A Comparison of Base Strengths Using Visual and EMF Observations
J. A. Campbell Horvey Mudd College Clarernont, California 9171 1
from your list of simple bases. ~=perimentallyverify the 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 estimation 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 concentration 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. Calculate KsDfor any insoluble products formed. VI) Measure E , and calculate [Cuz+l and K for the redox systems 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 1M solution of the potential 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 1ml of 0.1 M CuZ+ and 1ml of 1M 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 1M solution of a second base to the other tube. Compare the
Many instructors use the lecture experiment in which Cl-, 3 M NH3, Br-, 7 M NH,, I-, SzO$-, SZ-,and CNare added sequentially t o a dilute aqueous solution of silver nitrate. The alternate formation and solution of precipitates, 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, b e used t o compare the basicity of t h e species added with respect t o silver ion a s a n acid. We have developed a similar series, based on copper ion a s t h e acid. T h e colors are more varied, and reduction of Cu(II) is introduced a s a complication. Many other acids could be used h u t 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. T h e chemicals are stocked a s 1M aqueous solutions of: ZnSOd, CuSOn, NaC1, NaBr, KI, NaOH, NH3, KCN, Na2S, NaaP01, NazC03. Other hases could, of course, be added so t h a t the experiment can he a s open-ended a s the student and instructor wish. T h e following experiments are performed. Chemical equations are required for all observed reactions Prepare a 0.1 M CuSO* solution and measure, or calculate, the concentration of each principal species presence. (pH papers or pH meters are used.) Calculate K for copper(II) ion acting as an acid in water. (pH and atom balance give ( H i ) , (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 I)
Table 1. Some Typical Results Initial base system ~ d & dbase Orieinal
51PO,# -
COF NHa OH -
Br~o order of
m,a-
CI -
Br-
OH -
NHs
COP
It blue soh
gn blue soh
It blue ppt
dk blue soh
It blue ppt
--
bk P P ~ It blue ppt It blue ppt dk blue soh I t bl-e ppt an blve Boln
base drength is: S 2 -
Adding I - or CN - to Cu** (aq)
--r
It blue ppt +
4
no change
?
no change
> POa,- > OH- > NHI > Cox1- > Er- > C sives redox
I
> Hz0
Table 2. Some Typical Results 'Base
8
HO C1BrCOP
1.14
0.94
1 10-7
0.90
10-8
0.78 0.76 0.71 0.70 0.68
lo-" lo-"
[Cu(base)n"l M
lo-. lo-. lo"
10-4 10-2
lo-%
St-
-0.03
10.'
solid 10-3 redoa, solid solid did solid
CN -
-0.46
lod8
redox,
NHI
IOH -
p0.1-
q-hese mulas of
(Vl
[Cu2+(sqll M
In-"
lo-" lo-16
10-2
10-31
-
K
[CuOH+lm*l/[Cu~+l [Cu~*l[C1-l~/[C~Cl$*l [CuZilBr-l'/[CuBrrl-I
= = = [CU~+I[COI~-I = [Cu=*l[NHal~/[Cu(NH~)~~*l
-
lo-a* [Cu~+I~[I-lVl1~-1 lo-" = .[c~*l[OH-p = [cu~*I~Po,~-I' lo-= = [Cu~+l[s'-l M
[c~~+l[cN-l~~~[oH-I
K = [Cu(CN)n-l ICNO -1'"
obtained from 25 ml1 M bsse plus 0.25 ml 1 M Cua* using a c o p p r electmde against a zinc electmde in 1M ZnSO*.Errors in presumed forand 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
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 ~ ~ a1.) ble
(m)(preparing more only if needed) for thb measurements of Gin expenments (V) and (VI).
C) Use the remaining solutions from (II) and
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 constant for the system in which the concentration is being varied. The equation is used in this experiment to calculate [CuZ+] as a function of 8 for the zinc-copper cell: Zn(c) C U ~ +(aq) = ZnZ+ (aq) + CU(C).Q, of course,
+
186
/ Journal of Chemical Education
uses experimental concentrations,. not equilibrium concentrations 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 potential 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 published K's gives an experimental basis for a discussion of hydrolysis, i.e. competititive acid-base phenomena. Allowance for hydrolysis gives results in good agreement with published data.
