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Chapter 8 — Solutions and Their Behavior. Some Terminology. Solution A solution is a homogeneous mixture (not possible to see boundaries between components) that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent

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Chapter 8 — Solutions and Their Behavior

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## Chapter 8 — Solutions and Their Behavior

### Some Terminology

• Solution

• A solution is a homogeneous mixture (not possible to see boundaries between components) that consists of one or more solutes uniformly dispersed at the molecular or ionic level throughout a medium known as the solvent

• Solvent is present in the larger amount

• Solute is present in smaller amount than the solvent

• Examples:

Liquids: NaCL (sodium chloride + water)

0.89% 99.11%

(solute) (solvent)

Gases: Air is a solution of N2, O2, few other minor gases

### Units of Concentration 1

• Molality (m) aka “molal concentration”

Solvent= does not include solute

• Molarity (M) aka “molar concentration”

Most common concentration unit in Chemistry

Solution=solute+solvent

### Example

• What is the molarity of a solution prepared by dissolving 45.0 grams of NaCl in enough water to give a total volume of 489 mL?

Na = 23 a.m.u.

Cl = +35 a.m.u.

58 a.m.u. 1mol NaCl = 58g

45g X 1 mol = 0.7759 mol of NaCL

58g

0.7759 mol of NaCl = 1.6 M

0.489 L

### Molality vs. Molarity

• Molality is never equal to molarity

• But the difference becomes smaller as solutions become moredilute(denominators are very similar)

• Molarity is more useful when dealing with solution stoichiometry

• Molality is more appropriate for dealing with physical chemistry

• Question: which is temperature-dependent?

• Molarity depends on temperature. Molarity will decrease as temperature increases since the amount of the solution will decrease (from evaporation). Temp  M

• Molality does not depend on temperature since mass (kg) does not change with temperature

### Units of Concentration 2

• Mole Fraction ( )

• Percent by Volume (% w/v) AND Percent by Weight (%w/w)

### Example

• Calculate the percent by weight NaCl in a solution comprised of 45.0 g NaCl and 457 g of water.

% (w/w) = grams of solute x 100% = 45g x 100% = 8.96%

grams of solution 502g

Solute = 45g of NaCl

Solution = 457g of NaCl (solute) + H2O (solvent) = 502g

### Units of Concentration 3

• Parts per million (ppm): Extremely dilute solutions. Compares amount of solute to a million parts of solution (rather than 100 parts).

• Parts per billion (ppb) Even more extremely dilute solutions. Compares amount of solute to a billion parts of solution (rather than 1million parts).

### Converting Between Units

• Every concentration unit is a ratio of two quantities

• Pick a sample size

This fixes one of the two quantities

• Use the factor-label method (dimensional analysis a.k.a. “conversion factor”) to systematically convert the given quantities into the desired quantities

### Some More Terminology

• Solubility

• Solubility is the amount of solute that will dissolve in a given amount of solvent at a given temperature

• Saturated

• A saturated solution contains the maximum amount of a solute, as defined by its solubility (no more solute will dissolve in a solution that is already saturated with that solute)

• Supersaturated

• A solution contains more solute than allowed by the solubility. Not a stable system since there is more solute dissolved than the solvent can “accommodate” and the excess solute will come out of solution and will crystallize as a solid (e.g., calcium oxalate or CaPO4 will make kidney stone) or it will separate as a liquid or it will bubble out as a gas (e.g., soda)

### Some More Terminology

• Miscible

• Two liquids are miscible if they are soluble in each other in all proportions

• Question: Can a solution be both saturated and dilute?

Yes, just remove some of the solvent.

If you want to see this for yourself, mix a little salt and water together. Then leave it stand so most of the water evaporates. You have saturated a dilute solution.

### Solubility Guidelines

• Like dissolves like

• Polar solutes are more soluble in polar solvents

• Nonpolar solutes are more soluble in nonpolar solvents

### Heat of Solution (Energy change or Enthalpy of Solution)

ΔH°solution is defined as the enthalpy (energy) change that accompanies dissolving exactly 1 mole of solute in a given solvent

Enthalpy = Heat (as long as pressure remains constant)

• Some substances have positive (endothermic) ΔH°solution and some have negative (exothermic) ΔH°solution

• Endothermic: energy flows INTO the system (gain of energy, temperature of the system decreases)

• Exothermic: energy flows OUT of the system (loss of energy, temperature of system increases)

ΔH°solution is the sum of two terms:

Lattice Energy and Solvation Energy

### LE and Solvation

• Lattice energy is the energy released when molecules or ions settle into a crystalline lattice. It is inherently exothermic (opposite charges coming towards each other)

• The opposite: Pulling all the ions away from each other in an ionic substance in the solid state requires overcoming the lattice energy (endothermic)

• Solvation (hydration) energy is the energy released when an ion (or molecules) settles into a sphere of solvent molecules

• Solvation is inherently exothermic (opposite charges coming towards each other also)

• Both LE and SE result from IM forces

• ### Sodium Chloride: Exothermic DH°solution

The magnitude of the lattice energy is lower than the magnitude of the solvation energy, in going from solid sodium chloride to a solution of sodium chloride so the energy of this system has decreased (it has undergone an exothermic change)

### Ammonium Chloride: Endothermic DH°solution

The magnitude of the lattice energy is greater than the magnitude of the solvation energy, in going from solid ammonium chloride to a solution of ammonium chloride so the energy of this system has increased (it has undergone an endothermic change)

### Effect of Pressure on Solubility

• Gaseous solute

• As P increases, solubility increases

• Henry’s Law:

S = kHPgas

S=solubility, k=Henry’s constant, P=partial pressure of the gas

• Liquid and Solid Solutes

• P has negligible effect (liquids and solids are not very compressible)

### Example

• The Henry’s law constant for oxygen in water is 0.042g/L/atm at 25°C. What is the solubility (in grams per mL) of O2 in pure water when the atmospheric pressure is 740 torr. In air, the mole fraction of oxygen is 0.21.

Ptotal = 740 torr x 1 atm = 0.974 atm

760 torr

PO2 = 0.21 x 0.974 atm = 0.204 atm

S = 0.042g/L/atm x 0.204 atm = 0.0086 g/L = 8.6 mg/L

### Colligative Properties of Solutions

• A colligative property depends only on the number of solute particles, not the identity of the solute particles

• Colligative properties include:

• Vapor pressure of a solution decreases with increasing [solute]

• Boiling point of a solution increases with increasing [solute]

• Freezing point of a solution decreased with increasing [solute]

• Osmotic pressure of a solution increases with increasing [solute]

### Vapor Pressure Decreases

• Introduce a solute (solute will “plug escape sites” of molecules that are at the surface of the solution “ready to jump out” and become gas)

Raoult’s Law (vapor pressure of a volatile component of a solution (P) is equal to the vapor pressure of the pure substance Po) times the mole fraction (X) of that substance)

• The vapor pressure of a solution is given by Raoult’s Law

### Example (Textbook page 206)

• What is the vapor pressure of a 5% (by mass) sugar solution at 45°C? (P°water = 71.9 torr)

### Example (Textbook, page 206)

• What is the total vapor pressure of a solution comprised of 0.75 moles of benzene and 0.45 moles hexane at 60°C? (P°benzene= 390 torr; P°hexane= 580 torr)

### Boiling Point Elevation(BP: Temperature at which the vapor pressure of the material is equal to the ambient pressure)

• As vapor pressure goes down, boiling point goes up

ΔTbp = Kbpmsolute

ΔTbp is the boiling point elevation

Kbp is the boiling point (ebullioscopic) constant

msolute is the molality of all solute particles

Freezing Point Depression(Freezing point: temperature at which the liquid phase of the material is in equilibrium with the solid phase (aka melting point)

ΔTfp = Kfpmsolute

is ΔTfp the freezing point depression

Kfp is the freezing point (cryoscopic) constant

msolute is the molality of all solute particles

### Ionic Solutes

• When ionic solutes dissolve, they dissociate into solvated ions

• Each ion counts as a particle for colligative properties

### Osmosis

• Osmosis is diffusion of water through a semipermeable membrane

• Solute particles are too big (or too polar) to make it across the membrane

• This is how water gets moved around cells

### Tonicity

• Isotonic solutions have equal concentrations of solute particles

• A hypertonic solution has a greater concentration of solute

• A hypotonic solution has lower concentration of solute

### Example

Hypertonic solution

Less water

Hypotonic solution

More water

### Osmotic Pressure (Increases with increasing solute concentration)

• Osmotic pressure (P) results from the potential drive for the concentration of water to equalize

P V = nRT

Or

P = MRT

A 1.0 M solution of glucose exerts an osmotic pressure of

22.4 atm at 25°C

### Question

• What will happen to a red blood cell when it is placed into pure water? Cells are isotonic with normal saline (0.89% NaCl)

### Question

• What will happen to a red blood cell when it is placed into 10% aqueous sodium chloride? Cells are isotonic with normal saline

(0.89% NaCl)

### Colloids

• Colloids are not true solutions

• Particle size is on the order of 10 to 200 nm

• Might be super-sized molecules (e.g., proteins) or aggregates of ions

• Colloidal particles cannot be filtered and do not settle out of solution

• Colloids exhibit the Tyndall effect (whereas solutions don’t)

• The particles in a colloid are large enough to scatter light passing through)

• Examples:

• blood, milk, jelly, propofol

### Surfactants

• A surface active agent breaks the surface tension of water

• Surfactants improve a solvent’s ability to be a solvent

• Soaps and detergents are common surfactants

greasy tail

### Soaps and Detergents

• A soap is prepared by hydrolyzing fat with alkali

• A detergent is a synthetic chemical with a structure similar to soap

• Both have a polar (hydrophilic) head and a non-polar (hydrophobic, greasy) tail

### Monolayers

• The greasy tails stick out of the surface of water

• This breaks down the surface tension of water

### Bilayers

• The tails can dissolve in each other in a bilayer

• This structure is used in cell membranes

### Micelles

• The tails can dissolve in each other forming a sphere

• This creates a non-polar microenvironment in the water