TYPES OF CHEMICAL REACTIONS AND SOLUTION CHEMISTRY
1.
Much of the chemistry that affects each of us occurs among substances dissolved in water.
2.
Before we can understand solution reactions, we must discuss the nature of solutions in which
water is the dissolving medium, or solvent.
3.
These solutions are called _______________________________.
Water – The Common Solvent
1.
One of the most valuable properties of water is its ability to dissolve many different substances.
2.
Liquid water consists of a collection of H2O molecules. The H2O molecule is __________ or Vshaped with an H-O-H bond angle of about 1050.
3.
The O-H bonds in the water molecule are _______________ bonds formed by electron sharing
between the oxygen and hydrogen atoms. However, the electrons are not shred equally.
4.
The electrons tend to spend more time close to the oxygen than to the hydrogen because oxygen
has a greater attraction for the electrons.
5.
This gives the oxygen atom a slight negative charge and the hydrogen a slight positive charge.
6.
Because of this unequal charge distribution, water is said to be a _________________________.
7.
It is this polarity that gives water its great ability to dissolve compounds.
8.
When an ionic substance dissolves in water, the “positive ends” of the water molecule are attracted
to the anions while the “negative ends” of the water are attracted to the cations.
9.
The process is called _______________.
OVER
10. The hydration of its ions tends to cause a salt to “fall apart” in the water, or to dissolve. The
strong forces of attraction between the cations and the anions are replaced by strong water-ion
interactions.
11. When ionic substances dissolve in water, they break up into the individual cations and anions.
(aq) designates that the ions are hydrated by an unspecified number of water molecules.
12. The _______________ of ionic substances in water varies greatly. Some are very soluble while
others are barely soluble. The differences in solubilities of ionic compounds in water depend on
the relative attractions of the ions for each other and the attraction of the ions for water molecules.
13. Water also dissolves many nonionic substances.
14. Ethanol contains a polar O-H bond like those in water, which makes it very compatible with
water.
15. Pure water will not dissolve animal fat.
16. Polar and ionic substances are more soluble in water than nonpolar substances.
Strong and Weak Electrolytes
1.
What happens when a substance, the ____________, is dissolved in liquid water, the
____________?
2.
One useful property for characterizing a solution is its _________________________, its ability
to conduct electricity.
3.
Some aqueous solutions conduct electricity very efficiently. These are said to contain
_________________________.
4.
Some aqueous solutions conduct only a small amount of current. These are said to contain
_________________________.
5.
Some aqueous solutions permit no current flow and are said to contain
_________________________.
6.
The extent to which a solution can conduct an electric current depends directly in water.
Strong Electrolytes
1.
Strong electrolytes are substances are substances that are _______________ ionized when they are
dissolved in water.
2.
Several classes of strong electrolytes are:
3.
Arrhenius proposed that an __________ is a substance that produces H+ ions when it is dissolved
in water.
OVER
4.
When HCl, HNO3 and H2SO4 are placed in water, virtually every molecule ionizes. These
substances are strong electrolytes and are thus called _________________________.
5.
A strong acid is one that completely dissociates into its ions.
6.
Another important class of strong electrolytes consists of ____________________. These are
soluble ionic compounds containing the hydroxide ion __________.
Weak Electrolytes
1.
Weak electrolytes are substances that exhibit a small degree of ionization in water.
2.
The most common are weak acids and weak bases.
3.
Because it is a weak electrolyte, it is also called a weak acid.
4.
Another type of weak electrolyte is a weak base.
Nonelectrolytes
1.
Substances that do not produce ions.
Composition of Solutions
1.
Chemical reactions often take place when two solutions are mixed.
2.
To perform stoichiometric calculations, one must know:
a.
the nature of the reaction.
b.
the amount of chemicals present in the solution, called ____________________.
3.
One expression of concentration is _______________.
***** Determine the molarity of a solution in which a 0.4508g sample of iron is dissolved in enough water
to make 250.0 mL of solution.
4.
The conventional description of a solution’s concentration may not be accurate. One must look at
what is really happening when the solute dissolves. Then determine the concentration of each ion
present.
***** Determine the concentration of the ions present in 0.250 M K 2Cr2O7
OVER
***** Determine the mass, in grams, of NaOH contained in 250.0 mL of 0.400 M NaOH solution.
Dilution
1.
Routinely used solutions are often purchased or prepared in concentrated form. Water is then
added to achieve the desired molarity.
2.
This process is called __________
3.
Remember:
***** Determine the volume of 16.0M HNO3 needed to prepare 2.50 L of 0.100 M HNO3 solution.
Precipitation Reactions
1.
When two solutions are mixed, an insoluble substance sometimes forms. Such a reaction is called
a ____________________.
2.
The solid that forms is called a ___________________.
3.
To write an equation for a precipitation reaction, one must know the identities of the reactants and
the products.
4.
In virtually every case, when a solid containing ions dissolves in water, the ions separate and
move around independently.
***** Look at the reaction between lead (II) nitrate and potassium chloride:
A more accurate representation is:
5.
Now the chemist must __________ the products. This is hard to do. Only after identifying each
product ____________________ is the chemist sure what reaction has taken place.
6.
Some things help us predict products.
a.
When ions form a solid compound, the compound must have zero net charge.
b.
Most ionic materials contain only two types of ions – one type of cation and one type of anion.
***** From our earlier example, the possible product combinations are:
Since two possibilities are the reactants in solution, the possible products are:
OVER
7.
In order to predict products and determine whether a reaction takes place, one must remember the
solubility rules. The rules help one determine which ionic species will form a solid or remain in
solution.
***** According to the solubility rules, PbCl2 will form a precipitate because it is insoluble in water. Our
reaction becomes:
After balancing we have:
8.
The key to dealing with the chemistry of an aqueous solution is first to focus on the actual
components of the solution before any reaction occurs and then to figure out how these
components will react with each other.
***** Using the solubility rules, predict what will happen when each of the following pairs of solutions are
mixed:
a.
BaCl2 (aq) and Na2SO4 (aq)
b.
AgNO3 (aq) and Na3PO4 (aq)
c.
NaOH (aq) and Fe(NO3)3 (aq)
Describing Reactions in Solution
1.
The ___________________________________ gives the overall reaction stoichiometry but not
necessarily the actual forms of the reactants and products in solution.
***** From our previous example:
2.
The ______________________________ represents as ions all reactants and products that are
strong electrolytes.
3.
The _________________________ includes only those solution components undergoing a
change. Spectator ions are not included.
Stoichiometry of Precipitation Reactions
1.
The same stoichiometric principles apply.
2.
It is sometimes difficult to tell immediately what reaction will occur when two solutions are
mixed.
3.
To obtain the moles of reactants we must use the volume of the solution and its molarity.
***** What volume of 0.100M Na3PO4 is required to precipitate all the lead(II) ions from 150.0mL of
0.250 M Pb(NO3)2 ?
OVER
***** What mass of barium sulfate is produced when 100.0 mL of a 0.100 M solution of barium chloride is
mixed with 100.0 mL of a 0.100 M solution of iron (III) sulfate?
***** A 100.0 mL aliquot of 0.200 M aqueous potassium hydroxide is mixed with 100.0 mL of 0.200 M
aqueous magnesium nitrate.
a. Write a balanced chemical equation for any reaction that occurs.
b. What precipitate forms?
c. What mass of precipitate is produced?
d. Calculate the concentration of each ion remaining in solution after precipitation is complete.
ACID BASE REACTIONS
1. The Arrhenius definition of an acid and a base is that an acid produces _______ ions and a base
produces _______ ions.
2. A more generalized, and useful, definition was provided by Johannes Bronsted and Thomas Lowry. The
Bronsted-Lowry definition of an acid is that it is a substance that is a ______________________________.
A base is a ______________________________.
3. Predicting the results of an acid-base reaction can be difficult.
4. When aqueous hydrogen chloride mixes with aqueous sodium hydroxide, the combined solution
contains:
OVER
5. This is because HCl is a strong acid and NaOH is a strong base. Both will dissociate completely.
6. The recombination produces _________, which is soluble so it remains _____ and _____. These
species are spectator ions.
7. The _____ and the _____ cannot coexist in solution because water is a nonelectrolyte.
8. If one mixes acetic acid, a weak acid, with an aqueous solution of potassium hydroxide, the species
present are:
9. Because the OH- has such a strong affinity for protons, it can strip them from the HC 2H3O2:
10. The hydroxide ion is such a strong base that for purposes of stoichiometric calculations it can be
assumed to react completely with any weak acid it encounters.
11. An acid-base reaction is often called a ______________________________. When just enough base
is added with the acid in a solution, we say the acid has been neutralized.
***** What volume of 0.250 M KOH will react completely with 25.00 mL of 0.200 M HCl?
***** Hydrochloric acid (75.0 mL of 0.250 M) is added to 225.0 mL 0f 0.550 M Ba(OH)2 solution. What
is the concentration of the excess H+ or OH- left in the solution?
Acid Base Titrations
1. ______________________________ is a technique for determining the amount of a certain substance
by doing a titration.
2. A _________________________ involves delivery of a measured volume of a solution of known
concentration, ____________________, into a solution containing the substance being analyzed,
____________________.
3. The titrant contains a substance that reacts in a known manner with the analyte.
4. The point in the titration where enough titrant has been added to react exactly with the analyte is called
the ______________________________ or ______________________________.
OVER
5. This point is often marked by an ____________________, a substance added at the beginning of the
titration that changes color at (or near) the equivalence point.
6. The point where the indicator actually changes color is called the ____________________ of the
titration.
7. When the analyte is a base or an acid, the required titrant is a strong acid or strong base, respectively.
This procedure is called an ___________________________________.
OXIDATION – REDUCTION REACTIONS
1. 2 Na (s) + Cl2 (g) → 2 NaCl (s)
2. Both reactants have no charge, they are neutral. NaCl is an ionic compound containing Na + ions and Clions.
3. Reactions like this one, in which one or more electrons are transferred, are called
______________________________ reactions, or ____________________ reactions.
4. Most reactions used for energy production are redox reactions.
Oxidation States
1. The concept of oxidation states provides a way to keep track of electrons in redox reactions.
2. This is particularly useful for redox reactions involving covalent substance.
3. For a covalent bond between two identical atoms, the electrons are split equally between the two.
4. When two different atoms are involved, and the electrons are shared unequally, the electrons are
assigned to the atom that has the stronger attraction for the electrons.
5. Oxidation states are assigned according to the rules on page __________.
6. While uncommon, noninteger oxidation states can exist.
***** Assign oxidation states to all atoms in each of the following compounds:
UO22+
As2O3
HAsO2
Mg2P2O7
Characteristics of Oxidation – Reduction Reactions
1. In some cases the transfer of electrons is very obvious.
2. In other cases the transfer is less obvious.
3. ____________________ is the increase in oxidation state; a loss of electrons.
4. ____________________ is the decrease in oxidation state; a gain of electrons.
5. The ______________________________ is the ______________________________.
6. The ______________________________ is the ______________________________.
7. _____________ says _____________.
OVER
***** Identify the oxidizing agent, the reducing agent, the substance being oxidized, and the substance
being reduced.
Cu (s) + 2 Ag+ (aq) → 2 Ag (s) + Cu2+ (aq)
SiCl4 (l) + 2 Mg (s) → 2 MgCl2 (s) + Si (s)
BALANCING OXIDATION – REDUCTION EQUATIONS
1. Difficult to do by simple inspection.
Half – Reaction Method
1. Separate the reaction into two half-reactions, one involving oxidation and one involving reduction.
2. Balance them separately.
3. Add them together.
In Acidic Solution
1. Write separate equations for the oxidation and reduction half-reactions.
2. Balance all the elements except hydrogen and oxygen.
3. Balance oxygen using H2O.
4. Balance hydrogen using H+.
5. Balance the charge using electrons.
6. Multiply one, or both, balanced half-reactions by an integer to equalize the number of electrons
transferred in the two half-reactions.
7. Add the half-reactions and cancel identical species.
8. Check that the elements and charges are balanced.
***** Balance the following oxidation-reduction reaction that occurs in acidic solution using the halfreaction method.
Cr2O72- (aq) + Cl- (aq) → Cr3+ (aq) + Cl2 (g)
OVER
In Basic Solution
1. Use the half-reaction method specified for acidic solutions to obtain the final balanced equation as if H+
ions were present.
2. To both sides of the equation, add a number of OH - ions that is equal to the number of H+ ions.
3. Form H2O on the side containing both H+ and OH- ions, and eliminate the number of H 2O molecules
that appear on both sides of the equation.
***** Balance the following oxidation-reduction reaction that occurs in basic solution.
CN- (aq) + MnO4- (aq) → CNO- (aq) + MnO2 (s)
CHAPTER 4 HOMEWORK
1. Read the chapter and review your notes.
2. Answer the following questions; solve the following problems – remember answers without
supporting work will not receive credit.
1. Show how each of the following strong electrolytes dissociates upon dissolving in water:
(a) NaBr
(b) MgCl2
(c) (NH4)2SO4
(d) HNO3
(e) Al(NO3)3
2. Calculate the molarity of each of the following solutions.
(a) A 5.623g sample of NaHCO3 is dissolved in enough water to make a 250.0 mL solution.
(b) A 184.6 mg sample of K2Cr2O7 is dissolved in enough water to make a 500.0 mL solution.
OVER
(c) A 0.1025g sample of copper metal is dissolved in 35.0 mL of concentrated Cu2+ ions and then
water is added to make a total volume of 200.0 mL. Calculate the molarity of the Cu2+ ions.
3. Calculate the concentration of all ions present in each of the following solutions of strong
electrolytes.
(a) 0.15 M CaCl2
(b) 0.26 M Al(NO3)3
(c) 0.25 M K2Cr2O7
(d) 2.0 x 10-3 M Al2(SO4)3
4. Calculate the volume of a 0.100 M solution of NaHCO3 that contains 0.350 g of NaHCO3.
5. Describe how you would prepare 2.00 L of 0.250 M NaOH from solid NaOH.
6. Describe how you would prepare 1.00 L of 0.500 M HCl from a concentrated (12 M) stock
solution.
7. A solution is prepared by dissolving 10.8 g of ammonium sulfate in enough water to make a
100.0 mL stock solution. A 10.00 mL sample of this stock solution is added to 50.00 mL of
water. Calculate the concentrations of ammonium ions and sulfate ions in the final solution.
8. When the following solutions are mixed together, what precipitate (if any) will form?
(a) BaCl2 (aq) and Na2SO4 (aq)
(b) Pb(NO3)2 (aq) and KCl (aq)
(c) K2S (aq) and Ni(NO3)2 (aq)
OVER
9. Write complete, balanced molecular equations and net ionic equations for the reaction, if any,
which occurs when aqueous solutions of the following are mixed:
(a) silver nitrate and barium chloride
(b) iron (II) sulfate and potassium sulfide
(c) sodium hydroxide and potassium sulfate
(d) dimercury (I) nitrate and calcium chloride
10. What mass of solid aluminum hydroxide is produced when 50.0 mL of a 0.200 M solution of
Al(NO3)3 is added to 200.0 mL of 0.100 M KOH?
11. Determine the mass, in grams, of sodium carbonate (Na2CO3) required to completely react
with 25.0 mL of 0.155 M nitric acid (HNO3).
Na2CO3 (aq) + 2 HNO3 (aq) → 2 NaNO3 (aq) + CO2 (g) + H2O (l)
12. Calculate the mass of carbon dioxide produced when 75.0 mL of 0.350 M HCl reacts with
excess Na2CO3.
Na2CO3 (s) + 2 HCl (aq) → 2 NaCl (aq) + H2O (l) + CO2 (g)
13. You can dissolve an aluminum soft drink can in an aqueous base such as potassium
hydroxide:
2 Al (s) + 2 KOH (aq) + 6 H2O (l) → 2 KAl(OH)4 (aq) + 3 H2 (g)
If you place 2.05g of aluminum in a beaker with 185 mL of 1.35 M KOH, will any aluminum
remain? What mass, in grams, of KAl(OH)4 is produced?
OVER
14. How many grams of silver chloride can be prepared by the reaction of 100.0 mL of 0.200 M
silver nitrate with 100.0 mL of 0.150 M calcium chloride? Calculate the concentration of each ion
remaining in solution after the precipitation reaction is complete.
15. Determine the volume of 0.150 M HNO3 required to neutralize 50.00 mL of 0.200 M NaOH.
16. Determine the volume of 0.0200 M calcium hydroxide is required to neutralize 35.00 mL of
0.0500 M nitric acid?
17. Assign oxidation states to each of the atoms in the following ions/compounds:
(a) Cr(OH)4-
(b) K2SO4
(c) CH2Cl2
(d) H2PO4-
OVER
18. Balance the following oxidation-reduction reactions in an acidic solution:
(a) VO2+ (aq) + Zn (s) → VO2+ (aq) + Zn2+ (aq)
(b) Cr2O72- (aq) + Fe2+ (aq) → Cr3+ (aq) + Fe3+ (aq)
19. Balance the following oxidation-reduction reactions in a basic solution:
(a) CrO42- (aq) + SO32- (aq) → Cr(OH)3 (s) + SO42- (aq)
(b) NiO2 (s) + Zn (s) → Ni(OH)2 (s) + Zn(OH)2 (s)
OVER
AP CHEMISTRY CHAPTER 4 REVIEW
1. Determine whether the following compounds are soluble or insoluble in water:
a. Hg2Cl2
b. KI
c. Pb(NO3)2
d. NaBr
e. Ba(OH)2
f. CrCO3
2. Determine the precipitate that will form, if any, when the following solutions are mixed:
a. 2 Al(NO3)3 (aq) + 3 Ba(OH)2 (aq)
b. CuSO4 (aq) + HgCl2 (aq)
c. 3 AgNO3 (aq) + Na3PO4 (aq)
d. 3 NaOH (aq) + Fe(NO3)3 (aq)
3. Write net ionic equations for the reaction, if any, that occurs when aqueous solutions of the following
are mixed:
a. copper (II) nitrate and ammonium sulfate
b. silver nitrate and iron (III) chloride
c. ammonium phosphate and potassium carbonate
d. sodium hydroxide and manganese (II) nitrate
OVER
4. Determine the volume of 0.100M Na3PO4 required to precipitate all the lead (II) ions from 150.0 mL of
0.250M Pb(NO3)2.
5. Determine the mass of barium sulfate that is produced when 100.0 mL of 0.100 M BaCl 2 is mixed with
100.0 mL of 0.100 M Fe2(SO4)3.
6. Determine the molarity of a solution prepared by dissolving 11.85 g of KMnO 4 in enough water to make
750.0 mL of solution.
7. A 0.3025gsample of nickel metal is dissolved in 45.0 mL of concentrated nitric acid forming Ni 2+ ions
in solution. The solution is then diluted to a total volume of 600.0 mL. Determine the molarity of the Ni 2+.
8. Determine the volume of 0.850M Al(NO3)3 that contains 50.00g of Al(NO3)3.
9. Determine the volume of 18.0M H2SO4 solution needed to prepare 250.0 mL of
1.75M H2SO4.
10. Determine the concentration of a solution of NaOH after diluting 10.0 mL of 10.0 M stock solution to
250.0 mL.
OVER
11. Heme, obtained from red blood cells, binds oxygen, O 2. Calculate the number of moles of heme in
150.0 mL of 0.0019M heme solution.
12. Determine the volume of 0.100 M NaOH required to react completely with 25.0 mL
of 0.200 M H2SO4.
13. Determine the molarity of Fe3+ and SO42- ions in a solution prepared by dissolving 48.05g of Fe 2(SO4)3
in enough water to make a 800.0 mL solution.
14. The distinctive odor of vinegar is due to acetic acid, HC 2H3O2, which reacts with sodium hydroxide in
the following fashion:
HC2H3O2 (aq) + NaOH (aq) → H2O (l) + NaC2H3O2 (aq)
3.45 mL of vinegar requires 42.5 mL of 0.115 M NaOH to reach the equivalence point in a titration.
Determine the mass, in grams, of vinegar titrated by the NaOH solution.
15. A solution of 100.0 mL of 0.200 M KOH is mixed with a solution of 200.0 mL
of 0.150 M NiSO4 according to the following reaction:
2 KOH (aq) + NiSO4 (aq) → K2SO4 (aq) + Ni(OH)2 (s)
(a) Determine the mass, in grams, of Ni(OH) 2 that will precipitate.
(b) What is the identity of the limiting reactant?
(c) What is the identity of the excess reactant?
(d) Determine the concentration of each ion remaining in solution after the precipitation is complete.
OVER
16. Assign oxidation states to each of the atoms in the following compounds.
a. GaO3
b. KBrO4
c. NbO2
d. CrO3
e. Na3PO4
f. NH2-
g. IO3-
17. Balance the following oxidation-reduction reactions that occur in acidic solution.
a. Zn (s) + NO3- (aq) → Zn2+ (aq) + NH4+ (aq)
b. Br - (aq) + MnO4 - (aq) → Br2 (l) + Mn2+ (aq)
18. Balance the following oxidation-reduction reactions in basic solution.
a. MnO4- (aq) + SO32- (aq) → MnO2 (s) + SO42- (aq)
b. Co2+ (aq) + H2O2 (l) → Co(OH)3 (s) + H2O (l)
OVER