtain a single small sample of each: magnesium (Mg), calcium (Ca), copper (Cu), and zinc (Zn). Add 2-3 mL of deionized water to each tube. Record any observations at room temperature.

Identify the metal(s) which are weaker reducing agent than H2 (g). Identify the strongest reducing agent in this exercise. What information do you have that identifies it as the strongest reducing agent? Which metals did you observe not to react with water?

In the table below under the first heading write the symbols, in alphabetical order, for the elements which you did not observe reacting in the previous section. Obtain new samples of those metals that were not consumed in the water and place them again in labeled test tubes. Caution: Do this only with the samples that did NOT react with water. This time add 2-3 mL of 1 M HCl solution instead of pure + water and observe for any reaction. All metals that react with 1 M HCl (which contains H3O ions as the active oxidizing agent) produce a metal ion/chloride ion solution along with the hydrogen gas.

Which of these metals is the weakest reducing agent (least reactive)? + Which is the stronger oxidizing agent, H2O or H3O , and defend your answer with experimental evidence?

You should now be able to determine the activity series for Mg, Ca, Cu, H2, Na and Zn. Note that hydrogen gas from pure water and hydrogen gas from an acidic solution will appear in different places. (Use the > symbol to rank the metals.)

126 PART II Reactions of halide ion solutions with halogens A. Preliminary Observations: The visual clue you will have for this part of the investigation is the color that halogens exhibit when they are dissolved in hexane. First, you will investigate the properties (solubility and color) of the halogens in solution with water and with hexane. Next you will use this information to probe what happens when a halogen is mixed with a salt of a different halide.

CAUTION: The halogens are all good oxidizing agents and can cause burns. Take extra care with their solutions.

Label six small test tubes and place into them 1 mL each of the chlorine water (Cl2), bromine water (Br2),

iodine water (I2), 0.1M NaCl solution, 0.1M NaBr solution, and 0.1M NaI solution, respectively. Observe the appearance of each solution including its color and record this information under aqueous solution in the table below. Then add 0.5 mL of hexane to each tube, stopper, and shake well to mix (no thumbs!). Observe the appearance of both layers in each tube. Which layer is which? Try to figure this out. Be sure to follow instructions for the safe disposal of these solutions.

Record the colors of the halogen / halide ions in water and hexane. Based on color can you say something about the solubilities of the halogens, as a group, in water versus hexane? Based on color alone can you say something about the solubilities of the halides, as a group, in water versus hexane?

Now that you know what colors to expect for each of the halogens and halide ions in both water and hexane, you can determine the results of any of the six possible combinations where a reaction might occur. (Of course, three & additional combinations, such as mixing Br2 and Br , solutions will not result in any changes.) You will be given specific instructions below for one such combination; it is up to you to decide how to perform the rest of the tests.

B. Testing Halogen Oxidizing Strengths: Pour 1 mL of chlorine water into a test tube. Add 1 mL of 0.1M NaBr, stopper and mix. Note any changes. Next add 1 mL hexane, shake, and again note any changes. Which halogen is now present? Which is higher in the activity series, Cl& ion or Br& ion?

The reaction under investigation is + & ! + & Cl2 (aq) 2 Br (aq) Br2 (aq) 2 Cl (aq)

Label (in the spaces provided) the oxidizing agent and conjugate oxidizing agent in the reaction above. By observing and recording the color changes in both layers, decide which is the strongest oxidizing agent. Try all possible combinations given in the following table below, following a procedure similar to the one above.

Write the equation that describes the reaction for each system in which you observed a change. Rank the halogen molecules in order of decreasing oxidizing agent strength. (Use the > symbol to rank them.) Now write the activity series for the halide ions as reducing agents. (Use the > symbol to rank them.) & Would you expect a reaction to occur between At2 and I ? Explain.

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