Chapter 4:
Solution Chemistry and the
Hydrosphere
Problems: 4.1-4.80, 4.85-4.96, 4.99-4.100,
4.111-4.113, 4.119-4.120, 4.129, 4.131-4.132
Solutions on Earth and Other
Places
aqueous solution: a solution where water is the
dissolving medium (the solvent)
• For example, when table salt (NaCl) is dissolved in
water, it results in an aqueous solution of sodium
chloride, NaCl(aq), with Na+ and Cl- ions dissolved in
water.
• Note: The physical state aqueous,(aq), indicates an
element or compound dissolved in water while the
physical state liquid,(l), means a pure substance in the
liquid state.
– Thus, NaCl(aq)  NaCl(l), which is molten NaCl
requiring very high temperatures.
A solution consists of a solute dissolved in a solvent.
Solutions
solute: component present in smaller amount
solvent: component present in greater amount
The formation of a solution:
As a solute crystal is dropped into a solution, the water
molecules begin to pull apart the ionic compound ion by ion
Solvent molecules surround the solute
particles, forming a solvent cage
– the ions are now hydrated
(surrounded by polar water
molecules)
– solute is now dissolved in the
solvent and cannot be seen
because the ions are far apart, like
the particles in a gas
Unsaturated, Saturated and
Supersaturated Solutions
• In general, if a solid is soluble in a solvent, more solid
dissolves in the solvent at higher temperatures.
unsaturated: contains less than the maximum amount
of solute that a solvent can hold at a specific
temperature
saturated: contains the maximum amount of solute
that a solvent can hold at a specific temperature
supersaturated: contains more than the maximum
amount of solute that a solvent should be able to hold
at specific temperature
Unsaturated, Saturated and
Supersaturated Solutions
• A supersaturated NaC2H3O2 solution recrystallizing
after addition of more solute:
Unsaturated, Saturated and
Supersaturated Solutions
• How can a solution hold more solute than it should
be able to hold?
– If a given amount of solute is dissolved in a
solvent at a higher temperature, and the solution
is allowed to cool without being disturbed, the
solute will remain in solution.
• But the solution is unstable, and the solute will come
out of solution (i.e. recrystallize) if the solution is
disturbed (e.g. by adding more solute, scratching the
glass, etc.)
Unsaturated, Saturated and
Supersaturated Solutions
For some substances, recrystallization is exothermic,
releasing heat to the surroundings.
– Hot packs used to warm hands and feet in winter
(though some of these are oxidation reactions, which
we will discuss later)
For other substances, recrystallization is endothermic,
absorbing heat and making the surroundings colder.
– Cold packs used for sports injuries
How do we measure
concentration?
solution: homogeneous mixture of substances present
as atoms, ions, and/or molecules
solute: component present in smaller amount
solvent: component present in greater amount
Note: Unless otherwise stated, the solvent for most
solutions considered in this class will almost always be
water!
Aqueous solutions are solutions in which water is the
solvent.
How do we measure
concentration?
• A concentrated solution has a large quantity of
solute present for a given amount of solution.
• A dilute solution has a small quantity of solute
present for a given amount of solution.
amount of solute
amount of solvent
The more solute in a given amount of solution  the
more concentrated the solution
Example: Explain the difference between the
density of pure ethanol and the concentration of an
ethanol solution.
SOLUTION CONCENTRATION =
How do we measure
concentration?
Concentration can be measured a number of ways:
• ppm (parts per million) – one part in a million parts
• ppb (parts per billion) – one part in a billion parts
• g/kg (grams per kilogram) – one gram solute per one
kilogram of solvent
The chemical standard most used is Molarity
Molarity =
moles of solute
liters of solution
units: M (molar = mol/L)
We’ll come back to concentration later in the chapter…
Evidence of a Chemical
Reaction
a) A gas is produced.
b) A precipitate forms.
c) Heat is released or
absorbed
Types of Chemical Reactions
• Precipitation Reactions
• Acid-Base Neutralization Reaction
• Oxidation-Reduction (Redox) Reactions
– Further classified as:
• Combination
• Decomposition
• Combustion
• Single-replacement reactions
Precipitation Reactions
• Solubility Rules: Indicate if an ionic compound is soluble
or insoluble in water.
• Keep in mind that these are just general guidelines, and in
reality, some ionic compounds are slightly soluble, and
solubility may depend on the temperature.
Solubility Rules for Ionic Compounds in Water
Soluble if the ionic compound contains:
•
•
•
•
Li+, Na+, K+, NH4+ (ALWAYS!)
C2H3O2–, NO3–, ClO3–, ClO4–
Halide ions (X–): Cl–, Br–, or I–, but AgX,
PbX2, HgX, and Hg2X2 are insoluble
sulfate ion (SO42-), but CaSO4, SrSO4,
BaSO4, Ag2SO4, `and PbSO4 are
insoluble.
Insoluble if the ionic compound contains:
•
•
•
•
•
carbonate ion, CO32chromate ion, CrO42phosphate ion, PO43sulfide ion (S2–), but CaS, SrS, and BaS are
all soluble.
hydroxide ion (OH–), but Ca(OH)2, Sr(OH)2,
and Ba(OH)2 are soluble.
Precipitation Reactions
soluble = compound dissolves in water  exists as individual
ions in solution
 physical state is aqueous, (aq)
Insoluble = compound does not dissolve in water but remains
a solid
 physical state is shown as solid, (s)
Precipitation Reactions
• Example 1: Use the Solubility Rules and identify the
ionic compounds are soluble or insoluble by
indicating the physical state of each compound.
a. NaCl
d.LiOH
g.Mg(OH)2
j. Ag3PO4
b.MgS
e. CaS
h. SrSO4
k.
c. K3PO4
f. Li2CrO4
i. Na2CO3
l. (NH4)2CrO4
BaCO3
Precipitation Reactions
• Example 1: Use the Solubility Rules and identify the
ionic compounds are soluble or insoluble by
indicating the physical state of each compound.
a. NaCl
d.LiOH
g.Mg(OH)2
j. Ag3PO4
b.MgS
e. CaS
h. SrSO4
k.
c. K3PO4
f. Li2CrO4
i. Na2CO3
l. (NH4)2CrO4
Soluble
(aq)
Insoluble
(s)
BaCO3
Precipitation Reactions
• In a precipitation reaction, two solutions react to
form a precipitate (an insoluble solid):
AX(aq) + BZ(aq)  AZ(s) + BX(aq)
precipitate
For example:
KI (aq)+ Pb(NO3)2(aq)  PbI2 (s) + KNO3 (aq)
Precipitation Reactions
To balance and complete the precipitation reactions:
1. Exchange the anions, writing the formulas for the
products based on the charges of the ions!
2. Use the Solubility Rules to determine if each product is
soluble or insoluble.
– If at least one product is insoluble, a precipitation
reaction has occurred. Write the formulas for both
products, indicating the precipitate as (s), then balance
the equation.
– If both products are soluble, write NR (=no reaction).
3. Keep in mind that the charges on ions do NOT change
in precipitation reactions. For metals that can form more
than one charge, use the charge on the metal ion from the
reactant side of the equation.
Precipitation Reactions
Ex 1.
MgSO4(aq)
+
NaOH(aq)

Ex 2.
K2CO3(aq)
+
AlCl3(aq)

Ex 3.
SrBr2(aq)
+
Zn(NO3)2(aq) 
Ex 4.
CuSO4(aq)
+
NaOH(aq)
Ex 5.
KI(aq)
+
Pb(NO3)2(aq) 

Acid-Base Neutralization
Reactions: Proton Transfer
Properties of Acids and Bases
Acids
–produce hydrogen ions, H+
–taste sour
–turn blue litmus paper red
Bases
–produce hydroxide ions, OH–
–taste bitter; feel soapy,
slippery
–turn red litmus paper blue
Arrhenius Definitions
acid: A substance that releases H+ when dissolved in
water
– Some acids are monoprotic (release only H+ per
molecule)
• e.g. HCl, HBr, HI, HNO3, HClO4
– Some acids are polyprotic (release more than on
H+ per molecule)
• e.g. H2SO4 and H2CO3 are both diprotic; H3PO4
is triprotic.
base: A substance that releases OH– when dissolved in
water
Acid-Base Reactions
In an acid-base reaction,
• H+ from acid reacts with the OH– from base to form
water, H2O
• The cation (M+) from base combines with anion from acid
(X–) to form a salt.
A general equation for an acid-base neutralization reaction
is shown below:
HX(aq) + MOH(aq)  H2O(l) + MX
acid
base
water
salt
Because water is always produced, an acid always reacts
with a base!
Examples
Complete and balance each of the equations below:
a.
HCl(aq) +
NaOH(aq) 
b.
H2SO4(aq) +
KOH(aq) 
c.
H3PO4(aq) +
Ca(OH)2(aq) 
Acid-Base Reactions with Gas
Formation
Some acid-base reactions produce carbon dioxide gas, CO2(g),
along with water and salt.
When the base contains carbonate ion (CO32–) or hydrogen
carbonate ion (HCO3–), then the products of the acid-base
reaction are water, carbon dioxide gas, and a salt.
The general equations for the unbalanced acid-base reactions
are:
HX(aq) + MCO3(s)  H2O(l) +
CO2(g) +
acid
base
water
carbon dioxide
HX(aq) + MHCO3(s)  H2O(l) +
CO2(g)
+
acid
base
water
carbon dioxide
MX
salt
MX
salt
Because water is always produced, an acid always reacts
with a base!
Acid-Base Reactions with Gas
Formation
Complete and balance each of the equations below:
a.
HCl(aq) + Na2CO3(s) 
b.
HNO3(aq) + CaCO3(s) 
c.
H2SO4(aq) + KHCO3(s) 
d.
HClO4(aq) + Sr(HCO3)2(s) 
A double-replacement reaction that produces
NH4OH(aq) actually produces ammonia, NH3(g).
NH4OH(aq)

NH3(g) +
H2O(l)
Example: Complete and balance the equation below:
(NH4)2SO4(aq) +
KOH(aq) 
Brønsted-Lowry Definition of
Acids and Bases
• Brønsted-Lowry acid: A substance that donates a
proton (H+)—i.e., a proton donor
• Brønsted-Lowry base: A substance that accepts a
proton (H+)—i.e., a proton acceptor
• Unlike an Arrhenius base, a Brønsted-Lowry base
does not need to contain OH–.
Why is H+ called a proton?
Brønsted-Lowry Acids and
Bases
A Brønsted-Lowry acid-base reaction simply involves
a proton (H+) transfer.
NH3(aq) + H2O(l) ⇄ NH4+(aq) + OH–(aq)
Note: This reaction simply involves H2O donating a H+ ion to NH3 to produce NH4+
and OH–.
• In this reaction, H2O is the Brønsted-Lowry acid, and NH3 is the Brønsted-Lowry
base.
• The conjugate acid-base pairs differ only by a H+.
• In this reaction, the conjugate acid-base pairs are NH3 and NH4+ and H2O and
OH–.
Conjugate Acid-Base Pairs
Conjugate acid-base pairs: a Brønsted-Lowry acid/base and its
conjugate differ by a H+
HA(aq) + H2O(l) ⇄ H3O+(aq) + A–(aq)
• For the reaction above, when HA donates H+ to H2O, it leaves
behind A–, which can act as a base for the reverse reaction.
• An acid and base that differ only by the presence of H+ are
conjugate acid-base pairs.
• The general reaction for the dissociation (or ionization) of an
acid can be represented as above, where the double-arrow
indicates both the forward and reverse reactions can occur.
• Note: The double arrow (⇄) indicates the reaction is reversible
(goes in both directions).
Conjugate Acid-Base Pairs
Determine the Brønsted-Lowry acid and base in each of
the following reactions:
a. CH3COOH(aq) + NH3(aq)
⇄
NH4+(aq) + CH3COO–(aq)
b. NH3(aq) + H2O(l) ⇄ NH4+(aq) +
c. H2O(l) +
H2SO4(aq) ⇄ H3O+(aq)
OH–(aq)
+ HSO4–(aq)
Oxidation-Reduction (Redox)
Reactions
Types of Redox Reactions
•
•
•
•
Combination Reaction
Decomposition Reaction
Single-Replacement (or Displacement) Reaction
Combustion Reaction
Combination Reactions
A + Z  AZ
Usually a meal and a non-metal react to form a
solid ionic compound:
metal + nonmetal
Δ
ionic compound(s)
Combination Reactions:
A + Z  AZ
Complete and balance each of the equations below:
a. Na(s) + Cl2(g)
Δ
b. Al(s) + O2(g)
Δ
c. Zn(s) + S8(s)
Δ
d. Mg(s) + N2(g)
Δ
Decomposition Reactions:
AZ  A + Z
Be able to classify and balance decomposition reactions. (You
won’t need to predict products.)
Δ
a. ___ KHCO3(s)
b. ___ Al2(CO3)3(s)
c. ___ KClO3(s)
Δ, MnO2
___ K2CO3(s) + ___ H2O(l) + ___ CO2(g)
Δ
_____ Al2O3(s) + _____ CO2(g)
_____ KCl(s) + _____ O2(g)
Single-Replacement Reactions
and the Activity Series
Activity Series: Relative order of elements arranged by
their ability to replace cations in aqueous solution
Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn > Zn >
Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au
Note: The Activity Series will be given to you on quizzes and
exams.
Single Replacement Reactions:
A + BZ  AZ + B
metal A + aqueous solution B  aqueous solution A + metal B
Zn(s) + Sn2+(aq)  Sn(s) + Zn2+(aq)
Cu(s) + 2 Ag+(aq)  2 Ag(s) + Cu2+(aq)
Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn >
Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au
To balance and complete the following rxns:
• Check the Activity Series to see which metal is more active.
– The more active metal replaces the less active by going
into solution as an ion and the less active metal ion
comes out as a solid.
1.
2.
3.
4.
Mg(s)
Al(s)
Cd(s)
Ag(s)
+
+
+
+
AlCl3(aq)

CdSO4(aq) 
AgNO3(aq) 
Mg(NO3)2(aq) 
Activity Series: Li > K > Ba > Sr > Ca > Na > Mg > Al > Mn >
Zn > Fe > Cd > Co > Ni > Sn > Pb > (H) > Cu > Ag > Au
To balance and complete the following reactions:
• Check the Activity Series to see which metal is more active,
the metal or H.
– The more active metal replaces the less active by going
into solution as an ion and the H comes out as hydrogen
gas, H2(g).
metal A + acid solution  aqueous solution A + H2(g)
1. Zn(s) + HCl(aq) 
2. Al(s) + HNO3(aq) 
3. Cu(s) + HI(aq) 
Active Metals:
Li > K> Ba > Sr > Ca > Na
Active metals react directly with water:
The active metal replaces the less active by going into
solution as an ion with hydroxide, OH–, and the H comes out
as hydrogen gas, H2(g).
active metal + H2O(l)  metal hydroxide + H2(g)
1. Ca(s) + H2O(l) 
2. Na(s) + H2O(l) 
3. Fe(s) + H2O(l) 
Combustion Reactions
CxHy + O2(g)  H2O(g) + CO2(g)
CxHyOz + O2(g)  H2O(g) + CO2(g)
1. C3H8(g) + O2(g)
2. C6H6O(l) + O2(g)
3. C2H2(g) + O2(g)
Δ
Δ
Δ
Identify the Reactions
For each of the following,
1. Identify the type of reaction using the letters
designated below:
Combination (C)
Precipitation (P)
Decomposition (D)
Acid-Base Neutralization (N)
Combustion (B)
Single Replacement/Displacement (SR)
2. Balance the equations:
____a.
____b.
____c.
____d.
___ Mg(NO3)2(aq) + ___ K3PO4(aq)  ___ Mg3(PO4)2(s) + ___ KNO3(aq)
___ Ni(OH)3(s) + ___ HCl(aq)  ___ H2O(l) + ___ NiCl3(aq)
___ Al(HCO3)3(aq) Δ ___ CO2(g) + ___ H2O(g) + ___ Al2(CO3)3(s)
___ Fe(s) + ___ Pb(NO3)2(aq)  ___ Pb(s) + ___ Fe(NO3)3(aq)
Electrolytes and NonElectrolytes
Electrolytes and electrical conductivity
• If a solution conducts electricity, it contains ions
• A solution that contains many ions is a strong electrolyte.
• A solution that contains only a few ions is a weak
electrolyte.
• A solution that contains only a no ions is a nonelectrolyte.
Electrolytes and NonElectrolytes
• A solution that contains many ions is a strong
electrolyte.
 Light bulb burns brightly in a light bulb
conductivity apparatus.
• A solution that contains only a few ions is a weak
electrolyte.
 Light bulb burns dimly in a light bulb
conductivity apparatus.
• A solution that contains only a no ions is a
nonelectrolyte.
 Light bulb does not light in a light bulb
conductivity apparatus.
Strong Electrolytes
strong electrolytes: substances that are good conductors
of electricity
• These substances break up to produce many ions in water
– many ions present to move electrons/conduct
electricity  strong electrolyte
For example,
H2O
NaCl(s)
Na+(aq) + Cl–(aq)
HO
KOH(s) 2 K+(aq) + OH–(aq)
HO
HBr(aq) 2 H+(aq) + Br–(aq)
Examples: strong acids, strong bases, all soluble ionic
compounds
Weak Electrolytes
weak electrolytes: substances that are weak/poor
conductors of electricity
• These substances mostly remain intact as
compounds, producing very few ions in water
– only a few ions present to move electrons/conduct
electricity  weak electrolyte
For example,
H2O
Mg(OH)2(s)
Mg(OH)2(s)
HNO2(aq)
HNO2(aq)
Examples: weak acids, weak bases, insoluble ionic
compounds
Non-electrolytes
nonelectrolytes: substances that cannot conduct
electricity
• These molecules never break down into ions.
– They always remain intact as neutral molecules
that have no charge  no ions to move
electrons/conduct electricity
For example,
C12H22O11(s)
H2O
C12H22O11(aq)
Examples: sugar (e.g. sucrose), ethanol (C2H5OH), and all
other molecules that are not acids
Acids and Bases
Know the following acids and bases. All other acids and
bases are weak!
Strong Acids
Strong Bases
HCl, HBr , HI, HNO3, HClO4,
H2SO4
LiOH, NaOH, KOH, Ca(OH)2,
Sr(OH)2, Ba(OH)2
Strong acids and bases dissolve in water to form ions (or
species) in solution.
HNO3(aq)  H+(aq) + NO3–(aq)
Ca(OH)2(aq)  Ca2+(aq) + 2 OH–(aq)
Note: H2SO4(aq) is a strong acid and diprotic (able to release 2
H+ ions), but it generally ionizes to release only one H+ ion in
water: H2SO4(aq)  H+(aq) + HSO4–(aq)
Recognize that both protons are not released in water!
Molecular, Ionic and Net Ionic
Equations
• molecular equation: equation showing reactants and
products as compounds
• total/complete ionic equation: shows strong
electrolytes as individual ions while all solids, liquids,
gases, and weak electrolytes remain intact as compounds
• spectator ions: ions that do not form solids, liquids,
gases, weak electrolytes – appear on both sides of total
ionic equation as ions
• net ionic equations: show only solids, liquids, gases,
weak electrolytes (weak acids and weak bases), and ions
undergoing a chemical change/reaction – excludes
spectator ions
Net Ionic Equations
Guidelines for Writing Net Ionic Equations
1. Balance the chemical/molecular equation.
2. Convert the molecular equation to total ionic equation
– Leave solids, liquids, gases, and weak acids and bases as
compounds
– Show strong acids and all aqueous ionic compounds as
ions in solution.
3. Cancel spectator ions to get net ionic equation
– If canceling spectator ions eliminates all ions  NO
REACTION (NR)
– If coefficients can be simplified, do so to get the lowest
ratio.
4. Make sure total charges on both sides of the equation are
equal.