CHAPTER 14

Ions in Aqueous Solutions andColligative Properties

• When an ionic compound dissolves in water, the ions separate from one another in a process called dissociation.

• For example, when sodium chloride reacts with water, a sodium cation and chloride anions are released in solution.

• NaCl (s) Na+ (aq) + Cl- (aq)

• Another example of dissociation is seen when calcium chloride reacts with water to form a calcium cation and two chloride anions.

• CaCl2 (s) Ca2+ (aq) + 2Cl- (aq)

• For ionic compounds, formula units (think of these as “molecules” of ionic compounds) tell how many ions should be formed during dissociation.

Dissociation

• So, if 1 mole of NaCl completely dissolves, or undergoes 100% dissociation, 1 mol of Na+ ions and 1 mol of Cl- ions will result.

• Similarly, if 1 mol of CaCl2 dissociates completely, then 1 mol of Ca2+ ions and 2 mol of Cl- ions will result.

• All soluble ionic compounds will be assumed to undergo 100% dissociation, unless told otherwise.

Dissociation (continued)

• Write the equation for the dissolution of aluminum sulfate, Al2(SO4)3 in water. How many moles of aluminum ions and sulfate ions are produced by dissolving 1 mol of aluminum sulfate? What is the total number of moles of ions produced by dissolving 1 mol of aluminum sulfate?

Dissociation Example 1

• Write the equation for the dissolution of each of the following in water, and then determine the number of moles of each ion produced as well as the total number of moles of ions produced.

• A. 1 mol of ammonium chloride, NH4Cl

• B. 1 mol of sodium sulfide, Na2S

• C. 0.5 mol of barium nitrate, Ba(NO3)2

Dissociation Example 2

• Write the equation for the dissolution of magnesium chlorate in water. How many moles of ions are produced for every 1 mol of magnesium chlorate dissolved?

Dissociation Example 3

• Write the equation for the dissolution of NH4NO3 in water. If 3.5 mol of NH4NO3 are dissolved, how many moles of each type of ion are produced? How many total moles of ions are produced?

Dissociation Example 4

• Sometimes ionic compounds are combined together and solids form as a result. These solids are called precipitates.

• Prior to combining these ionic compounds, it may be necessary to predict whether a precipitate will form or not.

• Solubility rules will be useful in the prediction of precipitates. It is difficult to write solubility rules that cover all possible conditions.

• YOUR SOLUBILITY RULES CHART WILL BE YOUR BEST FRIEND!

Precipitation Reactions

• Using your solubility rules, determine whether or not each chemical compound will be a precipitate. If not, write an equation that shows its complete dissociation.

• A. Ca3(PO4)2

• B. (NH4)2S

• C. Cd(NO3)2

• D. PbSO4

• E. Al2(CO3)3

Precipitation Reactions (continued)

• Sometimes, we use the solubility rules to confirm the prediction of precipitates in a double-replacement reaction.

• If we were to react ammonium sulfide and cadmium nitrate, the two products that form are ammonium nitrate and cadmium sulfide. But the question must be asked, are there any precipitates formed?

• (NH4)2S (aq) + Cd(NO3)2 (aq) 2NH4NO3 (?) + CdS (?)

Precipitation Reactions (continued)

• Identify the precipitate that forms when aqueous solutions of zinc nitrate and ammonium sulfide are combined. Write the equation for the possible double-replacement reaction including all physical states of products.

Precipitation Reactions Example 1

• Will a precipitate form if solutions of potassium sulfate and barium nitrate are combined? If so, write the equation of the possible double-replacement reaction indicating the physical states of all products.

Precipitation Reactions Example 2

• Will a precipitate form if solutions of potassium nitrate and magnesium sulfate are combined? Write the equation of the double-replacement reaction indicating the physical states of all products formed.

Precipitation Reactions Example 3

• Reactions of ions in aqueous solution are typically shown using net ionic equations rather than traditional formula equations.

• In net ionic equations, we only use those compounds and ions that undergo a chemical change in a reaction in an aqueous solution.

• There is a systematic approach to arriving at a correct net ionic equation.

Net Ionic Equations

• 1. Write out the formula equation with all physical states of reactants and products.

• 2. Write out the overall ionic equation by breaking down all reactants and products that are not insoluble into their respective ions. All precipitates are shown using (s).

• 3. Remove all spectator ions. These are ions that appear on both sides of the overall ionic equation. These do not contribute to the formation of the precipitate.

• 4. The leftovers = net ionic equation!

How to Write a Net Ionic Equation

• In a reaction between strontium chloride, SrCl2, and sodium sulfate, Na2SO4, a white precipitate is observed. Write the net ionic equation for the precipitate observed in the reaction above.

Net Ionic Equations Example 1

• For the following pairs, identify the precipitate formed, if any, and write the net ionic equation for the reaction.

• A. KCl and AgNO3

• B. Na2CO3 and CaCl2

• C. Na2S and Fe(NO3)2

• D. K2SO4 and Ba(NO3)2

• E. Cu(CH3COO)2 and K2CO3

Net Ionic Equations Example 2

• Some molecular compounds can also form ions in solution. These compounds are considered to be polar.

• Ions are formed from solute molecules by the action of the solvent via ionization.

• Ionization and dissociation ARE NOT THE SAME!

• Dissociation of ionic compounds result in the release of ions that were already present in the compound.

• Ionization results in the formation of ions that were not already present.

Ionization

• Many molecular compounds contain a hydrogen atom bonded by a polar covalent bond.

• When these compounds undergo ionization in an aqueous solution, H+ ions are released.

• These H+ ions are so strongly attractive that they do not exist alone, but form the hydronium ion, H3O+.

• Consider the ionization of hydrogen chloride in water.

• H2O (l) + HCl (g) H3O+ (aq) + Cl- (aq)

The Hydronium Ion

• Electrolytes are substances that yield ions and conduct an electric current in solution.

• Nonelectrolytes are substances that do not yield ions and do not conduct electric current in solution.

• We can classify electrolytes as strong (where there is 100% dissociation) or weak (where there is partial dissociation).

• Strong = HCl, HBr, HI, along with all soluble ionic compounds

• Weak = HF, CH3COOH, molecular compounds

Strong and Weak Electrolytes

• The presence of solutes affect the properties of the solutions.

• Some of the properties are not dependent on the nature of the dissolved substance but only on how many dissolved particles are present.

• These properties are called colligative properties.

• For calculations that involve colligative properties, concentration is given in molality.

Colligative Properties of Solutions

• The boiling point and freezing point of a solution differ from those of the pure solvent.

• The freezing point of a 1 m solution of any nonelectrolyte solute in water is found by experiment to be 1.86C lower than the freezing point of water.

• In other words, when 1 mol of a nonelectrolyte is dissolved in 1 kg of water, the freezing point of the solution is -1.86C instead of 0.00C.

• Molal freezing-point constant (Kf): the freezing point depression of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute

Freezing Point Depression

• Each solvent has its own characteristic molal freezing-point constant.

• The freezing-point depression, Tf is the difference between the freezing points of the pure solvent and a solution of a nonelectrolyte in that solvent, and it is directly proportional to the molal concentration of the solution.*

• *for electrolytes, this is not necessarily true.

• Formula: Tf = i Kf m Kf = -1.86C/m

• i = 1(nonelectrolytes); 2 (cmpds that yield 2 ions); 3 (cmpds that yield 3 ions)

Freezing-Point Depression (continued)

• What is the freezing-point depression of water in a solution of 17.1 g of C12H22O11 in 200.0 g of water? What is the actual freezing point of the solution?

Freezing-Point Depression Example 1

• A water solution containing an unknown quantity of a nonelectrolyte solute is found to have a freezing point of -0.23C. What is the molal concentration of the solution?

Freezing-Point Depression Example 2

• A solution consists of 10.3 g of the nonelectrolyte glucose, C6H12O6, dissolved in 250.0 g of water. What is the freezing-point depression of the solution?

Freezing-Point Depression Example 3

• In a laboratory experiment, the freezing point of an aqueous solution of glucose is found to be -0.325C. What is the molal concentration of this solution?

Freezing-Point Depression Example 4

• The freezing point of an aqueous solution that contains a nonelectrolyte is -9.0C. A. What is the freezing-point depression of the solution? B. What is the molal concentration of the solution?

Freezing-Point Depression Example 5

• The freezing point of an aqueous NaCl solution is -0.20C. What is the molality of the solution?

Freezing-Point Depression Example 6

• What is the expected change in the freezing point of water in a solution of 62.5 g of barium nitrate, Ba(NO3)2, in 1.00 kg of water?

Freezing-Point Depression Example 7

• What is the expected freezing-point depression for a solution that contains 2.0 mol of magnesium sulfate dissolved in 1.0 kg of water?

Freezing-Point Depression Example 8

• The boiling point of a solution is higher than the boiling point of the pure solvent.

• Molal boiling-point constant (Kb): the boiling-point elevation of the solvent in a 1-molal solution of a nonvolatile, nonelectrolyte solute.

• The boiling-point elevation of a 1-molal solution of any nonelectrolyte solute in water has been found to be 0.51C.

• Kb = 0.51C/m

Boiling-Point Elevation

• The boiling-point elevation, Tb is the difference between the boiling points of the pure solvent and a nonelectrolyte solution of that solvent, and is directly proportional to the molal concentration of the solution.*

• *For electrolytes, this is not necessarily so.

• Formula: Tb = i Kb m Kb = 0.52C/m

• i = 1 (nonelectrolytes); 2 (if 2 ions produced); 3 (if 3 ions produced)

Boiling-Point Elevation (continued)

• If 90.0 g of nonionizing C6H12O6 are dissolved in 255 g of water, what will be the boiling point of the resulting solution?

Boiling-Point Elevation Example 1

• A solution contains 50.0 g of sucrose, C12H22O11, dissolved in 500 g of water. What is the boiling-point elevation?

Boiling-Point Elevation Example 2

• If the boiling-point elevation of an aqueous solution containing a nonvolatile electrolyte is 1.02C, what is the molality of the solution?

Boiling-Point Elevation Example 3

• The boiling point of an aqueous solution containing a nonvolatile nonelectrolyte is 100.75C. What is the boiling-point elevation? What is the molality of the solution?

Boiling-Point Elevation Example 4

• What is the expected boiling-point elevation of water for a solution that contains 150 g NaCl dissolved in 1.0 kg of water?

Boiling-Point Elevation Example 5

• Calculate the boiling point of an ionic solution containing 29.7 g Na2SO4 and 84.4 g of water.

Boiling-Point Elevation Example 6

• Salt is often added to water in order to raise the temperature of the boiling point and to heat food more quickly. If you add 120.0 g of salt to 2.75 kg of water, what will be the change in the boiling point of the water solution?

Boiling-Point Elevation Example 7