Chapter 15 Notes: Solutions
Nature of Solutions
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What is a solution?
 **ask students what kinds of substances they think of when they hear the word
solution and write down the responses** **some surprising examples are below
 Air
 Brass
 Window glass
 Definition: any substance that is evenly distributed throughout another substance
- Can be solid, liquid, or gas
 Also called homogenous mixtures because they are not pure substances
 The substance that is doing the dissolving is called the solvent
- the substance that is being dissolved is the solute
- sometimes liquids are so soluble in each other that at one concentration
one substance is the solute and at another concentration it is the solvent
 for example, alcohol and water
 not all solutions are homogenous
- at the beach, the sea water in a handful of sea water-sand mixture is not
evenly distributed throughout the sea water
 this is an example of a heterogeneous mixture
- in fact, sand itself is a heterogeneous mixture
The process of solvation
 Show the Dissociation of Water video
- In the video, when the charged ions are surrounded by water molecules,
the ion is hydrated
- The process of dissolving a solute in a solvent is called solvation
 **Then discuss how you would represent the dissociation of salt in an equation**
NaCl(s)  Na+(aq) +Cl-(aq)
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Molecular salvation
- Molecular substances (not just ions) will also dissolve in water
 For example, ethanol, which is a polar molecule
- The polar ends of the solute are attracted to the oppositely
charged polar end of the solvent
 This disperses the solute molecules throughout the
solvent
 These attractions can result in a smaller volume
than expected
 **Drawing 1**
- This occurs because ethanol has hydrogen
bonding just like water
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The additional hydrogen bonding that takes
place between ethanol and water molecules
causes the volume to decrease
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Miscibility
- Our previous example (water and ethanol) can be mixed in any proportion
and always result in a solution
 Many gaseous solutions have this same property
- Each of the components is said to be miscible
- Two liquids that do not mix are immiscible
 Like dissolves like
- In most cases, solvation depends on whether two substances are alike
(both polar or both nonpolar)
- There are exceptions, especially when dealing with trace amounts of
solute
Solubility
 A saturated solution is one that cannot dissolve any more solute at a given
temperature
 An unsaturated solution is able to dissolve more solute
 Definition: the amount of a substance needed to make a saturated solution at a
specified temperature
 In the same liquid, solubilities of different solids vary considerably
- So ‘soluble’ is a term that does not have an exact meaning
Factors affecting solubility
 Temperature
- For many substances, solubility increases as the temperature of the
solution increases
- One method of growing crystals uses this principle
 A heated saturated sugar solution is cooled
 Crystals start to form and grow as the solution gets cooler and
cooler
- A solution can be cooled so that more solute can be dissolved without
crystals forming
 The solution now holds more than it usually would
- the solution is considered supersaturated
 Pressure
- **have the students think about opening a can of pop that has been shaken
up**
- The solubility of a gas depends on the pressure
 If the pressure is reduced, the gas leaves solution
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Ionic Equations and Precipitation Reactions
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Reactions in Solution
 Reactions that take place in aqueous solutions can form a precipitate
- For example, the reaction between silver nitrate and potassium chloride
AgNO3(aq) + KCl(aq)  AgCl(s) + KNO3(aq)
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We know from last semester that this is a double replacement
reaction
- How can we explain this reaction in terms of solubility?
 **have students look at solubility table**
Writing Ionic Equations
 In aqueous solutions, ionic compounds (that are soluble) dissociate into ions
- From the previous reaction, we get two dissociation equations
AgNO3(s)  Ag+(aq) + NO3-(aq)
and
KCl(s)  K+(aq) + Cl-(aq)
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We can use ions instead of formulas to get:
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl(s) + K+(aq) + NO3-(aq)
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 This is called an ionic equation
We then look for spectator ions: ions that do not participate in the
reaction
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq)  AgCl(s) + K+(aq) + NO3-(aq)
Ag+(aq) + Cl-(aq)  AgCl(s)
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 `this is called a net ionic equation
Precipitates to do not always from; there can be other products, like water, gas or
metallic ions
sodium hydroxide and hydrogen chloride
zinc and hydrochloric acid
baking soda and nitric acid
carbon dioxide and water
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Colligative Properties of Solutions
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Boiling Point Elevation and Freezing Point Depression
 If you had a table sugar – water solution, it would boil at a temperature greater
than 100°C and freeze at a temperature below 0°C
 Boiling point elevation
- Definition: the temperature difference between the boiling point of a pure
solvent and the temperature at which a solutions begins to boil
 Freezing point depression
- Definition: the temperature difference between the freezing point of a pure
solvent and the temperature at which solution begins to freeze
 The higher the concentration of the solute, the greater the solution deviates from
the pure solvent’s boiling and freezing points
 Ions that dissociate in water can carry en electrical current and are called
electrolytes
- In general, electrolytes are more effective than molecular solutes in
changing the boiling and freezing points of a solution
 this is because electrolytes break into several parts whereas
molecular solutes do not
 general principle: the number of solute particles, not their size or whether
they are molecules or ions, determines to a large extent how the solute will
affect the boiling and freezing points of water
 both of these properties (BPE, FPD) are called colligative properties
Molality
 Recall that molarity is the number of moles in a given volume of solution
 When predicting temperature changes in initial boiling and freezing points,
molality is used
- It is the number of moles per kilogram of solvent
- Example: How would you prepare 0.50m solution of sucrose, C12H22,
using 500.0 grams of water?
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Calculations involving colligative properties
 The different in boiling point between the solution and pure solvent is represented
by:
ΔTb = kb x m, where kb = 0.51 °C·kg H2O / molsolute
 The difference in freezing point is represented by
ΔTf = kf x m, where kf = 1.86 °C·kg H2O / molsolute
 Example problems:
- At what temperature will a solution that is composed of 0.73 moles of
glucose in 650 mL of water being to boil?
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At what temperature will a saltwater solution freeze if the solution is
composed of 280 g of sodium chloride per 1000 g of water?
There are constants for solvents other than water in your book
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Calculating molar mass
 You can figure out the molar mass of a compound by knowing the freezing point
 Example problem: When 36.0 g of a nonvolatile molecular substance is dissolved
in 100 g of water, the solution begins to freeze at -3.72°C. What is the molar mass
of solute?
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