tested demonstrcrtions -,

GEORGE L. GILBERT Denison University

Granville, OH 43023

An Ex~anded Silver Ion Eauilibria Demonstration: including Use of the'~ernst Equation

and Calculation of Nine Equilibrium Constants

Robert Hunt Anderson Western Michigan University Kalamazoo, MI 49008

Amodification of the demonstration described by Shakh- ashiri, Dirreen, and Jeurgens (1, 2) and originally pro- posed in a slightly different form by Schwenck (3) can with very little additional effort yield a quantitative measure of the silver ion concentration a h r each of the nine steps. From this data the six solubility products and three forma- tion constants of complex ions can be readily calculated.

The entire demonstration can be done in 20 minor less depending on the amount of explanation. The calculation of the equilibrium constants and a description of conver- sion of millivolt readings from the concentration cell into silver ion concentrations could easily fill one class period. This modified demonstration gives more results with less effort than any other demonstration with which I am ac- quainted.

Modifications Equipment Setup

Use two 600-mL beakers instead of one. One is for the reference half-cell and contains 200 mL of distilled water and 10 mL of 0.1 M AgN03. They are wnnected by a %3- cm wide strip of filter paper previously dipped in 1.0 M NH&!03 and reaching into both solutions. This ads as a salt bridge. A pure silver wire electrode is mounted to dip into each heaker, and they are connected to a digital milli- voltmeter or pH meter on the millivolt scale.

The reference beaker of 0.00476 M AzNO- is connected to the negative or reference input. The s>mde or test bea- ker goes to the test or oositive input, making all the read- ings>egative. If the meter does not accept negative values, just interchange the two leads and attach a negative sign to all readings.

The test beaker sits on a magnetic stirrer and is mildly stirred to provide mixing after each addition. The refer- ence beaker is conveniently placed next to the test heaker on a laboratory jack. Thus, it can be raised to the same level as the test beaker as additions are made, keeping the levels approximately the same. However, this is not neces- sary because siphoning between the beakers does not seem to be significant during the 10 or 15 min of the demonstra- tion. If the experiment is repeated later, a new reference solution should be prepared because there can be diffusion of cyanide ion into the reference solution over a period of

30-60 min. This will significantly reduce the concentration of uncomplexed (free) silver ion in the reference solution.

The figure shows a diagram of the setup. A digital volt- meter with a display large enough to be seen by the class is desirable though not absolutely necessary.

Adjusting the Amount of Reagent

Some changes in the amounts of reagents are required in order to assure a known excess of the precipitant or com- plexing agent in each case. This is necessary in order to easily calculate the equilibrium wnstants. In five out of nine cases Shakhashiri (1,2) did not use an excess of the reactant. The amounts and concentrations of the reagents used are given in the table. They provide a known excess of each reagent and also satisfactorily precipitate Ag2C03 and Ag,0(AgOH(H20)).

I was not able to obtain any precipitate when 2 mL of 0.1 M NaHC03 was added to 200 mL deionized water and 10 mL 0.1 M AgN03, as called for in ref 2. Even using 36 mL of 0.1 M NaHC03 gave no immediate precipitate, though a slight turbidity appeared after 2 or 3 min. Theoretically, the ion product exceeds K, by about 100-fold, thus predict- ing a precipitate, but the result is nebulous and certainly not satisfactory for a demonstration.

Shakhashiri et al. (I) used 0.1 M Na2C03, which gives a precipitate, though it is slightly gray-brown when first formed from silver hydroxide or oxide due to the higher -

I Digital voltmeter I

Ag wir eF

600-mL ' Reaction

beaker

- Stirring bar

filter paper strip soaked in 1.0 M NH.NO,

for a salt bridge

Experimental setup.

940 Journal of Chemical Education

Experimental Data and Results

Reagent AgN03 NazCCb NaOH KC1 NH3 KBr NazSz03 KI NaCN NarS NaHC03

M 0.100 0.200 1 .OO 0.100 3.00 0.100 0.100 0.100 0.100 0.100 in each

Vadded 10 mL 10 50 20 20 15 40 15 30 10

h t a l 210 220 225 245 265 280 320 335 365 375

Run 1 +.I7 -99.5 -204 -351 -204 4 1 6 -534 -632 -907 -1163

R u n 2 +1.1 -117.3 -195 -353 -204 4 2 1 -538 -635 -916 -1158

Run 3 -13.0 -118.3 -194.6 -368 -204 4 3 6 -554 -647 -920 -1168

R u n 4 +5.9 -100.6 -177.3 -350 -204 4 1 5 -533 -204 -906 -1153

k l + I Run 4 ,00599 9.45 x lo4 4.76 x lo6 1.19 x lo-' 5.69 x lo4 4.52 x 4.56 x lo-'' 1.13 x 10-l3 2.23 x 1 0 ~ " 1.47 x

Run I 10.8 4.63 4.40 1.60

Run 2 2.88 4.25 3.29 1.55

Run 3 4.87 8.23 3.50 1 .46

Run 4 5.30 8.46 4.86 1.38

pH. He tried to avoid this by diluting the Na2C03 in the 200 mL of water before adding the AgN03. The resulting slightly brown precipitate is satisfactory, though not as white as a good silver carbonate precipitate obtained with a buffered carbonate solution. I used 10 mL of a buffer con- taining 0.2 M NaHC03 and 0.2 M Na2C03 and obtained a good white precipitate. The pH is 9.7, which agrees with a theoretical value of 9.6 when activity coefficients for an ionic strength of I = 0.8 are used, as given by Skoog and West (4).

The concentrations given in Shakhashiri's book of chem- ical demonstrations (2) gave no oxide precipitate. With my buffered carbonate even 30 mL of 0.1 M NaOH was not enough to change the white carbonate to the oxide, so I used 5.0 mL of 1.0 M NaOH. This gave a good black-brown precipitate.

Procedure and Results The amounts and concentrations used are given in the

table. Both 600-mL beakers start with 200 mL of distilled water. With 10 mL 0.1 M AgN03 in the reference beaker the millivolt reading is a measure of silver ion concentra- tion in the distilled water in the test beaker. (These read- ings are not reported in this paper.) When the 10 mL of 0.1 M AgN03 are added to the test beaker, the millivolt read- ing becomes a measure of the difference in silver ion con- centration in the two beakers. It should be close to zero. Because the volumes are measured with graduated cylin- ders, the values are a few millivolts either ~ l u s or minus. hen each reagent in turn is added to the test beaker, and the millivolt reading is recorded after mixing occurs. - -

The results of four runs are given in the table. The silver ion concentration is only shown for run 4, which will be used in sample calculations of four of the equilibrium con- stants.

Calculations The activity coefficients will he omitted in most of these

calculations. The Nernst equation for a silver ion half-cell is

where [MI should really be the activity, and n = 1 for sil- ver.

Silver Ion Concentration

Reference Beaker

In the reference beaker,

Neglecting the activity coefficient, which is 0.924 accord- ing to Skoog and West (41, we get

Concentration Cell

For the concentration cell, with % read in volts, subtract- ing eq 2 from eq 1, we get

where %, is the reading for the test half-cell from eq 1 when %=&.

Solving eq 3 for the silver ion concentration, we get

[Agl = 0.00476 x (4)

Volume 70 Number 11 November 1993 941

For KI addition in run 4, % =-0.628 V Thus, substituting into eq 4, we get

[Ag+] = 0.00416 x 1 0 ~ ~ ~ ~ ~ ~ ~ ~ ~ ~ = 1.13 x 10-l~ (5)

In my experiments I used a small BASIC program to cal- culate the silver ion concentrations.

Sample Calculations

Solubility Product Constants

The K.;s for AgOH, AgCl, AgBr, and AgI are the easiest to calculate and will be shown for AgI in run 4.

1.5 mmol r + 1.0 mmol Ag+

leaves

1.5 - 1 = 0.5 mmol I~

in excess. Thus,

1.5 - 1.0 117 =-=0.00149 M 335

From the Nernst equation as in eq 5, [Ag+] = 1.13 x lo-"

K,, = [Agtl[17 = 1.13 x loA3 x 0.00149 = 16.9 x 10-l7

Formation Constants of Complex Ions

The formation .constants of the complex ions Ag(NH&, Ag(CN);, and Ag(SzO& will be calculated for only the lat- ter in run 4 because the method is similar in all three cases.

From eq 4, [Ag+] = 4.56 x lo-''

1.0 mmolAg+ + 1.0 mmol Ag(SzO3)f

In the complex, there are

4 mmol S20;-- 2 mmol S20;

in 320 mL. Thus,

The formation constant of Ag(S20&is thus calculated as follows.

pH-Dependent Calculations

The calculations for AgzS and Ag2C03 are both depen- dent on the pH. The calculation for silver sulfide is shown first. In an earlier run, pH paper indicated a pH of 11.5 for the solution in equilibrium with the Ag2S. This was con- firmed with a pH meter reading in run 3, which gave a pH of 11.40. This value was then used for calculations in runs 3 and 4.

From the pH,

Also, 1.0 mmol of S" was added, and 0.5 mmol of S" is present in the Ag2S solid. Thus, we get

The calculation for Ag2C03 in run 4, where the pH mea- sured with a pH meter was 8.90, is given below.

From the pH,

[Htl = 1.26 x loJJ

where 2 mmol HC0i and 2 mmol CO; are in the reagent while 0.5 mmol of CO$- are in the Ag2C03 assuming that all of the silver precipitates.

From eq 6,

Substituting this into eq 7, we get

So the solubility product is found

Discussion of Results The calculated equilibrium constants are all within a

factor of 4 of the literature values except perhaps the thio- sulfate values. In the 1964 compilation of stability con- stants by Sillen (5) there is some scatter of these K's cen- tering around 20.0 x 10" in agreement with our results, though the 1971 supplement (6) gives one value of 4.3 x 10" at 25 T. Our results seem quite satisfactory because activity coefficients have been neglected, and the experi- ment is a demonstration-not a highly precise quantita- tive measurement.

Versatility in Applications

This demonstration can be used at several different lev- els of souhistication. The aualitative observation of ureciu- itates aAd complexes along with chemical equations'can 6e done as oro~osed bv Schwenck (3). Shakhashiri and co- workers'adktional"ly used literature equilibrium con- stants to derive the eauilibrium constants of the various conversions taking place (1,2).

By adding the reference solution, salt bridge, and milli- voltmeter, you can introduce a good example of a wncen- tration cell along with its Nernst equation and then relate them to this type of popular analytical method. This makes it suitable for an analytical chemistry or physical chemis- try course. Perhaps some second-semester general chemis- try courses could use this to illustrate the use of the Nernst equation.

This demonstration can be also extended by using the results to calculate equilibrium constants in qualitative or quantitative analysis courses or wherever equilibrium constants are extended to solubility products or complex ion formation. Finally, activity coefficients can be intro- duced into the calculations to show their effect in changing and improving the results. Perhaps by altering the ionic strength and extrapolating to zero this could be used as an experiment in the physical chemistry laboratory to obtain activity coeflicients.

Using [Ag+l from run 4, we get Literature Cited 1. Shakhashiri B.Z.: i3reen.G. E.:Juelgens, F. J Chem. Educ. 1980,57,81%814.

942 Journal of Chemical Education

2. ShaWlashui, B. 2. C h m i m l &monsfmtions. Vol 1; University of Wiseonsin: Mad% eon. WI. 1983: pp 307313.

3. Schwenek. J. R. J Ckem. Edue. 1859.36.45. 4. Skmg. D. A.; West, D. M.Auly t iml choAistry--~" Infrduelion, 4th ed.: Saundera:

Explosion of Nitrogen Triiodide: A Safer and Cleaner Demonstration

Gordon F. Hambly and Robert Peters John Abbon College P.O. Box 2000 Ste. Anne de Bellevue, QC, H9X 3L9, Canada

The explosive decomposition of nitrogen triiodide, as de- scribed bv Shakashiri ( I ) . is an effective demonstration, particula;ly for large clasbes, but i t does have two poten- tial drawbacks noted bv the authors and experienced on numerous occasions by us. Because a relatively large amount of iodine is used ( 2 3 g) to prepare the explosive, which is subsequently divided/nto tno-or three the nitrogen triiodlde may not explode totally, particularly when it is still slightly damp. Secondly, cleanup can be te-

dious particnlarly when the nitrogen triiodide sample has not exploded completely.

Both of these problems are circumvented by carrying out the reaction on a much-reduced scale. Students are sur- prised by the decibel level generated from the detonation of such a small amount of sample, and the small-scale preparation leads to rapid drying and complete destmc- tion of the sample. Cleanup merely involves discarding what is left on the filter DaDers. S~ecificallv. 0.2-0.3 e of iodine are placed in a 3 0 - k i beaker with 5 ;L of concen- trated aqueous ammonia and stirred briefly. The mixture is allowed to to stand for 5 min, and the supernatant liquid is decanted from the brown solid. I t is then washed 5 times with water that is decanted off each time after allowing most of the brown solid to settle. The brown solid is then scraped onto a few pieces of filter paper and patted to ab- sorb most of the water, then scraped onto a new filter paper. In our hands the nitrogen triiodide always has ex- ploded totally 45 min later when touched with a long pole. The chemical equations involved in the preparation and decomposition of the nitrogen triiodide are documented thoroughly elsewhere ( I ) .

Literature Cited: 1. ShaLashiri, B. S. CkemimlDemonsimfions: Uniu. Wimonah Press, 1983, Vol.1, p 96.

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