+
Ions in Aqueous
Solutions and
Colligative
Properties
Chemistry 1
(Chapter 13)
+
Exam Analysis

Next Exam…
REMINDER-Tests #2 and #3 will NOT replace Test #1 if…
You
have a 3 or less on your packet
the day of Test #2!!
+ Unit Objectives:
1. I can use a solubility table to predict
products in chemical reactions.
2. I can explain and correctly construct
molecular, complete ionic, and net
ionic equations (spectator ions).
3. I understand and can apply and
calculate the four colligative properties.
+ Questions From Readings
 Explain
what happens when an ionic substance is
dissolved in water. (examples to strengthen
response)
+
Dissociation reactions …
AgNO3
KBr


BaSO4
+ Questions from Reading
 What
is a precipitation reaction? How can
you determine if the reaction occurs?
+

Precipitate Reactions (ppt)
Precipitate: An insoluble solid
compound is formed during a reaction.

Anions are exchanged between two cations.


To be a ppt. rxn, both must occur:



(This is a double displacement reaction)
1. Both reactants must be aqueous (aq)
2. At least one product must be a solid (s)
You MUST include phase labels with your
equation. (See solubility table)
+ Predicting Solubility of Compounds

Use a solubility table to determine if a
substance is going to be soluble (aqueous) or
insoluble (solid) in water
a) Hg2Cl2
b) KI
c) lead (II) nitrate
+
Practice Solubility
8 Minutes!
+ Rules for ppt. reactions
1.Write a balanced chemical equation.
2.Use the solubility table to place phase
labels to each formula.
3.If one of the products is a solid and the
reactants are aqueous the reaction is
classified as a precipitate reaction.
4.If all of the products are (aq) then the
reaction is NOT a ppt rxn and is
classified as double displacement.
+Q: For each of the following decide
if a ppt. will occur.
A) Aqueous solutions of sodium chloride
and iron (II) nitrate are mixed.
B) Aqueous solutions of aluminum sulfate
and sodium hydroxide are mixed.
+
Check for Understanding
For the following reactions, predict the identity of the
precipitate formed. Write the correct formula of the
precipitate on the space. If no precipitate is likely, write
No Reaction.
 BaCl2
and K2SO4 ______________________
 CuCl2
and AgNO3 ______________________
 (NH4)3PO4
 KCl
and CaS
______________________
and Ca(NO3)2 ______________________
+ Ionic Equations
 Molecular
Equation: Chemical equation in which
the reactants and products are written as if they
were molecular substances, even though they may
exist in solutions as ions.
 Must include phase labels (s, l, g, aq)
 This provides you with the big picture
Example:
Al2(SO4)3(aq)+ 6NaOH(aq) ----- 2Al(OH)3(s)+ 3Na2SO4(aq)
+ The next 2 types of equations are only
completed for precipitate reactions!!!
 Complete
ionic equation: Shows all the
particles in the solution as they realistically
exist.
 Break apart the aqueous substances into
their ions.
 Do NOT break apart s, l, or g!!
 When writing a complete ionic equation
include the Amount, Symbol and CHARGE!
+ Example of a Complete Ionic
Equation
Molecular Equation:
Al2(SO4)3(aq)+ 6NaOH(aq)--- 2Al(OH)3(s)+3Na2SO4(aq)
Complete Ionic Equation: Amount, Symbol, Charge
+ Last step for precipitate reaction:
 Net
ionic equation: Ionic equations that
include only the particles that participate in the
reaction.
 This tells us what substances actually formed
something new in the reaction.
 Cross out the spectator ions
+ Writing Net Ionic Equation
Molecular Equation:
Al2(SO4)3(aq)+ 6NaOH(aq)--- 2Al(OH)3(s)+3Na2SO4(aq)
Complete Ionic Equation: Amount, Symbol, Charge
2Al+3 + 3SO4-2 + 6Na+1 + 6OH-1 ---- 2Al(OH)3 + 6Na+1 + 3SO4-2
Net Ionic Equation:
What formed during the reaction?
+ One more example of a precipitate
reaction
Calcium hydroxide reacts with sodium carbonate
to produce calcium carbonate and sodium
hydroxide
Molecular Equation:
Complete Ionic Equation:
Net Ionic Equation:
+ Check for Understanding

Why are complete ionic equations more informative than
molecular equations for reactions of ions in aqueous solutions?

What is the difference between a complete/total ionic equation
and a net ionic equation?

Why are spectator ions left out of the net ionic equation?

What substance is designated with an (s) in the net ionic
equation? What state designation do the other substances have?

Why is it necessary to balance the molecular equation before
writing the total and net ionic equation?
+
Practice, Practice, Practice
 Unit
Packet
+
Check for Understanding
 Explain
the difference between ionization and
dissociation.
 What
determines how much a solute ionizes in
solution?
 Explain
how to tell the difference between a strong
electrolyte and a weak electrolyte.
+
Check for Understanding
 Explain
the difference between ionization and
dissociation.

Ionization occurs when ions are formed from the solute particles due to the
action of the solvent.

Dissociation occurs when an ionic compound dissolves into ion.

Difference is ionization’s ions are formed from molecular compounds, not ionic
compounds.
 What
determines how much a solute ionizes in
solution?

The strength of the solute molecules

The strength of attraction between solute and solvent
 Explain
how to tell the difference between a strong
electrolyte and a weak electrolyte.

The degree of ionization or dissociation is what determines strength of electrolyte, not the amount of solute
dissolved.
+ Colligative
Properties
1.
2.
3.
4.
Vapor Pressure Lowering
Freezing-Point Depression
Boiling-Point Elevation
Osmotic Pressure
+
Colligative Properties
 Properties
that depend
upon the concentration of
solute particles, but not the
properties of the solute.
 Vapor
Pressure Lowering
 Boiling Point Elevation
 Freezing Point
Depression
 Osmotic Pressure
+ Vapor Pressure: the pressure exerted in a closed
container by liquid particles that have escaped the
liquid’s surface and entered the gaseous state.
 Adding
a solute to a solvent lowers the
solvent’s vapor pressure.
Vapor Pressure Lowering
+
Vapor Pressure Lowering

Fewer liquid molecules are available to escape from the
liquid to become gaseous in any solution vs. the pure solvent.

Does not depend upon the properties of the solute, only the
concentration=colligative.

This lowers the vapor pressure in solutions as compared to
pure solvents

When the vapor pressure is lower, the solution remains liquid
over a larger temperature range (solutions have lower
freezing points and higher boiling points than pure solvents).

True for nonvolatile substances-substances with little
tendency to become a gas under existing conditions.
+
Freezing Point Depression
+ Freezing-Point Depression

Molal (m) freezing-point constant (Kf) is the freezing-point
depression of the solvent in a 1-molal solution of a nonvolatile,
nonelectrolyte solution

For water, this value is -1.86°C/m

This means the freezing point of a 1m solution of any
nonelectrolyte solute in water is 1.86°C lower than the freezing
point of water (0°C).

Using this number, freezing-point depression of any solution can
be determined.

Freezing-point depression/ Δtf is the difference between the
freezing point of the pure solvent and a non-electrolyte solution
+ Freezing-Point Depression
Δtf = Kfmi
 Formula
can be used to solve for the freezingpoint of the molal concentration of a solution.
 Δtf = the change in the freezing point temperature
 Kf= the molal freezing point constant (pg. 448 in
text)
 m = the molality of the solution
 i = the number of ions present (covalent = 1, ionic
= total number of ions)
+ Solute
Ions
NaCl
HCl
MgCl2
Ca3(PO4)2
K2SO4
C12H22O11
Solute Impact on
Colligative Properties
# Particles
+
Example Problem 1
 Calculate
the freezing point depression of adding
150 g of NaCl into 250 g of water.
 Calculate
 Apply
molality
freezing-point depression equation
+
Example Problem 2
 Calculate
the freezing point depression of adding
100. g of CH3OH into 500. g of the non-electrolyte,
camphor.
 Calculate
 Use
molality
the freezing-point equation
+
Boiling Point Elevation
+ Boiling-Point Elevation

The boiling point is the temperature at which the vapor
pressure of a liquid is equal to the prevailing atmospheric
pressure.

The boiling point of a solution is higher than the boiling point
of the pure solvent (due to lower vapor pressure, less
particles available to become gaseous).

The molal boiling point constant, Kb, is the boiling-point
elevation of the solvent in 1-molal solution of a nonvolatile,
nonelectrolyte solution

In water, Kb= 0.51°C/m

The boiling point elevation, Δtb, is the difference between the
boiling point of the solution and the boiling point of the pure
solvent
+
Boiling-Point Elevation
Δtb = Kbmi

Use this equation to solve for boiling-point elevation or molal
concentration of solutions

Δtb = the change in the boiling point temperature

Kb = the molal boiling point constant (pg 448 text)

m = molality of the solution

i = the number of ions present (covalent =1, ionic = total
number of ions)
+
Example 1
 Calculate
the boiling point elevation of adding
150.g of NaCl into 250. g of water.
 Calculate
 Use
molality
boiling-point elevation equation
 New
boiling point?
+
Example 2
 Calculate
the boiling point elevation of adding 100.
g of CH3OH into 500.g of the non-electrolyte,
camphor.
 Calculate
 Use
molality
the boiling-point elevation equation
+
Osmotic Pressure

The movement of a solvent through a semipermeable
membrane from the side of lower solute concentration to the
side of higher solute concentration is osmosis.

In a U-tube, the side with higher concentration would move
up as water moves to that side

Only occurs in solution if each side contains a different
concentration

Osmotic pressure is the external pressure that must be
applied to stop osmosis

The greater the concentrations of solution (regardless of
solute), the greater the osmotic pressure of the solution.
+ Electrolytes and Colligative
Properties

Electrolytes behave differently than non-electrolytes and
result in changes in colligative properties.

The freezing point is lower and the boiling point is higher.

Due to the separation of the ions when electrolytes are dissolved
in a solution.

1 mole of NaCl, for example, has 2 particles (Na+ and Cl-), thus
results in 2 moles of solute

Nonelectrolyte molecules stay together in solution.

Some electrolytes may form more than 2 particles, thus change
the colligative properties more.
+ Electrolytes and Colligative
Properties
 Actual
values for electrolyte solutions
 The actual values are not quite those that are
calculated.
 The freezing point/boiling point are not quite
twice the expected or three times the expected.
 Due to attractive forces between the ions–
interferes with movement and affects
freezing/boiling
 Compounds with ions that have higher charges
show less change than those with lower charges
(+2, +3, -2 ….vs -1, +1)
+
Video Review

Boiling Point Elevation

Freezing Point Depression

Osmotic Pressure

Video 1

Video 2

Vapor Pressure Lowering

The Chemistry of Cars
+ Individual Practice
 Unit
Packet
Ice Cream Anyone?
Mini-Lab
20 Minutes!
+

Here's what happens. When ice cream is made the old-fashioned
way, rock salt (big chunks of salt crystals) is mixed with ice.
Only a little water melts before some of the dissolved salt lowers
its freezing point. Now when the ice wants to melt, it can absorb
lots of heat from the water to do it, and the water still will not
freeze. (In the recipe you will try, the temperature of the water
may decrease to almost minus 20 degrees C and still be liquid!
Pretty cool!) The water is very much colder than the cream
mixture. Because heat flows from hot things to cold things, the
cream now loses its heat to the water and rapidly cools down.

http://www.eduplace.com/kids/hmxs/g6/weatherwater/cricket/s
ect2cc.shtml

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