4.4: Solubility and Ionic Equations

Transcription

1 4.4: Solubility and Ionic Equations Solubility Curves Graphs of solubility (maximum concentration) against temperature allow quick and easy reference, and are very useful for a wide variety of questions involving solution concentrations. immiscible: miscible:

2 Saturation Recall that solubility is the quantity of solute that can dissolve in a given quantity of solvent. Quantitatively, solubility is often expressed as the maximum mass of solute that can be dissolved per 100 g of water at a given temperature. Table I gives the solubilities of some common solutes. According to these data, g of sucrose (table sugar) dissolves in 100 ml of water at20 "C. This is the maximum quantity of sucrose that will dissolve at this temperature. saturated solution unsaturated solution supersaturated solution Using Solubility Rules to Predict Precipitation Formation Solubility is a measure of how a substance dissolves in water at a given temperature and pressure. A substance that does not dissolve well in water is called insoluble. Chalk is a substance that has a low solubility. Substances like sodium chloride, NaCl, that dissolve well in water are called soluble. Predicting The Formation of A Precipitate Use the following table to determine if a precipitate would be form in the following examples.

3 Cations SCH3U: Solutions and Solubility Soluble (aq) Insoluble (s) Solubility of Ionic Compounds Anions Cl, Br, I S 2 OH 2 CO SO 3, 4 3 PO 4 C 2 H 3 O 2 Li +, Na +, most K +, Mg 2+ Li +, Na +, Li +, Na +,, Ca 2+ K +, Sr 2+ most K +, most Sr 2+ Ag +, Pb 2+ Cu + most most * Note that all acids and all nitrates are highly soluble in water. Ag +, Pb 2+, Ca 2+, most Ag Sr 2+ + Example 1: Determine the products (if any) when a solution of sodium sulfate is mixed with a solution of lead(ii) nitrate. If a reaction occurs, summarize the reaction as a balanced chemical equation.

4 Example 2: Determine the products (if any) when a solution of sodium chloride is mixed with a solution of silver nitrate. If a reaction occurs, summarize the reaction as a balanced chemical equation. Ionic and Net Ionic Equations ionic equations or total ionic equations net ionic equations (the spectators ions) How to Create Net ionic Equations Following these steps the same why each time will create a net ionic equation Step 1: Identify the type of reaction and decide on the possible products. Step 2: Write out the word equation for the reaction Step 3: Look up the solubility of the products using the solubility table. Step 4: Indicate the states of the reactants and products, either (aq) for aqueous or (s) for precipitate. Step 5: Write the correct formulas for the reaction. Step 6: Balance the equation Step 7: Write out the total ionic equation leaving any precipitates in compound or undissolved form Step 8: Cancel out any ions that are the same on both sides (the spectator ions) Step 9: Write out the equation showing only what is left. This is your net ionic equation.

5 Example 1 - Write the total ionic equation and net ionic equation between barium sulphide and sodium sulfate. Example 2 - Write the total ionic equation and net ionic equation between sodium chloride and lead(ii) nitrate. Qualitative Chemical Analysis Chemists use the information in the solubility table to create a diagnostic tool called the. Chemists can use the solubility rules to determine if certain ions are present in a solution by conducting double displacement reactions. They use solution of ions that will cause the ions they are looking for to precipitate out. If you suspect that a solution may have acetate ions in it, then you'd add silver ions in the form of silver nitrate to it. If there are acetate ions in the solution they will precipitate out because silver acetate is insoluble. If you add the silver nitrate and nothing happens then you can infer that the acetate ions are not present. precipitate - supernate -

6 Worksheet 4.4: Solubility and Ionic Equations 1. Refer to the Solubility Curves in notes. (a) What mass of K 2 SO 4 can be dissolved in 100 ml of water at 70 C? (b) At what temperature are the solubilities of KNO 3 and KCl approximately equal? (c) Why is it not possible to determine the molar concentration of a saturated solution of KClO 3 at 25 C from its solubility curve? What additional information would be required? (d) Which substance forms a saturated solution in which the mass ratio of solute to solvent is 1:1, and at what temperature? 2. A solution containing equal masses of NaCl and KCl is slowly cooled. (a) At what temperatures might sodium chloride precipitate first, without precipitating potassium chloride? (b) At what temperatures might potassium chloride precipitate first, without precipitating sodium chloride? 3. For any solute, what important condition must be stated in order to report the solubility? 4. A common diagnostic test for CO 2(g) involves the use of a saturated solution of calcium hydroxide, Ca(OH) 2(aq), historically called limewater. Calcium hydroxide has low solubility but is actually soluble enough to have a noticeable effect. The solution is easily prepared by placing a spoonful of solid Ca(OH) 2(s) in pure water, and stirring or shaking for a while. A small amount of the solid dissolves. The excess solute settles to the bottom, leaving a clear, saturated solution. Bubbling carbon dioxide gas through this saturated solution causes reactions that finally precipitate calcium carbonate, which is roughly 10 times less soluble than calcium hydroxide. (a) The first reaction is the apparent dissolving of a tiny amount of carbon dioxide in water. What we believe really occurs is that some carbon dioxide gas reacts with water to form carbonic acid in solution. Write a balanced chemical equation to represent this reaction. (b) The acid reacts with dissolved calcium hydroxide to form calcium carbonate precipitate. Write a balanced chemical equation to represent this reaction. (c) Use the reaction equations to explain why this is a test for the presence of CO 2(g). (d) If the calcium hydroxide solution were not saturated, or nearly so, would this diagnostic test still work? Explain the reasoning behind your answer. (e) Assuming you wished to prepare 1.0 L of saturated limewater solution to be stored at a normal room temperature of 22 C, what minimum mass of solid calcium hydroxide would you require? Why should you actually use considerably more than the minimum required? 5. Strontium compounds are often used in flares because their flame colour is bright red. One industrial process to produce low-solubility strontium compounds (that are less affected by getting wet) involves the reaction of aqueous solutions of strontium nitrate and sodium carbonate. Write the balanced chemical equation, total ionic equation, and net ionic equation for this reaction. 6. Placing aluminum foil in any solution containing aqueous copper(ii) ions will result in a reaction. The reaction is slow to begin with, but then proceeds rapidly. (a) Referring to the solubility table (Appendix C), name at least four ionic compounds that could be dissolved in water to make a solution containing aqueous copper(ii) ions. (b) Write a balanced chemical equation for the reaction of aluminum with one of the compounds you suggested in (a). (c) Write the total ionic equation for the reaction. (d) Write the net ionic equation for the reaction.