Reactions in Aqueous
Solutions
Chapter 7
Aqueous Solutions
Water is the dissolving
medium, or solvent.
Driving Forces for Chemical
Reactions
1. Formation of a precipitate (solid).
2. Formation of molecular substance
(water).
3. Formation of a gas.
4. Transfer of electrons (Redox).
Types of Solution Reactions
-
Precipitation reactions
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
-
Acid-base reactions
NaOH(aq) + HCl(aq)  NaCl(aq) + H2O(l)
Electrolytes
Strong - conduct current efficiently - many ions in solution.
NaCl, KNO3, HNO3, NaOH
Weak - conduct only a small current - few ions in solution,
HC2H3O2, aq. NH3, tap H2O
Non - no current flows - no ions in solution.
pure H2O, sugar solution, glycerol
In what two ways can a solid ionic
compound be made to conduct electricity?
•Dissolve it in water.
•Melt or fuse it.
Ions must be free to move
(mobile) in order to conduct
electricity!
Dissociation
ionic compounds
metal + nonmetal (Type I & II)
metal + polyatomic anion
Ammonium compounds
Acids
When ionic compounds dissolve in water the anions and
cations are separated from each other; this is called
dissociation
We know that ionic compounds dissociate when they
dissolve in water because the solution conducts
electricity
Dissociation
potassium chloride dissociates in water into
potassium cations and chloride anions
KCl(aq) = K+ (aq) + Cl- (aq)
K
Cl
K+
Cl-
copper(II) sulfate dissociates in water into copper(II)
cations and sulfate anions
CuSO4(aq) = Cu+2(aq) + SO42-(aq)
Cu SO4
Cu+2
SO42-
Dissociation
potassium sulfate dissociates in water into
potassium cations and sulfate anions
K2SO4(aq) = 2 K+ (aq) + SO42-(aq)
K
SO4 K
K+
SO42K+
Ionic Compounds in Solution
In aqueous solution, soluble ionic compounds
exist in the form of ions.
K2CrO4(aq) + Ba(NO3)2(aq) ----> BaCrO4(s) + 2KNO3(aq)
Figure 7.1: The
precipitation reaction
that occurs when
yellow potassium
chormate, K2CrO4 (aq),
is mixed with a
colorless barium nitrate
solution, Ba(NO3)2 (aq)
Solubility
1. A soluble solid readily dissolves in water-designated with (aq).
2. A slightly soluble solid only dissolves to a
tiny extent in water--designated with (s).
3. An insoluble solid does not dissolve to any
appreciable extent in water--designated
with (s).
Simple Rules for Solubility
1. Most nitrate (NO3), acetate (C2H3O2-), & chlorate
(ClO3-) salts are soluble.
2. Most alkali (group 1A) salts and NH4+ are soluble.
3. Most Cl, Br, and I salts are soluble (NOT Ag+,
Pb2+, Hg22+)
4. Most sulfate salts are soluble (NOT BaSO4, PbSO4,
HgSO4, CaSO4)
5. Most OH salts are only slightly soluble (NaOH,
KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally
soluble)
6. Most S2, CO32, CrO42, PO43 salts are only
slightly soluble.
Precipitation Reactions
In all precipitation reactions, the ions of one substance
are exchanged with the ions of another substance when
their aqueous solutions are mixed
At least one of the products formed is insoluble in water
KI(aq) + AgNO3(aq)  KNO3(aq) + AgIs
K+
Ag+
K+
Ag
I-
NO3-
NO3-
I
Describing Reactions in
Solution
1. Molecular equation (reactants and
products as compounds)
AgNO3(aq) + NaCl(aq)  AgCl(s) + NaNO3(aq)
2. Complete ionic equation (all strong
electrolytes shown as ions)
Ag+(aq) + NO3(aq) + Na+(aq) + Cl(aq) 
AgCl(s) + Na+(aq) + NO3(aq)
Describing Reactions in
Solution (continued)
3. Net ionic equation (show only
components that actually react)
Ag+(aq) + Cl(aq)  AgCl(s)
Na+ and NO3 are spectator ions.
Solutions of:
AgNO3
Na2SO4
If AgNO3 is mixed with Na2SO4 what ions are most
abundant in the solution?
AgNO3
Na2SO4
With what ions is the solution saturated?
2AgNO3(aq) + Molecular
Na2SO4(aq)
 2 NaNO3 (aq) +
Equation
Ag2SO4(s)
2Ag+(aq) + 2 NO3-(aq) + Overall
2Na+(aq) +Ionic
SO42-(aq)
 2Na+(aq) + 2 NO3- (aq) + Ag2SO4(s)
Equation
+
2AgNet
SO42-(aq)
 Ag2SO4(s)
Equation
(aq) +Ionic
Silver Sulfate
Precipitate
Acids
The nature of acids was discovered by Svante
Arrhenius.
Acids are characterized by:
• a sour taste (lemons -- citric acid).
• producing H+ ions (protons) in aqueous
solution.
• turn blue litmus red.
Acids
Strong acids - dissociate completely (nearly
100 %) to produce H+ in solution
HCl, H2SO4, HNO3, HBr, HI, & HClO4
Weak acids - dissociate to a slight extent
(approximately 1 %) to give H+ in solution
HC2H3O2, HCOOH, HNO2, & H2SO3
Figure 7.5: When gaseous HCl is dissolved in
water, each molecule dissociates to produce H+
and Cl- ions
Bases
Bases are characterized by:
• bitter taste (soap).
• feel slippery.
• produce hydroxide (OH-) ions in aqueous
solution.
• turn red litmus blue.
A basic solution is said to be alkaline since
they often contain one of the alkali metals
--Na, K, Li, etc.
Bases
Strong bases - react completely with water to
give OH ions.
sodium hydroxide
Weak bases - react only slightly with water to
give OH ions.
ammonia
Acid-Base Reactions
Acid + Base ----> Salt + water
Molecular Equation
HCl(aq) + NaOH(aq) ----> NaCl(aq) + HOH(l)
Complete Ionic Equation
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) -----> Na+(aq) + Cl-(aq)
+ HOH(l)
Net Ionic Equation
H+(aq) + OH-(aq) -----> HOH(l)
Acid-Base Neutralization
The net ionic equation for the
neutralization of any strong acid and
base will always be:
H+(aq) + OH-(aq) -----> HOH(l)
Salts
Salts are ionic compounds consisting of:
a. metal & nonmetal -- KCl
b. metal & polyatomic ion -- CuSO4
c. polyatomic ion & nonmetal -- NH4Cl
d. two polyatomic ions -- (NH4)2SO4
Oxidation-Reduction Reactions
(Redox Reactions)
Redox reaction -- involves the transfer of
electrons
• loss and gain of electrons must be
exactly equal.
• loss and gain of electrons must be
simultaneous.
Oxidation
•
•
•
•
•
•
•
loss of electrons
metal atoms -- Na, Ca, & K
Nao ---> Na1+ + 1eCao ---> Ca2+ + 2enonmetal ions -- Cl-, S2-, & O2Cl1- ---> Clo + 1eS2- ---> So + 2 e-
Reduction
•
•
•
•
•
•
•
gain of electrons.
nonmetal atoms.
Oo + 2e- ---> O2Fo + 1e- ---> F1metal ions.
K1+ + 1e- ---> Ko
Ba2+ + 2e- ---> Bao
Figure 7.7: When powdered aluminum and iodine (shown in the
foreground) are mixed (and a little water added), they react
vigorously
Oxidation & Reduction HalfReactions
•
•
•
•
Always add electrons (negative) to the
more positive side of the equation.
The charge on both sides of an equation
must be equal.
Oxidation -- Nao ----> Na1+ + 1 eReduction -- Cl2 + 2 e- ----> 2 Cl-
OIL RIG
Oxidation Is Loss.
Reduction Is Gain.
Redox Reactions
•
•
•
•
Metal-nonmetal reactions are always
redox reactions.
Any reaction that has a free element (such
as O2) as a reactant or product are redox.
All single replacement reactions are redox.
All combustion reactions are redox.
Synthesis I (Composition)
A + X ----> AX ALWAYS REDOX
Element + Element -----> Binary Compound
Fe(s) + S(s) ----> FeS(s)
4 Al(s) + 3 O2(g) ----> 2 Al2O3(s)
Synthesis II (Composition)
A + X ----> AX NOT REDOX
Compound + Compound ----> Compound (3
or more elements)
Ammonia + Acid ----> Ammonium Salt
NH3(g) + HCl(g) ----> NH4Cl(s)
2 NH3(aq) + H2SO4(aq) ----> (NH4)2SO4(aq)
Synthesis II (Composition)
Continued
Water + An Oxide
Rule # 1 --Water + Metal Oxide ----> Metal
Hydroxide (Base)
HOH(l) + CaO(s) ----> Ca(OH)2(s)
HOH(l) + Na2O(s) ----> 2 NaOH(aq)
Rule # 2 --Water + Nonmetal Oxide ----> Acid
HOH(l) + SO3(g) ----> H2SO4(aq)
HOH(l) + N2O5(g) ----> 2 HNO3(aq)
Single Replacement I
A + BX ----> AX + B ALWAYS REDOX
Element + Compound ----> Different Element
+ Different Compound
Metal Reactivities:
K > Na > Ca > Mg > Al > Zn > Cr > Fe > Cd
> Co > Ni > Sn > Pb > H > Sb > Cu > Hg
> Ag > Pt > Au
Single Replacement I
(Continued)
A is a metal.
A is more reactive than B
A + BX ----> AX + B
Cu(s) + 2 AgNO3(aq) ---->Cu(NO3)2(aq) + 2 Ag(s)
Zn(s) + 2 HCl(aq) ----> ZnCl2(aq) + H2(g)
Rule # 3 -- Active Metal + Water ----> Metal
Hydroxide + Hydrogen
Ca(s) + 2 HOH(l) ----> Ca(OH)2(s) + H2(g)
Single Replacement I
(Continued)
A is a metal.
A is less reactive than B
A + BX ----> No Reaction (NR)
Cu(s) + HCl(aq) ----> NR
Single Replacement II
Y + BX ----> BY + X ALWAYS REDOX
Element + Compound ----> Different Element
+ Different Compound
Nonmetal Reactivities:
F > O > Cl > Br > I
Single Replacement II
(Continued)
Y is a nonmetal.
Y is more reactive than X
Y + BX ----> BY + X
Cl2(g) + 2 KBr(aq) ----> 2 KCl(aq) + Br2(aq)
Y is less reactive than X
Y + BX ----> No Reaction
Cl2(g) + KF(aq) ----> NR
Single Replacement
(Continued)
Remember!!
Metals replace metals.
Mg(s) + 2 Ag(NO3)2(aq) ----> 2 Ag(s) + Mg(NO3)2(aq)
Nonmetals replace nonmetals.
Cl2(g) + 2 KBr(aq) ----> Br2(aq) + 2 KCl(aq)
Figure 7.6: The
thermite reaction
gives off so
much heat that
the iron formed
is molten
Decomposition I
AX ----> A + X ALWAYS REDOX
AX is a binary compound.
Binary Compound ----> Element + Element
elect
2 NaCl(l) ----> 2 Na(l) + Cl2(g)
elect
2 HOH(l) ----> 2 H2(g) + O2(g)
Decomposition II
AX -----> A + X MAY OR MAY NOT BE
REDOX
AX is a ternary compound.
Rule # 4 -- Metal Chlorates ----> Metal Chlorides
+ Oxygen

2 KClO3(s) ----> 2 KCl (s) + 3 O2(g) REDOX
Decomposition II
(Continued)
Rule # 5 -- Metal Carbonates ----> Metal Oxides
+ Carbon Dioxide

CuCO3(s) ----> CuO (s) + CO2(g) NOT REDOX
Decomposition II
(Continued)
Rule # 6 -- Metal Hydroxides ----> Metal Oxides
+ Water

Ca(OH)2(s) ----> CaO (s) + HOH(g) NOT REDOX
Decomposition II
(Continued)
Rule # 7 -- Acid ----> Nonmetal Oxide + Water

H2SO4(aq) ----> SO3(g) + HOH(g) NOT REDOX
Double Displacement
(Double Replacement)
AX + BY ----> AY + BX NEVER REDOX
Compound + Compound ----> 2 New
Compounds
Acid-Base & Precipitation Reactions (Chapter 7)
H2SO4(aq) + 2 NaOH(aq) ----> Na2SO4(aq) + 2 HOH(l)
NaCl(aq) + AgNO3(aq) ----> NaNO3(aq) + AgCl(s)
Combustion
AX + O2 ALWAYS REDOX
AX is a hydrocarbon.

CxHy + O2(g) ----> CO2(g) + HOH(g)

CH4(g) + 2 O2(g) ----> CO2(g) + 2 HOH(g)

2 C6H6(l) + 15 O2(g) ----> 12 CO2(g) + 6
HOH(g)