SOLUTIONS and SOLUBILITY
- Classifying Solutions
o A solution is a homogeneous mixture of 2 or more substances
 Homogeneous refers to a uniform composition throughout
o Once a solution is fully mixed the substances within remain evenly distributed
throughout
 Ex. Each drop in a cup of tea contains the same proportions of water and
dissolved substances that give the tea its colour and flavour
o The substance that is present in the greatest amount in a solution is called the solvent
 Any other substance in the solution is the solute
 Ex. In a cup of coffee, water is the solvent and the coffee beans, milk,
and sugar would be the solute
o Solutions that contain water are called aqueous solutions
 May be coloured or colourless, but they are always clear
o Types of solutions:
 A solution can be formed from a solvent and a solute in any state
 Solid dissolved in solid
 Ex. Structural steel that is used to construct buildings usually contains a
small % of carbon dissolved in iron
 Solid dissolved in liquid
 Ex. An average, 1 L of sea water contains about 19 g of chloride ions, 11
g of sodium ions, and 5 g of other solutes (Mg+2, SO4-2, Ca+2, K+)
 Solid dissolved in a gas
 Ex. Mothballs form a solution with air which is toxic for moths and other
organisms
 Liquid dissolved in solid
 Ex. Liquid mineral spirits or toluene are dissolved in solid wax to make it
easier to apply
 Liquid dissolved in liquid
 Ex. Antifreeze liquids are mixed with water in car radiators
o Solutions are described as immiscible (insoluble) [2 liquids that
form separate layers instead of dissolving] or miscible (soluble)
[liquids that mix in all proportions]
 Liquid dissolved in gas
 Ex. Humidity results from water dissolving in the air
 Gas dissolved in solid
 Ex. Most ice contains a small amount of dissolved air
 Gas dissolved in liquid
 Ex. When a bottle of a carbonated beverage is opened, some carbon
dioxide bubbles out of the solution
 Gas dissolved in gas

o
o
o
o
o
o
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Ex. Natural gas is a solution of methane gas, ethane, nitrogen, carbon
dioxide, and other gases dissolved in it
The solubility of a substance is defined in terms of the maximum amount of solute that
will dissolve in a given quantity of solvent at a specific temperature
Solubility in water is stated in terms of the mass in grams of solute that will dissolve in a
deciliter (100 mL) of water at 20 C
 Ex. NaCl (s) in water is 35.9 g/100 mL
Qualitatively:
 Solute is soluble if more than 1 g dissolves in 100 mL of solvent
 Solute is slightly (sparingly) soluble if it has a solubility between 0.1 g - 1 g per
100 mL of solvent
 Solute is insoluble if less than 0.1 g will dissolve in 100 mL of solvent
 Ex. O2 (g) in water is 0.0009 g/100 mL
o Some still dissolves in water which is needed for organisms in
oceans, lakes, and rivers
A saturated solution cannot dissolve any more solute
 You will see a very small amount of solid solute left in the solvent
An unsaturated solution could dissolve more solute
A supersaturated solution contains more dissolved solute then a saturated solution at
the same temperature
 A change in temperature can cause a saturated solution to become
supersaturated
 Ex. A solid dissolved in water becomes supersaturated when cooled
 Ex. A saturated gas in a liquid becomes supersaturated when heated
 Solutions are unstable
 Ex. If a crystal of solute is added to a supersaturated liquid solution of
solute, the excess solute will precipitate leaving the solution saturated
Factors that affect Solubility
o The formation of most solutions depends on the relative strength of 3 categories of
forces
 Forces that attract particles of the solute to each other
 Forces that attract particles of the solute to particles of the solvent
 Forces that attract particles of the solvent to each other
o
o
When a solution forms, particles of the solute are attracted to particles of the solvent
All the intermolecular forces between the particles of the solute are broken and some of
the intermolecular forces between the particles of the solvent

o
However when gases are mixed together they form readily [unless the gases
react chemically together]
 The forces of attraction between gases is very small and therefore there
influence on solubility is negligible
Intramolecular forces are attractive force within a molecule
o
Solubility in water
 The polar nature of water enables it to dissolve a wide range of solutes
 1 oxygen molecule can be bonded to as many as 6 H's at a time
o Ex.
o
Oxygen is more EN so the shared electrons in the covalent bond are displaced towards
the oxygen
 Makes them slightly negative and H slightly positive
 Creates a dipole between the O-H bond
Its strength is due to H bonding which again is a type of dipole-dipole attraction
between molecules
o
o
o
o
o
Most ionic compounds are soluble in water
 Attraction between ions of ionic compound and dipoles on water is great
enough to pull ions away
Water molecules surround each ion in a process called hydration
 Causes the ions to separate and disperse through the water, and therefore
dissolve
The stronger the attraction between ions in ionic compounds the less likely it will
dissolve
Some molecular compounds due to the dipole-dipole attraction being less than that of
the attraction between ions in ionic compounds can dissolve in water
 Again due to the polarity of the structure
o
What factors determine which ionic compounds will be soluble in water?
 Ion charge
 The greater the charge on each ion, the less soluble
o Ex. O-2 less soluble than F-1
 Ion size
 As ion size (ionic radius) increases, so does solubility
 Force of attraction between opposite charges decreases as distance
between charges increases
o
The conductivity of an aqueous solution can indicate if the compound that dissolved in
water is ionic or molecular
 Ionic compounds dissociate and its ions move freely to produce a charge due to
electrolyte
 Molecular compounds remain intact and are neutral so they have no
charge
o
Examine solubility of common Ionic compounds in water chart


o
Soluble will dissolve, i.e. aqueous solution
 The soluble compounds have solubility’s of at least 1 g/100 mL of water
at 25 C and 100 kPa
Low solubility will not dissolve completely, therefore state is solid
Overall we can use the basic concept that "Like dissolves Like"
 In other words a polar solute will more often than none dissolve in a polar
solvent
 A non-polar solute more often than none will dissolve in a non-polar
solvent

However certain, larger, molecules can have polar and non-polar properties
 Ex. CH3COOH - acetic acid


o
o
o
CH3 non-polar
COOH polar
Recall: changes in temperature can make a saturated solution become supersaturated
 This shows that the solubility of many substances changes with temperature
At higher temperature, solvent particles have greater energy resulting in more frequent
and more energetic collisions with a solute
 Therefore, solubility’s are often stated at standard ambient temperature and
pressure, SATP, which is 25 C and 100 kPa
For a given solid solute and liquid solvent, solubility usually increases with temperature
 Ex. Boiling water is better than hot water for making tea or coffee
 Therefore more mass can dissolve for the same volume of water
o
o
o
o
o
o
o
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When 2 liquids at the same temperature mix to form a solution, there is usually little or
no energy exchanges between particles
 Therefore solubility of 1 liquid in another is usually not affected by temperature
 Similar for a gas mixed with another gas
The solubility of a gas in a liquid depends on both temperature and pressure
 Due to a large change in the kinetic energy (energy of motion) of gas molecules
when they enter or leave a solution
Molecules in gaseous state have greater K.E. than the same molecules dissolved in a
solvent
Increased temperature of solution allows dissolved gas molecules to escape
 Therefore solubility of most gases in most liquid solvents decreases with an
increase in temperature
 Ex. This is why carbonated drinks (coke, pepsi) go flat
Pressure is force per unit area and has very little effect on solubility of liquids and solids
but a significant effect on the solubility of gas
The solubility of a gas in a liquid is directly proportional to the pressure of that particular
gas above the liquid
 Not affected by pressure of any other gas
Both temperature and pressure can affect the rate at which a solute will dissolve in a
solvent
 Also mixing the solution increases the number of collisions and increase the
surface area, which then increases the number of collisions
Concentrations of Solutions
o Concentration is the quantity of solute per unit of solution or solvent
o Qualitatively a solution can be described as concentrated (high ratio of solute to
solution) or dilute (low ratio of solute to solution)
o Percent concentration:
 These ratios refer to the solution as a whole
 Mass/volume [m/v] percent
 % m/v = mass of solute (in grams) x 100
volume of solution (in mL)
o
Ex. An intravenous solution for a patient was prepared by
dissolving 17.5 g of glucose in distilled water to make 350 mL of
solution. Find the % m/v concentration of the solution.

Mass %
 Solution that contains a solid state dissolved in a liquid solvent
 % m/m = mass of solute x 100
mass of solution
 Sometimes referred to as weight/weight (w/w) which is equivalent to
m/m
o Ex. A cooking pan with 18/8 stainless steel contains 18 % m/m
chromium and 8 % m/m nickel in iron
o Ex. Find the mass of pure silver in a sterling silver ring that has a
mass of 6.45 g and concentration is 92.5 % m/m.

% Volume
 When 2 liquids are mixed to form a solution
 % v/v = volume of solute x 100
volume of solution
 Note the volume of solution produced by dissolving 1 liquid in another
is usually not equal to the sum of the volumes of the 2 separate liquids
o Ex. Acetic acid is a liquid at room temperature. How much pure
water should be added to 15.0 mL of pure acetic acid to make a
5.00 % v/v solution of acetic acid?

Very small concentrations
 Very dilute solutions have concentrations that are much less than 1 %
m/m
o Expressed in parts per million, ppm (a ratio of solute to solution
multiplies by 106) or parts per billion, ppb (a ratio of solute to
solution multiplied by 109)
 ppm = mass of solute x 106
mass of solution
 ppb = mass of solute x 109
mass of solution
o Ex. Health Canada's guideline for the maximum mercury
content in commercial fish is 0.5 ppm. When a 1.6 kg salmon
was tested, it was found to contain 0.6 mg of mercury. Would
this salmon be safe to eat?
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Molar concentrations
o Molarity is the amount, in moles, of solute dissolved in 1 L of solution
 molar concentration = amount of solute
volume of solution
 c = n/V
o where n is the amount of solute in moles
o V is the volume of solution in litres
o c is molarity in mol/L or M
o
When calculating molarity of ions make sure to balance the equation showing the
dissociation of the compound into its ions
 Ex. CaCl2 (aq) → Ca+2 (aq) + 2Cl- (aq)
 Ex. A student dissolved 0.212 mol of iron (III) chloride to make 175 mL solution.
Find the molar concentration of the solution and the concentrations of the ions
in the solution.

-
Ex. The most commonly used intravenous solution, normal saline, contains
0.154 mol/L of NaCl (aq). Calculate the mass of Na in a 500 mL bag of normal
saline solution
Preparing Solutions
o A solution with an accurate, known concentration is called a standard solution
o 2 ways to prepare:
 Dissolve a measured mass of pure solute in water and then dilute the solution
to a known volume
 Dilute a standard solution by adding a known volume of water
o
A stock solution can be diluted to prepare a standard aqueous solution
 Adding more solvent does not remove or add any particles of solute
 Amount of solute (moles) is the same before and after dilution
o Therefore, C1V1= C2V2
 C1V1 is the molarity and volume of the concentrated
solution
 C2V2 is the molarity and volume of the diluted solution
 Ex. A stock solution of 12 M HCl is used in an experiment that requires
2.0 L of 0.10 mol/L HCl solution. What volume of concentrated solution
should be used to make the dilute solution for the experiment?
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Net Ionic Equations
o A chemical reaction between 2 aqueous solutions that contains ions is always a double
displacement reaction
o Water dissociates ionic substances into their component ions, allowing reactant ions to
mix and react more readily
 However water is not a reactant and neither are some of the ions in solution
 Non-reacting ions in an aqueous solution are called spectator ions
o Spectator ions usually form soluble compounds
o An ionic equation replaces the formulas of soluble ionic compounds with the ions that
these compounds form in water
 Ex. AgNO3 (aq) → Ag+ (aq) + NO3- (aq)
 Note the total charge on the left and right side of the ionic equation
equal
o Cancelling the spectator ions leaves the net ionic leaves just the net ionic equation
 Shows only the ions that react and the insoluble product or precipitate
o Rules:
 Write the complete chemical equation for the reaction
 Rewrite the soluble ionic compounds (aqueous) as ions
 Leave insoluble ionic compounds (solids) alone
 Leave molecular compounds alone since they produce relatively few ions in
aqueous solution
 Write all acids as compounds, except the following:
 HCl, HBr, HI, HNO3, and HClO4
o Because they are stronga cids that completely ionize in water
 Cancel spectator ions
 Keep only covalent compounds, the ions that react, and the precipitate
 Any gases present must appear in the net ionic equation
 Check that both charges and atoms are balanced
 Ex. What substance will precipitate when an aqueous solution of sodium
sulfide is mixed with an aqueous solution of silver nitrate?

o
Ex. Write the net ionic equation for aqueous hydrobromic acid and
aqueous potassium hydroxide
A qualitative analysis identifies substances in a sample, for example:
 Flame test
 Many metal ions produce a distinctive colour when they are heated



-
A small sample of a solid or a drop of solution is used
o Ex. Na+ is yellow-orange, Cs+ is blue, Cu+3 is bluish-green
A wire loop made of platinum or an alloy (mixture of metals) of nickel
and chromium is dipped in hydrochloric acid or nitric acid and then used
to collect a sample to place in the flame
o The electrons in the atoms of the sample absorb energy from
the flame
 The electrons then re-emit some of the energy as visible
light
 i.e. Recall Bohr - excited and ground state
electrons
 Note: The wavelength of light emitted by different
elements is so unique that astronomers analyze light
from distant stars to determine the elements present
Aqueous solutions of ionic compounds of certain cation and anions also
have characteristic colours
o Ex. Cu+2 - blue, Cu+1 - green, Fe+3 - yellow, Fe+2 - green

Precipitates
 Another way to identify an unknown ion in a solution is to add a known
reactant to the solution and observe whether a precipitate forms
 The solubility guidelines can be used to infer which ion must be present
 Each time a precipitate forms ions are removed from the solution
o Can filter out the precipitate and continue to test the liquid
filtrate remaining
o Flame tests can also be used on the precipitate

The colour of a solution
Solution Stoichiometry
o Ex. A student carefully measured 100 mL of a silver nitrate aqueous solution of unknown
concentration. The solution was poured into a beaker and a coil of copper was added.
The solution was left for a few days. Later the student scraped the silver that attached
to the copper and filtered the solution. The dry precipitate had a mass of 1.65 g. What
was the molarity of AgNO3 (aq)?
-
o
Ex. When excess aqueous lead (II) nitrate was added to 125 mL aqueous solution of
sodium iodid, a bright yellow precipitate of lead (II) iodide formed. The dry precipitate
has a mass of 4.13 g. What was the concentration of the iodide ions in the solution of
NaI (aq)?
o
Ex. In one process for water purification, aluminum sulfate (aq), reacts with calcium
hydroxide (aq). Find the mass of the precipitate when 20.0 mL of 0.0150 mol/L
aluminum sulfate is mixed with 30.0 mL of 0.0185 mol/L calcium hydroxide.
o
Ex. A chemist wants to react 60.0 mL of 0.135 mol/L sodium oxalate, Na2C2O4 (aq), with
0.226 M calcium chloride (aq), to precipitate calcium oxalate. What is the minimum
volume of calcium chloride solution required? What mass of calcium oxalate will be
precipitated?
Acids and Bases
o Arrhenius principle:
 Assumes that all acids form H3O+ (hydronium) and all bases form OH(hydroxide) in water
o Brønsted -Lowry Theory
 Defines acids and bases according to the donation and acceptance of a proton
 Acids and bases don’t need to be in aqueous solutions and their activity in water
is not part of the definition [unlike Arrhenius]
 A Brønsted-Lowry acid is a proton donor or any substance that donates H ions,
ex. HCl, HNO3
A Brønsted-Lowry base is a proton acceptor or any substance that accepts a H

ion, ex. NaOH, Mg(OH)2
 A base must have a lone pair of electrons to bind with H ion, ex. NH3
o
o
A substance behaves as an acid only if another substance behaves as a base (and vice
versa)
Gives a conjugate base of an acid (particle that remains when a proton is removed from
acid) and a conjugate acid of a base (particle formed when the base receives the proton
from the acid)

Ex.
o
A molecule or ion that can accept or donate a proton can act as an acid or a base; is said
to be amphiprotic, ex. H2O, HSO4o Acids are molecular compounds that are held together by covalent bonds
 When dissolved in water they form ions and therefore undergo ionization
o Bases are ionic compounds and when in water the ions in the base break apart in a
process called dissociation
o Properties of acids and bases:
Property
Acids
Bases
Taste
Sour
Bitter
Texture
None
Slippery
Conductivity in aqueous
Conducts electricity
Conducts electricity
solution
pH
Less than 7
Greater than 7
Corrosion
Corrode tissue and metals
Corrode tissue
Reaction with metals
Produce hydrogen gas
No reaction
Reaction with carbonate
Produce carbon dioxide
No reaction
Example
Citrus fruit, vinegar, vit C,
Soap, baking soda, oven
carbonated drinks
cleaner, household
ammonia
o
The pH scale is used to describe the acidity or basicity of a solution based on the
concentration of H+ ions in solution
 i.e. Acids have a high concentration of H+ ions and bases have a low
concentration of H+ ions and a high concentration of OH- ions
 pH = - log [H+]
pOH = - log [OH-]
o pH + pOH = 14
 A neutral solution has an equal balance of H ions and OH- ions, and therefore
has a pH equal to 7
o
A pH meter or indicators are used to help measure and identify if a solution is acidic or
basic
Types of indicators:
 Litmus paper
o


 blue litmus turns red in acid; no colour change means solution is basic
 red litmus turns blue in base; no colour change means solution is acidic
Universal indicator
 Is a mixture of chemicals that turn a specific colour to give an
approximate value for pH and determine if acidic or basic
Chemical indicators
 Placed in a solution and if colour changes pH is within a specific range
o Ex. Methyl violet pH range: 0-1.6, colour changes: yellow-blue;
bromocresol green pH: 3.8-5.4, colour: yellow-blue;
phenolphthalein pH: 8.2-10, colour: colourless-pink
o
Strong acids
 Completely ionize in water
 Ex. Hydrochloric acid, HCl; hydrobromic acid, HBr; hydroiodic acid, HI;
perchloric acid, HClO4; nitric acid, HNO3; sulfuric acid, H2SO4
 The stronger the EN the harder it is to ionize
 In oxoacids the H atom that ionizes is always attached to the O atom
 Acids with 2 H's are diprotic, 3 H's triprotic; ex. H2SO4
o
Weak acids
 Does not ionize completely in water and most acids are weak
 The number of molecule that ionize depends on concentration and temperature
of solution
 Ex. Acetic acid, CH3COOH; hydrocyanic acid, HCN; hydrofluoric acid, HF;
phosphoric acid, H3PO4
o
Strong bases
 Dissociate completely in water
 All group 1 and 2 metals with OH, except Be (due to size of atom and its bond
with O is strong)
 Ex. lithium hydroxide, LiOH; sodium hydroxide, NaOH; potassium
hydroxide, KOH; calcium hydroxide, Ca(OH)2; barium hydroxide, Ba(OH)2
o
Weak Bases
 Produces relatively few OH- in water, most bases are weak; ex. NH3
o
In a concentrated solution of strong acid or base there are many H+ or OH- and acid
anions or base cations
 A concentrated solution of weak acid or base has many acid or base molecules
and some H+ or OH-, as well as, acid anions or base cations
In a dilute solution of a strong acid or base there are few H+ or OH-
o

-
Even less for weak acids and bases
Neutralization Reactions and Titration
o A neutralization reaction is between acids and bases
o The acid counter acts the properties of the base and the base counteracts the properties
of the acid
o The H+ from acid combines with OH- from the base to form water
o The metal cation from the base and the anion from the acid combine to form an ionic
compound called salt
 Ex. HCl (aq) + KOH (aq) → H2O (l) + KCl (aq)
o Not all reactions result in a pH = 7, i.e. neutralization, however different mixtures of
acids and bases still neutralize at different pH values
o To determine the concentration of an acid or base we use a titration
 In a titration the concentration of a solution is determined by reacting a known
volume of that solution with a measured volume of a solution with a known
concentration
o The solution with the known concentration is the titrant and it is placed into a burette
o Indicators are used to identify the neutralization by reaching an end point when
changing colour
 Ex. Phenolphthalein is colourless between pH 0-8, and turns pink between pH 810
o The aim of a titration is to know when the amount of titrant that has been added to the
sample is just enough to react with all the acid or base
 This point is referred to as the equivalence point
 Very slight colour change
o Ex. What volume of 1.25 mol/L hydrobromic acid will neutralize
75.0 mL of 0.895 mol/L magnesium hydroxide?

Ex. If 15.0 mL of phosphoric acid completely neutralize 38.5 mL of 0.150
M of NaOH (aq), what is the concentration of the phosphoric acid?