Chapter 14
Ions in Aqueous Solutions and Colligative Properties
Compounds in Aqueous Solutions
Solids compounds can be ionic or molecular
Differ in the basic particles - ions v atoms
Behavior is different when dissolved in water
Dissociation
applies to ionic compounds
dissociation - definition
NaCl(s) ---> Na+(aq) + Cl-(aq)
1 mol ---> 1 mol + 1 mol = 2 mol of ions
N.B. charge of each ion times coefficient = cancellation of charges
CaCl(s) ---> Ca+2(aq) + 2Cl1-(aq)
1 mol ---> 1 mol + 2 mol = 3 mol of ions
N.B. charge of each ion times coefficient = cancellation of charges
Number of ions produced from each formula.
100% dissociation implies 1 mole of compound breaks down totally into its ions in solution. No part of the compound remains - only ions in solution.
Sample Problem 14-1 page 426
a) Write the equation for the dissociation of aluminum sulfate,
Al2(SO4)3.
b) How many moles of each ion are produced when you start with one mole of aluminum sulfate?
c) What is the total number of ions?
Al2(SO4)3 (s) ---> 2 Al+3(aq) + 3 SO4-2(aq)
1 mol ---> 2 mol + 3 mol = 5 mol of ions
N.B. charge of each ion times coefficient = cancellation of charges
Homework: 14.1
Precipitation Reactions
No compound is completely insoluble - some small amount dissolves.
Table 14-1 page 427
Can’t write dissociation reactions for insoluble compounds.
Can use table 14-1 to predict solubility of products when two solutions are mixed -- double replacement reactions we studied earlier in the year.
If one of the compounds formed is insoluble we get a precipitate formed.
Precipitate - definition
Forms because the attraction between the ions of the precipitate is stronger than the attraction of the water molecules for each ion.
(NH4)2S(aq) + Cd(NO3)2(aq) ---> 2 NH4NO3(aq) + CdS(s)
First write the products of the double replacement reaction -- two solutions reacting -- then use table 14-1 to predict if the products individually will be soluble.
Net Ionic Equations
Used instead of formula equations for reactions that happen in solutions.
Net ionic equation - definition
Steps:
a) take the balanced chemical equation and covert it into an ionic equation;
b) soluble compounds are shown as ions, insoluble compounds are not shown as ions;
c) cancel those ions that appear on both sides of the equation, exactly;
d) rewrite the equation showing what is left -- usually two or more ions on the left and a compound on the right
Example using equation above.
Spectator Ions - definition
Sample Problem 14-2 page 430
Zinc nitrate solution and ammonium sulfide solution are mixed. Identify the precipitate after you write the formula equation, the ionic equation, and the net ionic equation and identify the spectator ions.
a) write chemical equation
Zn(NO3)2(aq) + (NH4)2S(aq) ---> ZnS(s) + NH4NO3(aq)
b) balance the equation
Zn(NO3)2(aq) + (NH4)2S(aq) ---> ZnS(s) + 2 NH4NO3(aq)
c) write the ionic equation
Zn+2(aq) + 2 NO3 1-(aq) + 2 NH41+(aq) + S2-(aq) ---> ZnS(s) +
2 NH41+(aq) + 2 NO31-(aq)
d) check to see that all charges cancel
1(+2) + 2(-1) + 2(+1) + 1(-2) = 0 for the reactants
zero + 2(+1) + 2(-1) = 0 for the products
Homework: 14.2
Ionization
The term applies to molecular compounds that form ions in solution.
Ionization - definition
ionization v dissociation
formation v separation
the extent of ionization depends on a) strength of the bonds within the molecule of the solute; b) strength of the attraction between the solute and solvent molecules
if a above is greater than b above, the solution process will not occur; if a above is weaker than b above,the solution process and the formation of ions will occur
HCl ---> H1+(aq) + Cl1-(aq)
The Hydronium Ion
H+ does not exist alone; it attracts other ions or molecules -- it is a bare proton.
a better way to represent when HCl is added to water is:
H2O(l) + HCl(g) ---> H3O1+(aq) + Cl1-(aq)
figure 14-4 page 431
H3O1+ is known as the hydronium ion -- a hydrogen ion attached to a water molecule or a hydrated hydrogen ion
Homework: 14.3
Strong and Weak Electrolyte
If a substance in solution produces ions and conduct electricity it is known as an electrolyte -- ionic compounds that are soluble or molecular compounds that ionize.
hydrogen halides -- all gases, all except HF conduct electricity well
the strength of an electrolyte is directly related to the fact that whatever dissolves in water exists completely as hydrated ions in solution -- a strong electrolyte
more ions means better electrical conductivity
Figure 14-5 page 432
Strong Electrolytes
strong electrolyte - definition
HCl, HBr, HI are 100% ionized and are strong electrolytes; all form acids in water solutions.
The acids above, several other acids, and all ionic compounds are strong electrolytes.
AgCl is considered insoluble: 0.000 089g AgCl/100 g water and yet the small amount that dissolves is complete dissociated into hydrated ions, therefore, it is considered a strong electrolyte.
Weak Electrolytes
In HF, there are dissolved ions and some molecules that are not ionized. The solution is called hydrofluoric acid.
The bond between H and F is stronger than the bond between H and other halogens, therefore, hydrogen fluoride solution has fewer ions than the same amount of the other hydrogen halogen compounds in water.
This can be represented as
HF(aq) + H2O(l) = H3O1+(aq) + F1-(aq)
where the concentration of unionized HF is high compared to the concentration of the hydronium and fluoride ions.
weak electrolyte - definition
nonelectrolyte - definition
HC2H3O2(aq) + H2O(l) ---> C2H3O21-(aq) + H3O1+(aq)
strong and weak v concentrated and dilute
Colligative Properties of Solutions
Whenever we add a solute to a solvent, the resulting solution has different properties than the pure solvent would have.
The properties of the solution depend on the number of solute particles.
Colligative Properties - definition
Vapor-Pressure Lowering
Boiling and freezing points depend on the vapor pressure of the liquid.
A solution has a
a) lower vapor pressure than the pure solvent and so
b) the boiling point of the solution is higher than the boiling point of the pure solvent and the
c) freezing point of the solution is lower than the freezing point of the pure solvent.
This assumes a nonvolatile solute is added to a pure solvent.
Nonvolatile substance - definition
Figure 14-6 page 436
Figure 14-7 page 437
Boiling and freezing points depend on vapor pressure.
The nonvolatile solute prevents the solvent from creating vapor pressure as readily and so lowers the vapor pressure and as a result raises the boiling point and lowers the freezing point.
The more solute you have the lower the vapor pressure of the solvent will be and the effect on the bp and fp will be even greater.
A 1m aqueous solution of glucose, a nonelectrolyte, lowers the vapor pressure eo water 5.5 x 10-4 atm at 25oC.
A 1 m aqueous solution of sucrose, a nonelectrolyte, lowers the vapor pressure eo water 5.5 x 10-4 atm at 25oC.
The lowering of the vapor pressure depends on the number of solute particles (concentration of the nonelectrolyte solute) which makes the lowering of vapor pressure a colligative property.
Figure 14-6 page 436
After the solute is added the solution remains liquid over a larger temperature range because of the bp increasing and the fp decreasing.
Freezing Point Depression
For water a 1m solution of a nonelectrolyte lowered the freezing point 1.86oC i.e. the freezing point of the solution is
-1.86oC.
A 2m solution of a nonelectrolyte in water lowers the freezing point by 2(-1.86) or -3.72oC.
For water the molal freezing point constant, Kf, is -1.86oC/m
Molal freezing point constant - definition
Each solvent has its own unique molal freezing point constant.
Table 14-2, page 438
Freezing point depression has the symbol delta Tf
Freezing point depression - definition
Formula: delta Tf = Kf m
Sample Problem 14-3 page 439
a) What is the freezing point depression of water in asolution of 17.1 g of sucrose,
C12H22O11, and 200. g of water?
b) What is the actual freezing point of the solution?
Homework: 14.4
Boiling Point Elevation
boiling point - definition
adding a solute to a solvent lowers the vapor pressure and thus raises the boiling point. The solution boils at a temperature higher than the pure solvent.
molal boiling point constant (Kb) - definition
for water it is 0.51oC/m
the constant is unique to each solvent
page 438
boiling point elevation (delta tb) - definition
delta tb = Kb m
Homework: 14.5
Osmotic Pressure
Figure 14-8 page 442
Semipermeable membranes - definition
Osmosis - definition
Osmotic pressure - definition
Cells and cell membranes
Homework: 14.6
Electrolytes and Colligative Properties
The colligative properties depend on the number of solute particles added to a pure solvent. We measure that number by using molality (m).
An electrolyte breaks down into more particles in solution than a nonelectrolyte.
e.g. C12H22O11 (s) --->C12H22O11 (aq)
one molecule of the solid (for covalent substances) when added to water yields one molecule of the hydrated substance.
CaCl2 ---> Ca2+ (aq) + 2 Cl1- (aq)
Calculated Values for Electrolyte Solutions
(substances which when dissolved in water form ions)
one mole of CaCl2 produces 3 moles of ions ( one mole of calcium ions and two moles of chloride ions) -- one formula produces 3 ions in solution.
Figure 14-9 page 443
Thus, we would expect a solution of calcium chloride to lower the vapor pressure, raise the boiling point, and lower the freezing point three times more than the same concentration of sucrose solution.
We need to take this into account when calculating the colligative properties of solutions.
Actual Values for Electrolyte Solutions
The values mentioned above for solutions of electrolytes are expected values. This is not always the case.
Table 14-3 page 445
Attractive forces need to be considered. These forces become more important as the concentration of the solution increases because of the increase in the number of solute particles and the closer these hydrated ions are.
They also become less important as the concentration of the solution decreases (becomes more dilute) since the number of solute particles decreases and the hydrated solute particles are farther apart.
Also, the higher the charge on the individual ions, the less the colligative properties are affected -- Debye-Huckel theory.
Compare MgSO4 and KCl in table 14-3 page 445.
Homework: 14.7
end of notes
The boiling point is the temperature at which the equilibrium vapor pressure of a liquid is equal to the prevailing atmospheric pressure
A nonelectrolyte is a substance whose water solution does not conduct electricity
A precipitate is an insoluble solid that results from mixing two solutions
Dissociation is the separation of ions that occurs when an ionic compound dissolves.back
Net ionic equation includes only those compounds and ions that undergo a chemical change in a reaction in an aqueous solution
Spectator ions are ions that do not take part in a chemical reaction and are found in solution both before and after the reaction
Ionization is the process by which ions are formed from solute molecules by the action of the solvent
The hydronium ion is a hydrated proton
A strong electrolyte is any compound o which all or almost all of the dissolved compound exists as ions in an aqueous solution
A weak electrolyte is a compound of which a relatively small amount of the dissolved compound exists as ions in an aqueous solution
Colligative properties are those properties that depend on the concentration of solute particles but not on their identify
A nonvolatile substance is a substance that has little tendency to become a gas under existing conditions
Molal freezing point constant (Kf) is the freezing point depression of the solvent ina 1 molal solution of a nonvolatile nonelectrolyte solute
The freezing point depression (delta tf) is the difference between the freezing points of the pure solvent and a solution of a nonelectrolyte in that solvent and it is directly proportional to the molal concentration of the solution
The molal boiling point constant (Kb) is the boiling point elevation of the solvent in a 1 molal solution of a nonvolatile, nonelectrolyte solute
The boiling point elevation, (delta tb) is the difference between the boiling points of the pure solvent and a nonelectrolyte solution of that solvent
Semipermeable membranes allow the movement of some particles while blocking the movement of others
Osmosis is the movement of solvent through a semipermeable membrane from the side of lower solute concentration to the side of higher solute concentration
Osmotic pressure is the external pressure that must be applied to stop osmosis