Ch. 13: Solutions

Dr. Namphol Sinkaset Chem 201: General Chemistry II

I. Chapter Outline

I. Introduction II. Types of Solutions and Solubility III. Energetics of Solution Formation IV. Solution Equilibrium V. Solution Concentration VI. Colligative Properties

I. Solutions Everywhere

• solution: homogeneous mixture of 2 or more substances or components

• solute: substance present in smaller amount • solvent: substance present in greater amount

II. Types of Solutions

• Aqueous solutions are most common. • solubility: amount of a substance that will

dissolve in given amount of solvent • Solubility depends on IM forces and entropy.

III. Why Do Solutions Form?

• What is the reason most things happen?*

• Is the PE lowered when the barrier is removed?

• No! So why does this gas/gas solution form?

III. Entropy • In addition to lowering PE, nature’s

desire to disperse energy is a reason why we observe things happening.

• entropy: measure of energy of randomization or energy dispersal in a system

• KE originally confined now spread out over a larger volume.

• Entropy also explains heat flow. • Much more on entropy later…

III. What About IM Forces? • In most solutions, PE is a factor. • In solution, new particle-particle interactions form. • IM forces can help or hinder solution formation.

III. Solution Interactions

• Remember that “like dissolves like.” • If liquids are soluble in all proportions, then

they are said to be miscible.

III. When Will a Solution Form?

• If solvent molecules and solute molecules interact more strongly with their own kind, a solution tends not to form.

III. Sample Problem 13.1

• Explain the trends in solubilities in terms of intermolecular forces.

III. Energetics of Solution Formation

• Although not a reaction, we see energy changes when a solution forms.

• The energy change observed depends on relative strengths of IM forces.

• The enthalpy of solution, ΔHsoln, can be exo or endo and can be estimated using a Hess’s Law type of calculation.

III. Steps of Solution Formation

1) Separating solute into its constituent particles, ΔHsolute.

III. Steps of Solution Formation

2) Separating solvent particles to make room for solute particles, ΔHsolvent.

III. Steps of Solution Formation

3) Mixing solute particles with solvent particles, ΔHmix.

III. Steps of Solution Formation

• Hess’s Law allows us to break up a process into a series of steps.

• Therefore, the 3-step path for solution formation yields the following equation.*

ΔHsoln = ΔHsolute + ΔHsolvent + ΔHmix

III. Hess’s Law Applied

III. Aqueous Ionic Solutions

• Many solutions are ionic, water-based. • For these, ΔHsolvent and ΔHmix can be combined

into ΔHhydration. • heat of hydration: enthalpy change when 1

mole of gaseous solute ions are dissolved in water

III. Heat of Hydration

• Why is ΔHhydration always very exo for ionic compounds?*

• Compare relative strengths of ΔHsolvent and ΔHmix.

III. KF Aqueous Solution

• For ionic compounds, ΔHsolute = -ΔHlattice. • Thus, whether or not formation of an aqueous

ionic solution is exo or endo depends on the relative values of ΔHsolute and ΔHhydration.

IV. Solution Equilibrium

• Initially, when solid is placed in a solvent, solid particles rapidly go into solution.

• Eventually, there will be excess solute particles in solution, and they will begin to redeposit.

• At dynamic equilibrium, the rates of dissolution and deposition are equal.

IV. Dissolution of NaCl

IV. Solution Terminology

• Equilibrium between solid and solvated particles can only occur if solution is saturated.

• If there’s less than this equilibrium amount, the solution is unsaturated.

• Some solutions can be made to be supersaturated.

IV. Solubility of Solids

• In general, the solubility of solids increases with temperature.

IV. Solubility of Gases

• Gas solubility in water depends on temperature and external pressure.

• Gas solubility decreases with temp. • Gas solubility increases with higher

external pressure.

IV. Henry’s Law

• Dependence of gas solubility on external pressure is described by Henry’s Law.

• Henry’s constant (kH) usually in M/atm.

𝑆𝑆𝑔𝑔𝑔𝑔𝑔𝑔 = 𝑘𝑘𝐻𝐻𝑃𝑃𝑔𝑔𝑔𝑔𝑔𝑔

IV. Henry’s Constants

• Why is the value for ammonia (NH3) so high?*

V. Solution Concentration

• One of the most important aspects about a solution is its concentration.

• There are many different units of concentration, and we need to be able to convert between all of them.

• Which unit we use depends on what type of problem we are solving.

V. Units of Concentration

V. Sample Problems 13.2 1) How many grams of KNO3 are needed to

make 200.0 mL of a 0.025 M solution? 2) Calculate the weight fraction of NaCl in a

solution consisting of 12.5 g NaCl and 75.0 g H2O.

3) Calculate the number of grams of CH3OH needed to make a 0.250 m solution w/ 2000 g of H2O.

4) What’s the molality of a 37.0% HCl solution? 5) A 40.0% HBr solution has a density of 1.38

g/mL. What is molarity?

VI. Colligative Properties

• Salt is added to ice in an ice cream maker. • Icy roads are salted in winter. • Antifreeze is mixed with water. • These are practical uses of colligative

properties. • colligative properties: properties that

depend on number of dissolved particles, not the type particle.

VI. Colligative Properties

• We will look at 3 colligative properties: Vapor pressure lowering Freezing point depression Boiling point elevation Osmotic pressure

VI. Vapor Pressure

• Vapor pressure is the pressure of a gas in dynamic equilibrium w/ its liquid.

• It’s a measure of how many molecules go into the gas phase.

VI. Vapor Pressure of Solutions

• All liquid solutions of nonvolatile solutes have lower vapor pressures than the pure solvents.

• Why?*

VI. Striking Example

• Why is the volume on the right increasing?*

VI. Raoult’s Law

• The vapor pressure of a solution can be calculated using Raoult’s Law.

VI. Sample Problem 13.3

• What’s the vapor pressure at 20 °C of a solution of 20.0 g dibutyl phthalate (MW = 278 g/mole) in 50.0 g of octane (MW = 114 g/mole) if pure octane has a vapor pressure of 10.5 torr at 20 °C?

VI. Ionic Solutes

• An aqueous solution of MgCl2 will have a lower vapor pressure than an aqueous solution of NaCl – why?*

• Recall the definition of colligative properties to understand why!

VI. Sample Problem 13.4

• A solution contains 0.115 mole H2O and an unknown number of moles of sodium sulfate. The vapor pressure of the solution at 30 °C is 25.7 torr. If the vapor pressure of pure water at 30 °C is 31.8 torr, how many moles of sodium sulfate are in solution?

VI. Ideal and Nonideal Liquid/Liquid Solutions

• Ideal solutions obey Raoult’s law at all concentrations for both solute and solvent.

• In an ideal solution, ΔHsoln = 0, which means forces of attraction between all molecules are identical.

• We can calculate vapor pressures of each component…

VI. Rauolt’s Law for 2 Liquids

• Stronger or weaker solute-solvent interactions will cause deviations as shown in the curves in the next slide…

VI. Raoult’s Law Curves

VI. Effect of Vapor Pressure Lowering

• Vapor pressure lowering occurs at all temperatures.

• This lowering alters the appearance of phase diagrams of a solution when compared to a pure solvent.

VI. Altered Phase Diagram

VI. Changes in FP and BP

• We see that there’s a freezing point depression and a boiling point elevation.

• Since they are caused by vapor pressure lowering, these are also colligative properties.

VI. Changes in FP and BP

• Equations for calculating freezing point depression and boiling point elevation are very simple.

VI. Sample Problem 13.5

• At what temperature will a 10% aqueous solution of sugar (C12H22O11) boil if the Kb for water is 0.51 °C/m?

VI. Sample Problem 13.6

• A solution made by dissolving 3.46 g of an unknown compound in 85.0 g of benzene froze at 4.13 °C. What’s the molar mass of the compound if Kf for benzene is 5.07 °C/m and the normal freezing point of benzene is 5.45 °C?

VI. Osmosis • osmosis: the flow of solvent from a solution of

lower solute concentration to one of higher solute concentration

• Why does osmosis occur?*

VI. Osmotic Pressure

• The flow of solvent can be stopped by applying external pressure. This pressure is called the osmotic pressure.

• The “gas” value of R is used.

VI. Isotonic Solutions • The concept of osmotic pressure is important

in medicine. • IV solutions need to have the same osmotic

pressure as body fluids; if it’s off, there’s a “burning” sensation.

VI. Ionic Solutions

• Since colligative properties only depend on number, we need to consider how many particles form when ionics form a solution.

• We have to adjust the molality/molarity values used in the equations.

• Problem – ionics do not dissociate completely!

VI. Ion Pairs

• In an aqueous NaCl solution, some Na+ and Cl- pair up and behave as one entity.

• These are known as ion pairs.

VI. Accounting for Ion Pairs

• We use the van’t Hoff factor (i) to account for ion pairs.

VI. van’t Hoff Factors

VI. Equations for Ionic Solutions

• Incorporating van’t Hoff factors is easy…