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6. How do the properties of a solution differ from that of pure solvent?

Colligative properties - properties that depend on the number of dissolved particles in a given mass of solvent

does not depend on the chemical properties

1.0M NaCl and 1.0M KBr have the same colligative properties because there solutions contain the same number of ions.

a. Vapor pressure lowering - solute added to a solvent makes the v.p. of a solution less than solvent

decrease in vapor pressure is directly proportional to the number of solute particles in the solution

b. boiling point elevation - because vapor pressure is lowered, must go to higher temperature to get solution boiling.

Covalent compounds:

1M ethanol(aq) boils at 100.7oC

Ionic compounds:

dissolve with the formation of ions, number of ions is important

1M NaCl(aq) boils at 101.4oC  (NaCl forms Na+ and Cl- when it dissolves in solution so the affect is twice as much as with the molecule ethanol)

1M KBr(aq) boils at 101.4oC (colligative properties do not depend on what the material is, just the number of particals)

c. freezing point depression - the addition of solute lowers the freezing point of water. 

1M NaCl(aq) freezes at -3.72oC

Demo: string on ice cube

d. Osmotic pressure

Water seeks solute

consider two solutions of different solute concentrations separated by a osmotic membrane.

osmotic membranes allow only water molecules to pass through (no solute passes through)

 As if a pressure was pushing down on the side with the pure water.  The amount of pressure on the left tube it would require to bring the two water levels equal is the osmotic pressure.

e. dialysis

solute seeks water

dialyzing membranes - semipermeable membranes that permit the passage of small molecules and ions, large molecules are retained.


Hemodialysis is where a dialyzing membrane Is used to remove toxic waste products from the blood.

Acids and Bases

Certain types of compounds share similar chemical and physical properties.  Acids and bases are two types of such classes of compounds:


taste sour (Latin "acidus" meaning tart)

react with metals to give off H2 gas 

cause litmus indicator dye to turn red


taste bitter

feel slippery

cause litmus to turn blue

Molecular definition of acids and bases:

Acids and bases react in characteristic ways with water.  Because so much of chemistry (in particular the chemistry of life) is carried out in water, the presence of acids or bases can strongly influence chemical reactivity.

1. acids - any species that dissolves in water to produce hydronium ions: H3O+ (Arrhenius definition)

Ex: HCl

HCl + H2O  --> H3O+ + Cl-

For every mole of HCl we put in water, it forms one mole of H3O+.  The reaction goes to completion (all the way to the right).  Acids for which this is true are termed strong acids.

For simplicity, this reaction can be written as:

HCl(aq) ----> H+(aq) + Cl-(aq)

Need to remember, however, that this is a simplified form and that the equation above showing the reaction with water is closer the the truth.  From the simplified reaction, it is clearer these reactions are termed dissociation reactions.  HCl "falls apart" or dissociates to form H+ and Cl- just like the dissolution of NaCl

Demo: How do we know?  conducts electricity

Weak acids only partially dissociate in water

Ex: acetic acid CH3COOH


H+ +



equilibrium <----


Demo: how do we know? doesn't conduct electricity as well.

As we have learned, equilibria are described by equilibrium constants. 

The equilibrium constant for the dissociation of an acid is termed Ka = the acid dissociation constant (note that in this special dissociation constant the water molecule is not included.


[H3O+][ CH3COO-]

Ka =



Remember all concentrations are in molarity!

Ka very large (>1) - near complete dissociation - strong acid. 

Ka small (<1) - little dissociation - weak acid





1 x 107



1 x 10-5


2. Bases - Any species that dissolve in water to produce OH- ions (Arrehnius definition)


in water



Na+ + OH-

For every mole of NaOH we put in water, one mole of OH- forms. Bases for which this is true are termed strong bases.  Strong bases completely dissociate.

Weak bases only partially dissociate - ammonia (NH3) for instance is a weak base

NH3 + H2O

NH4+ +



equilibrium <----


3. Water - can act as both an acid and base

Can react with itself to form both H3Oand OH-

H2O + H2O -> H3O+ + OH-

Kw = [H3O+][OH-] = 1 x 10-14  ( ion product constant of water)

This reaction, termed the autoionization of water, maintains the product of the H3O+ and OH- concentrations constant.  It provides the link between [H3O+] and [OH-] in aqueous solutions.  Cannot affect one without affecting the other.  Add an acid, the H3O+ concentration goes up so the OH- concentration must go down.

If we do not add any additional acid or base, the [H3O+] = [OH-], so in the so-called "neutral" water:

[H3O+] = [OH-] = (1 x 10-14)1/2 = 1 x 10-7 M


Acid + base --> water + salt

NaOH + HCl --> H2O  + NaCl (neutral, not acidic or basic)

The acid and base in effect cancel each others influence.  No production of OH- or H+ just H2O

Bronsted-Lowry definition-

Does not have water in the definition so more general.

Acid - proton donor

Base - proton acceptor

Using this definition, we can look at a reaction such as:

NH3 + HCl --> NH4+ + Cl-

and see that NH3 accepted a proton so it is the base and HCl donated a proton so it is the acid.