Physical Properties of Physical Properties of Solutions

 Types of Solutions (13.1)  A Molecular View of the Process (13.2)  Concentration Units (13.3)  Effect of Temperature on Solubility (13.4)  Effect of on the Solubility of Gases (13.5)  (13.6) General Chemistry I – Concepts

 Mixtures and representations of matter (1.3 and 1.4)  Chemical formulas and nomenclature (2.6 and 2.7)  Formula calculations particularly mole calculations and percent composition (3.3 and 3.5)  Properties of solutions, introduction to hydration and concentrations of solutions (4.1 and 4.5)  Enthalpy and heats of reaction (6.6)  Intermolecular forces and vapor pressure (12.2 and 12.6) 13.1 Types of Solutions

 (Review 1.3 and 4.1) ◦ What is the definition of a mixture and how is this different from a pure substance? ◦ What are the definitions of solute, and solution? ◦ How can you differentiate between a heterogeneous and homogeneous solution? 13.1 Types of Solutions  Key Definitions:  Unsaturated solutions: ◦ A solution that contains less solute than it has the capacity to dissolve.  Saturated solutions: ◦ A solution that contains the maximum amount of a solute in a given solvent, at a specific temperature.  Using sodium chloride in water describe how these are different.  How would this be different for a solution of silver chloride? 13.1 Types of Solutions

 Key Definition:  Supersaturated solutions: ◦ A solution that contains more solute than is present in a saturated solution.

 What is happening in the supersaturated solution shown in Figure 13.1?

Figure 13.1 p. 440 13.2 A Molecular View of the Solution Process

 Consider the intermolecular forces between: ◦ solute/solute ◦ solvent/solvent ◦ solvent/solute Discussed in Chapter 12 – intermolecular forces (12.2) Review: 12.2 Intermolecular Forces  Interactions between: Dispersion ◦ solute/solute forces ◦ solvent/solvent Dipole forces

Hydrogen bonding IMF – Substances Pure Review:12.2 Intermolecular Forces

Interactions between solute/solute and solvent/solvent: 1. Determine the intermolecular forces of the substances (usually considering the solubility of two pure substances) Approximating solubility 2. Determine the relative size of the two molecules (same or different)

IMF – Mixtures If both (IMF and size) are the same, then we would approximate that the substances are soluble If both (IMF and size) are different, then we would approximate that the substances are insoluble If one (IMF or size) are the same, then we would approximate that the substances are partially soluble Review: 12.2 Intermolecular Forces

Interaction between solvent/solute: ◦ Dispersion forces (or induced dipole-induced dipole) between all substances in a solution. ◦ A nonpolar and a polar substance would have induced dipole-dipole forces ◦ A polar and a polar substance would have dipole-dipole forces

IMF – Mixtures ◦ A ionic and a nonpolar substance would have ion-induced dipole forces ◦ A ionic and a polar substance would have ion-dipole forces 13.2 A Molecular View of the Solution Process

Incorporating energy into the discussion…

Figure 13.2 p. 441 13.3 Concentration Units

 (Review 3.3, 3.5 and 4.5) What is molarity?

 Other types of concentration units: ◦ percent by mass ◦ ◦ ◦ ppm 13.3 Concentration Units Concentrated sulfuric acid used in the lab is

96% H2SO4 by mass. What is the molarity, molality and mole fraction of the acid solution? The density of the solution is 1.83 g·mL–1. 13.4 Effect of Temperature on Solubility

 Consider two types of solutes: ◦ Solids (ionic) ◦ Gases

 (Review, 6.6) What is enthalpy of solution? 13.4 Effect of Temperature on Solubility – ionic solids

ΔHsoln Solute kJ∙mol–1

KNO3 36.5 NaNO3 21.56 NaBr –0.61 KBr 19.87 KCl 17.51 NaCl 3.87

Na2SO4 –5.02

Figure 13.3 p. 444 13.4 Effect of Temperature on Solubility – ionic solids

What does 100 g of potassium nitrate in 100 g water look like at 40oC?

What does this same mixture look like at 70oC?

Figure 13.3 p. 444 13.4 Effect of Temperature on Solubility – gases

Figure 13.4 p. 444 13.5 Effect of Pressure on the Solubility of Gases  Pressure and solubility of liquids and solids  Pressure and solubility of gases ◦ Using oxygen gas in solution (use water) describe what is the balanced equation? ◦ What would happen if we increase the pressure? ◦ What is the relationship between pressure and the solubility of gases? 13.5 Effect of Pressure on the Solubility of Gases

Figure 13.5 p. 445 13.5 Effect of Pressure on the Solubility of Gases  Pressure and solubility of gases ◦ What is the equation for Henry’s law?

 Practice: The solubility of carbon dioxide in water at 25oC and 1 atm is 0.034 mol·L–1. What is the mass of carbon dioxide found in a 355 mL can of soda at 25oC, assuming that the manufacturer used a pressure of 2.0 atm of carbon dioxide to carbonate the beverage? Does Henry’s law always hold? When does it not? 13.6 Colligative Properties

 What are colligative properties? ◦ Vapor pressure lowering (Raoult’s Law) ◦ Boiling-point elevation ◦ Freezing-point depression ◦ Osmotic pressure  Why could these be useful?

 Will first discuss nonelectrolyte solutions 13.6 Colligative Properties

 (Review, 12.6) What is vapor pressure?  (Review 13.3) What is mole fraction?  If something is nonvolatile, what does that mean?

 What is vapor pressure lowering? Vapor-pressure lowering (Raoult’s lowering Law) Vapor-pressure 13.6 Colligative Properties

 Practice (nonvolatile solutions)

The solubility of sugar in water at 25oC is 67.47 mass fraction of sugar. What is the vapor pressure of water in this solution? Vapor-pressure lowering (Raoult’s lowering Law) Vapor-pressure

Table 5.2 p. 156 13.6 Colligative Properties

 (Review, 12.6) What is vapor pressure?  (Review 13.3) What is mole fraction?  If something is volatile, what does that mean?

 How does combining two volatile substances to form a solution affect the vapor of each substance? Vapor-pressure lowering (Raoult’s lowering Law) Vapor-pressure 13.6 Colligative Properties Vapor-pressure lowering (Raoult’s lowering Law) Vapor-pressure Figure 13.6 p. 449 Figure 13.6 p. 449 13.6 Colligative Properties

 Practice (volatile solutions)

Using figure 13.6, what is the vapor pressure of benzene at 80oC? What is the vapor pressure of toluene at 80oC?

What is the total pressure when 50 g of toluene and 50 g of benzene are mixed? Vapor-pressure lowering (Raoult’s lowering Law) Vapor-pressure

Figure 13.6 p. 449 13.6 Colligative Properties

 (Review, 12.7) On a phase diagram where is the normal boiling point located?  What happens at this point? Boiling-Point Elevation Boiling-Point

Figure 12.32 p. 427 13.6 Colligative Properties  (Review, 12.7) On a phase diagram where is the normal boiling point located?  What happens at this point?  (Review 13.3) What is molality?  Why is molality better than molarity?

 What is boiling-point elevation? Boiling-Point Elevation Boiling-Point  How is boiling-point elevation quantified? 13.6 Colligative Properties Boiling-Point Elevation Boiling-Point

Figure 13.7 p. 449 13.6 Colligative Properties

 Practice

What is the boiling point of a saturated sugar solution (saturated at 25oC)?

At 25oC, 67.47 mass fraction of the

Boiling-Point Elevation Boiling-Point solution is sugar. 13.6 Colligative Properties Boiling-Point Elevation Boiling-Point

Table 13.2 p. 450 13.6 Colligative Properties

 (Review, 12.1-12.4) What is the definition of a liquid versus a solid?  Which has a greater disorder?  (Review 13.3) What is molality?  Why is molality better than molarity?

 What is freezing-point depression? Freezing-Point Depression Freezing-Point  How is freezing-point depression quantified? 13.6 Colligative Properties Freezing-Point Depression Freezing-Point

Figure 13.7 p. 449 13.6 Colligative Properties

 Practice

What is the freezing point of a saturated sugar solution (saturated at 0oC)?

At 0oC, 64.447 mass fraction of the solution is sugar. Freezing-Point Depression Freezing-Point 13.6 Colligative Properties Freezing-Point Depression Freezing-Point

Table 13.2 p. 450 13.6 Colligative Properties

 (Review, 13.1) What is vapor pressure?

 (Review, 12.6) What is ? Osmosis and Osmotic Pressure Osmosis and 13.6 Colligative Properties Osmosis and Osmotic Pressure Osmosis and

Figure 13.9 p. 453 13.6 Colligative Properties

 (Review, 13.1) How do we model solutions on the particle level?  (Review, 12.6) How do we model vapor pressure on the particle level?  What happens if a semi-permeable membrane is placed between a solvent and a solution? Osmosis and Osmotic Pressure Osmosis and 13.6 Colligative Properties Osmosis and Osmotic Pressure Osmosis and

Figure 13.8 p. 452 13.6 Colligative Properties  (Review, 13.1) How do we model solutions on the particle level?  (Review, 12.6) How do we model vapor pressure on the particle level?  What happens if a semi-permeable membrane is placed between a solvent and a solution?

 What is osmosis and osmotic pressure Osmosis and Osmotic Pressure Osmosis and  How is osmotic pressure quantified? 13.6 Colligative Properties

 Practice

The molar mass of a type of hemoglobin was determined by osmotic pressure of 4.60 mmHg for a solution at 20oC containing 3.27 g of hemoglobin in 0.200 L of solution. What is the molar mass of hemoglobin? Osmosis and Osmotic Pressure Osmosis and 13.6 Colligative Properties

 What are colligative properties? ◦ Boiling-point elevation ◦ Freezing-point depression ◦ Osmotic pressure  What do colligative properties depend on? Electrolyte Solutions Electrolyte

 Will now discuss electrolyte solutions 13.6 Colligative Properties

 How are nonelectrolytes and electrolytes different?  (Review 2.6 and 2.7) How is a solution of calcium chloride different from sodium chloride?  How is this incorporated into the relationships for colligative properties? Electrolyte Solutions Electrolyte 13.6 Colligative Properties

Electrolyte solutions – van’t Hoff Factors Electrolyte Solutions Electrolyte

Table 13.3 p. 457 13.6 Colligative Properties

 Why are the measured values not equal to the calculated values?  What are ion pairs? Electrolyte Solutions Electrolyte

Figure 13.11 p. 457 13.6 Colligative Properties

 Practice:

What is the freezing point of a solution made by dissolving 10 g of sodium chloride in 100 g of water?

What mass of magnesium chloride would be Electrolyte Solutions Electrolyte needed for the same change in freezing point?