Reactions in
Aqueous Solutions
AP Chemistry
Ms. Grobsky
PROPERTIES OF AQUEOUS SOLUTIONS
What is a Solution?
• Defined as homogeneous
mixtures of two or more
pure substances
• The solvent is present in
greatest abundance
• Does the dissolving
• All other substances are
solutes
• Thing being dissolved
Water: The Universal Solvent
• One of the most valuable
properties of water is its
ability to dissolve other
substances
• An individual water
molecule has a bent
shape with a H-O-H bond
angle of approximately
105 degrees
• Water is polar
• Thus having positive &
negative partial charges on
its ends
Ionic Compounds in Water
• When an ionic substance
dissolves in water, the
solvent pulls the individual
ions from the crystal and
solvates them
• The positive ends of a
water molecule are
attracted to negative
anions and the negative
ends are attracted to
positive cations in an
ionic compound
• This is called hydration
• The ions can then
move around
independently
Covalent Compounds in Water
• Water also dissolves many nonionic substances such as
ethanol (C2H5OH)
• The reason for this is that ethanol is also polar
LIKE DISSOLVES LIKE
• Substances with similar types of intermolecular forces dissolve
in each other
• Polar solvent is used to dissolve polar or ionic solute
• Nonpolar solvent is used to dissolve a nonpolar solute
Solutions as Electrolytes
• An electrolyte is a
substance that
dissociates into ions
when dissolved in water
• A nonelectrolyte may
dissolve in water, but it
does not dissociate into
ions when it does so
Electrolytes and Nonelectrolytes
• Soluble ionic
compounds tend to
be electrolytes
• Molecular
compounds tend to
be nonelectrolytes
• Does not conduct
electricity
© 2012 Pearson Education,
Inc.
Electrolytes and Nonelectrolytes
Strong vs. Weak Electrolytes
• A strong electrolyte dissociates completely when dissolved in
water
• Conducts electricity
• A weak electrolyte only dissociates partially when dissolved in
water
• Conducts electricity poorly
Strong Electrolytes Are…
• Strong acids
• Strong bases
Strong Electrolytes Are…
© 2012 Pearson Education,
Inc.
• Strong acids
• Strong bases
• Soluble ionic salts
Ways to Measure Solution
Composition
• Molarity
𝑎𝑚𝑜𝑢𝑛𝑡 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒
𝑣𝑜𝑙𝑢𝑚𝑒 𝐿 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
• Parts by Mass
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
• Mass percent
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
× 100
𝑚𝑎𝑠𝑠 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
Ways to Measure Solution
Composition
• Mole fraction
• Molality
• Parts by
Volume
𝑎𝑚𝑜𝑢𝑛𝑡 𝑚𝑜𝑙 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝑎𝑚𝑜𝑢𝑛𝑡 𝑚𝑜𝑙 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒 + 𝑎𝑚𝑜𝑢𝑛𝑡 𝑚𝑜𝑙 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
𝑎𝑚𝑜𝑢𝑛𝑡 𝑚𝑜𝑙 𝑠𝑜𝑙𝑢𝑡𝑒
𝑎𝑚𝑜𝑢𝑛𝑡 𝑘𝑔 𝑜𝑓 𝑠𝑜𝑙𝑣𝑒𝑛𝑡
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑒
𝑣𝑜𝑙𝑢𝑚𝑒 𝑜𝑓 𝑠𝑜𝑙𝑢𝑡𝑖𝑜𝑛
Molarity Example Problem
• The most common method of expressing
solution concentration is molarity:
M = moles solute
volume (L) solution
• Example Problem
• Calculate the molarity of a solution prepped by
dissolving 1.56 g of gaseous hydrochloric acid in
enough water to make 26.8 mL of solution.
• 1.60 M HCl
Concentration of Ions in Solution
•
When an ionic salt dissolves in water, ions are
form in solution
•
•
Example Problem
•
(a)
(a)
The moles of ions formed must be considered in
concentration!
What is the concentration each ion in (a) 0.50M
cobalt II nitrate (b) 2.0M iron III perchlorate?
Co(NO3)2 (s) + H2O  Co2+ (aq) + 2NO3- (aq)
[Co2+] = 1 x 0.50M = 0.50M
[NO3-] = 2 x 0.50M = 1.0M
Fe(ClO4)3 (s) + H2O  Fe3+ (aq) + 3ClO4- (aq)
[Fe3+] = 2.0M
[ClO4-] = 6.0M
Concentration of Ions in Solution
• Example Problem
• Calculate the number of moles of chloride ions in 1.75 L of 1.0 x
10-3M zinc chloride.
(Step 1) ZnCl2 (s)  Zn2+(aq) + 2Cl-(aq)
(Step 2) [Cl-] = 2 x 1.0 x 10-3 = 2.0 x 10-3M
(Step 3) 1.75 L x 2.0 x 10-3 mole Cl- = 3.5 x 10-3 moles of Cl- ions
1L
Example Problems for
Concentration & Volume
• Typical blood serum is about 0.14M sodium
chloride. What volume of blood contains 1.0 mg
of sodium chloride?
• 0.12 mL of blood
• To analyze the alcohol content of a certain wine,
a chemist needs 50.00 mL of an aqueous 0.200M
potassium dichromate solution. How much solid
potassium dichromate must be weighed out to
make this solution?
• 2.94 g K2Cr2O7
Dilution Formula
• This formula allows a chemist to prepare a diluted solution
from a concentrated one:
𝑀1 𝑉1 = 𝑀2 𝑉2
𝑀𝐶 𝑉𝐶 = 𝑀𝐷 𝑉𝐷
• Example Problem
• What volume of 16 M sulfuric acid must be used to prepare 1.5 L
of a 0.10 M sulfuric acid solution?
• 9.4 mL of H2SO4 must be diluted with 1.5 L of water
TYPES OF REACTIONS IN
AQUEOUS SOLUTIONS
Types of Chemical Reactions in
Aqueous Solutions
Precipitation Reactions
• When 2 aqueous solutions are mixed an insoluble
precipitate sometimes forms – also known as double
replacement or metathesis reactions.
• It is important to remember that some ions are the
key players and some are just spectators:
• The formula equation gives the overall reaction.
• The complete ionic equation represents all ions involved in
the reaction.
• The net ionic equation includes only those solution
components undergoing a change, spectator ions are not
included.
The reaction of Pb(NO3)2 & NaI.
Write the formula equation, ionic & net ionic equations:
Stoichiometry of Precipitation Reactions
• Calculate the mass of solid sodium chloride that must be
added to 1.50 L of a 0.100 M silver nitrate solution to
precipitate all the silver ions in the form of silver chloride.
• First find the # of moles Na necessary for the # of moles of
silver ions already present, then convert to grams.
• 8.77 g NaCl
Example: Determine Mass of Product Formed
• When aqueous solutions of sodium sulfate and lead II nitrate
are mixed a precipitate forms. Give the net ionic equation for
the reactions and calculate the mass of this precipitate when
1.25 L of 0.0500 M lead II nitrate and 2.00 L of 0.0250 M
sodium sulfate.
15.2 g PbSO4
Acid-Base Reactions
• An acid is a substance that produces H+ ions when dissolved in
water.
• A base is a substance that produces OH− ions when dissolved
in water.
Acids and bases as Electrolytes
A, Strong acids and bases are strong electrolytes, as indicated by the
brightly lit bulb. B, Weak acids and bases are weak electrolytes.
Neutralization Reactions
Strong Acid—Strong Base:
Because both ionize completely, the H+ ions and OH- ions
react with each other to form water molecules.
Basic: HNO3 (aq) + NaOH (aq) 
Ionic: H+ (aq) + NO3- (aq) + Na+ (aq) + OH- (aq) 
Net Ionic:
H+ (aq) +
OH- (aq)

H2O (l)
Weak Acid – Strong Base:
A two-step process occurs:
• 1. The ionization of the weak acid
HB (aq)  H+ (aq) + B- (aq)
• 2. The neutralization of the H+ ion by the OH- from the strong
base
H+ (aq) + OH- (aq)  H2O (l)
Net Ionic: HX
+
OH-

X-
+
H2O
Strong Acid – Weak Base:
This is also a two-step process:
•
1.The first step occurs when the weak base reacts with water
to produce OH- ions.
NH3 + H2O  NH4+ + OH-
•
2. In the second step the H+ ions from the strong acid
neutralize the OH- ions to form water.
H+ + OH-  H2O
• Net Ionic: H+ +
X

XH+
Example: Neutralization Reactions
• 1)What volume of a 0.100 M HCl is needed to
neutralize 25.0 mL of 0.350 M NaOH?
8.75 x 10-2 L
• 2)In a certain experiment, 28.0 mL of 0.250 M
nitric acid and 53.0 mL of 0.320 M potassium
hydroxide are mixed. Calculate the moles of
water formed in the resulting reaction. What is
the [H+] and [OH-] after the reaction goes to
completion?
0.024 moles H2O, [H+] = 0, [OH-] = 0.123 M
Acid-Base Titration
• This is a volumetric analysis technique for
determining the amount (usually concentration) of a
substance.
• This involves the delivery (from a buret) of a
measured volume of a solution of known
concentration (the titrant) into a solution containing
the substance being analyzed (the analyte).
• The neutralization point is known as the equivalence
point. This point is marked by an indicator.
• The point when the indicator changes color is known
as the end point.
• The goal is to choose an indicator which has a similar
endpoint as the equivalence point your reaction.
Titration Example #1
You perform an acid-base titration to standardize an HCl solution by
placing 50.00 mL of HCl in a flask with a few drops of indicator solution.
You put 0.1524 M NaOH into the buret, and the initial reading is 0.55 mL.
At the end point, the buret reading is 33.87 mL. What is the concentration
of the HCl solution?
Titration Example #2
In a titration, it is found that 25.0 mL of 0.500 M NaOH is
required to react with
• (a) a 15.0-mL sample of HCl. What is the molarity of HCl?
0.833 M
• (b) a 15.0-mL sample of a weak acid, H2A. What is the
molarity of H2A, assuming the reaction to be
H2A(aq) + 2OH-(aq)  2H2O + A2-(aq)?
0.417 M
• (c) an aspirin tablet weighing 2.50 g. What is the percentage
of acetylsalicylic acid, HC9H7O4, in the aspirin tablet? The
reaction is
HC9H7O4 (s) + OH- (aq)  H2O + C9H7O4 - (aq)
90.0%
Oxidation-Reduction (Redox)
• Oxidation means the losing of electrons (an increase in the
oxidation #) and reduction means the gaining of electrons (a
decrease in the oxidation #). The 2 occur together, they are
opposite sides of the same coin.
• A good way to remember LEO the lion goes GER – losing electrons
oxidation…..gaining electrons reduction OR OIL RIG – oxidation is
losing, reduction is gaining.
Oxidizing & Reducing Agents
• An oxidizing agent is the species that accepts the electrons i.e.
the H+ ion above.
• Non-metals tend to be oxidizing agents.
• A reducing agent is the species that donates the electrons i.e.
the Zn above.
• Metals tend to be reducing agents.
Oxidation Numbers
• The first step to balancing any redox is assigning
oxidation numbers to reactants and products in the
equation – please reference pg. 89 in text:
• The oxidation # of an element in its elemental state is 0.
• The oxidation # of an element in a monatomic ion is equal
to the charge of that ion.
• Certain elements have the same oxidation in all or almost
all their compounds.
• The sum of the oxidation numbers in a neutral species is 0;
in a polyatomic ion, it is equal to the charge of that ion.
Assigning Oxidation #’s
Practice
• What is the oxidation number of phosphorus
• in sodium phosphate?
• P = +5
• In the dihydrogen phosphate ion?
• P = +5
Balancing Redox Reactions
•
•
Assign oxidation numbers (Rules – Pg. 89)
Identify the oxidation and reduction reactions.
Split in 2 half reactions.
Balance the element being oxidized and reduced.
Balance the elements (that is being reduced or oxidized) oxidation # by
adding electrons. Oxidation adds to the right, reduction adds to the left.
Balance the oxygens by adding water molecules.
Balance the hydrogens by adding H+ ions.
If the electrons on both sides are not the same you must find the least
common multiple between the 2 electrons. Multiply each reaction to
get the same number of electrons on both sides.
•
•
•
•
•
•
•
•
It is important to check that all atoms and charges balance at this point.
THE MISTAKES ARE LIKELY TO HAPPEN HERE!!!
Combine the reactions and simplify if necessary.
If in basic solution add OH- ions to both sides to produce water
molecules on the side with H+ ions. Simplify the water molecules if
necessary.
Zn (s) + 2HCl (aq)  ZnCl2 (aq) + H2 (g)
• Oxidation: Zn  Zn+2 + 2e• Reduction: 2H+ + 2e-  H2
Net ionic equation:
Zn (s) + 2H+ (aq)  H2 (g) + Zn+2 (aq)
Balance the following redox…
• (a) Fe2+(aq)+ MnO4- (aq) Fe3+(aq) Mn2+(aq) (acidic solution)
• (b) Cl2(g) Cr(OH)3(s)  Cl- (aq) CrO42- (aq)
(basic solution)
MC #1
• When 70. milliliter of 3.0-molar Na2CO3 is added to 30.
milliliters of 1.0-molar NaHCO3 the resulting concentration of
Na+ is
(A) 2.0 M
(B) 2.4 M
(C) 4.0 M
(D) 4.5 M
(E) 7.0 M
MC #2
• A student wishes to prepare 2.00 liters of 0.100-molar
KIO3 (molecular weight 214 g/mol). The proper
procedure is to weigh out
(A) 42.8 grams of KIO3 and add 2.00 kilograms of H2O
(B) 42.8 grams of KIO3 and add H2O until the final
homogeneous solution has a volume of 2.00 liters
(C) 21.4 grams of KIO3 and add H2O until the final
homogeneous solution has a volume of 2.00 liters
(D) 42.8 grams of KIO3 and add 2.00 liters of H2O
(E) 21.4 grams fo KIO3 and add 2.00 liters of H2O
MC #3
• A 20.0-milliliter sample of 0.200-molar K2CO3
solution is added to 30.0 milliliters of 0.400molar Ba(NO3)2 solution. Barium carbonate
precipitates. The concentration of barium ion,
Ba2+, in solution after reaction is
(A) 0.150 M
(B) 0.160 M
(C) 0.200 M
(D) 0.240 M
(E) 0.267 M
MC #4
• The weight of H2SO4 (molecular weight 98.1 g/mol) in 50.0
milliliters of a 6.00-molar solution is
(A) 3.10 grams
(B) 12.0 grams
(C) 29.4 grams
(D) 294 grams
(E) 300. grams
MC #5
• Given that a solution is 5 percent sucrose by
mass, what additional information is necessary
to calculate the molarity of the solution?
I. The density of water
II. The density of the solution
III. The molar mass of sucrose
(A) I only
(B) II only
(C) III only
(D) I and III
(E) II and III
MC #6
• A yellow precipitate forms when 0.5 M NaI(aq) is added to a
0.5 M solution of which of the following ions?
A) Pb2+ (aq)
B) Zn2+ (aq)
C) CrO42¯ (aq)
D) SO42¯ (aq)
E) OH¯ (aq)
MC #7
• When 100 mL of 1.0 M Na3PO4 is mixed with 100
mL of 1.0 M AgNO3, a yellow precipitate forms
and [Ag+] becomes negligibly small. Which of
the following is a correct listing of the ions
remaining in solution in order of increasing
concentration?
A) [PO43¯] < [NO3¯] < [Na+]
B) [PO43¯] < [Na+] < [NO3¯]
C) [NO3¯] < [PO43¯] < [Na+]
D) [Na+] < [NO3¯] < [PO43¯]
E) [Na+] < [PO43¯] < [NO3¯]
MC #8
• The volume of distilled water that should be added to 10.0 mL
of 6.00 M HCl(aq) in order to prepare a 0.500 M HCl(aq) solution
is approximately
A) 50.0 mL
B) 60.0 mL
C) 100. mL
D) 110. mL
E) 120. mL
MC #9
• The net ionic equation for the reaction between silver carbonate and
hydrochloric acid is
(A) Ag2CO3 + 2H+ + 2 Cl¯ ---> 2 AgCl + H2O + CO2
(B) 2Ag+ + CO32¯ + 2 H+ + 2 Cl¯ ---> 2 AgCl + H2O + CO2
(C) CO32¯ + 2 H+ ---> H2O + CO2
(D) Ag+ + Cl¯ ---> AgCl
(E) Ag2CO3 + 2H+ ---> 2Ag+ + H2CO3
MC #10
5 Fe2+ + MnO4¯ + 8 H+ <===> 5 Fe3+ + Mn2+ + 4H2O
• In a titration experiment based on the equation above,
25.0 milliliters of an acidified Fe2+ solution requires 14.0
milliliters of standard 0.050-molar MnO4¯ solution to
reach the equivalence point. The concentration of Fe2+ in
the original solution is
(A) 0.0010 M
(B) 0.0056 M
(C) 0.028 M
(D) 0.090 M
(E) 0.14 M
FRQ #1
• A 1.2516 gram sample of a mixture of CaCO3 and Na2SO4 was analyzed by
dissolving the sample and completely precipitating the Ca2+ as CaC2O4. The
CaC2O4 was dissolved in sulfuric acid and the resulting H2C2O4 was titrated with
a standard KMnO4 solution.
(a)
Write the balanced equation for the titration reaction, shown unbalanced
below:
MnO4- + H2C2O4 + H+  Mn2+ + CO2 + H2O
(i) Indicate which substance is the oxidizing agent and which substance is
the reducing agent.
(b)
The titration of the H2C2O4 obtained required 35.62 milliliters of 0.1092
molar MnO4- solution. Calculate the number of moles of H2C2O4 that reacted
with the MnO4(c)
Calculate the number of moles of CaCO3 in the original sample.
(d)
Calculate the percentage by weight of CaCO3 in the original sample.
FRQ #2
• Permanganate ion, MnO4-, oxidizes sulfite ions to
sulfate ion. The manganese product depends
upon the pH of the reaction mixture. The mole
ratio of oxidizing to reducing agent is two to five
at pH 1 (acidic), and is two to one at pH 13
(basic). For each of these cases, write a balanced
equation for the reaction, and indicate the
oxidation state of the manganese in the product
containing manganese.
FRQ #3
• A 0.150 g sample of solid lead(II) nitrate is added to 125 mL of 0.100
M sodium iodide solution. Assume no change in volume of the
solution. The chemical reaction that takes place is represented by the
following equation:
Pb(NO3)2(s) + 2 NaI(aq)  PbI2(s) + 2NaNO3(aq)
(a) List an appropriate observation that provides evidence of a
chemical reaction between the two compounds.
(b) Calculate the number of moles of each reactant.
(c) Identify the limiting reactant. Show calculations to support your
identification.
(d) Calculate the molar concentration of NO3–(aq) in the mixture after
the reaction is complete.
FRQ #4
• Answer the following questions about acetylsalicylic acid,
the active ingredient in aspirin.
(a) The amount of acetylsalicylic acid in a single aspirin
tablet is 325 mg, yet the tablet has a mass of 2.00 g.
Calculate the mass percent of acetylsalicylic acid in the
tablet.
(b) A student dissolved 1.625 g of pure acetylsalicylic
acid in distilled water and titrated the resulting solution
to the equivalence point using 88.43 mL of 0.102 M
NaOH(aq). Assuming that acetylsalicylic acid has only one
ionizable hydrogen, calculate the molar mass of the acid.
FRQ #5 (part I)
5 Fe2+(aq) + MnO4–(aq) + 8 H+(aq)  5 Fe3+(aq) + Mn2+(aq) + 4H2O(l)
• The mass percent of iron in a soluble iron(II) compound is measured using
a titration based on the balanced equation above.
(a) What is the oxidation number of manganese in the permanganate
ion, MnO4–(aq)?
(b) Identify the reducing agent in the reaction represented above.
Explain your reasoning.
• The mass of a sample of the iron(II) compound is carefully measured
before the sample is dissolved in distilled water. The resulting solution is
acidified with H2SO4(aq). The solution is then titrated with MnO4–(aq) until
the end point is reached.
(c) Describe the color change that occurs in the flask when the end
point of the titration has been reached. Explain why the color of the
solution changes at the end point.
FRQ #5 (part II)
(d)
Let the variables g, M, and V be defined as follows:
g = the mass, in grams, of the sample of the iron(II) compound
M = the molarity of the MnO4–(aq) used as the titrant
V = the volume, in liters, of MnO4–(aq) added to reach the end point
• In terms of these variables, the number of moles of MnO4–(aq) added to
reach the end point of the titration is expressed as M x V. Using the
variables defined above, the molar mass of iron (55.85 g mol-1), and the
coefficients in the balanced chemical equation, write the expression for
each of the following quantities
(i) The number of moles of iron in the sample
(ii) The mass of iron in the sample, in grams
(iii) The mass percent of iron in the compound
(e) What effect will adding too much titrant have on the experimentally
determined value of the mass percent of iron in the compound? Justify
your answer.
Equations #1
(a)
A sample of solid iron(III) oxide is reduced completely with solid
carbon.
(i)
Balanced equation:
(ii) What is the oxidation number of carbon before the reaction, and
what is the oxidation number of carbon after the reaction is complete
(b)
Equal volumes of equimolar solutions of ammonia and hydrochloric
acid are combined.
(i)
Balanced equation:
(ii) Indicate whether the resulting solution is acidic, basic, or neutral.
Explain.
(c)
Solid mercury(II) oxide decomposes as it is heated in an open test
tube in a fume hood.
(i)
Balanced equation:
(ii) After the reaction is complete, is the mass of the material in the test
tube greater than, less than, or equal to the mass of the original sample?
Explain.
Equations #2
(a) A small piece of sodium is placed in a beaker of distilled water.
(i) Balanced equation:
(ii) The reaction is exothermic, and sometimes small flames are
observed as the sodium reacts with the water. Identify the product of the
reaction that burns to produce the flames
(b)
Hydrogen chloride gas is oxidized by oxygen gas.
(i) Balanced equation:
(ii) If three moles of hydrogen chloride gas and three moles of oxygen
gas react as completely as possible, which reactant, if any, is present in
excess? Justify your answer.
(c)
Solid potassium oxide is added to water.
(i) Balanced equation:
(ii) If a few drops of phenolphthalein are added to the resulting solution,
what would be observed? Explain.