The purpose of this experiment is to investigate soluble and insoluble ionic compounds and to make predictions regarding solubility.
Solutions consist of two or more substances combined to form a homogeneous mixture. The components present in lesser quantities are solutes dissolved in the larger quantity component, the solvent. The most common solutions involve solid solutes dissolved in water to form aqueous solutions.
When a solute dissolves within a solvent, the solute can be described as “soluble” within that solvent. Solubility plays an important role within chemistry, as it is a significant method for adding or removing chemicals from a system. Understanding solubility allows hospital pharmacists to correctly determine if a particular medication will dissolve in an IV solution, or help an environmental chemist remove certain chemicals from water by precipitation. In this lab, students explore a bit about solutions that contain ions that, while individually soluble, may combine to form insoluble substances.
A precipitation reaction is a common example of a double displacement reaction in which two ionic compounds react to form two new compounds. The initial two ionic compounds are soluble in water. When mixed, the cation of each reactant combines with the anion of the other reactant as shown in Reaction IS.2. Remember that cations are written first in ionic formulas (A and B in the example below) followed by the anions (X and Y in the example below).
\begin {equation} \ce{ AX + BY -> AY + BX} \end {equation}
During a double displacement reaction, there is no change in the charges of individual ions. However, the subscripts for each ion may change because the new compounds must contain the correct ratio of cations and anions to form a neutral compound (total charge of zero). One driving force of a double displacement reactions is the formation of an insoluble product, or precipitate. This is known as a precipitation reaction. Two soluble ionic compounds combine in solution to exchange ions in such a way that an insoluble ionic compound forms and precipitates out of solution. A precipitate will appear cloudy (can be white or colored) within the water, and depending on the size of the particles it will stay suspended or settle to the bottom of the solution. In Reaction IS.4, aqueous (aq) barium chloride, \(\ce {BaCl2}\), reacts with aqueous sodium sulfate, \(\ce {Na2SO4}\), to form solid barium sulfate, \(\ce {BaSO4}\) (the precipitate) and aqueous sodium chloride, NaCl. Barium sulfate is an insoluble solid in water and will be written with an (s) for solid as its physical state. Sodium chloride is soluble in water and will be written with an (aq) for aqueous to represent that it is dissolved in water.
\begin {equation} \ce{ BaCl2(aq) + Na2SO4(aq) -> BaSO4(s) + 2NaCl(aq)} \end {equation}
Before the reaction occurs, we have two solutions that contain dissolved ions. In a solution of \(\ce {BaCl2}\), there are \(\ce {Ba^{2+}}\) ions and \(\ce {Cl^{-}}\) ions in a ratio of 1:2, respectively. In a solution of \(\ce {Na2SO4}\), there are \(\ce {Na^{+}}\) ions and \(\ce {SO4^{2-}}\), in a ratio of 2:1, respectively. When these two solutions are mixed, specifically \(\ce {Ba^{2+}}\) ions and \(\ce {SO4^{2-}}\) ions bond together to make the insoluble solid precipitate of \(\ce {BaSO4}\). Note that even though sodium and chlorine have subscripts of “2” on the reactant side, the two ions \(\ce {Na^{+}}\) and \(\ce {Cl^{-}}\) will form a compound in a ratio of 1:1 to make a neutral compound. However, in order to balance the entire chemical equation, a 2 must be placed in front of NaCl.
It is possible to combine two solutions of ions and observe no sign of a reaction. In these cases, the reactant ions simply exist in solution together, forming a sort of “ion soup” with no formation of a precipitate, water molecules, or a gas. An example of this seen in Reaction IS.6, where no reaction has occurred because the products are all water soluble.
\begin {equation} \ce {NaCl(aq) + KNO3(aq) -> KCl(aq) + NaNO3(aq)} \end {equation}
How can you predict whether a combination of ions will form an insoluble compound? The solubility rules in Table IS.1 describe some general patterns of solubility of ionic compound as observed by chemists. These rules or guidelines can be used to predict the solubility of an ionic compound in water. Terms such as ”nitrates” and ”sulfates” refer to all ionic compounds that contain nitrate ions ( \( \ce {NO3^-} \) ) and sulfate ions ( \( \ce {SO4^2-} \) ), respectively as their anions.
| Rule | Applies to | Statement | Exceptions |
|---|---|---|---|
| 1 | Group 1 and ammonium (\(\ce{NH4^{+}}\)) ions | All compounds are soluble. | -- |
| 2 | Acetates (\(\ce{CH3COO^{-}}\), nitrates (\(\ce{NO3^{-}}\)) | All compounds are soluble. | -- |
| 3 | Halides (\(\ce{Cl^-, Br^-, I-}\)) | Most halides are soluble. | Silver (\(\ce{Ag^{+}}\)), mercury (\(\ce{Hg2^{2+}}\)), and lead (\(\ce{Pb^{2+}}\)) halides |
| 4 | Sulfates (\(\ce{SO4^{2-}}\)) | Most sulfates are soluble. | Calcium (\(\ce{Ca^{2+}}\)), strontium (\(\ce{Sr^{2+}}\)), barium (\(\ce{Ba^{2+}}\)), silver (\(\ce{Ag^{+}}\)), mercury (\(\ce{Hg2^{2+}}\)), and lead (\(\ce{Pb^{2+}}\)) sulfates |
| 5 | Carbonates (\(\ce{CO3^{2-}}\)) | Most carbonates are insoluble. | Group IA and ammonium carbonates |
| 6 | Phosphates (\(\ce{PO4^{3-}}\)) | Most phosphates are insoluble. | Group IA and ammonium phosphates |
| 7 | Sulfides (\(\ce{S^{2-}}\)) | Most sulfides are insoluble. | Group IA and ammonium sulfides |
| 8 | Hydroxides (\(\ce{OH^{-}}\)) | Most hydroxides are insoluble. | Group IA and ammonium hydroxides |
Below is an example of a precipitation reaction. When aqueous solutions of potassium sulfate, \( \ce {K2SO4} \), and silver (I) nitrate, \( \ce {AgNO3} \) are mixed, the results will be silver (I) sulfate, \( \ce {Ag2SO4} \), and potassium nitrate, \( \ce {KNO3} \), after the reactant compounds exchange ions. See Reaction IS.8 below. Silver (I) sulfate will be the insoluble precipitate solid because according to the solubility rules most sulfates are soluble with the exception of silver. All nitrates are soluble with no exceptions, and therefore \( \ce {KNO3} \) will remain aqueous (soluble).
\begin {equation} \ce{ K2SO4(aq) + 2 AgNO3(aq) -> 2 KNO3 (aq) + Ag2SO4(s)} \end {equation}
Example IS.1
Use the solubility rules to determine if any of the following compounds is likely to be
soluble in water: sodium phosphate, calcium carbonate, and lead(II) chloride.
Start by writing the chemical formulas for each compound and applying the solubility
rules.
All the solutions used in this experiment are the same concentrations. Each will be in a dropper bottle. The chemicals include:
1.
Obtain a spot plate and make sure it is clean. Contaminants left from previous
use or ions found in tap water can cause false reactions in this procedure. If
necessary, use soap and water to wash it first, rinsing three times with DI water.
It is not essential to have it completely dry.
2.
Divide the metal ions into two groups of four as you set up your mixtures. A
spot plate has six rows and four columns. Thus, the six anions will each be
added to one row on both spot plates while the metal ions will each be added to
one column.
3.
Your first spot plate will contain all six anions:
\( \ce {CO3^2-, Cl^-, OH^-, NO3^-, PO4^3-, and SO4^2-} \)
and these four cations:
\( \ce {K+, Ca^2+, Ba^2+, Al^3+} \)
4.
Draw your spot plate in your notebook and label the paper for each row and
column similarly to the way your report sheet is set up.
5.
In each well (or spot) of the spot plate you will place one pipet full of each of the
two specified ions. Observe and note any signs of a reaction.
6.
Record the results in your notebook. Use ppt to indicate the formation of a solid
precipitate. Describe the color and appearance of the precipitate, such as
“yellow, chunky ppt” or “white, fluffy ppt.” If no precipitate forms, enter NR for
“no reaction.”
7.
Once you have completed all 24 combinations and recorded your observations,
clean the spot plate by pouring the contents carefully into a large waste beaker.
Rinse the plate twice with DI water and pour the rinses into the waste
beaker. Wash the plate with soap and water and rinse thoroughly with DI
water.
8.
Draw a second spot plate, which contains all six anions:
\( \ce {CO3^2-, Cl^-, OH^-, NO3^-, PO4^3-, and SO4^2-} \)
and these four cations:
\( \ce {Ag+, Cu^2+, Fe^2+ and Pb^2+} \)
9.
Complete another 24 combinations of two chemical solutions and record your
observations for each, as you did for the first spot plate.
10.
Clean the spot plate by pouring the contents carefully into a large waste beaker.
Rinse the plate twice with DI water and pour the rinses into the waste
beaker. Wash the plate with soap and water and rinse thoroughly with DI
water.
11.
Pour the contents of the waste beaker into the metal ion waste container in the
hood.
Write the correct ionic compound formulas and names for all the precipitates that you observed in the first spot plate. For example, if you observed a solid and wrote “ppt” for the first well that shows a combination of \( \ce {CO3^{2-}} \) and \( \ce {K^{+}} \), write the formula and name for this compound. If you wrote “NR” for no reaction in this well, then no formula and name will be written for this ion combination.
Write the correct ionic compound formulas and names for all the precipitates that you observed in the second spot plate. For example, if you observed a solid and wrote “ppt” for the first well that shows a combination of \( \ce {CO3^{2-}} \) and \( \ce {Ag^{+}} \), write the formula and name for this compound. If you wrote “NR” for no reaction in this well, then no formula and name will be written for this ion combination.
Post-Lab Questions: Be sure to answer the questions that are listed at the end of the online lab submission.