In the Classroom
Experimentation and Group Discussion as a Means of Determining Solubility Rules Karen E. Stevens Department of Chemistry, Whitworth College, Spokane, WA 99251;
[email protected] Students often dislike studying solubility rules and precipitation equations in part because of the memorization required. Several authors (1, 2) have suggested various sets of generalized solubility rules intended to aid students in their memorizing. This paper presents a technique used to teach this topic in a manner that shows the experimental basis of these “facts”. We have found that this significantly adds to student interest, enjoyment, and comprehension. In our method, the emphasis is on having the students see that the solubility “rules” are based on experimental observations, rather than merely being a jumble of complicated and strange facts they are expected to know. This series of experiments is presented in class as a demonstration and followed by group discussion, which leads to a determination of the solubility rules. In this activity, we will be able to classify anions as either generally soluble or generally insoluble, while cations will be categorized as always soluble or as a cation that follows the rule of the anion. This differs significantly from standard methods used to present solubility rules, in which memorized rules are applied to predict precipitation equation products.
(from the end of the alphabet) is the anion. For example, in AW, A must be a cation, while W is an anion. This represents the same notation used for all ionic formulas: for example, in KBr, K+ is the cation and Br᎑ is the anion. Examining our first reaction above, we find that we have two cations (A and B) and two anions (W and X). The products must be written so that the cations and anions are paired in the way opposite to their pairing in the reactants. Thus we must write our reaction as AW + BX → AX + BW. Note that we still don’t know whether AX or BW is the substance responsible for the green precipitate formed. Writing full “equations” for all four reactions above gives: Expt. 1: AW + BX → AX + BW Expt. 2: CW + BX → CX + BW Expt. 3: AW + BY → AY + BW Expt. 4: AZ + BX → AX + BZ
Now we can summarize our results this far as follows: Based on Expt. 1, either AX or BW is a green solid. Based on Expt. 2, either CX or BW is a white solid.
Method
Based on Expt. 3, neither AY or BW is a solid.
Several simple reactions are presented. Prepared 1.0 M aqueous solutions are mixed together and the resulting mixture is examined to see whether a precipitate has formed. (The solutions are labeled with letter codes to avoid complicating the issue with formulas and to show that the classification can be determined without any knowledge other than the current experimental observations. In case some students have already memorized some solubility rules, this technique requires them to think through the results with everyone else.) The experiment is then repeated several times with a slight change in ions used in order to see how the products differ. For this initial set of data, solutions were chosen so that some of the precipitates are a different color than the others, making the identifications somewhat more simple, although it is not required for the analysis. One set of example results is shown here: Expt. 1: AW + BX produces green precipitate upon mixing,
Based on Expt. 4, either AX or BZ is a green solid.
We can now fairly easily see that the green solid must be AX and the white solid must be CX. Similarly, BW and BZ must be soluble ionic species because we see no difference between the results of the two experiments producing green solids. The four reactions involve seven different ions: A, W, B, X, C, Y, and Z. We can quickly separate them into cations (A, B, C) and anions (W, X, Y, Z) on the basis of their position in each formula. Furthermore, we can categorize them into one of the four groupings below: Cation 1 Cation 2
Anion 1
Expt. 2: CW + BX produces white precipitate upon mixing, Expt. 3: AW + BY produces no precipitate, Expt. 4: AZ + BX produces green precipitate upon mixing,
where BX, CW, and BY are all clear, colorless aqueous stock solutions and AW and AZ are clear, green aqueous stock solutions. To analyze these results, one must recall that the first species in each pair (represented by a letter from the beginning of the alphabet) is the cation and the second half of each pair
Anion 2
Always soluble (such as alkali metal cations or ammonium ion) Follows rule of the anion, i.e. forms solids with "insoluble anions" (such as most alkaline earth or transition metal cations) Generally soluble (such as nitrate, sulfate, chloride, bromide, or iodide anions) Generally insoluble (such as hydroxide, sulfide, carbonate, or phosphate anions)
The reasoning for assignment of each ion is as follows: Anion X : AX was the precipitate formed in reaction 1. Thus, anion X must be categorized in group 2 (generally insoluble). Cation A : Cation A must be of type 2. If A were of cation type 1, it would always be soluble, which would not have allowed solid AX to form.
JChemEd.chem.wisc.edu • Vol. 77 No. 3 March 2000 • Journal of Chemical Education
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In the Classroom Anion W : Because cation A will form solids with any anion of type 2, we know that anion W must be of type 1 because the initial aqueous solution of AW did not form a solid. Cation B : As a reactant, BX is an aqueous solution, and the product BW was also determined to be aqueous (nonsolid). Since anion W is of type 1, no solid is expected to form. However, anion X was determined to be generally insoluble, which “should” have caused the reactant BX to be a solid. The fact that BX was instead an aqueous solution tells us that cation B must have “overridden” the insolubility of X. Cation B is then of type 1, which is always soluble, no matter the type of anion. Cation C : In reaction 2, the white solid was determined to be CX. This allows us to quickly categorize cation C as type 2 because it formed a solid with the generally insoluble anion X. Anion Y : Anion Y is determined by examining the properties of the two species in which it appears. In reaction 3, BY is an aqueous reactant and AY is an aqueous product. Cation B has already been determined to be of type 1 (always soluble) and cation A of type 2 (follows rule of anion). Since cation B will “force” any compound to be soluble, this tells us nothing about the character of Y. However, because cation A will follow the rule of the anion and we know that AY is an aqueous solution, we can assign Y as an anion of type 1 (generally soluble). Anion Z: Finally, anion Z appears in experiment 4 as a soluble reactant (AZ). We recall again that A is a type 2 cation, which will form solids if the anion is generally insoluble. Since AZ was a soluble aqueous reactant, we can classify Z as a type 1 anion (generally soluble). The product BZ is soluble for two “reasons”: B is an always-soluble cation (type 1) and Z is a generally soluble anion (type 1).
We have now categorized all seven of the ions: Type 1 Cations B
Type 2 Cations A, C
Type 1 Anions W, Y, Z
Type 2 Anions X
After the classifications above have been discussed, a real chemical system can be presented that fits the above data. In this set of reactions, we have used the assignments A = Ni2+, W = Cl ᎑, B = Na+, X = OH ᎑, C = Mg2+, Y = NO3᎑, Z = SO42᎑, so that the four equations are: 1. NiCl2(aq) + 2NaOH(aq) → Ni(OH)2(s) + 2NaCl(aq) 2. MgCl2(aq) + 2NaOH(aq) → Mg(OH)2(s) + 2NaCl(aq)
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3. NiCl2(aq) + 2NaNO3(aq) → Ni(NO3)2(aq) + 2NaCl(aq) 4. NiSO4(aq) + 2NaOH(aq) → Ni(OH)2(s) + Na2SO4(aq)
In this example, the green solid AX is nickel(II) hydroxide, Ni(OH)2, and the white solid CX is magnesium hydroxide, Mg(OH)2. (Note that, for simplicity, the reactions presented earlier do not show correct stoichiometry. Attention is focused on the outcome of understanding the basis of the solubility rules, not on balancing equations or on stoichiometry.) Discussion This classification can be stopped after an example such as the one presented above. In this case, just four reactions were required to classify seven ions. Further steps can be included to show as much detail as desired. For example, some exceptions to the solubility guidelines can be presented. We have used this technique as a lecture demonstration of the four reactions. This is followed by discussion time in which the students break into small groups to try to classify the ions as explained above. Student involvement and interest in this activity is very high, often resulting in raised voices as they excitedly assign each ion to its proper classification. This activity could easily be modified into a laboratory activity in which the students perform the reactions themselves. This could perhaps be followed by more detailed reactions or questions to solve. For example, a more challenging set of solutions that produce only white precipitates when mixed (such as MgBr2, K2CO3, CaCl2, Na2CO3, and MgSO4) could be given to the students disguised as before by letter codes such as AX, BY, etc. The students can then mix pairs of solutions and observe the products of the reactions. This exercise allows students to see the rationale for assigning the hydroxide ion as an insoluble anion, for example, based on experimental evidence that they have actually seen and analyzed. As a result, solubility rules become less a tangle of special “rules” for students to know and more an experimentally based classification system that they have personally helped to develop. Literature Cited 1. Heimerzheim, C. J. J. Chem. Educ. 1941, 18, 377. 2. Monroe, M.; Abrams, K. J. Chem. Educ. 1984, 61, 885.
Journal of Chemical Education • Vol. 77 No. 3 March 2000 • JChemEd.chem.wisc.edu
