Chapter 8A Solutions 1 CHAPTER OUTLINE Type of Solutions Electrolytes & Non-electrolytes Equivalents of Electrolytes Solubility & Saturation Soluble & Insoluble Salts Formation of a Solid Precipitation Reactions 2 TYPE OF SOLUTIONS A solution is a homogeneous mixture of two substances: Solute: substance being dissolved present in smaller amount Solvent: substance doing the dissolving present in larger amount Solutes and solvents may be of any form of matter: solid, liquid or gas. 3 TYPE OF SOLUTIONS Type Example Solute Solvent Gas in gas Air Oxygen Nitrogen CO2 Water Gas in liq. Soda water Liq. in liq. Vinegar Acetic acid Water Solid in liq. Seawater Salt Water 4 TYPE OF SOLUTIONS Type Example Solute Solvent Liquid in solid Dental amalgam Mercury Silver Solid in solid Brass Zinc Copper 5 SOLUBILITY Solutions form between solute and solvent molecules because of similarities between them. Like dissolves Like Ionic solids dissolve in water because the charged ions (polar) are attracted to the polar water molecules. Non-polar molecules such as oil and grease dissolve in non-polar solvents such as kerosene. 6 SOLUBILITY Water (polar) CH2Cl2 (non-polar) Ni(NO3 )2 (polar) I2 (non-polar) 7 ELECTROLYTES & NON-ELECTROLYTES Solutions can be characterized by their ability to conduct an electric current. Solutions containing ions are conductors of electricity and those that contain molecules are non-conductors. Substances that dissolve in water to form ions are called electrolytes. The ions formed from these substances conduct electric current in solution, and can be tested using a conductivity apparatus. 8 STRONG ELECTROLYTES Electrolytes are further classified as strong electrolytes and weak electrolytes. In water, a strong electrolyte exists only as ions. Only ions Strong electrolytes make the light present after bulb on the conductivity solution apparatus glow brightly. Ionic substances such as NaCl are strong electrolytes. NaCl (s) Na+ (aq) + Cl (aq) 9 WEAK ELECTROLYTES Solutions containing weak electrolytes contain only a few ions. These solutions make the light bulb on the conductivity Few ions apparatus glow dimly. present after Weak acids and basessolution that dissolve in water and produce few ions are weak electrolytes. HF (aq) H+ (aq) + F (aq) 10 NONELECTROLYTES Substances that do not form any ions in solution are called non-electrolytes. No ions present With these solutions the bulb on theafter conductivity solution apparatus does not glow. Covalent molecules that dissolve in water but do not form ions, such as sugar, are non-electrolytes. C12H22O11 (s) C12H22O11 (aq) 11 ELECTROLYTES & NON-ELECTROLYTES 12 Example 1: Identify the predominant particles in each of the following solutions and write the equation for the formation of the solution: NH4Br NH4Br (s) Strong electrolyte (only ions) NH4+ (aq) + Br (aq) 13 Example 1: Identify the predominant particles in each of the following solutions and write the equation for the formation of the solution: CH4 N2O Non-electrolyte CH4N2O (s) (only molecules) CH4N2O (aq) 14 Example 1: Identify the predominant particles in each of the following solutions and write the equation for the formation of the solution: HClO Weak electrolyte HClO (aq) (few ions) H+ (aq) + ClO (aq) 15 EQUIVALENTS OF ELECTROLYTES Body fluids typically contain a mixture of several electrolytes, such as Na+, Cl–, K+ and Ca2+. Each individual ion is measured in terms of an equivalent (Eq), which is the amount of that ion equal to 1 mole of positive or negative electrical charge. For example, 1 mole of Na+ ions and 1 mole of Cl– ions are each 1 equivalent (or 1000 mEq) because they each contain 1 mole of charge. 16 EQUIVALENTS OF ELECTROLYTES An Some ionexamples with a charge of ionsofand 2+ or their 2– contains equivalents 2 are equivalents shown below: per mole. Ion Electrical Charge No. of Equivalents in 1 Mole Na+ 1+ 1 Eq Ca2+ 2+ 2 Eq Fe3+ 3+ 3 Eq Cl– 1– 1 Eq SO42– 2– 2 Eq 17 EQUIVALENTS OF ELECTROLYTES In body, the charge of the positive ion is always balanced by the charge of the negative ion. For example, a solution containing 25 mEq/L of Na+ and 4 mEq/L of K+ must have 29 mEq/L of Cl– to balance. Shown next are examples of some common intravenous solutions and their ion concentrations. 18 EQUIVALENTS OF ELECTROLYTES 19 Example 1: Indicate the number of equivalents in each of the following: 2 mol K+ 1 Eq 2 mol x 1 mol = 2 Eq Mg2+ 2 Eq 0.5 mol x 1 mol = 1 Eq 3 mol CO32 2 Eq 3 mol x 1 mol = 6 Eq 0.5 mol 20 Example 2: A typical concentration for Ca2+ in blood is 8.8 mEq/L. How many moles of Ca2+ are present in 0.50 L of blood? Liter blood 8.8 mEq 0.50 L x 1L moles Ca2+ mEq x 1 Eq 103 mEq x 1 mol 2 Eq = 0.0022 mol 21 Example 3: An IV solution contains 155 mEq/L of Cl–. If a patient received 1250 mL of the IV solution, how many moles of chloride were given to him? 1L 1250 mL x 103 mL x 155 mEq 1L x 1 Eq 103 mEq x 1 mol 1 Eq = 0.194 mol 22 Example 4: A sample of Ringer’s solution contains the following concentrations (mEq/L) of cations: Na+ 147, K+ 4, and Ca2+ 4. If Cl– is the only anion in the solution, what is the concentration of Cl– in mEq/L? Total cation mEq/L = 147 + 4 + 4 = 155 mEq/L Total cation mEq/L = Total anion mEq/L Total Cl mEq/L = 155 mEq/L 23 SOLUBILITY Solubility refers to the maximum amount of solute that can be dissolved in a given amount of solvent. Many factors affect the solubility of a solute in a solution. Type of solute Type of solvent Temperature Solubility is measured in grams of solute per 100 grams of solvent at a given temperature. 24 SATURATION A solution that does not contain the maximum amount of solute in it, at a given temperature, is called an unsaturated solution. A solution that contains the maximum amount of solute in it, at a given temperature, is called a saturated solution. Un-dissolved solid in solution 25 SOLUBILITY Solubility of most solids in water increases as temperature increases. Using a solubility chart, the solubility of a solute at a given temperature can be determined. For example, KNO3 has a solubility of 80 g/100 g H2O (80%) at 40 C. 26 SOLUBILITY OF GASES Solubility of gases in water decreases as temperature increases. At higher temperatures more gas molecules have the energy to escape from solution. Henry’s law states that the solubility of a gas is directly proportional to the pressure above the liquid. For example, a can of soda is carbonated at high pressures in order to increase the solubility of CO2. Once the can is opened, the pressure is reduced and the excess gas escapes from the solution. 27 SOLUBLE & INSOLUBLE SALTS Many ionic solids dissolve in water and are called soluble. However, some ionic salts do not dissolve in water and do not form ions in solution. These salts are called insoluble salts and remain solid in solution. Chemists use a set of solubility rules to predict whether a salt is soluble or insoluble. These rules are summarized next. 28 SOLUBILITY RULES S O L U B L E NO3 No exceptions Na+, K+ NH4+ No exceptions Cl, Br, I Except those containing Ag+ , Pb2+ SO4 2 Except those containing Ba2+ , Pb2+, Ca2+ 29 SOLUBILITY RULES I N S O L U B L E S2, CO32 PO43 Except those containing Na+ , K+, NH4+ OH Except those containing Na+ , K+, Ca2+, NH4+ 30 Example 1: Use the solubility rules to determine if each of the following salts are soluble or insoluble: K3PO4 soluble CaCO3 insoluble (all salts of K are soluble) (most carbonates are insoluble) 31 Example 2: Using the solubility chart, determine if each of the following solutions is saturated or unsaturated at 20C: 25 g NaCl in 100 g water Solubility of NaCl at 20 C is 40% Solution is unsaturated 32 Example 2: Using the solubility chart, determine if each of the following solutions is saturated or unsaturated at 20C: 11 g NaNO3 in 25 g water Solubility of NaNO3 at 20 C is 85% Solution is unsaturated 33 Example 2: Using the solubility chart, determine if each of the following solutions is saturated or unsaturated at 20C: 400. g of glucose in 125 g water Solubility of glucose at 20 C is 80% Solution is saturated 34 FORMATION OF A SOLID Solubility rules can be used to predict whether a solid, called a precipitate, can be formed when two solutions of ionic compounds are mixed. A solid is formed when two ions of an insoluble salt come in contact with one another. For example, when a solution of K2CrO4 is mixed with a solution of Ba(NO3)2 a yellow insoluble salt BaCrO4 is produced. 35 AQUEOUS REACTIONS K2CrO4 (aq) + Ba(NO3)2 (aq) BaCrO4 (s) + 2 KNO3 (aq) precipitate Solid BaCrO4 forms 36 PRECIPITATION REACTIONS Double replacement reactions in which a precipitate is formed are called precipitation reactions. To predict a precipitate, follow the steps outlined next. 37 PRECIPITATION REACTIONS 1. Write the reactant ions that form after dissolution. 2. Write the product combinations possible when reactant ions combine. 3. Use solubility rules to determine if any of the products are insoluble. 4. If a precipitate forms, write the formula for the solid. Write other ions that form soluble salts as ions. If no precipitate forms, write “NO REACTION” after the arrow. 5. Cancel ions that appear the same on both sides of the equation (spectator ions), to form Net Ionic Equation. 38 PRECIPITATION REACTIONS The reaction of K2CrO4 and Ba(NO3)2 can be predicted as shown below ??? Step 1: 2 K+ +CrO42 + Ba2+ + 2 NO3 Step 2: 2 K+ +CrO42 +Ba2+ + 2 NO3 BaCrO4(?) + 2K+NO3 (?) Net Ionic Step 3: 2 K+ +CrO42 +Ba2+ + 2 NO Equation BaCrO4(s) + 2K+NO3 (aq) 3 Step 4: 2 K+ +CrO42 +Ba2+ + 2 NO3 BaCrO4(s) + 2K+ +2NO3 Step 5: Ba2+ + CrO42 BaCrO4 (s) spectator ions precipitate 39 Example 1: Predict the products, if any, for the reaction of AgNO3 and NaCl and write the net ionic equation. Step 1: Ag+ +NO3 + Na+ + Cl Step 2: Ag+ +NO3 +Na+ +ClNetIonic Ag+Cl (?) + Na+ NO3 (?) Step 3: Ag+ +NO3 +Na+ +ClEquation Ag+Cl (s) + Na+ NO3 (aq) Step 4: Ag+ +NO3 +Na+ +Cl AgCl (s) + Na+ + NO3 (aq) Step 5: ??? + +Cl Agspectator ions AgCl (s) spectator ions precipitate 40 Example 2: Predict the products, if any, for the reaction of Na2SO4 and Pb(NO3)2 and write the net ionic equation. Step 1: 2 Na+ +SO42 + Pb2+ + 2 NO3 Step 2: 2Na+ +SO42 +Pb2+ +2NO3 2Na+NO3(?) + Pb2+SO42(?) Step 3: 2Na+ +SO42 +Pb2+ +2NO3 2Na+NO3(aq) + Pb2+SO42(s) Step 4: 2Na+ +SO42 +Pb2+ +2NO3 2Na+ + 2NO3(aq) + PbSO4 (s) Step 5: Pb2+ +SO42 ??? PbSO4 (s) (NIE) precipitate 41 Example 3: Predict the products, if any, for the reaction of PbCl2 and KI and write the net ionic equation. Step 1: Pb2+ + 2 Cl Step 2: Pb2+ + 2 Cl + K+ + I Pb2+I2(?) + K+Cl(?) Step 3: Pb2+ + 2 Cl + 2 K+ +2 I Pb2+I2(s) + 2 K+Cl(aq) Step 4: Pb2+ + 2 Cl + 2 K+ +2 I PbI2 (s) + 2 K+ + 2 Cl Step 5: + K+ + I ??? Pb2+ + 2 I PbI2 (s) (NIE) precipitate 42 Example 4: Predict the products, if any, for the reaction of NH4Cl and KNO3 and write the net ionic equation. Step 1: NH4+ + Cl + K+ + NO3 ??? No Reaction Step 2: NH4+ + Cl + K+ + NO3 NH4+NO3(?) + K+Cl(?) Step 3: NH4+ + Cl + K+ + NO3 NH4+NO3(aq) + K+Cl(aq) Step 4: NH4+ + Cl + K+ + NO3 NH4+ + NO3 + K+ + Cl 43 THE END 44