Chapter 12 SOLUTIONS

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Chapter 12SOLUTIONS
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Chapter 12

• Section 1 Types of Mixtures
• Section 2 The Solution Process
• Section 3 Concentration of Solutions

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12.1 - Types of Mixtures

• Distinguish between electrolytes and
  nonelectrolytes.
• List three different solute-solvent combinations.
• Compare the properties of suspensions, colloids,
  and solutions.
• Distinguish between electrolytes and
  nonelectrolytes.

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Introduction

• Sugar Water is a mixture

• Soil is a mixture

• 2 or more different substances

• Homogeneous Mixture
• (uniform properties)

• Heterogeneous Mixture
• (nonuniform properties)

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Solutions

• Solution homogeneous mixture of two or more
  substances
• Solute substance dissolved in a solution
• Solvent substance doing the dissolving (usually
  in a large quantity)
• Can be any combination of solid, liquid, or gas
• A homogeneous mixture of solids is called an
  alloy
• Ex brass solution of zinc in copper

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Examples of Solutions
In these examples, the substance with the greater
volume is the solvent
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Particle Models for Gold and a Gold Alloy
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Solutions

• Soluble capable of being dissolved
• Insoluble not capable of being dissolved
• What substances are soluble?
• Depends on your solute your solvent (and other
  factors, too)
• In general, follow the rule
• like dissolves like

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Polarity Solubility

• Polar substances dissolve in polar substances
• Ionic compounds
• Polar molecules
• Nonpolar substances dissolve in nonpolar
  substances
• Nonpolar molecules
• Organic compounds

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Like Dissolves Like
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Heterogeneous vs. Homogeneous
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Suspensions

• Suspension particles in a solvent are so large
  they settle out unless the mixture is constantly
  stirred
• What is an example of a suspension?
• Oil and water
• Sand and water

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Colloids

• Colloid particles are small enough to be
  suspended throughout the solvent by constant
  movement of the surrounding molecules
• (particles that are intermediate in size)
• Colloidal particles make up the dispersed phase,
  and water is the dispersing medium.
• What are some examples of Colloids?
• Paint (solid in liquid)
• Whipped cream (gas in liquid)
• fog

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Emulsions A Specific Type of Colloid
a mixture of two or more immiscible
(un-blendable) liquids
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Tyndall Effect

• Tydnall Effect when light is scattered by
  colloidal particles dispersed in a transparent
  medium (like water)
• This is used to distinguish between a solution
  and a colloid

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Solutions, Colloids, and Suspensions
Solutions Colloids Suspensions
Type of Mixture Homogeneous Heterogenous Heterogenous
Particle Size 0.01 1 nm 1 1000 nm gt 1000 nm
Separation? No separation No separation Particles settle
Filtration Possible? No No Yes
Tyndell Effect? No Yes Yes No
Property Uniformity? Uniform nonuniform nonuniform
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Electrolytes

• Electrolyte a substance that dissolves in water
  to give a solution that conducts electric current
• What would make a good electrolyte?
• Ionic compounds the positive and negative ions
  separate from each other in solution and are free
  to move making it possible for an electric
  current to pass through the solution
• Polar covalent molecules

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Nonelectrolyte

• Nonelectrolyte a substance that dissolves in
  water to give a solution that does not conduct
  electricity
• Neutral solute molecules do not contain mobile
  charged particles so they do not conduct electric
  current
• Example sugars, alcohols, organic hydrocarbons

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Electrical Conductivity of Solutions
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Strong and Weak Electrolytes

• What is an electrolyte?
• The strength of an electrolyte depends on the
  ability of a substance to form ions, or the
  degree of ionization or dissociation.

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Strong Electrolytes

• Strong Electrolyte conducts electricity well due
  to presence of all or almost all of the dissolved
  compound as ions
• What is an example of a strong electrolyte?
• Strong acids (HCl, HNO3)
• Strong bases (NaOH)
• ionic compounds (assuming their soluble)

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Weak Electrolytes

• Weak Electrolyte conducts electricity poorly due
  to presence of a small amount of the dissolved
  compound as ions
• What is an example of a weak electrolyte?
• Weak acids and bases
• Ex Acetetic Acid, H3PO4 NH3 (weak base)

• HF gtgt H and F

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Models for Strong and Weak Electrolytes and
Nonelectrolytes
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12.2 - The Solution Process
List and explain three factors that affect the
rate at which a solid solute dissolves in a
liquid solvent. Explain solution equilibrium,
and distinguish among saturated, unsaturated, and
supersaturated solutions. Explain the meaning of
like dissolves like in terms of polar and
nonpolar substances. List the three interactions
that contribute to the enthalpy of a solution,
and explain how they combine to cause dissolution
to be exothermic or endothermic. Compare the
effects of temperature and pressure on solubility.
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Factors Affecting the Rate of Dissolution

• 1. Increase the surface area
• -- because dissolution occurs at the surface
• 2. Stirring or shaking
• -- increases contact between solvent and solute
• 3. Increase the temperature
• -- increases collisions between solute and
  solvent and are of higher energy

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Dissolving Process Video
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Solubility

• If you add spoonful after spoonful of sugar to
  tea, eventually no more sugar will dissolve.
• There is a limit to the amount of solute that can
  dissolve in a solvent.

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Solubility Values

• Solubility measure of how much solute can be
  dissolved in a specific amount of solvent at
  specific temperature
• Specifically, how much is required to form a
  saturated solution
• Example solubility of sugar is 204 g per 100 g
  of water at 20OC
• Solubilities vary widely and must be determined
  experimentally

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Solutions

• Saturated Solution contains the maximum amount
  of dissolved solute
• How does this happen?
• When a solute is added, the molecules leave the
  solid surface and move about at random.
• As more solute is added, more collisions occur
  between dissolved solute particles. Some of the
  solute molecules return to the crystal.
• When maximum solubility is reached molecules are
  returning to the solid form at the same rate they
  are going into solution.
• This is called solution equilibrium

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Solutions

• Unsaturated Solution contains less solute than a
  saturated solution under the same condition
• Supersaturated Solution contains more dissolved
  solute than a saturated solution
• Usually created by cooling a saturated solution
  slowly.
• If disturbed, a supersaturated solution will
  precipitate out crystals of excess solute

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Saturated Solution and Temperature
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Mass of Solute Added Versus Mass of Solute
Dissolved
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Solubility of Common Compounds
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Solubility of Common Compounds
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Dissolving Ionic Compounds in Aqueous Solutions

• What type of forces are present in water?
• These forces attract the ions in ionic compounds
  and surround them, separating them from the
  crystal surface and drawing them into solution.
• Hydration when water is the solvent
• Ions are said to be hydrated
• A solute particle that is surrounded by solvent
  molecules is solvated. (like hydrated, but not
  water)

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Hydration of Lithium Chloride
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Liquid Solutes and Solvents

• Oil and water do not mix because oil is nonpolar
  and water is polar. The hydrogen bonding squeezes
  out whatever oil molecules may come between them.
• Two polar substances or two nonpolar substances
  easily form solutions.
• Immiscible liquids that are not soluble in each
  other
• Miscible liquids that dissolve freely in one
  another in proportion

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Comparing Miscible and Immiscible Liquids Video
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Temperature on Solubility

• In general, an increases in temperature usually
  increases the solubility of solids in liquids
• A few solid solutes are actually less soluble at
  higher temperatures.

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Pressure and Solubility

• Increases in pressure, increase gas solubilities
  in liquids
• An equilibrium is established between a gas above
  a liquid solvent and the gas dissolved in a liquid

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Pressure and Solubility Cont.

• Increasing the pressure
• causes gas particles to collide with the liquid
  surface and forces more gas into the solution
• Decreasing the pressure
• allows more dissolved gas to escape from solution

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Soda Carbonation Video
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Henrys Law

• Henrys Law solubility of a gas in a liquid is
  directly proportional to the partial pressure of
  that gas on the surface of the liquid
• In carbonated beverages, the solubility of carbon
  dioxide is increased by increasing the pressure.
  The sealed containers contain CO2 at high
  pressure, which keeps the CO2 dissolved in the
  beverage, above the liquid.
• When the beverage container is opened, the
  pressure above the solution is reduced, and CO2
  begins to escape from the solution.
• Effervescence the rapid escape of a gas from a
  liquid

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Henrys Law Video
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Effervescence
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Temperature and Solubility of Gas

• Increasing the temperature
• Increases the average kinetic energy and allows
  more solute to escape the attraction of the
  solvent and escape to the gas phase
• At higher temperatures, equilibrium is reached
  with fewer gas molecules in solution

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Enthalpies of Solution

• The formation of a solution is accompanied by an
  energy change.
• The formation of a solid-liquid solution can
  either absorb energy (KI in water) or release
  energy as heat (NaOH in water)

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Enthalpies of Solution Cont

• Before dissolving begins, solute particles are
  held together by intermolecular forces. Solvent
  particles are also held together by
  intermolecular forces.
• Energy changes occur during solution formation
  because energy is required to separate solute
  molecules and solvent molecules from their
  neighbors.

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Enthalpy changes during the formation of a
solution
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Enthalpies of Solution Cont.

• Enthalpy of Solution net amount of energy
  absorbed as heat by the solution when a specific
  amount of solute dissolves in a solvent
• Negative energy is released (exo)
• (Steps 1/2 attractions is less than Step 3)
• Positive energy is absorbed (endo)
• (Steps 1/2 attractions is greater than Step3)

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12.3 - Concentration of Solutions
Given the mass of solute and volume of solvent,
calculate the concentration of solution. Given
the concentration of a solution, determine the
amount of solute in a given amount of
solution. Given the concentration of a solution,
determine the amount of solution that contains a
given amount of solute.
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Concentration

• Concentration measure of the amount of solute in
  a given amount of solvent or solution
• Concentration is a ratio any amount of a given
  solution has the same concentration
• Dilute the opposite of concentrated
• These terms DO NOT relate to saturation!

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Concentration Units
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Concentration
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Molarity

• Molarity(M) moles of solute in one liter of
  solution

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Preparing a Solution

• NOTE 1 M solution is NOT made by adding 1 mol of
  solute to 1 L of solvent

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Example Problem A

• You have 3.50 L of solution that contains 90.0 g
  of sodium chloride. What is the molarity of that
  solution?
• Given solute mass 90.0 g sodium chloride
• solution volume 3.50 L
• Unknown molarity of NaCl solution

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Example Problem B

• You have 0.8 L of a 0.5 M HCl solution. How many
  moles of HCl does this solution contain?
• Given volume of solution 0.8 L
• concentration of solution 0.5 M HCl
• Unknown moles of HCl in a given volume

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Example Problem C

• To produce 40.0 g of silver chromate, you will
  need at least 23.4 g of potassium chromate in
  solution as a reactant.
• If you have 5 L of a 6.0 M K2CrO4 solution. What
  volume of the solution is needed to give you the
  23.4 g K2CrO4 needed for the reaction?

• Given volume of solution 5 L
• concentration of solution 6.0 M K2CrO4
• mass of solute 23.4 K2CrO4
• mass of product 40.0 g Ag2CrO4
• Unknown volume of K2CrO4 solution in L

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Example Problem C Solution

• Find Molar Mass of Solute
• Find Moles of Solute
• Find Volume of Solution

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Molality

• Molality(m) moles of solute per kilogram of
  solvent
• Molality is used when studying properties of
  solutions related to vapor pressure and
  temperature changes, because molality does not
  change with temperature.

Molality (m) mol of solute kg of
solvent
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Making a Molal Solution
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Example Problem A

• A solution was prepared by dissolving 17.1 g of
  sucrose (table sugar, C12H22O11) in 125 g of
  water. Find the molal concentration of this
  solution.
• Given solute mass 17.1 C12H22O11
• solvent mass 125 g H2O
• Unknown molal concentration

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Sample Problem A Solution

• Convert the Given
• Find Molality

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Example Problem B

• A solution of iodine, I2, in carbon
  tetrachloride, CCl4, is used when iodine is
  needed for certain chemical tests. How much
  iodine must be added to prepare a 0.480 m
  solution of iodine in CCl4 if 100.0 g of CCl4 is
  used?
• Given molality of solution 0.480 m I2
• mass of solvent 100.0 g CCl4
• Unknown mass of solute

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Molality, continued

• Convert Solvent
• Find moles
• Convert to grams

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Chapter 13Ions in solution and Colligative
Properties
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Chapter 13

• Section 1 Compounds in Aqueous Solutions
• Section 2 Colligative Properties

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13.1 Compounds in Aq Solutions

• Write equations for the dissolution of soluble
  ionic compounds in water.
• Predict whether a precipitate will form when
  solutions of soluble ionic compounds are
  combined, and write net ionic equations for
  precipitation reactions.
• Compare dissociation of ionic compounds with
  ionization of molecular compounds.
• Draw the structure of the hydronium ion, and
  explain why it is used to represent the hydrogen
  ion in solution.
• Distinguish between strong electrolytes and weak
  electrolytes.

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Dissociation

• Dissociation separation of ions that occurs when
  an ionic compound dissolves

1 mol 1 mol 1 mol
1 mol 1 mol 2 mol
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Dissociation of NaCl
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Example Problem 1

• 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?

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Example Problem 1 Solution

• Equation
• How many moles of each ion?
• What is the total number of ions present?

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Precipitation Reactions

• No ionic compound is completely insoluble, but
  ones with very low solubility can be considered
  to be insoluble and to form a precipitate.
• REMEMBER THE SOLUBILITY RULES!!!

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General Solubility Guidelines
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Which Ionic Compounds are INSOLUBLE?
NiCl2 KMnO4 CuSO4 Pb(NO3)2 AgCl
CdS
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Particle Model for the Formation of a Precipitate
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Net Ionic Equations

• Total Ionic Equation contains all compounds and
  ions including any substances that are not
  involved in the reaction (spectator ions)
• Spectator ions do not take part in a chemical
  reaction and are found in solution before and
  after the reaction
• Net Ionic Equation includes only those compounds
  and ions that undergo a chemical change in a
  reaction in an aqueous solution

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Example Problem

• 1. Write the balanced reaction
• 2. Identify precipitate

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• 3. Write the ionic equation
• 4. Write the net ionic equation

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2
2
2
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Example Reaction
Starting with two solutions of Potassium Sulfate
and Barium Nitrate
At the end of the reaction, there is one solution
and one precipitate
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Net Ionic Equation Video
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Writing Net Ionic Equaitons
The ions crossed out in blue are the spectator
ions and do not take part in the reaction.
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Ionization

• Ionization formation of ions from solute
  molecules by the action of the solvent
• When a molecular compound dissolves and ionizes
  in a polar solvent, ions are formed where none
  existed in the undissolved compound
• Example HCl is molecular but ionizes in water

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Dissociation and Ionization Video
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Hydronium Ion

• When the H ion is formed, it attracts other
  molecules or ions so it does not exist alone.
• It forms the hydronium ion H3O

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13.2 Colligative Properties

• List four colligative properties, and explain why
  they are classified as colligative properties.
• Calculate freezing-point depression,
  boiling-point elevation, and solution molality of
  nonelectrolyte solutions.
• Calculate the expected changes in freezing point
  and boiling point of an electrolyte solution.
• Discuss causes of the differences between
  expected and experimentally observed colligative
  properties of electrolyte solutions.

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Colligative Properties

• Colligative Properties properties that depend on
  the concentration of solute particles but not on
  their identity
• Vapor Pressure Lowering
• Freezing Point Depression
• Boiling Point Elevation
• Osmotic Pressure

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Colligative Properties

• Freezing Point Depression
• Freezing Point of a solution is always lower than
  the pure solvent.
• Boiling Point Elevation
• Boiling Point of a substance is always higher
  than the pure solvent.

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Freezing Point Depression

• Freezing Point Depression (?tf )
• Difference between the freezing points of the
  pure solvent and a solution
• Molal freezing point constant (Kf)
• Freezing point depression of a solvent in 1 molal
  solution (nonvolatile, nonelectrolyte)
• ?tf Kfm

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• If the solution contains an electrolyte it will
  produce a higher concentration of mols than the
  undissolved solute
• ?tf Kf x m x d.f.
• d.f is the dissociation factor (how many
  particles the solute break up into)
• Ex NaCl produces 2 mols of ions so d.f. 2

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Molal Freezing Point and Boiling Point Constants
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Freezing-Point Depression
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Example Problem 3

• What is the freezing-point depression of water in
  a solution of 17.1 g of sucrose, C12H22O11, in
  200. g of water? What is the actual freezing
  point of the solution?
• Given 17.1 g C12H22O11 of solute 200.
  g H2O of solvent
• Unknown a. freezing-point depression b.
  freezing point of the solution

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Example Problem 3 Solution

• Calculate Moles
• Calculate Molality

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Example Problem 3 Cont.

• Calculate Freezing Point Depression
• ?tf Kfm
• ?tf 0.250 m (-1.86C/m) -0.465C
• Calculate New Freezing Point
• Freezing Pointsolution Freezing Pointsolvent
  ?tf
• FPsolution 0.000C (-0.465C) -0.465C

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Boiling Point Elevation

• Boiling Point Elevation (?tb)
• Difference between the boiling points of the pure
  solvent and a solution
• Molal boiling point constant (Kb)
• Boiling point elevation of a solvent in 1 molal
  solution solution (nonvolatile, nonelectrolyte)
• ?tb Kbm

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Boiling-Point Elevation and the Presence of
Solutes
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Example Problem 4

• What is the boiling-point elevation of a solution
  made from 20.1 g of a nonelectrolyte solute and
  400.0 g of water? The molar mass of the solute is
  62.0 g.
• Given solute mass 20.1 g solute molar
  mass 62.0 g/mol solvent mass and
  identity 400.0 g H2O
• Unknown boiling-point elevation

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Example Problem 4 Solution

• Calculate Moles
• Calculate Molality
• Calculate Boiling Point Elevation

?tb 0.51C/m 0.810 m 0.41C
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Electrolytes and Colligative Properties

• Electrolytes depress the freezing point and
  elevate the boiling point of a solvent MORE than
  expected.
• Electrolytes produce MORE THAN 1 mol of solute
  particles for each mole of a compound
• Moles of Solute?
• ___
• ___
• ___

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Calculations for Electrolytes

• Colligative properties depend on the total
  concentration of solute particles.
• The changes in colligative properties cause by
  electrolytes is proportional to the TOTAL
  molality of all dissolved particles, not to
  formula units.
• For the same molal concentrations of sucrose and
  sodium chloride
• What would you expect the effect on the
  colligative property to be? By what proportion is
  there a difference?

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• Sample Problem F
• 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?

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• Sample Problem F Solution
• Given solute mass and formula 62.5 g Ba(NO3)2
• solvent mass and identity 1.00 kg water
• ?tf Kfm
• Unknown expected freezing-point depression
• Solution

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• Sample Problem F Solution, continued
• Solution

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• Sample Problem F Solution, continued
• Solution

Each formula unit of barium nitrate yields three
ions in solution.
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Actual Values for Electrolytes
The actual values of the colligative properties
for all strong electrolytes are almost what would
be expected based on the number of particles they
produce in solution.
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Actual Values for Electrolytes

• The differences between the expected and
  calculated values are caused by the attractive
  forces that exist between dissociated ions in
  aqueous solution.
• According to Debye and Hückel a cluster of
  hydrated ions can act as a single unit rather
  than as individual ions, causing the effective
  total concentration to be less than expected.
• Ions of higher charge have lower effective
  concentrations than ions with smaller charge.