Bài giảng hoá phân tích oxidation–reduction reactions

Experiment 27Oxidation–Reduction Reactions • To observe and predict products of oxidation–reduction reactions • To determine the relative reactivity of a series of metallic elements The

Experiment 27

Oxidation–Reduction

Reactions

• To observe and predict products of oxidation–reduction reactions

• To determine the relative reactivity of a series of metallic elements

The following techniques are used in the Experimental Procedure:

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)

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.

Objectives

Techniques

Introduction

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

dis-placement, double disdis-placement, 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

re-works, 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 (H) are transferred; in a redox reaction,

elec-trons (e) are transferred from one substance to another, resulting in changes in

oxida-tion numbers (see Dry Lab 2A) of two or more elements in the chemical reacoxida-tion.

In a redox reaction the substance that experiences an increase in oxidation number

is said to be oxidized—to do 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

sub-stance that is reduced (or gains the electrons) Therefore, the subsub-stance 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):

(27.1)

Zn(s)  Cu2(aq) l Zn2(aq)  Cu(s)

1 Synthesis and decomposition reactions were observed in Experiment 7 and double displacement

reactions in Experiment 6.

In the reaction, zinc has increased its oxidation number from 0 to 2 by releasing

2 moles of electrons per mole of zinc; zinc has therefore been oxidized:

(27.2)

In addition, copper(II) ion has decreased in oxidation number from 2 to 0 by accepting

2 moles of electrons per mole of copper(II) ion; copper(II) ion has therefore been reduced:

(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

A common laboratory reducing agent (in addition to zinc metal) is the thiosulfate ion, Its half-reaction for oxidation in an aqueous solution is

(27.4) Reducing agents with their corresponding half-reactions used in this experiment are:

Cu(s, copper-colored) l Cu 2(aq, blue)  2 e

3 I(aq, colorless) l (aq, purple)  2 e

Fe2(aq, light green to colorless) l Fe 3(aq, red-brown)  e

(aq, colorless) l 2 CO 2(g, colorless)  2 e

(aq, colorless) l (aq, colorless)  2 e

A common laboratory oxidizing agent is the permanganate ion, Its half-reaction for reduction in an acidic solution is

(27.5) Oxidizing agents with their corresponding half-reactions used in this experiment are:

(aq, colorless)  2 H (aq)  e l NO 2(g, red-brown)  H 2O(l ) (aq, colorless)  4 H (aq)  3 el NO(g, colorless)  2 H2O(l ) (aq, purple)  8 H (aq)  5 e l Mn 2(aq, light pink to colorless)  4 H 2O(l)

O 2(g, colorless)  4 H (aq)  4 e  l 2 H 2O(l)

H 2 O 2(aq, colorless)  2 H (aq) + 2 el 2 H 2O(l)

ClO(aq, colorless)  2 H (aq)  2 e l Cl (aq, colorless)  H 2O(l )

In Part A, several redox reactions involving these reducing and oxidizing agents are ob-served 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

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

MnO 4 

NO 3 

NO 3 

MnO4 (aq)  8 H(aq)  5 el Mn2(aq)  4 H2O(l)

MnO4 

S 4 O 6 

2 S 2 O 3 2

C 2 O 4 2

I 3 

2 S2O3 2(aq) l S4O6 2(aq)  2 e

S2O3 2

Cu2(aq)  2 el Cu(s) a reduction half-reaction Zn(s) l Zn2(aq)  2 e an oxidation half-reaction

Figure 27.1 The relationship between oxidation,

reduction, oxidizing agent, and reducing agent

Reducing Agents and

Oxidizing Agents

Displacement Reactions

(Activity Series)

Activity series: a listing of (generally)

metals in order of decreasing

chemical reactivity

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 Bnfrom an aqueous solution Metal A is oxidized to

Amin solution, and Bnis reduced to metal B

(27.6)

A specific example is the relative reactivity of iron versus lead:

experimen-tally, when iron metal is placed in a solution containing lead(II) ion, iron is

oxi-dized (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

(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):

(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 superscript1–9 in the

pro-cedure, stop and record your observation on the Report Sheet Discuss your

observa-tions with your lab partner and your instructor

1 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.1

2 Oxidation of copper.In three separate small test tubes add approximately 1 mL

of 6 M HCl, 1 mL of 6 M HNO3, and 1 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.2

3 A series of redox reactions.Refer to Table 27.1, page 312 Clean and label eight

small test tubes Place about 1 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

1 mL of solution B until a permanent change is observed For test tubes 5 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.3

CLEANUP: Rinse the test tubes twice with tap water and twice with deionized

water Discard the rinses in the Waste Salts container

1 Reactivity of Ni, Cu, Zn, Fe, Al, and Mg with H 3 Oⴙ Obtain 1-cm strips of Ni,

Cu, Zn, Fe, Al, and Mg Place about 1.5 mL of 6 M HCl into each of six small

Disposal: Dispose of the test solutions in the Waste Salts container

Pb(s)  Fe2(aq) l no reaction

Fe(s)  Pb2(aq) l Fe2(aq)  Pb(s) A(s)  Bn(aq) l Am(aq)  B(s)

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

B Reactions with Hydronium Ion

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!4

1 Reactivity of Cu, Zn, and Fe with Ni 2ⴙ Obtain 1-cm strips of Cu, Zn, Fe, and Ni

Place about 1 mL of 0.1 M NiSO4in each of three test tubes Place a short strip of freshly polished Cu, Zn, and Fe in successive test tubes (Figure 27.5) Tarnishing or dulling of the metal or color change of the solution indicates that a reaction has occurred Allow 5–10 minutes to observe the reaction Record.5

2 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)2,6 Zn(NO3)2,7 and Fe(NH4)2(SO4)2.8 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

3 Relative activity of the metals.List the four metals in order of decreasing activity.9

CLEANUP: Rinse the test tubes twice with tap water and twice with deionized water Discard the rinses in the sink

(1) Oxidation–reduction reactions are involved in respiration and photosynthesis,

cor-rosion, 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?

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

Table 27.1 Preparation of Several Oxidation–Reduction Reactions

Test

1 Chlorine bleach and 2 drops of 6 M HCl Drops of 0.1 M Fe(NH4 ) 2 SO 4

(the active ingredient in chlorine bleach is ClO)

2 Chlorine bleach and 2 drops of 6 M HCl Add 1 drop of starch solution followed by

drops of 0.1 M KI

3 0.01 M KMnO4and 2 drops of 6 M H2 SO 4 Drops of 0.1 M Fe(NH4 ) 2 SO 4 (Figure 27.3)

4 0.01 M KMnO4and 2 drops of 6 M H2 SO 4 Drops of 1 M K2 C 2 O 4

5 Deionized water and 2 drops of 6 M H2 SO 4 Drops of 0.1 M Fe(NH4 ) 2 SO 4

6 Deionized water and 2 drops of 6 M H2 SO 4 Add 1 drop of starch solution followed by

drops of 0.1 M KI

7 0.1 M H2 O 2and 2 drops of 6 M H2 SO 4 Drops of 0.1 M Fe(NH4 ) 2 SO 4

8 0.1 M H2 O 2and 2 drops of 6 M H2 SO 4 Add 1 drop of starch solution followed by

drops of 0.1 M KI

Figure 27.3 The deep purple

permanganate ion, added to an

acidic iron(II) solution, is reduced

to the nearly colorless

manganese(II) ion.

Figure 27.5 Setup for

observing the relative reactivity of

metals and metal ions

C Displacement Reactions

between Metals and Metal

Cations

2 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.

M

Figure 27.4 Setup for observing

the reactivity of six metals in hydrochloric acid

The Next Step

Experiment 27 Prelaboratory Assignment

Oxidation–Reduction Reactions

Date Lab Sec Name Desk No

1 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

2. 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, Fe2, to iron metal

3. 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:

b. Co3l Co2 _

e. Hl H2 _

4 Cite the part of the Experimental Procedure for three Cautions in this experiment Identify the reason or response for

each caution

S2O8 l SO4 

CrO4 

Cr2O7 2

ClO2 l ClO

5. The following equation is not balanced for both mass and charge! Explain.

What is the correct balanced equation? See the list of half-reactions that appear under “Reducing Agents and Oxidiz-ing Agents” in the Introduction

6. 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

7. 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  Tl X T

R  Xl no reaction

R  Tl R 

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

MnO4(aq)  H2C2O4(aq)  6 H(aq) l Mn2(aq)  2 CO2(g)  4 H2O(l)

Corrosion of the Alyeska (Alaskan) pipeline is inhibited by the presence of zinc metal.

Experiment 27 Report Sheet

Oxidation–Reduction Reactions

Date Lab Sec Name Desk No

A Oxidation–Reduction Reactions

1. 1Oxidation 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 the oxidizing agent in the equation

2. 2Oxidation of copper.Describe your observations of each test tube

Comment on the relative oxidizing strengths of the three acids in the test tubes

3. 3A series of redox reactions.On a separate sheet of paper, organize your data to record the test tube number and your observations for the 8 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: the oxidizing agent in each written equation Submit this with the completed Report Sheet.

B Reactions with Hydronium Ion

1. 4Reactivity of Ni, Cu, Zn, Fe, Al, and Mg with H 3 Oⴙ

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 do react in order of decreasing activity.

Write a balanced equation for the reaction that occurs between Mg and H3O

Box Box

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

bal-anced net ionic equation Use additional paper if necessary the oxidizing agent in each written equation

4HCl

5NiSO4 NR

6Cu(NO3)2 NR

7Zn(NO3)2 NR

8Fe(NH4)2(SO4)2 NR

9List 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

1. 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?

2. 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

3. Part A.3 Test tube 7 Does the ferrous ion in the Fe(NH4)2SO4solutions function as an oxidizing agent or a reducing agent? Explain

4. Part A.3, Test tube 8

a. What was the color change of the Iin the reaction?

b. Does the Isolution function as an oxidizing agent or a reducing agent? Explain

5. Part B.1 Eliseo couldn’t nd the 6 M HCl and so used 6 M HNO3for testing the metals instead His logic? Both are strong acids Explain how the results of the experiment would have been different

6. Part C Single displacement, double displacement, and decomposition reactions may all be redox reactions Identify the type of redox reactions in Part C Explain

7. 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

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