2
Many instructors use the lecture experiment in which Cl-, 3 M NH3, Br-, 7 M NH,, I-, SzO$-, SZ-, and CN- are added sequentially to a dilute aqueous solution of sil- ver nitrate. The alternate formation and solution of pre- cipitates, with some variation in color, provide excellent evidence for different degrees of dissociation of a series of complexes and precipitates. The results can, of course, be used to compare the basicity of the species added with re- spect to silver ion as an acid. We have developed a similar series, based on copper ion as the acid. The colors are more varied, and reduction of Cu(II) is introduced as a complication. Many other acids could be used hut we settled on Cu(I1) because of low cost, the varied colors involved, and the ease with which the concentration of copper(II) ions can be measured using a copper electrode, a constant EMF zinc electrode, and a low-cost voltmeter. The chemicals are stocked as 1 M aqueous solutions of: ZnSOd, CuSOn, NaC1, NaBr, KI, NaOH, NH3, KCN, Na2S, NaaP01, NazC03. Other hases could, of course, be added so that the experiment can he as open-ended as the student and instructor wish. The following experiments are performed. Chemical equations are required for all observed reactions I) Prepare a 0.1 M CuSO* solution and measure, or calcu- late, the concentration of each principal species presence. (pH papers or pH meters are used.) Calculate K for cop- per(II) ion acting as an acid in water. (pH and atom bal- ance give (Hi), (CuOH+), (Cue+), (SO?), (OH-), (HzO).) II) Design and perform a series of experiments with aqueous 1 M C1-, Br-, I-, OH-, NHa, and NazS, to arrive at a table of relative bsse strengths toward aqueous Cu(II). Note that one of these species causes redox with Cu(I1). Eliminate it J. A. Campbell Horvey Mudd College Clarernont, California 9171 1 from your list of simple bases. ~=perimentally verify the A Comparison of Base Strengths Using Visual and EMF Observations oxidized product. III) Try one or more possible other bases and add them to the table. (We provide 1 M NazCOs, NaaP01, KCN (to be used only with great care and never added to a strongly acid solution) .) Watch out for redoa systems. IV) Set up an electric cell with 0.1 M ZnZ+ (aq) around a zinc electrode (in a pomus cup) and 1 M CU~+ around a copper electrode in a small beaker. Measure F. Use the same zinc half-cell in all future measurements. Change the C U ~ + con- centration to 0.1 M, to 0.01 M, and compare your results with those predicted from the Nernst equation (allom esti- mation of experimental uncertainty, not accurate value). V) Measure E (against the standard zinc half-cell) for a dilute Cu(I1) solution (of known total Cu(II) concentration) in a 1 M solution of one of the bases you used in (11) or (III) above. Repeat for some other bases. Calculate the concen- tration of Cu2+ (aq) in each solution, and the equilibrium constant between Cu2+ (aq) and the complex between Cu(II) and each base you use. Calculate AG" values. Cal- culate KsD for any insoluble products formed. VI) Measure E, and calculate [Cuz+l and K for the redox sys- tems you identified. Hints A) Your results in (I) will be most visible if you start with about 1 ml of 0.1 M Cua+ and add a 1 M solution of the po- tential base to it dropwise, noting changes with each drop after shaking. After adding about 1 ml of the 1 M hase, pour in enough of the 1 M base to give a total volume of 10 ml. Color comparison is easier if you look down the test tube at a white surface with a piece of white paper wrapped around the tube. Start with 1 M Cu2+ sdution~ if color changes are hard to see. B) Add 1 ml of 0.1 M CuZ+ and 1 ml of 1 M solution of s single base to each of two tubes. The results should match thnse in (I). Now add 1 ml of H20 to one tube and 1 ml of a 1 M solution of a second base to the other tube. Compare the Table 1. Some Typical Results Initial base CI - Br- OH - NHs COP m,a- Orieinal system It blue soh gn blue soh It blue ppt dk blue soh It blue ppt It blue ppt ~d&d base 51- bk PP~ + PO,# - It blue ppt 4 COF It blue ppt - ? no change NHa dk blue soh --r no change OH - It bl-e ppt - Br- an blve Boln ~o order of base drength is: S2- > POa,- > OH- > NHI > Cox1- > Er- > CI > Hz0 Adding I - or CN - to Cu** (aq) sives redox Table 2. Some Typical Results [Cu2+(sql l [Cu(base)n"l 'Base 8 (Vl M M K HO 1.14 1 10-4 lo-. - [CuOH+lm*l/[Cu~+l C1- 0.94 10 -7 10-2 lo-. = [Cu~*l[C1-l~/[C~Cl$*l Br- 0.90 10-8 lo-% lo" = [CuZilBr-l'/[CuBrrl-I COP 0.78 lo-" solid 10-31 = [CU~+I[COI~-I NHI 0.76 lo-" 10-3 = [Cu=*l[NHal~/[Cu(NH~)~~*l I - 0.71 In-" redoa, solid lo-a* - [Cu~+I~[I-lVl1~-1 OH - 0.70 lo-" solid lo-" = .[c~*l[OH-p p0.1- 0.68 lo-16 did = [cu~*I~Po,~-I' St- -0.03 10.' solid lo-= = [Cu~+l[s'-l lod8 redox, 10-2 M K = [c~~+l[cN-l~~~[oH-I CN - -0.46 [Cu(CN)n-l ICNO -1'" q-hese obtained from 25 ml1 M bsse plus 0.25 ml 1 M Cua* using a coppr electmde against a zinc electmde in 1 M ZnSO*. Errors in presumed for- mulas of and in degree of hydrolysis of bases account for deviations from pblished values. Gwd analytical data are needed to get good K's. Volume 52, Number 3, March 1975 / 185

A comparison of base strengths using visual and EMF observations

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Page 1: A comparison of base strengths using visual and EMF observations

Many instructors use the lecture experiment in which Cl-, 3 M NH3, Br-, 7 M NH,, I-, SzO$-, SZ-, and CN- are added sequentially t o a dilute aqueous solution of sil- ver nitrate. The alternate formation and solution of pre- cipitates, with some variation in color, provide excellent evidence for different degrees of dissociation of a series of complexes and precipitates. The results can, of course, be used t o compare the basicity of the species added with re- spect t o silver ion as an acid.

We have developed a similar series, based on copper ion as the acid. The colors are more varied, and reduction of Cu(II) is introduced as a complication. Many other acids could be used hut we settled on Cu(I1) because of low cost, the varied colors involved, and the ease with which the concentration of copper(II) ions can be measured using a copper electrode, a constant EMF zinc electrode, and a low-cost voltmeter.

The chemicals are stocked as 1 M aqueous solutions of: ZnSOd, CuSOn, NaC1, NaBr, KI, NaOH, NH3, KCN, Na2S, NaaP01, NazC03. Other hases could, of course, be added so tha t the experiment can he as open-ended as the student and instructor wish.

The following experiments are performed. Chemical equations are required for all observed reactions

I) Prepare a 0.1 M CuSO* solution and measure, or calcu- late, the concentration of each principal species presence. (pH papers or pH meters are used.) Calculate K for cop- per(II) ion acting as an acid in water. (pH and atom bal- ance give (Hi) , (CuOH+), (Cue+), (SO?), (OH-), (HzO).)

II) Design and perform a series of experiments with aqueous 1 M C1-, Br-, I-, OH-, NHa, and NazS, to arrive at a table of relative bsse strengths toward aqueous Cu(II). Note that one of these species causes redox with Cu(I1). Eliminate it

J. A. Campbell Horvey Mudd College

Clarernont, California 9171 1

from your list of simple bases. ~=perimentally verify the

A Comparison of Base Strengths Using Visual and EMF Observations

oxidized product. III) Try one or more possible other bases and add them to the

table. (We provide 1 M NazCOs, NaaP01, KCN (to be used only with great care and never added to a strongly acid solution) .) Watch out for redoa systems.

IV) Set up an electric cell with 0.1 M ZnZ+ (aq) around a zinc electrode (in a pomus cup) and 1 M C U ~ + around a copper electrode in a small beaker. Measure F . Use the same zinc half-cell in all future measurements. Change the C U ~ + con- centration to 0.1 M, to 0.01 M, and compare your results with those predicted from the Nernst equation (allom esti- mation of experimental uncertainty, not accurate value).

V) Measure E (against the standard zinc half-cell) for a dilute Cu(I1) solution (of known total Cu(II) concentration) in a 1 M solution of one of the bases you used in (11) or (III) above. Repeat for some other bases. Calculate the concen- tration of Cu2+ (aq) in each solution, and the equilibrium constant between Cu2+ (aq) and the complex between Cu(II) and each base you use. Calculate AG" values. Cal- culate KsD for any insoluble products formed.

VI) Measure E , and calculate [Cuz+l and K for the redox sys- tems you identified.

Hints A) Your results in (I) will be most visible if you start with

about 1 ml of 0.1 M Cua+ and add a 1 M solution of the po- tential base to it dropwise, noting changes with each drop after shaking. After adding about 1 ml of the 1 M hase, pour in enough of the 1 M base to give a total volume of 10 ml. Color comparison is easier if you look down the test tube a t a white surface with a piece of white paper wrapped around the tube. Start with 1 M Cu2+ sdution~ if color changes are hard to see.

B) Add 1 ml of 0.1 M CuZ+ and 1 ml of 1 M solution of s single base to each of two tubes. The results should match thnse in (I). Now add 1 ml of H20 to one tube and 1 ml of a 1 M solution of a second base to the other tube. Compare the

Table 1. Some Typical Results

Initial base CI - Br- OH - NHs COP m,a- Orieinal system It blue soh gn blue soh It blue ppt dk blue soh I t blue ppt It blue ppt ~ d & d base

51- bk P P ~ + PO,# - It blue ppt 4

C O F It blue ppt - ? no change NHa dk blue soh --r no change OH - I t bl-e ppt - Br- an blve Boln ~o order of base drength is: S 2 - > POa,- > OH- > NHI > Cox1- > Er- > C I > Hz0 Adding I - or CN - to Cu** (aq) sives redox

Table 2. Some Typical Results

[Cu2+(sql l [Cu(base)n"l 'Base 8 (Vl M M K

HO 1.14 1 10-4 lo-. - [CuOH+lm*l/[Cu~+l C1- 0.94 10 -7 1 0 - 2 lo-. = [Cu~*l[C1-l~/[C~Cl$*l Br- 0.90 1 0 - 8 lo-% lo" = [CuZilBr-l'/[CuBrrl-I C O P 0.78 lo-" solid 1 0 - 3 1 = [CU~+I[COI~-I NHI 0.76 lo-" 10-3 = [Cu=*l[NHal~/[Cu(NH~)~~*l I - 0.71 In-" redoa, solid lo-a* - [Cu~+I~[I-lVl1~-1 OH - 0.70 lo-" solid lo-" = .[c~*l[OH-p p0.1- 0.68 lo-16 d i d = [cu~*I~Po ,~- I ' St- -0.03 10.' solid lo-= = [Cu~+l[s'-l

lod8 redox, 10-2 M K = [ c ~ ~ + l [ c N - l ~ ~ ~ [ o H - I

CN - -0.46 [Cu(CN)n-l ICNO -1'"

q-hese obtained from 25 ml1 M bsse plus 0.25 ml 1 M Cua* using a coppr electmde against a zinc electmde in 1 M ZnSO*. Errors in presumed for-

mulas of and in degree of hydrolysis of bases account for deviations from pblished values. G w d analytical data are needed to get good K's.

Volume 52, Number 3, March 1975 / 185

Page 2: A comparison of base strengths using visual and EMF observations

two tubes to see which base is atr0ngt.r. Repeat for other pain of hases. If you plan and observe pruperly you will nur have to run through all possible pairs. Systematize ynur data in a table. ( ~ e ~ ~ a b l e 1.)

C) Use the remaining solutions from (II) and (m) (preparing more only if needed) for thb measurements of Gin expen- ments (V) and (VI).

We provide a derivation of the Nemst equation of the form

6, - 6 , = (0.@592/n) log (Q,/QJ

where Q has the same form as K, the equilibrium con- stant for the system in which the concentration is being varied. The equation is used in this experiment to calcu- late [CuZ+] as a function of 8 for the zinc-copper cell: Zn(c) + C U ~ + (aq) = ZnZ+ (aq) + CU(C). Q, of course,

uses experimental concentrations,. not equilibrium con- centrations as does the corresponding equation for K.

The experiment can be written with any desired degree of guidance for the students; it provides a very wide po- tential scope for investigating both acid-base and redox changes (and relating the two to one another), it involves interesting color changes (some of which are dramatic and easy to see, othen of greater subtlety), and it gives a deep insight into the power of thermodynamics to measure numbers varying over a great range. Table 2 gives some typical results. Note that many of the values of [ C U ~ + ] are too ridiculously low to be estimated in any other way. The lack of direct agreement between the calculated and pub- lished K's gives an experimental basis for a discussion of hydrolysis, i.e. competititive acid-base phenomena. Al- lowance for hydrolysis gives results in good agreement with published data.

186 / Journal of Chemical Education