Bài giảng hoá phân tích oxidation–reduction reactions

Experiment 27 Oxidation–Reduction Reactions Zinc metal is placed into a blue copper(II) ion solution (before, left) Copper metal collects on the zinc strip and the colorless zinc(II) ion goes into solution (after, right) Zinc has a greater activity than does copper • To observe and predict products of oxidation–reduction reactions • To determine the relative reactivity of a series of metallic elements Objectives The following techniques are used in the Experimental Procedure: Techniques Most chemical reactions are classi ed as being either acid–base ( Experiment 6) or oxidation–reduction (redox) Reactions may also be classi ed as synthesis, single displacement, double displacement, or decomposition.1 Redox reactions are often accompanied by spectacular color changes, generally more so than what is observed in acid–base reactions The changing of the color of the leaves each fall season, the rusting of our automobiles, the detonation of reworks, the formation of “brown” smog, and the commercial production of copper, aluminum, and iron are all examples of redox reactions In an acid–base reaction, protons (Hϩ) are transferred; in a redox reaction, electrons (eϪ) are transferred from one substance to another, resulting in changes in oxidation numbers (see Dry Lab 2A) of two or more elements in the chemical reaction In a redox reaction the substance that experiences an increase in oxidation number is said to be oxidized—to so it must have donated or lost electrons, a process of oxidation Conversely, the substance that experiences a decrease in oxidation number is said to be reduced—as a result, it must have accepted or gained electrons, a process of reduction In all redox reactions, the substance that is oxidized must lose its electrons to the substance that is reduced (or gains the electrons) Therefore, the substance being oxidized is causing a reduction and therefore is called a reducing agent Conversely, the substance that is reduced in the reaction must gain electrons from the substance that is oxidized, thereby causing oxidation; it is called an oxidizing agent (Figure 27.1, page 310) Electrons are never considered a reactant or product in a chemical reaction but are merely transferred—the total number of electrons donated must equal the total number of electrons gained in a redox reaction For example, zinc metal reacts with copper(II) ion (see opening photo): Introduction Zn(s) ϩ Cu2ϩ(aq) l Zn2ϩ(aq) ϩ Cu(s) (27.1) Oxidation: a process whereby a substance loses electrons but increases in oxidation number Reduction: a process whereby a substance gains electrons but decreases in oxidation number Reducing agent: a substance that donates electrons, causing reduction of another substance to occur (the reducing agent is therefore oxidized) Oxidizing agent: a substance that accepts electrons, causing oxidation of another substance to occur (the oxidizing agent is therefore reduced) Synthesis and decomposition reactions were observed in Experiment and double displacement reactions in Experiment Experiment 27 309 Figure 27.1 The relationship between oxidation, reduction, oxidizing agent, and reducing agent In the reaction, zinc has increased its oxidation number from to ϩ2 by releasing moles of electrons per mole of zinc; zinc has therefore been oxidized: Zn(s) l Zn2ϩ(aq) ϩ eϪ an oxidation half-reaction (27.2) In addition, copper(II) ion has decreased in oxidation number from ϩ2 to by accepting moles of electrons per mole of copper(II) ion; copper(II) ion has therefore been reduced: Cu2ϩ(aq) ϩ eϪ l Cu(s) a reduction half-reaction (27.3) An overview of the reaction suggests that the presence of zinc caused the reduction of copper(II) ion—zinc is the reducing agent; the presence of copper(II) ion caused the oxidation of zinc—copper(II) ion is the oxidizing agent The total balanced redox reaction (equation 27.1) is a sum of the two half-reactions—the oxidation half-reaction and the reduction half-reaction Reducing Agents and Oxidizing Agents A common laboratory reducing agent (in addition to zinc metal) is the thiosulfate ion, S2O32Ϫ Its half-reaction for oxidation in an aqueous solution is S2O32Ϫ(aq) l S4O62Ϫ(aq) ϩ eϪ (27.4) Reducing agents with their corresponding half-reactions used in this experiment are: Cu(s, copper-colored) l Cu2ϩ(aq, blue) ϩ eϪ IϪ(aq, colorless) l I3Ϫ(aq, purple) ϩ eϪ Fe2ϩ(aq, light green to colorless) l Fe3ϩ(aq, red-brown) ϩ eϪ C2O4 2Ϫ(aq, colorless) l CO2(g, colorless) ϩ eϪ S2O3 2Ϫ(aq, colorless) l S4O62Ϫ(aq, colorless) ϩ eϪ A common laboratory oxidizing agent is the permanganate ion, MnO4Ϫ Its halfreaction for reduction in an acidic solution is MnO4Ϫ(aq) ϩ Hϩ(aq) ϩ eϪ l Mn2ϩ(aq) ϩ H2O(l) (27.5) Oxidizing agents with their corresponding half-reactions used in this experiment are: Ϫ NO3 (aq, colorless) ϩ Hϩ(aq) ϩ eϪ l NO2(g, red-brown) ϩ H2O(l ) NO3Ϫ(aq, colorless) ϩ Hϩ(aq) ϩ eϪ l NO(g, colorless) ϩ H2O(l ) MnO4Ϫ(aq, purple) ϩ Hϩ(aq) ϩ eϪ l Mn2ϩ(aq, light pink to colorless) ϩ H2O(l) O2(g, colorless) ϩ Hϩ(aq) ϩ eϪ l H2O(l) H2O2(aq, colorless) ϩ Hϩ(aq) + eϪ l H2O(l) ClOϪ(aq, colorless) ϩ Hϩ(aq) ϩ eϪ l ClϪ(aq, colorless) ϩ H2O(l ) In Part A, several redox reactions involving these reducing and oxidizing agents are observed and studied, and reactants and products are recorded You may have to consult with your instructor when writing the balanced redox equation for the reactions that are studied Displacement Reactions (Activity Series) Activity series: a listing of (generally) metals in order of decreasing chemical reactivity 310 Parts B and C of the experiment seek to establish the relative chemical reactivity of several metals and hydrogen The metals and hydrogen, listed in order of decreasing activity (decreasing tendency to react), constitute an abbreviated activity series The result of one metal being placed into a solution containing the cation of another metal establishes the relative reactivity of the two metals A displacement Oxidation–Reduction Reactions redox reaction occurs if there is evidence of a chemical change For example, if metal A is more reactive (has a greater tendency to lose electrons to form its cation) than metal B, then metal A displaces Bnϩ from an aqueous solution Metal A is oxidized to Amϩ in solution, and Bnϩ is reduced to metal B A(s) ϩ Bnϩ(aq) l Amϩ(aq) ϩ B(s) (27.6) A specific example is the relative reactivity of iron versus lead: experimentally, when iron metal is placed in a solution containing lead(II) ion, iron is oxidized (loses two electrons) to form the iron(II) ion, and the lead(II) ion is reduced (gains two electrons) to form lead metal (Figure 27.2, right) The equation for the reaction is Fe(s) ϩ Pb2ϩ(aq) l Fe2ϩ(aq) ϩ Pb(s) (27.7) On the other hand, when lead metal is placed into a solution containing iron(II) ion, no reaction occurs (Figure 27.2, left): Pb(s) ϩ Fe2ϩ(aq) l no reaction (27.8) Consequently, iron prefers to be ionic, whereas lead prefers the metallic state— iron loses electrons more easily than lead Iron therefore is more easily oxidized than lead and is said to have a greater activity (chemical reactivity) than lead Procedure Overview: Observations of a number of redox reactions are analyzed and equations are written The relative reactivity of several metals is determined from a series of oxidation–reduction or displacement reactions Perform the experiment with a partner At each circled superscript 1–9 in the procedure, stop and record your observation on the Report Sheet Discuss your observations with your lab partner and your instructor Oxidation of magnesium Half- ll a 20-mL beaker with deionized water and carefully heat to near boiling—test the water with litmus or pH paper Grip a 2-cm piece of magnesium ribbon with crucible tongs Heat it directly in a Bunsen burner ame until it ignites ( Caution: Do not look directly at the ame ) The white-ash product is magnesium oxide Allow the ash to drop into the 20-mL beaker Swirl and again test the solution with litmus or pH paper at the end of the laboratory period Compare your observations for the two litmus or pH paper tests Record your observations on the Report Sheet Figure 27.2 Lead metal does not react in a solution containing iron(II) ion (left) Iron metal does react in a solution containing lead(II) ion (right) Experimental Procedure A Oxidation–Reduction Reactions Oxidation of copper In three separate small test tubes add approximately mL of M HCl, mL of M HNO3, and mL of conc HNO3 (Caution: Avoid skin contact Flush affected areas with large amounts of water.) Place the three test tubes in a fume hood and add a 1-cm wire strip of copper metal to each Record your observations on the Report Sheet A series of redox reactions Refer to Table 27.1, page 312 Clean and label eight small test tubes Place about mL of each solution listed as solution A in Table 27.1 into the test tube (Caution: Avoid skin contact with all solutions!) Slowly add up to mL of solution B until a permanent change is observed For test tubes and 6, after the addition of solution B, stopper the test tube and shake vigorously Some reactions may be slow to develop Record your observations Disposal: Dispose of the test solutions in the Waste Salts container CLEANUP: Rinse the test tubes twice with tap water and twice with deionized water Discard the rinses in the Waste Salts container Reactivity of Ni, Cu, Zn, Fe, Al, and Mg with H3O؉ Obtain 1-cm strips of Ni, Cu, Zn, Fe, Al, and Mg Place about 1.5 mL of M HCl into each of six small B Reactions with Hydronium Ion Experiment 27 311 Table 27.1 Preparation of Several Oxidation–Reduction Reactions Test Tube No Figure 27.3 The deep purple permanganate ion, added to an acidic iron(II) solution, is reduced to the nearly colorless manganese(II) ion Solution A Addition of Solution B Chlorine bleach and drops of M HCl (the active ingredient in chlorine bleach is ClOϪ) Drops of 0.1 M Fe(NH4)2SO4 Chlorine bleach and drops of M HCl Add drop of starch solution followed by drops of 0.1 M KI 0.01 M KMnO4 and drops of M H2SO4 Drops of 0.1 M Fe(NH4)2SO4 (Figure 27.3) 0.01 M KMnO4 and drops of M H2SO4 Drops of M K2C2O4 Deionized water and drops of M H2SO4 Drops of 0.1 M Fe(NH4)2SO4 Deionized water and drops of M H2SO4 Add drop of starch solution followed by drops of 0.1 M KI 0.1 M H2O2 and drops of M H2SO4 Drops of 0.1 M Fe(NH4)2SO4 0.1 M H2O2 and drops of M H2SO4 Add drop of starch solution followed by drops of 0.1 M KI test tubes (Figure 27.4) Polish each metal (with steel wool or sandpaper) to remove any oxide coating and immediately place it in a test tube of hydrochloric acid.2 Allow 10–15 minutes for any reactions to occur Record what you see—look closely! M Reactivity of Cu, Zn, and Fe with Ni2؉ Obtain 1-cm strips of Cu, Zn, Fe, and Ni Place about mL of 0.1 M NiSO4 in each of three test tubes Place a short strip of freshly Figure 27.4 Setup for observing polished Cu, Zn, and Fe in successive test the reactivity of six metals in tubes (Figure 27.5) Tarnishing or dulling of hydrochloric acid the metal or color change of the solution indicates that a reaction has occurred Allow 5–10 minutes to observe the reaction Record C Displacement Reactions between Metals and Metal Cations Reactivity of Cu, Zn, Fe, and Ni with other cations Follow the procedure in Part C.2 (and correspondingly on the Report Sheet) to test the reactivity of the metals in the following 0.1 M test solutions: Cu(NO3)2, Zn(NO3)2, and Fe(NH4)2(SO4)2 This series of tests will ll the table on the Report Sheet The metal strip may be reused if it is unreacted after the previous test and is rinsed with deionized water or if it is again freshly polished Record your observations Relative activity of the metals List the four metals in order of decreasing activity Figure 27.5 Setup for observing the relative reactivity of metals and metal ions Disposal: Dispose of the waste test solutions and one rinse of the test tubes in the Waste Salts container Place the unreacted metals in the Waste Solids container CLEANUP: Rinse the test tubes twice with tap water and twice with deionized water Discard the rinses in the sink The Next Step (1) Oxidation–reduction reactions are involved in respiration and photosynthesis, corrosion, bleaching (Experiment 29), water purification, electroplating, and so on Research an area of interest where oxidation–reduction reactions are an integral part of the study (2) What metals are used to inhibit the corrosion of underground (iron/steel) storage tanks, of metal structures below water in harbors? The polishing of the metal and quick placement in HCl are especially critical for aluminum and magnesium as they form quick, tough protective oxide coatings 312 Oxidation–Reduction Reactions Experiment 27 Prelaboratory Assignment Oxidation–Reduction Reactions Date Lab Sec Name Desk No a Oxygen is a common oxidizing agent in nature What change (increase or decrease) in the oxidation number of oxygen must occur if it is to be an oxidizing agent? Explain b If oxygen gas were to oxidize copper metal, what change (increase or decrease) in oxidation number must occur for the copper metal? Write an appropriate half-reaction for the copper metal Zinc metal is a common reducing agent in analytical chemistry a What does it mean for a substance to be a reducing agent? b Write a balanced oxidation–reduction equation showing how zinc metal reduces the ferrous ion, Fe2ϩ, to iron metal Each of the following processes is a likely change in a redox reaction Label the chemical change as an oxidation process, a reduction process, or as neither: a ClO2Ϫ l ClOϪ _ b Co3ϩ l Co2ϩ _ c Cr2O7 2Ϫ l CrO42Ϫ d S2O82Ϫl SO4 2Ϫ Ϫ e H l H2 _ _ _ Cite the part of the Experimental Procedure for three Cautions in this experiment Identify the reason or response for each caution Experiment 27 313 The following equation is not balanced for both mass and charge! Explain MnO4Ϫ(aq) ϩ H2C2O4(aq) ϩ Hϩ(aq) l Mn2ϩ(aq) ϩ CO2(g) ϩ H2O(l) What is the correct balanced equation? See the list of half-reactions that appear under “Reducing Agents and Oxidizing Agents” in the Introduction Experimental Procedure, Part A.3 Write a balanced equation for the oxidation–reduction reaction that occurs when copper metal reacts with the permanganate ion in an acidic solution See the list of half-reactions that appear under “Reducing Agents and Oxidizing Agents” in the Introduction Experimental Procedure, Part C List the generic chemicals R, T, and X in order of deceasing activity on the basis of the following reactions: X ϩ Tϩ l Xϩ ϩ T R ϩ Xϩ l no reaction R ϩ Tϩ l Rϩ ϩ ⌻ most active _, _, _, least active *8 Along the 800-mile Alyeska (Alaska) pipeline transporting oil from Prudhoe Bay (north) to the Valdez Marine Terminal (south), zinc ribbon is buried to inhibit the corrosion of the below-ground sections of the (iron/steel) pipeline Explain how zinc serves in this function Corrosion of the Alyeska (Alaskan) pipeline is inhibited by the presence of zinc metal 314 Oxidation–Reduction Reactions Experiment 27 Report Sheet Oxidation–Reduction Reactions Date Lab Sec Name Desk No A Oxidation–Reduction Reactions 1 Oxidation of magnesium Write a description of the reaction What did the litmus tests reveal? Write a balanced equation for the reaction of magnesium in air Box the oxidizing agent in the equation 2 Oxidation of copper Describe your observations of each test tube Comment on the relative oxidizing strengths of the three acids in the test tubes 3 A series of redox reactions On a separate sheet of paper, organize your data to record the test tube number and your observations for the reaction mixtures in Table 27.1 Use the appropriate half-reactions (one for solution A and another for solution B) that appear under Reducing Agents and Oxidizing Agents in the Introduction Add the two half-reactions such that the electrons gained by one solution equal the electrons lost by the other—this sum provides the overall balanced equation Write a balanced redox equation for each observed reaction: Box the oxidizing agent in each written equation Submit this with the completed Report Sheet B Reactions with Hydronium Ion Reactivity of Ni, Cu, Zn, Fe, Al, and Mg with H3O؉ Which metals show a de nite reaction with HCl? Record this information on the table in Part C of the Report Sheet Arrange the metals that react in order of decreasing activity Write a balanced equation for the reaction that occurs between Mg and H3Oϩ Experiment 27 315 C Displacement reactions between metals and metal cations Complete this table with NR (no reaction) or R (reaction) where appropriate For all observed reactions, write a balanced net ionic equation Use additional paper if necessary Box the oxidizing agent in each written equation Ni Cu Zn Fe Al Mg NR HCl NiSO4 Cu(NO3)2 NR Zn(NO3)2 NR Fe(NH4)2(SO4)2 NR List the four metals along with Al, Mg, and hydrogen in order of decreasing activity , , , , , , Balanced net ionic equations Use equations 27.1 and 27.7 as models Laboratory Questions Circle the questions that have been assigned Part A.1 Sodium metal is also readily oxidized by oxygen If the product of the reaction were dissolved in water, what would be the color of the litmus for a litmus test? Explain What is the product? Part A.2 Oxygen gas has an oxidizing strength comparable to that of nitric acid Patina is a green or greenish-blue coating that forms on copper metal in the environment Account for its formation Part A.3 Test tube Does the ferrous ion in the Fe(NH4)2SO4 solutions function as an oxidizing agent or a reducing agent? Explain Part A.3, Test tube a What was the color change of the IϪ in the reaction? b Does the IϪ solution function as an oxidizing agent or a reducing agent? Explain Part B.1 Eliseo couldn’t nd the M HCl and so used M HNO3 for testing the metals instead His logic? Both are strong acids Explain how the results of the experiment would have been different Part C Single displacement, double displacement, and decomposition reactions may all be redox reactions Identify the type of redox reactions in Part C Explain Part C a On the basis of your intuitive understanding of the chemical properties of sodium and gold, where in your activity series would you place sodium and gold? b Will hydrochloric acid react with gold metal to produce gold(III) ions and hydrogen gas? Explain 316 Oxidation–Reduction Reactions