Chapter 13 Notes – Solution Chemistry


1.  Definitions of Solutions:

Substance:  One element - for example:  Gold, copper, silver...etc.

Solution:  A solid, liquid or gas which has one substance evenly distributed in another.

Miscible (Immiscible):  Like dissolves like.  Non polar substances will dissolve more easily in non polar liquids.  And polar substances in polar liquids (p. 406)

Hydration:  The process of polar water molecules pulling ionic compounds apart so that their ions float around When you crystallize compounds which were once in water, you get “hydrates”.  Hydrates have a specific ratio between the number of water molecules and ionic compounds. (p. 405). 

Homogeneous mixture:  another name for a solution – usually when two solids are combined.

Alloy:  A homogeneous mixture of 2 or more elements with metallic properties (ex. Brass = Cu + Zn)

Heterogeneous mixture: When you can visibly see particles in solution.  (sand/water)

Solvent:  The dissolving medium.

Solute:  The substance dissolved in the solvent.

Solubility:  The amount of a miscible substance that can be dissolved in another.  See

      table A-12 on p. 900 for a listing of solubilites of ionic compounds in water.   This

      will let you know if a single or double displacement reaction will form a precipitate or not.  If it is insoluble, that means it forms a precipitate. (Example:  Sodium Chloride is soluble – NO precipitate)

Precipitate:  An insoluble substance in solution. 

Saturated/Unsaturated:  A solution that cannot hold any more solute is saturated.

Suspensions:  If the particles in a solvent are so large that they settle out unless the mixture is constantly stirred (example:  Muddy water)

Colloids:  Particles that are intermediate in size between those in solutions and suspensions.  (Example:  The “cloudy” part of muddy water left over after the big chunks settle to the bottom.  Shaving cream, mayonnaise, smoke, fog, cheese, butter)

Electrolytes:  Substances which dissolve into water and allow it to carry an electric current (example:  Salt water conducts electricity better than “pure” water).  See example on p. 400

Solution Equilibrium:  Where dissolution and crystallization are occurring at the same rate (p. 402)

Saturated solution:  Contains the maximum amount of solute in a solvent. (p. 403)

Supersaturated solution:  Contains more than the maximum.  This is unstable and if “seeded” crystals will form (p. 403)


Read the section:  Heats of Solution (p. 409-410) aloud in class


2.      Changing the Rate of Dissolution:

          1.  Increasing the Surface area of the solute (p. 401)

          2.  Agitating the solution (p. 401)

          3.  Heating the solvent (p. 402)

          4.  How Temperature and Pressure Affect Solubility? (p. 407-409)


Temperature:  As the temperature of a liquid increases, the solubility of a solid in that liquid increases.  Example.  Hot water will dissolve more sugar than cold water.


Temperature II:  As the temperature of a liquid increases, the solubility of a gas in that liquid decreases.  Example:  Warm soda will lose its fizz but a cold soda won’t (as fast).


Pressure:  As the pressure on a liquid increases, the solubility of both solids or gases in that liquid will increase.  (pressure doesn’t affect it very much, but it does increase solubility a little bit) (p. 407)

Example:  When you open a can of soda to the reduced pressure of outside air, the bubbles escape.




Chapter 14 Notes Ions in Aqueous Solution and Colligative Properties


1.  Net Ionic Equations – (p. 429)

     A net ionic equation lists the ionic formula for substances which are insoluble in water and just the individual ions for those substances which are soluble in water (see p. 900 for a listing of what ions are soluble in water and which are not).  IMPORTANT – NOT ALL IONS ARE WATER SOLUBLE!  There are some general rules for solubility on p. 427 Table 14-1


Spectator ions:  Ions which are soluble in water stay surrounded by water and thus are not part of the reaction and are called spectators.  Spectator ions are usually dropped from the net ionic equation. See example below where spectator ions K and NO3 are dropped.

Ionic Eq: Ag+(aq) + NO3- (aq) + K+ (aq) + Cl-(aq) --à AgCl(s) + K+(aq) + NO3-(aq)

Net Ionic Eq:  Ag+(aq) + Cl-(aq) -----à AgCl(s)

Making Crystals:  Crystals are made from changes in solubility due to temperature.  You can dissolve a lot of sugar in hot water, but not in cold.  So, you supersaturate a hot water solution with sugar and then slowly cool it to bring out the crystals.


2.  Ionization (p. 431-433)

     When you put ionic compounds into water they will do one of three things:  1) nothing  2) dissolve somewhat 3) dissolve a lot!  Electrolytes are dissolved ionic compounds in water which can then carry an electric charge (remember chapter 13).  Weak electrolytes don’t dissolve much (such as HF).  But strong electrolytes (like NaCl mostly break apart completely in water).  COVALENT MOLECULES USUALLY NEVER BREAK APART LIKE IONIC ONES DO!!!!!!!! (an exception is the hydrogen halides like HCl, HBr and HI which are covalent and do break apart)

       Question:  AgCl is practically insoluble in water and yet it is considered to be a strong electrolyte.  What’s the deal?  The deal is that the small part of AgCl which actually does dissolve in water is completely dissociated into Ag+ and Cl- ions which can then carry an electric charge. (p. 432-433)


3.  Freezing point/Boiling Point Changes - AKA Colligative Properties. (p. 436-441)


Colligative properties - A property of solutions that doesn’t depend on the size or type of molecule or atom in solution, just the concentration.  For example, the Freezing point and melting point of a liquid are determined by the number of particles in solution, not the type.  1 mole of salt in 1L of water will do the same as 1 mole of sugar in 1L of water.


Ions vs. molecules in determining F.P and B.P. changes: 

     Ions separate into more than one particle in water.  For example, NaCl breaks up into Na+ and Cl- ions in water.  But C12H22O11 (sucrose) remains in its molecular form in water.  Therefore, 1 mole of NaCl makes 2 moles of solute in water and 1 mole of sucrose makes 1 mole of solute in water.  Calcium Chloride (CaCl2) turns into 3 ions and is therefore triple the effectiveness as sucrose in changing b.p and f.p.

                Ions in solution are called electrolytes because they can make a solution carry an electric charge due to the charges on the ions. (see p. 431-433)


5.  Molality: (p. 416-418)  m = molality = # of moles of solute/Kg of solvent.

     Molality is used to predict temperature changes between b.p. and f.p. of liquids.


     M = Molarity = moles of solute/L of solution (which would include L of the solute if applicable).  Molarity is used do determine the number of moles of a substance in solution.


Molality problem:  How would you prepare a 0.50 m solution of sucrose in 500.0 g of water?

Remember that the density of water at room temp is about 1g/mL or 1 g/cm3.  Therefore, if you have 1000 mL of water, you also have 1000g or 1Kg.


6.  By adding a solute to any liquid, you increase the boiling point and reduce the freezing point. 

  For example: adding salt to water will increase its boiling point and decrease its freezing point. 

              Why does it increase the boiling point?:  You lower the vapor pressure of the liquid when you add a solute because the solute prevents the pure solvent molecules from escaping.  (Recall that vapor pressure is the internal desire of a liquid to boil away).  A lower vapor pressure means you have to heat water hotter to get it to vaporize. Salt water will boil at a temperature greater than 100oC.

                Why does it lower the freezing point?:  By adding a solute to a liquid solvent, you decrease the amount of intermolecular interactions between molecules of the pure solvent.  Thus, if pure water can be frozen at 0oC, salt water will freeze at a colder temperature because the salt ions get in the way of the intermolecular attraction between water molecules.  When salt water freezes, the frozen substance contains water only, not salt water – want proof?  Icebergs floating in the ocean are pure water!  More proof?  How about this one:  Spreading salt on snow melts it because the outside air temperature is cold enough to freeze pure water, but not cold enough to freeze salt water.

                What does this have to do with making ice cream?  Ice cream is mostly frozen milk and sugar.  You can think of ice cream as mostly water with some dissolved solute in there (lactose, protein…etc).  This means that ice cream will freeze at a temperature BELOW 0oC (the normal freezing point of pure water.).  Your freezer is kept at about –10oC.  Cold enough to keep ice cream from melting.

            Pure ice (made from water) will draw heat out of the surrounding air to begin melting into liquid water.  As ice melts, its temperature remains constant until all of the ice has melted (we will learn more about this interesting phenomenon of solids keeping their temperature constant until they have completely melted in Chapter 16 – Thermodynamics)  Normally, it will draw in enough heat to reach its

melting/freezing point of 0oC.  BUT…..If you add salt to the ice, this reduces the melting/freezing point of ice to about –10oC (remember adding a solute to a solvent will reduce the freezing point).  Now, the ice will draw in enough heat to make its temperature –10oC.  This is cold enough to freeze milk into ice cream.

7.  Calculating F.P. depression or B.P. elevation: 

      DTb = kb x m x # of particles (see VERY IMPORTANT below to understand “# of particles”)

      DTf = kf x m x # of particles 

     k is a constant used for the solvent you are talking about – for example, water.  (see table 14-2 p. 438)

         Kb for water is = 0.51 oC(kg H2O)/molsolute

              Kf for water is = -1.86 oC(kg H2O)/molsolute   

    VERY IMPORTANT:  Molality of a solution needs to be doubled or tripled when you are using an ionic substance as your solute.  For example, a molality value of 1m for a solution of NaCl is actually 2m.  But a molality value of 1m for a solution of sucrose is still 1m.  NONELECTROLYTE solutes (such as insoluble ionic compounds or insoluble covalent compounds) do NOT break apart!


8.  Alloys:

     An alloy is a homogeneous mixture of two or more elements to make a substance that is metallic in nature.  For example, if you mix copper and zinc together you will get brass.  But there is no chemical bond created between the zinc and the copper and therefore brass is NOT a chemical compound (like water or salt or sugar…etc).  It is an alloy.

     Alloys generally have lower electrical conductivity than their original metals.  They also generally have a lower melting point.  Generally, alloys are stronger than their original metal and are less likely to corrode.  Pound for pound, alloys are stronger than their original metals and therefore a smaller amount of alloy can be used than an original metal.

     There are two basic types of alloys:  Substitutional alloys and Interstitial alloys:

                Substitutional alloys are where one metal atom is replaced by another (example:  brass)

                Interstitial alloys are where one metal has another atom in the spaces between its atoms

(example: steel)


Example uses of alloys:

Steel is an alloy of carbon and iron.  Uses:  Making cars.  Steel is stronger than either carbon or iron. So a thinner amount of steel can be used in the door of your car than a piece of iron.  Also, steel will not rust as easily as pure iron.


Special metal alloys are used in spacecraft because they are lighter, stronger and more heat resistant than their original elemental parts.


Brass is used in faucet fixtures because it will not corrode as easily as a zinc or copper fixture (remember brass is an alloy of zinc and copper).


Bronze is used in making statues because it is more bendable and corrodes less than its original atoms (which are Cu, Zn, Pb and Sn).


Solder is used to hold electrical contacts together because it has a lower melting point than its original atoms (which are Pb and Sn).  You can use a soldering iron to melt solder, but it won’t melt Pb or Sn.


Sterling silver (and likewise stainless steel) is used to make silverware which won’t tarnish as much as pure silver.  Sterling silver is made out of Ag and Cu.


Dentists use a special alloy called an “Amalgam” which is an alloy of Hg and Au for fillings.  Pure gold is not as strong or as pliable as it is when mixed with Hg.  Any alloy with Hg is called an amalgam.


Homework for Chapter 13 and 14:

20 points Total


Predicting Solubility of compounds:  p. 447 #2


Net Ionic Equations:  Practice Problems 1-4 p. 430; p. 448 #13, 16, 17, 18


Molality:  Practice Problems # 1-4 p. 418, p. 421 # 24, 25


Molarity (our old friend): Practice Problems p. 415 #1-3;  p. 421 # 17, 18 (skip 18b), 20


Freezing Point/Boiling Point Change Problems:  Practice Problems 1-4 p. 440; Practice Problems 1-4 p. 441; Practice Problems 1-3 p. 445; p. 449 #19, 20, 25, 26, 27, 28, 30

VERY IMPORTANT:  Molality of a solution needs to be doubled or tripled when you are using an ionic substance as your solute.  For example, a molality value of 1m for a solution of NaCl is actually 2m.  But a molality value of 1m for a solution of sucrose is still 1m.  NONELECTROLYTE solutes (such as insoluble ionic compounds or insoluble covalent compounds) do NOT break apart!