Chapter 4: Chemical Reactions &
Solution Stoichiometry
•OWL
Deadline: 15-October-12 (7:50 AM)
•Exam #2 on 15-Oct with Dr. Vining.
•Types of Reactions
•Solubility Rules
•Aqueous Reactions and Stoichiometry
•Oxidation-Reduction Reactions
Adding lead(II) nitrate to potassium iodide gives a yellow precipitate of lead(II) iodide.
Precipitation reactions are important in biological organisms— for example, in the
production of bone (calcium phosphate) and seashell (calcium carbonate).
Combination Reaction
A reaction in which two or more substances combine
to form a third substance.
For example:
2Na(s) + Cl2(g)  2NaCl(s)
Decomposition Reaction
A reaction in which a single compound
reacts to give two or more substances.
For example:
2HgO(s)  2Hg(l) + O2(g)
Let’s look at some reactions:
Sodium metal reacts with oxygen gas to produce sodium oxide.
-What kind of reaction ?
-Write the reaction (equation)
-Balance the equation
Aluminum Carbonate is heated and forms aluminum oxide and
carbon dioxide gas.
-What kind of reaction ?
-Write the reaction (equation)
-Balance the equation
Displacement Reaction
A reaction in which an element reacts
with a compound, displacing another
element from it.
For example:
Zn(s) + 2HCl(aq) 
H2(g) + ZnCl2(aq)
Acid-Base Reaction (usually
a type of Double
Displacement Reaction that
happens in solution)
An acid that reacts with a base
will usually for a salt and water.
For example:
CH3COOH (aq) + NaOH
(aq) CH3COONa (aq) +
HOH (l)
Precipitation Reaction (usually a
type of Double Displacement
Reaction in Solution)
A reaction in which an or ion reacts with
a compound or ion, displacing another
element from it. A solid precipitate
forms from the reactants in solution.
For example:
MgCl2 (aq) + 2 AgNO3 (aq) 
Mg(NO3)2 (aq) + 2 AgCl(s)
Gas Forming Reaction
(usually a type of Double
Displacement Reaction
in Solution)
A reaction in which an or
ion reacts with a compound
or ion, displacing another
element from it. A solid
precipitate forms from the
reactants in solution.
For example:
Na2CO3(aq) + 2HCl(aq) 
2NaCl(aq) + H2O(l) + CO2(g)
Combustion Reaction
A reaction in which a substance
reacts with oxygen, usually with the
rapid release of heat to produce a
flame.
For example:
4Fe(s) + 3O2(g)  2Fe2O3(s)
Let’s look at some reactions:
Calcium carbonate reacts with hydrochloric acid to produce
water, carbon dioxide and calcium chloride .
-What kind of reaction ?
-Write the reaction (equation)
-Balance the equation
A solution of barium chloride reacts with a solution of sodium
sulfate. Barium sulfate precipitates. Sodium chloride remains in
solution.
-What kind of reaction ?
-Write the reaction (equation)
-Balance the equation
What happens in aqueous solution?

Molecules and ions
become solvated.
◦ The solvent is the larger
component, usually a
liquid
◦ The solute is the minor
component, usually
starting as a solid or a
liquid (but it can be a
gas)

When water is the
solvent, then the
solute becomes
hydrated

Nonelectrolytes can dissolve in water, but do not
conduct charge
◦ They don’t form ions
◦ CH3OH (l)  CH3OH (aq) methanol dissolves in water
 Where is this reaction important for your car?

Weak electrolytes dissolve, but only partially dissociate
into their component ions. The solutions are weak
conductors
◦ CH3COOH (l) + H2O (l)  CH3COO- (aq) + H3O+ (aq)

Strong electrolytes dissolve and completely dissociate
into their component ions. The solutions are good
conductors
◦ NaCl (s)  Na+ (aq) + Cl- (aq)
Non-, Strong- or Weak-Electrolytes?
Solubility of Ionic Compounds

Soluble: to dissolve. A
solute that is soluble,
dissolves in the solvent.
◦ Say that 10 times really
fast. 
There are specific rules
for predicting the
solubility of ionic
compounds.
 When a solid forms
from a solution, it is
called a precipitate

Solubility Rules (know them)
1. Group IA and ammonium compounds are soluble.
2. Acetates and nitrates are soluble.
3. Most chlorides, bromides, and iodides are soluble.
Exceptions:
AgCl, Hg2Cl2, PbCl2;
AgBr, Hg2Br2, HgBr2, PbBr2;
AgI, Hg2I2, HgI2, PbI2
4. Most sulfates are soluble.
Exceptions:
CaSO4, SrSO4, BaSO4,
Ag2SO4, Hg2SO4, PbSO4
Reactions in Aqueous Solution

Acid-Base Reactions
◦ An acid + base produces a salt
+ water
◦ “Salt” refers to an ionic
compound.

Acids produce H+ ions
when dissolved in water.
Bases increase the amount
of OH- ions when dissolved
in water.
◦ Simplest definitions
◦ What else do we know?
 Taste
 pH
 Slippery?



Strong acids are strong
electrolytes
Weak acids are weak
electrolyotes
Acids can also be:
◦ Monoprotic (one proton
produced)
◦ Diprotic (two protons
produced)
◦ Triprotic….
Let’s look at some reactions….
Autoprotolysis of water
Nitric acid plus sodium hydroxide (SA + SB)
Oxalic acid plus sodium hydroxide (WA + SB)
Acetic acid plus water (WA + WB)
Acetic acid plus lithium hydroxide (WA + SB)
Ammonia plus water (WB + WA)
Precipitation Reactions


You MUST know solubility rules
Ordinarily double displacement reactions
3 CaCl2 (aq) + 2 Na3PO4 (aq)  Ca3(PO4)2 (s) + 6 NaCl (aq)
If the polyatomic ions are constant, you can balance them
as “groups”

Ammonium chloride + lead (II) nitrate solutions are mixed.
◦ Does a reaction occur?
◦ If so, is there a precipitate?
◦ Write and balance the chemical equation

Lithium phosphate and potassium iodide are mixed.
Net Ionic Equations (NIE)

In many reactions, there are important
compounds/ions and then there are spectators
◦ Spectator ions are present in the solution, but don’t
actually react or change what they are associated with
◦ They are sometimes an ion in a strong electrolyte

Net Ionic Equations eliminate spectators
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
Ag+(aq) + NO3-(aq) + Na+(aq) + Cl-(aq) 
AgCl(s) + Na+(aq) + NO3(aq)
Ag+(aq) + Cl-(aq)  AgCl(s)
Potassium iodide reacts with lead (II) nitrate.
Does a reaction take place?
Write the balanced equation.
Identify strong electrolytes, soluble compounds.
Identify spectator ions.
Write Net Ionic Equation.
Net Ionic Equations & Acid/Base Reactions

Acids and bases are strong or weak electrolytes.

Lithium hydroxide reacts with phosphoric
acid to produce lithium phosphate and
water.
Net Ionic Equations & Gas Forming Reactions

Even though non-electrolytes are a big part of
the reaction (the gas), there are still spectators.

Calcium carbonate reacts with
hydrochloric acid to produce carbon
dioxide, water and calcium chloride.
Oxidation Reduction (REDOX)
Reactions

Now we start looking at when electrons
move from one atom to another
◦ Charges Change
Oxidation: when electrons are lost
 Reduction: when electrons are gained

◦ Because the charge is reduced
Half-Reactions

Half-reactions are like mini net ionic
equations.
◦ Oxidation half-reaction
◦ Reduction half-reaction

They are used to identify the two
processes and balance the overall
equation.
◦ Plus, lots of times a chemist is only worried
about one half of the reaction.
Oxidation Numbers (charges)

The charge on some ion or species
◦ Cations (+) or anions (-)
◦ Now we are looking at:
 Elements becoming ions
 Ions changing their charge
2 HCl(aq) + Zn(s)  ZnCl2(aq) + H2(g)
Write the half-reactions
Identify oxidizing and reducing agents
Chromate, dichromate, sulfite, sulfate, nitrite, nitrate. What are the oxidation numbers in each?
What is the oxidation number of Cr in
dichromate, Cr2O72-?
Cr
O
2(oxidation number of Cr) + 7(-2) = -2
2(oxidation number of Cr) + (-14) = -2
2(oxidation number of Cr) = -12
Oxidation number of Cr = +6
Solution Stoichiometry……
In the titration above, the indicator changes
color to indicate when the reaction is just
complete.

Molarity (M): moles solute/liter of solution
◦ The solution volume is the FINAL volume (includes
everything.
◦ Molarity changes with temperature, which can be a
problem because solutions expand and contract.
◦ One of the more common units we will use

Molality (m): moles of solute / kg of solvent
◦ The mass of a solvent used is constant, regardless of
temperature
◦ A significant unit in geochemistry, especially when
dealing with various melted rocks (hot boiling
magma)….

How much sodium chloride do I need to make
a 1.0 liter solution that is 0.35 M NaCl?

A solution is made from 45.0 g of Calcium
Nitrate dissolved in 2.0 liters of water. What is
the concentration (M) of each ion in the
solution?

Describe how you would make a 2.0 molar
solution of phosphoric acid (500 mL).
Stoichiometry Calculations

250 mL of a 1.0 M calcium chloride
solution is mixed with a solution of
sodium phosphate (in excess). After the
reaction is complete, the resulting volume
is 1.0 liters.
◦ Write the balanced equation
◦ Identify any precipitates
◦ Calculate the amount of any precipitate that is
formed
◦ What is the resulting concentration of
calcium ion in solution?
Acid Base Titrations
The titrant is added to the buret.
 The equivalence point is when the moles
of titrant are exactly the amount needed
to react completely with the analyte
 An indicator is a substance that is used to
show the equivalence point (a dye of
some sort)

◦ The indicator changes color at the end pointthis is sometimes a little different from the
equivalence point.
In the titration above, the indicator changes color
to indicate when the reaction is just complete.

2.00 grams of oxalic acid dihydrate are titrated
to the endpoint with 25.0 mL of sodium
hydroxide. What is the molarity of the NaOH
solution?

The standardized NaOH solution above is used
to titrate 10.0 mL of sulfuric acid.
◦ What is the molarity of the sulfuric acid solution
◦ How many grams of sulfuric acid were in the
solution?