CHAPTER 4:
Types of Chemical Reactions
•Dissolution
•Precipitation
•Acids and Bases and their reactions
•Oxidation-Reduction Reactions
CHEM 1310 A/B Fall 2006
Dissolution of Ionic Compounds
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Dissolution of Molecules
CHEM 1310 A/B Fall 2006
Electrolytes
• Substances which
increase the conductivity
of water when they
dissolve
• Usually ionic compounds
(e.g., NaCl), but some
molecules can also
dissolve into ions (e.g.,
acids like HCl, HNO3).
Strong bases can also be
electrolytes (NH3 can split
H2O into ions).
CHEM 1310 A/B Fall 2006
Solubility
• Tells how much of something (the “solute”)
will dissolve into a given solvent
• Example: incomplete dissolution of sugar
in iced tea
• Observation: hot tea dissolves sugar much
easier than iced tea. This implies…
• Another observation: Coke goes flat faster
at higher temperatures. Why? Solubility
of gases change …
CHEM 1310 A/B Fall 2006
Rules about solubility
• Most general rule: “Like
dissolves like”
• Polar solvents (water,
methanol, etc) usually
dissolve polar molecules
and ionic substances;
nonpolar solutes not very
soluble
• Nonpolar solvents (CCl4,
benzene, etc) usually
dissolve nonpolar
molecules; don’t dissolve
polar molecules well
Polar ethylene glycol dissolves in water
Oil and water do not mix
CHEM 1310 A/B Fall 2006
More rules about solubility
• No gases or solids are infinitely soluble (“miscible in all
proportions”) with water; many liquids are
• Common inorganic acids soluble in water; organic acids
which are small molecules soluble in water
• Small organic compounds with –OH and –NH2 groups
usually soluble in water
• Gases are less soluble as temperature increases.
Liquids and soluds usually more soluble as temperature
increases
• Solubilities of ionic compounds varies in water; see
Table 4-1
CHEM 1310 A/B Fall 2006
Precipitation Reactions
• Opposite of dissolution --- a
solid comes out of solution
Na+ (aq) + Cl- (aq) → NaCl(s)
• Can cause precipitation by (a)
lowering temperature (e.g.,
“rock candy”), (b) evaporating
solvent, (c ) mixing 2 or more
solutions, e.g,
BaCl2 (aq) + K2SO4 (aq) →
BaSO4 (s) + 2 KCl (aq)
all of these are quite soluble in
water except BaSO4
CHEM 1310 A/B Fall 2006
Precipitation of lead iodide, PbI2,
from KI (aq) + Pb(NO3)2 (aq)
Ionic equations and
net ionic equations
• The previous example,
BaCl2 (aq) + K2SO4 (aq) → BaSO4 (s) + 2 KCl (aq),
can be written in terms of ionic equations as
Ba2+ (aq) + 2 Cl- (aq) + 2 K+ (aq) + SO42- (aq) →
BaSO4 (s) + 2 K+ (aq) + 2 Cl- (aq)
• Notice that 2 Cl- (aq) and 2 K+ (aq) appear on
both sides of the equation. They “cancel” and
we are left with
Ba2+ (aq) + SO42- (aq) → BaSO4 (s)
the net ionic equation. The Cl- and K+ ions are
spectator ions
CHEM 1310 A/B Fall 2006
Predicting precipitation
• In the previous example, how did we know that BaSO4
would precipitate out, and KCl would not?
• Need to consider all possible precipitates, and consult
solubility tables to see if they remain dissolved or if they
precipitate
• Table 4-1 in book gives rules of thumb
• Could look up “solubility constants” to do this more
quantitatively --- more on this later in the course!
• Example: what happens when an aqueous solution of
beryllium nitrate is mixes with an aqueous solution of
sodium acetate?
CHEM 1310 A/B Fall 2006
Example
CHEM 1310 A/B Fall 2006
Example, cont’d
CHEM 1310 A/B Fall 2006
Acids & Bases
• Acids
–
–
–
–
Citric acid (lemons)
Acetic acid (CH3COOH, in vinegar)
Hydrochloric (HCl), sulfuric (H2SO4), nitric (HNO3)
Taste sour, can be very reactive
• Bases
–
–
–
–
In soaps (KOH, NaOH)
Drain openers (NaClO, NaOH)
Bleach (NaClO; ClO- is hypochlorite)
Taste bitter, are corrosive, degrade organic matter
CHEM 1310 A/B Fall 2006
Chemistry of acids & bases
• All our example acids contained hydrogen
• Several of our example bases contained hydroxide
• An “Arrhenius acid” gives H+ ions when dissolved in
water
• An “Arrhenius base” gives OH- ions when dissolved in
water
• Water itself is both an Arrhenius acid and Arrhenius
base. Small amounts undergo this reaction:
H2O (l) → H+ (aq) + OH- (aq)
• Note: the H+ ions stay closely surrounded by at least a
few H2O molecules
CHEM 1310 A/B Fall 2006
Strong acids and bases
• Dissociate completely in water (and are
therefore also strong electrolytes)
• HCl (g) → H+ (aq) + Cl- (aq) [demo]
• NaOH (s) → Na+ (aq) + OH- (aq)
• Strong acids: HCl, HNO3, H2SO4, HClO3, HClO4,
HBr, HI
• Strong bases: NaOH, KOH
• Neutralization reaction: H+ from acid combines
with OH- from base, e.g.,
HCl (aq) + NaOH (aq) → H2O (l) + NaCl (aq)
What’s the “net ionic equation” for this reaction?
CHEM 1310 A/B Fall 2006
Weak acids
• Do not completely dissociate in water; weak electrolytes
• E.g., HF, acetic acid, CH3COOH, formic acid HCOOH,
phosphoric acid H3PO4, … [demo]
• Although they don’t normally completely dissociate in
water, a strong base can still “force” them to fully
dissociate, e.g.,
CH3COOH (aq) + NaOH (aq) →
H2O (l) + NaCH3COO (aq)
most of the acetic acid remains undissociated before the
reaction, so the net ionic equation is
CH3COOH (aq) + OH- (aq) →
H2O (l) + CH3COO- (aq)
• See book about naming acids
CHEM 1310 A/B Fall 2006
Weak bases
• The only strong bases are alkali metal
hydroxides (e.g., NaOH, KOH) and
Ba(OH)2
• Everything else is a weak base
(incomplete dissociation in water unless
forced by an acid)
CHEM 1310 A/B Fall 2006
Generalization of Arrhenius Model
• Ammonia (NH3) tastes bitter, feels soapy,
neutralizes acids, just like a base --- but it does
not contain OH- ions!
• Revised definition: An Arrhenius base, when
dissolved in water, increases the amount of
hydroxide ion over that present in pure solvent
• Definition of Arrhenius acid expanded similarly
for H+ ion
• NH3 (aq) + H2O (l) → NH4+ (aq) + OH- (aq)
[demo]
CHEM 1310 A/B Fall 2006
Acid & base anhydrides
• Substances which become H+ or OH- donators
when or after they react with water
– Acid anhydride: SO3 (g) + H2O (l) → H2SO4 (aq)
– Base anhydride: Na2O (s) + H2O (l) → 2 NaOH (aq)
• Usually oxides of nonmetals are acid anhydrides
• Usually oxides of metals are base anhydrides
• (More electronegative → more acidic,
more electropositive → more basic)
CHEM 1310 A/B Fall 2006
Other reactions of acids
• Besides neutralization reactions, acids…
• React with carbonates and hydrogen carbonates
to liberate gaseous CO2, producing a salt and
water, e.g., CaCO3 (s) + HCl (aq) →
CaCl2 (aq) + H2O (l) + CO2 (g) [demo]
• act with oxides of metals to form salts and water,
e.g.,
CuO (s) + H2SO4 (aq) → ???
• React with various metals to liberate H2 and
form a salt, e.g.,
Mg (s) + 2 HCl (aq) → ??
[demo]
CHEM 1310 A/B Fall 2006
Other reactions of bases
• Bases (except NH3) act on ammonium salts to
generate gaseous ammonia, a salt, and water.
E.g.,
NH4Cl (s) + NaOH (aq) →
NaCl (aq) + NH3 (g) + H2O (l)
Note: gaseous ammonia can be dangerous!
• Bases react with the oxides of non-metals to
produce salts and water, e.g.,
SO3 (g) + 2 NaOH (aq) → Na2SO4 (aq) + H2O (l)
CHEM 1310 A/B Fall 2006
Oxidation-Reduction (“Redox”)
Reactions
• Example: Cu wire in AgNO3 (aq) [demo].
Why does this Rx happen? Why does the
solution turn blue?
• Goals: to understand the twin prcesses of
oxidation and reduction; assign “oxidation
numbers”; become familiar with the basic
types of redox reactions
CHEM 1310 A/B Fall 2006
A simple redox reaction
• Recall periodic table, electronegativity,
octets
• 2 Na (s) + Cl2 (g) → 2 Na+ Cl- (s)
each Na atom
loses an e“oxidation”
(“LEO”)
each Cl atom
gains an e“reduction”
(“GERtrude”)
Na+ and Cl- both
have filled shells
like a noble gas…
happy
• Transfer of electrons. To do bookkeeping
on electrons transferred, use “oxidation
numbers”
CHEM 1310 A/B Fall 2006
Oxidation numbers
A bookkeeping device, like formal charges (but not the same)
Book has a set of 5 rules…
1. Sum of oxidation numbers must equal total charge (=0 if neutral
molecule)
2.
Group I has oxidation number of +1 always
Group II has +2 always
Group III has +3 always
3. Halogens have o.n. -1 usually (always for F) except in compounds
with O or other halogens
4. H has +1 except for metal hydrides like LiH, where it has -1
5. O has -2 except (a) in compounds with F, use rule 3; (b) in O-O
bonds, use rules 2, 4
CHEM 1310 A/B Fall 2006
Examples
• What are the oxidation numbers for H2O2,
NaO2, ClO- ?
CHEM 1310 A/B Fall 2006
Oxidation and Reduction
• Use oxidation numbers to tell if an atom is
oxidized or reduced by a reaction
• “Oxidizing agents” promote the oxidation of
other compounds. The oxidizing agent itself
must be reduced (e.g., MnO4-)
• Redox reactions can change a lot with
conditions (acidity, basicity, etc)
• Some strong oxidizing/reducing agents can
react slowly, some weak ones can react rapidly
CHEM 1310 A/B Fall 2006
Example of reaction conditions
In acidic solution…
5 Fe2+ (aq) + MnO4- (aq) + 8 H+ (aq) →
5 Fe3+ (aq) + Mn2+ (aq) + 4 H2O (l)
Which is oxidized and which is reduced?
In neutral solution…
3 Fe3+ (aq) + MnO4- (aq) + 4 H+ (aq) →
3 Fe2+ (aq) + MnO2 (s) + 2 H2O (l)
CHEM 1310 A/B Fall 2006
Types of redox reactions
• Combination reactions: 2 elements combine to
form a compound. Metals usually lose electrons
(are oxidized)
2 Fe (s) + 3 Cl2 (g) → 2 FeCl3 (s) [demo]
• Decomposition reactions: Compound breaks up
into elements or simpler compounds
2 Ag2O → 4 Ag (s) + O2 (g)
• Oxygenation: Reaction of element or compound
with oxygen (either O2 or O3, both oxidants)
4 Li (s) + O2 (g) → 2 Li2O (s)
CHEM 1310 A/B Fall 2006
Redox reaction types, cont’d
• Hydrogenation: H2 is a good reducing agent
2 Na (l) + H2 (g) → 2 NaH (s)
In organic chem, H2 often adds to double or triple bonds,
e.g.,
H2C=O + H2 → H3C-OH
• Displacement:
2 AgNO3 (aq) + Cu(s) →
Cu(NO3)2 + 2 Ag (s) [from the demo!]
Can predict ability to displace using “activity series”
(Table 4-2). Ag is a weaker reducing agent than Cu (not
as good at giving up electrons)
CHEM 1310 A/B Fall 2006
More on displacement reactions
• The famous sodium in water Rx is also a
“displacement” Rx..
2 Na (s) + 2 H2O (l) → 2 NaOH (aq) + H2 (g)
Na “displaces” hydrogen
• Halogens displace “backwards”; they are
oxidizing agents. Increasing oxidizing power:
I2 < Br2 < Cl2 < F2
Cl2 (g) + 2 KI (aq) → I2 (g) + 2 KCl (aq)
CHEM 1310 A/B Fall 2006
Redox reaction types, cont’d 2
• Disproportionation: A single ion undergoes
both oxidation and reduction
2 H2O2 (l) → 2 H2O (l) + O2 (g)
O reduced
O oxidized
CHEM 1310 A/B Fall 2006
Stoichiometry of Reactions in
Solution
• Quantitative problems like in chapter 2, but now
with the solution phase
• Usually easiest to measure volumes of solutions.
Need to know how much of a compound for a
given volume. “Concentration”
Csolute = nsolute / Volume
molarity = moles of solute / liters of solution
1.0 M = 1.0 mol L-1 = “1 molar”
CHEM 1310 A/B Fall 2006
Example of molarity
• What molarity is a solution made by
dissolving 10.0 g of Al(NO3)3 in enough
water to make 250. mL of solution?
CHEM 1310 A/B Fall 2006
Solution stoichiometry, cont’d
• Note the wording of the last example.
Don’t dissolve solute in exactly 250 mL of
solvent. Dissolve in less solvent, then add
solvent until 250 mL of solution.
• Dilution: See book. Basically, to get
solution 1/3 as concentrated, need to
increase volume by a factor of 3.
CHEM 1310 A/B Fall 2006
Stoichiometry example
• When treated with acid, lead (IV) oxide is
reduced to a lead (II) salt, liberating
oxygen
2 PbO2 (s) + 4 HNO3 (aq) →
2 Pb(NO3)2 (aq) [reduced Pb!] + 2 H2O (l)
+ O2 (g)
• What volume of a 7.91 M solution of HNO3
is needed to react with 15.9 g of PbO2?
CHEM 1310 A/B Fall 2006
Example, cont’d
• 2 PbO2 (s) + 4 HNO3 (aq) →
2 Pb(NO3)2 (aq) + 2 H2O (l) + O2 (g)
CHEM 1310 A/B Fall 2006
Titrations
• A quantitative way to
perform a reaction with
solutions by measuring
volume of one solution
added to another using a
buret
• Frequently add base from
buret to an acidic solution
until the indicator in the
flask just barely turns
color when it turns basic
• “End point” – when base
has just neutralized the
acid
CHEM 1310 A/B Fall 2006
Titrations, cont’d
• Can know the amount of base added (from
accurate concentration plus the known
volume added) and get acid concentration
(from volume of acid in solution, plus the
knowledge that at endpoint, moles of acid
= moles of base)
• Or, sometimes “standardize” base
concentration by titrating against known
amount of acid
CHEM 1310 A/B Fall 2006
Example of titration
A 100. mL solution of hydrochloric acid is just neutralized
(indicator turns just pink) when 20. mL of a 1.5 M
solution of NaOH is added. What is the concentration of
the HCl solution?
CHEM 1310 A/B Fall 2006