CHAPTER 5
Chemical Reactions
Objectives
You will be able to do the following.
1. Given a description of a solution of two components, identify the solute and the solvent.
2. Write a description of the process for dissolving an ionic compound in water. Your
description should include mention of the nature of the particles in solution and the
attractions between the particles in the solution.
3. Given the formula for an ionic compound, predict whether it is soluble in water or not.
4. Given formulas for two ionic compounds, (1) predict whether a precipitate will form when
the water solutions of the two are mixed, (2) if there is a reaction, predict the products of
the reaction and write their formulas, and (3) if there is a reaction, write the complete,
complete ionic, and net ionic equations that describe the reaction.
5. Describe precipitation reactions. Your descriptions should include mention of the nature
of the particles in the system before and after the reaction, a description of the cause of
the reaction, and a description of the attractions between the particles before and after the
reaction.
6. Write a description of the process for the solution of a strong monoprotic acid, sulfuric
acid, and a weak monoprotic acid in water.
7. Identify the following as strong monoprotic acids: HCl, HBr, HI, HNO3, and HClO4.
8. Identify sulfuric acid, H2SO4, as a strong diprotic acid.
9. Given a formula for any acid, identify it as a strong or weak acid.
10. Write a description of the process for the solution of ammonia, NH3, in water.
11. Given the formula or name of a compound, classify it as a strong electrolyte, weak
electrolyte, or nonelectrolyte.
12. Given a formula for an ionic compound, identify it as representing a weak base, a strong
base, or a substance that is neutral in the Arrhenius acid-base sense.
13. Identify the following as anions that are neutral in the acid-base sense: Cl−, Br−, I−,
NO3−, and ClO4−.
14. Identify the following anions as weak acids: hydrogen sulfate ion, HSO4−, and dihydrogen
phosphate ion, H2PO4−.
15. Identify ammonia as an uncharged, weak base.
16. Identify ionic compounds containing hydroxide ions as strong bases.
17. Given the pH of a solution, identify the solution as acidic, basic, or neutral.
18. Given the pH of two acidic solutions, identify which solution is more acidic.
19. Given the pH of two basic solutions, identify which solution is more basic.
20. Describe how litmus paper can be used in the laboratory to identify whether a solution is
acidic or basic.
21. Write a description of the process that takes place at the molecular level when a strong acid
reacts with an aqueous strong base. Your descriptions should include mention of the nature
of the particles in the system before and after the reaction, a description of the cause of
the reaction, and a description of the attractions between the particles before and after the
reaction.
73
74
Chapter 5
Chemical Reactions
22. Given the formulas for an acid and a base, predict the products that would form
from the reaction between them, and write the balanced complete, complete
ionic, and net ionic equations that describe the reaction.
23. Identify H2O(l ) and CO2(g) as the products of the reaction of an acid with
carbonate, CO32−, or hydrogen carbonate, HCO3−.
24. Write an explanation for why a substance can be a Brønsted-Lowry acid in one
reaction and a Brønsted-Lowry base in a different reaction. Give an example to
illustrate your explanation.
25. Write an explanation for why the Arrhenius definitions for acid and base and
not the Brønsted-Lowry definitions are used to describe whether an isolated
substance is an acid or base.
26. Write an explanation for why it is useful to have two sets of definitions for acids
and bases (the Arrhenius definitions and the Brønsted-Lowry definitions).
27. Given a Brønsted-Lowry neutralization equation, identify the Brønsted-Lowry
acid and Brønsted-Lowry base.
28. Given an equation for a chemical reaction, (1) identify whether the equation
represents a redox reaction or not, (2) identify the substance that is oxidized and
the substance that is reduced, and (3) identify the substance that is the reducing
agent and the substance that is the oxidizing agent.
29. Given the mass of substance dissolved in a solution and the total volume of the
solution, calculate the molarity of the solution.
30. Convert between moles of solute and volume of solution using molarity as a
conversion factor.
31. Convert between amount of one substance and amount of another substance,
both involved in a chemical reaction, when the amounts of these substances are
described by (1) mass of pure substance, (2) mass of a mixture that contains one
of the substances in the reaction with the percentage of that substance described,
or (3) volume of a solution that contains one of the substances with the molarity
of the solution given.
32. Given the volume of a substance titrated, the molarity of the titrant used in the
titration, and the volume of titrant necessary to reach the equivalence point in
the titration, calculate the molarity of the substance titrated.
33. Given a volume and molarity for a solution, write a description of how the
solution could be made from pure solid and water.
34. Write a description of why when making a solution from solid and water, the
solid must be dissolved in water before more water is added to bring the volume
to the desired total.
35. Write an explanation for why pure or almost pure acid should always be added to
water and not water to pure or almost pure acid.
36. Given the volume and molarity of a dilute solution and the molarity of a more
concentrated solution used to make the dilute solution, calculate the volume of
the concentrated solution necessary to make the dilute solution.
37. Write a description of the procedure for making a more dilute solution of an acid
from a pure or almost pure acid (e.g. 18 M H2SO4).
38. Given three of the following four, calculate the fourth: volume of a more dilute
solution, molarity of the more dilute solution, volume of a more concentrated
solution of the same substance, and molarity of the more concentrated solution.
75
39. Given the volume and molarity of a more dilute solution of a substance and the
molarity of a more concentrated solution of the same substance (other than a
pure or almost pure acid) used to make the dilute solution, write a description of
the procedure for making the more dilute solution from the more concentrated
solution.
40. Convert between the definition and the term for the following words or phrases.
Skip Section 5.5 of the text.
Chapter 5 Glossary
Solution A mixture whose particles are so evenly distributed that the relative
concentrations of the components are the same throughout. Solutions can also be
called homogeneous mixtures.
Aqueous solution A solution in which water is the solvent.
Solute The gas in a solution of a gas in a liquid. The solid in a solution of a solid in a
liquid. The minor component in other solutions.
Solvent The liquid in a solution of a gas in a liquid. The liquid in a solution of a solid
in a liquid. The major component in other solutions.
Hydrated Bound to one or more water molecules.
Hydration The binding of one or more water molecules to an ion or molecule.
Electrolyte A substance that ionizes or dissociates in water to form an electrically
conducting solution.
Strong electrolyte A substance that ionizes or dissociates completely in an aqueous
solution.
Nonelectrolyte A substance that ionizes or dissociates incompletely in an aqueous
solution.
Ionize To form ions (often as a substance dissolves in water).
Dissociate The separation of ions (often as a substance dissolves in water).
Precipitation reaction A reaction in which one of the products is insoluble in water
and comes out of solution as a solid.
Precipitate A solid that comes out of solution.
Precipitation The process of forming a solid in a solution.
Crystals Solid particles whose component atoms, ions, or molecules are arranged in
an organized, repeating pattern.
Complete ionic equation A chemical equation that describes the actual form for
each substance in solution. For example, ionic compounds that are dissolved in
water are described as separate ions.
Spectator ions Ions that play a role in delivering other ions into solution to react but
that do not actively participate in the reaction themselves.
Complete equation or molecular equation A chemical equation that includes
uncharged formulas for all of the reactants and products. The formulas include the
spectator ions, if any.
Net ionic equation A chemical equation for which the spectator ions have been
eliminated, leaving only the substances actively involved in the reaction.
Solubility The maximum amount of solute that can be dissolved in a given amount
of solvent.
Hydronium ion H3O+
76
Chapter 5
Chemical Reactions
Arrhenius acid According to the Arrhenius theory, any substance that generates hydronium
ions, H3O+, when added to water.
Monoprotic acid An acid that donates one hydrogen ion per molecule in a reaction.
Polyprotic acid An acid that can donate more than one hydrogen ion per molecule in a
reaction.
Diprotic acid An acid that can donate two hydrogen ions per molecule in a reaction.
Triprotic acid An acid that can donate three hydrogen ions per molecule in a reaction.
Strong acid An acid that donates its H+ ions to water in a reaction that goes completely
to products. Such a compound produces close to one H3O+ ion in solution for each
acid molecule dissolved in water.
Completion reaction A reaction that shifts completely to products, that is, a reaction that
is not significantly reversible.
Reversible reaction A reaction in which the reactants are constantly forming products
and, at the same time, the products are reforming the reactants.
Weak acid A substance that is incompletely ionized in water due to the reversibility of the
reaction that forms hydronium ions, H3O+, in water. Weak acids yield significantly less
than one H3O+ ion in solution for each acid molecule dissolved in water.
Strong monoprotic acids HCl(aq), HBr(aq), HI(aq), HNO3, HClO4.
Strong diprotic acids H2SO4.
Arrhenius base A substance that produces hydroxide ions, OH−, when added to water.
Strong base A substance that generates at least one hydroxide ion in solution for every
unit of substance added to water.
Weak base A substance that produces fewer hydroxide ions in water solution than particles
of the substance added.
Miscible Able to be mixed in any proportion, that is, infinitely soluble.
Acidic solution A solution with a significant concentration of hydronium ions, H3O+.
Basic solution A solution with a significant concentration of hydroxide ions, OH−.
Neutralization reaction A chemical reaction between an acid and a base.
Brønsted-Lowry acid-base reaction A chemical reaction in which a proton, H+, is
transferred.
Brønsted-Lowry Acid A substance that donates protons, H+, in a Brønsted-Lowry
acid-base reaction.
Brønsted-Lowry Base A substance that accepts protons, H+, in a Brønsted-Lowry
acid-base reaction.
Amphoteric substance A substance that can act as either a Brønsted-Lowry acid and a
Brønsted-Lowry base, depending on the circumstances.
Oxidation Any chemical change in which at least one element loses electrons, either
completely or partially.
Reduction Any chemical change in which at least one element gains electrons, either
completely or partially.
Oxidation-reduction reactions The chemical reactions in which there is a complete or
partial transfer of electrons, resulting in oxidation and reduction. These reactions are
also called redox reactions.
Half-reactions Separate oxidation and reduction reaction equations in which electrons
are shown as a reactant or product.
Reducing agent A substance that loses electrons, making it possible for another substance
to gain electrons and be reduced.
77
Oxidizing agent
A substance that gains electrons, making it possible for another
substance to lose electrons and be oxidized.
Oxidation number A tool for keeping track of the flow of electrons in redox reactions
(also called oxidation state).
Figure 5.1
Sodium Chloride Dissolving in Water
This shows the mixture immediately after sodium chloride has been added to water.
Certain water molecules are highlighted to draw attention to their role in the process.
78
Chapter 5
Chemical Reactions
Figure 5.2
Aqueous Sodium Chloride
This image shows a portion of the solution that forms when sodium chloride
dissolves in water. Certain water molecules are highlighted to draw attention to their role in the process.
You can find an animation that shows the process by which sodium chloride dissolves at the following web
address:
http://www.mpcfaculty.net/mark_bishop/NaCl_dissolves.htm
Figure 5.3
Aqueous Calcium Nitrate
Notice that there are twice as many −1 nitrate ions as +2 calcium ions.
79
Figure 5.4
Mixture of Ca(NO3)2(aq) and Na2CO3(aq) at the Instant They Are Combined
Figure 5.5
Product Mixture for the Ca(NO3)2(aq) and Na2CO3(aq) Reaction
80
Chapter 5
Chemical Reactions
You can find an animation that shows this precipitation reaction at the following web
address:
http://www.mpcfaculty.net/mark_bishop/precipitation.htm
Table 5.1 Ionic Solubility
Category
Ions
Except with these ions
Examples
Soluble
Cations
Group 1A metal ions
and ammonium, NH4+
All soluble
Na2CO3, LiOH, and (NH4)2S are
soluble
Soluble
Anions
NO3− and C2H3O2−,
All soluble
Bi(NO3)3, and Co(C2H3O2)2 are
soluble
Usually
Soluble
Anions
Cl−, Br−, and I−
Soluble except with Ag+
and Pb2+
CuCl2 is water soluble,
but AgCl is insoluble.
SO42−
Soluble except with Ba2+
and Pb2+
FeSO4 is water soluble,
but BaSO4 is insoluble.
CO32−, SO32−, PO43−,
and OH−
Insoluble except with group
1A elements and NH4+
CaCO3, ZnSO3, Ca3(PO4)2, and
Mn(OH)2 are insoluble in water,
but (NH4)2CO3, K2SO3, Li3PO4,
and CsOH are soluble.
S2−
Insoluble except with group
1A and 2A metal ions and
NH4+
CoS is insoluble in water,
but MgS is soluble.
Usually
Insoluble
Anions
EXERCISE 5.1 - Ionic solubility
Predict whether each of the following are soluble or insoluble in water.
a. Hg(NO3)2
b. FeCO3
c. SnS
d. K3PO4
e. PbCl2
f. Al(OH)3
81
TIP-OFF You are asked to predict whether a precipitation reaction will take place
between two aqueous solutions of ionic compounds, and if the answer is yes, to write
the complete equation for the reaction.
GENERAL STEPS
STEP 1 Determine the formulas for the possible products using the general
double-displacement equation. (Remember to consider ion charges when
writing your formulas.)
AB + CD → AD + CB
STEP 2 Predict whether either of the possible products is water insoluble. If
either possible product is insoluble, a precipitation reaction takes place, and you
may continue with Step 3. If neither is insoluble, write “No reaction”.
STEP 3 Follow these steps to write the complete equation.
•
Write the formulas for the reactants separated by a “+”.
•
Separate the formulas for the reactants and products with a single arrow.
•
Write the formulas for the products separated by a “+”.
•
Write the physical state for each formula.
The insoluble product will be followed by (s).
Water-soluble ionic compounds will be followed by (aq).
•
Balance the equation.
EXERCISE 5.2 - Precipitation Reactions
Predict whether a precipitate will form when each of the following pairs of water
solutions are mixed. If there is a precipitation reaction, write the complete, complete
ionic, and net ionic equations that describe the reaction.
a. Li2CO3(aq) + Al(NO3)3(aq)
b. KOH(aq) + Fe(NO3)3(aq)
c. NaC2H3O2(aq) + CaS(aq)
d. K2SO4(aq) + Pb(NO3)2(aq)
Sample Study
Sheet 5.1:
Predicting
Precipitation
Reactions
and Writing
Precipitation
Equations
82
Chapter 5
Chemical Reactions
You can find a description of the procedure for writing complete and net ionic
equations at the following web address:
http://www.mpcfaculty.net/mark_bishop/precipitation_equations.htm
Having Trouble?
Are you having trouble with this chapter? People often do. To successfully complete
each of the tasks in this chapter, you need to have mastered the skills from previous
sections. If you missed anything along the way, you will have trouble. Here is a list of
the tasks you need to be able to do in order to work the problems in this chapter. You
should go through the list in order and be sure you have mastered each skill before you
go on to the next one.
1. Convert between names and symbols for the common elements.
2. Identify whether an element is a metal or a nonmetal.
3. Determine the charges on many of the monatomic ions.
4. Convert between the name and formula for polyatomic ions.
5. Recognize a name or formula as ionic or molecular.
6. Recognize a name or formula as representing a binary acid or an oxyacid.
7. Convert between the name and formula for ionic compounds, binary acids
and oxyacids.
8. Balance chemical equations.
9. Predict the products of double displacement reactions.
10. Predict ionic solubility.
11. Predict the states of ionic compounds and H2O.
12. Identify strong, weak, and nonelectrolytes.
13. Write complete ionic and net ionic equations.
83
Figure 5.6
Hydrochloric Acid in Water
Figure 5.7
Acetic Acid in Water
84
Chapter 5
Chemical Reactions
You can recognize strong and weak acids by using the following criteria.
♦ You can recognize acids either from their formula or their name.
∗ Acid formulas can be either HX(aq) like HCl(aq), H2S(aq), or
HaXbOc, such as HNO2.
∗ Acid names end in acid, such as hydrochloric acid, hydrosulfuric acid
and nitrous acid.
♦ The strong monoprotic acids you should know are HCl, HBr, HI, HNO3
and HClO4.
♦ Sulfuric acid, H2SO4, is a strong diprotic acid.
You will recognize a substance as a weak acid if it is not on the list of strong acids.
Figure 5.8
Ammonia in Water
85
Table 5.2 Classifying Arrhenius acids
Classification
of Arrhenius
Acids
Monoprotic
Diprotic
Triprotic
Weak or
Strong
Characteristics
How to Recognize
Strong
Completely
ionized in water
Memorize HCl(aq), HBr(aq), HI(aq),
HNO3, and HClO4
Weak
Incompletely
ionized in water
General formula of HX(aq) or HXbOc
and not on the list of strong acids*
Strong
First H+ ion
lost completely:
second
incompletely
H2SO4
Weak
Both H+ ions lost
incompletely
H2S(aq) or General formula of H2XbOc
(not H2SO4)
Weak
All three H+ ions
lost incompletely
General formula H3XbOc
*Some acids have formulas that do not fit our general formulas. Acetic acid, HC2H3O2,
is the only example that you need to recognize at this point.
Table 5.3 Classifying Arrhenius bases
Arrhenius Bases
Strong or
Weak
Characteristics
How recognize?
Anions
strong
Each unit added leads to
at least one OH− ion in
solution
Soluble Metal hydroxides*
(in ionic
compounds)
weak
Reversible reaction with
water to yield fewer OH−
ions than units added
Anions in ionic compounds except
with OH−, Cl−, Br−, I−, NO3−,
ClO4−, HSO4−, and H2PO4−
Some uncharged
molecules
weak
Reversible reaction with
water to yield fewer OH−
ions than molecules added
NH3
* Note: Since a metal hydroxide must be soluble in water to yield a significant
concentration of hydroxide in solution, it is common to consider only water soluble
metal hydroxides like NaOH to be strong bases. Insoluble metal hydroxides like
aluminum hydroxide, Al(OH)3, react like the soluble hydroxides in neutralization
reactions, so for some purposes they can be considered strong bases.
86
Chapter 5
Sample Study
Sheet 5.2:
Identification
of Strong And
Weak Acids
and Bases
Chemical Reactions
TIP-OFF You can use the following steps to identify a name or chemical formula as
representing either (1) an Arrhenius strong acid, (2) an Arrhenius weak acid, (3) an
Arrhenius strong base, (4) an Arrhenius weak base, or (5) not acidic or basic in the
Arrhenius sense (neutral).
GENERAL STEPS
STEP 1 Identify the substance as an acid, a base, or neither. You can identify
acids in the following ways.
• If you are given a name:
a. the names of the uncharged acids end in acid.
For example, hydrochloric acid is an acid.
b. the names for the only ionic compounds you are expected to
recognize as acidic end in hydrogen sulfate or dihydrogen phosphate.
Sodium hydrogen sulfate is acidic.
• If you are given a molecular formula:
a. Remember that acids have one of these forms: HX(aq) or HaXbOc.
HCl(aq) and H2SO4 are acids.
b. Acidic ionic compounds have formulas that include HSO4− or
H2PO4−.
NaH2PO4 is acidic.
You can identify bases in the following ways.
• We expect ionic compounds to be basic, except those containing the
following anions.
a. Cl−, Br−, I−, NO3−, ClO4− - neutral
b. HSO4−, H2PO4− - acidic
NaF is basic, NaCl is neutral, and NaHSO4 is acidic.
• Ammonia, NH3, is a base.
STEP 2 If you have an acid or base, determine whether it is strong or weak.
♦ We will consider all acids except HCl(aq), HBr(aq), HI(aq), HNO3,
HClO4, and H2SO4 to be weak.
HF is a weak acid.
♦ We will consider all bases except metal hydroxides to be weak.
NaF is a weak base.
EXERCISE 5.3 - Acid and Base Classification
Identify HNO2, lithium hydroxide, NaCN, sodium iodide, NaHSO4, nitric acid,
CH3OH, hydrofluoric acid, and KC2H3O2 as either (1) an Arrhenius strong acid, (2)
an Arrhenius weak acid, (3) an Arrhenius strong base, (4) an Arrhenius weak base, or
(5) not acidic or basic in the Arrhenius sense (neutral).
87
The following are characteristics of acids.
1. Acids have a sour taste.
2. Acids turn litmus from blue to red.
3. Acids react with bases. (When the base includes carbonate or hydrogen
carbonate ions, carbon dioxide gas is released.)
The following are characteristics of bases.
1. Bases have a bitter taste.
2. Bases feel slippery on your fingers.
3. Bases turn litmus from red to blue.
4. Bases react with acids.
�
�
�
���
��
�� �
�
�
�
�� ��
���
�
���
�� �
��
���
�
�
�
�
�
�� � ���
�
�
�
�
�
�
�� �
� � �� ���
�
��
��
�
�
��
��
���
�
�
�
�� ���� ����
�
�
�� ��
��
���
�
�
�
�
�������
�����������
�������������������������
��������������������
��
��
��
��
��
��
��
����������
������������������������
�����������������������
Figure 5.9
pH of Common Substances
Table 5.4 Strong, weak and nonelectrolytes
Ions in
solution?
Completely
ionized?
Types
Strong Electrolytes
Yes
Yes
Strong acids
Water soluble ionic compounds
Weak Electrolytes
Yes
No
Uncharged weak acids
Ammonia
Nonelectrolytes
No
Not ionized at all
Alcohols
Sugars
88
Chapter 5
Chemical Reactions
Table 5.5 Identification of Strong, Weak, and Nonelectrolytes
Category
General
Type of
Substance
How to
Recognize
Formula
How to Recognize Name
Examples
Strong
electrolytes
Ionic(aq)
MaAb
MaHbXcOd or
(NH4)aAb
(NH4)aHbXcOd
metal (root nonmetal)ide,
metal polyatomic ion,
ammonium (root
nonmetal)ide,
or ammonium polyatomic
ion
NaCl
or Li2HPO4
or NaNO3
or NH4NO3
Strong
acids
HCl(aq),
HBr(aq),
HI(aq), HNO3,
HClO4, and
H2SO4
hydrochloric acid,
hydrobromic acid,
hydroiodic acid,
nitric acid, perchloric acid,
and sulfuric acid
All are listed to
the left.
Uncharged
weak acids
HX(aq),
H2S(aq), or
HaXbOc
and not on the
list of strong
acids
hydrosulfuric acid
per(root)ic acid
(root)ic acid
(root)ous acid
hypo(root)ous acid
and not on the list of
strong acids
hydrosulfuric
acid
periodic acid
chloric acid
chlorous acid
hypochlorous
acid
NH3
NH3
ammonia
Alcohol
CaHbOH
name ends in -anol or
alcohol
methanol
or methyl
alcohol
Sugar
C6H12O6
or C12H22O11
name ends in -ose
glucose or
sucrose
Weak electrolytes
Nonelectrolytes
M=symbol of metal, A=symbol of nonmetal, X=some element other than H or O, and a, b, c & d indicate
subscripts
EXERCISE 5.4 - Strong, Weak, and Nonelectrolytes
Identify Al(NO3)3, acetic acid, NH3, ammonium acetate, HCl, glucose, CH3OH,
barium chloride, K2Cr2O7, nitric acid, HBrO2, 2–propanol (or isopropyl alcohol)
as either a (1) strong electrolyte, (2) weak electrolyte, or (3) nonelectrolyte, and in
parentheses write the type of compound each name or formula represents. Each
dissolves in water.
89
Figure 5.10
Aqueous Nitric Acid
Figure 5.11
Water Solution of Nitric Acid and Sodium Hydroxide before Reaction
90
Chapter 5
Chemical Reactions
Figure 5.12
After Reaction of Nitric Acid and Sodium Hydroxide
Table 5.6 Prediction of whether neutralization reactions are completion reactions or
reversible reactions
Strong Acid
Weak Acid
Strong Base
Completion
Completion
Weak Base
Completion
Reversible or Completion*
*When you need to know whether a neutralization reaction is a completion reaction
or reversible, you will be told.
Sample Study
Sheet 5.3:
Steps for
Writing
Neutralization
Equations
TIP-OFF You will be given formulas or names of two substances and asked to predict
whether a neutralization reaction will take place between them. If it does, you will be
asked to write the complete, complete ionic, and net ionic equations for the reaction.
The following steps allow you to do this.
GENERAL STEPS
STEP 1 Ask, “Do you have an acid and a base?” If yes, go to Step 2. If no, say “No
reaction”. See Previous Study Sheets.
Any one of the following conditions will lead you to predict “no reaction”.
1) One of the substances is neutral in the Arrhenius acid-base sense.
2) Both substances given are acids.
3) Both substances given are bases.
91
STEP 2 Describe a completion reaction with a single arrow, (→). Describe a reversible
reaction with a double arrow, ( ). Determine whether the reaction is reversible or
goes to completion by applying the following guidelines.
a. When one or both of the acid and base are strong, the reaction is a completion
reaction. See Previous Study Sheets.
b. If you have a weak acid and a weak base, you can assume that the reaction is
reversible.
STEP 3 Write the formulas and states of the products.
a. You must first decide on the correct form of the equation.
1) Most neutralization reactions that you will see in this chapter will be
double displacement. If the base is ammonia, NH3, the reaction will
have the second form.
Double Displacement
AB + CD → AD + CB
NH3(aq) + HX(aq) → NH4X(aq)
2) For double displacement reactions, follow these steps.
a. Identify the A, B, C and D.
(1) For acids the positive component is H+.
(2) Split ionic compounds into cation and anion.
b. Write the AD and CB formulas. Be careful to balance charges.
c. Remember that carbonic acid, H2CO3(aq), decomposes to form
CO2(g) and H2O(l ).
d. We will assume that all of the H+ ions react for polyprotic acids,
e.g. H2SO3 loses two H+ ions, and H3PO4 loses three H+ ions.
b. Write the states of reactants and products.
1. The states are usually given for the original reactants.
2. For the ionic products:
1). Most ionic products of neutralization reactions are water soluble
and described with (aq).
2). If they are insoluble, describe them with (s).
3. Describe acids as aqueous, (aq), unless stated otherwise.
4. Ammonia is aqueous, (aq), when in water and gaseous, (g), when pure.
5. Carbon dioxide is insoluble in water and described as a gas, (g).
6. Water is a liquid, (l ).
STEP 4 Balance the complete equation.
92
Chapter 5
Chemical Reactions
STEP 5 Write the complete ionic and net ionic equations.
a. To get the complete ionic equation, describe each reactant and product as ions
or with a complete (not ionized) formula.
1. Strong acids and ionic(aq) are strong electrolytes and completely ionized.
a. For this task, describe strong acids, HA(aq), as H+(aq) and A−(aq).
(The A−(aq) represents an anion)
b. Even though describing strong acids, HA(aq), as H3O+(aq) and
A−(aq) is a better description of the nature of the ions in solution, it
is probably more common to use H+(aq) and A−(aq) for the strong
acids.
2. Everything else is described with a complete (not ionized) formula.
b. To get the net ionic equation, eliminate spectator ions. Remember that
these are the ions that are in an identical form on both sides of the equation.
Rewrite what is left and balance.
c. You can check that you have written the net ionic equation correctly by
making sure the following are true.
1. Strong monoprotic acids are described as H+(aq) in a net ionic equation.
(Even though H3O+(aq) provides a better description of the ions in
solution, it is probably more common to use H+(aq).)
2. The strong diprotic acid sulfuric acid, H2SO4, is described as H+ and
HSO4−.
3. Weak, uncharged acids, like acetic acid, HC2H3O2(aq), are described
with complete formulas.
a. A very common mistake is to describe the weak acids like the strong
acids, as H+ in the net ionic equation.
b. Ionic formulas that contain the anions that are weak acids are
described as ions. For instance, NaHSO4(aq) is described as
Na+(aq) and HSO4− (aq).
4. Any ions in strong electrolytes on both sides of the equation are
spectator ions and should be eliminated in a net ionic equation.
5. Formulas with (g), (l ), and (s) are described with complete formulas.
93
EXERCISE 5.5 - Writing Neutralization Equations
For each of the following pairs of possible reactants, predict whether a neutralization
reaction will take place between them. If there is no reaction, write, “No Reaction”.
If there is a reaction, write complete, complete ionic, and net ionic equations for the
reaction. (The reactions between weak acids and weak bases given here are reversible
reactions. If an acid has more than one acidic hydrogen, assume that there is enough
base to remove all of them. Assume that there is enough acid to add as many protons
to the base as possible)
a. HCl(aq) + NaOH(aq)
b. HF(aq) + LiOH(aq)
c. NH3(aq) + HNO3(aq)
d. NH3(aq) + HClO(aq)
e. HC2H3O2(aq) + LiF(aq)
f. Na2CO3(aq) + HBr(aq)
g. HCl(aq) + HNO2(aq)
h. H3PO3(aq) + LiOH(aq)
i. Fe(OH)3(s) + HNO3(aq)
j. NaI(aq) + HCl(aq)
94
Chapter 5
Chemical Reactions
���������
��������������
���������
��������������
�������������������������������������������������������������������������������������������������
�
�
�������������������������������������������������������������������������������������������������
H+
H+
Figure 5.13
Brønsted-Lowry Conjugate Acid-Base Pairs
EXERCISE 5.6 - Brønsted-Lowry Acids and Bases
For each of the following equations, identify the Brønsted-Lowry acid and base.
a. HNO2(aq) + NaBrO(aq) → HBrO(aq) + NaNO2(aq)
b. H2PO4−(aq) + HNO2(aq)
H3PO4(aq) + NO2−(aq)
c. H2PO4−(aq) + 2OH−(aq) → PO43−(aq) + 2H2O(l)
d. H2SO3(aq) + 2NaOH(aq) → Na2SO3(aq) + 2H2O(l)
Table 5.7 Key definitions for Redox Reactions
Term
Definition
oxidation-reduction
reaction
An electron transfer reaction
oxidation
Loss of electrons
reduction
Gain of electrons
reducing agent
The substance that loses electrons (is oxidized) and makes it
possible for something else to gain electrons (be reduced)
oxidizing agent
The substance that gains electrons (is reduced) and makes it
possible for something else to lose electrons (be oxidized)
half-reaction
One-half of a redox reaction; just the oxidation or just the
reduction
95
TIP-OFF You are asked to determine the oxidation number of an atom, or you need to Sample Study
assign oxidation numbers to atoms to determine whether a reaction is a redox reaction, Sheet 5.4:
and if it is, to identify which element is oxidized, which is reduced, what the oxidizing
Assignment
agent is, and what the reducing agent is.
GENERAL STEPS
Use the following guidelines to assign oxidation numbers to as many atoms
as you can. (The following table provides a summary of these guidelines with
examples.)
of Oxidation
Numbers
• The oxidation number for each atom in a pure element is zero.
• The oxidation number of a monatomic ion is equal to its charge.
• When fluorine atoms are combined with atoms of other elements, their
oxidation number is −1.
• When oxygen atoms are combined with atoms of other elements, their
oxidation number is −2, except in peroxides, like hydrogen peroxide,
H2O2, where their oxidation number is −1.
• The oxidation number for each hydrogen atom in a molecular compound
or a polyatomic ion is +1.
If a compound’s formula contains one element for which you cannot assign
an oxidation number using the guidelines listed above, calculate the oxidation
number according to the following rules.
• The sum of the oxidation numbers for the atoms in an uncharged
formula is equal to zero.
• The sum of the oxidation numbers for the atoms in a polyatomic ion is
equal to the overall charge on the ion.
Table 5.8 Oxidation Numbers for Some Elements
Oxidation
Number
Examples
Exceptions
Pure, uncharged
element
0
Each atom is 0 in
Zn, H2, P4, and Cl2
None
Monatomic ions
charge on
ion
Cd in CdCl2 is +2.
Cl in CdCl2 is −1.
H in LiH is −1.
None
Fluorine in the
combined form
−1
F in AlF3 is −1.
F in CF4 is −1.
None
Oxygen in the
combined form
−2
O in ZnO is −2.
O in H2O is −2.
O is −1 in peroxides, such as
H2O2 and O22−.
Covalently
bonded hydrogen
+1
H in H2O is +1.
None
96
Chapter 5
Chemical Reactions
Equations for redox reactions can be difficult to balance, but your ability to
determine oxidation numbers can help. You can find a description of the process for
balancing redox equations at the following web address.
www.mpcfaculty.net/mark_bishop/redox_balancing.htm
EXERCISE 5.7 - Oxidation Numbers
Determine the oxidation number for the atoms in P4, PF3, PH3, P2O3, H3PO4, N2,
N3−, K3N, Co3O2, NaH, Na, HSO3−, Cu(NO3)2, K2O2, and Fe2(SO4)3.
Table 5.9 Questions Answered by the Determination of Oxidation Numbers
Question
How to answer the question
Is the reaction redox?
If any atoms change their oxidation number, the
reaction is redox.
What’s oxidized?
The element that increases its oxidation number is
oxidized.
What’s reduced?
The element that decreases its oxidation number is
reduced.
What’s the reducing agent?
The substance that contains the element that is oxidized
is the reducing agent.
What’s the oxidizing agent?
The substance that contains the element that is reduced
is the oxidizing agent.
97
EXERCISE 5.8 - Redox Reactions
Identify whether the following equations describe redox reactions or not. For each
of the redox reactions, identify what is oxidized, what is reduced, what the reducing
agent is, and what the oxidizing agent is.
a. Ca(s) + F2(g) → CaF2(s)
∆
b. CaCO3(s) → CaO(s) + CO2(g)
c. 2Al(s) + 3H2O(g) → Al2O3(s) + 3H2(g)
d. Cr2O72−(aq) + 6Cl−(aq) + 14H+(aq) → 2Cr3+(aq) + 3Cl2(g) + 7H2O(l )
Figure 5.14
Single-Displacement Reaction Between Copper(II) Sulfate and Solid Zinc
98
Chapter 5
Relating Quantities of Reactants and Products
Sample Study
Sheet 5.5:
TIP-OFF - You are asked to convert from amount of one substance in a chemical
reaction to amount of another substance in the reaction.
General
Equation
Stoichiometry
GENERAL PROCEDURE - Use the unit analysis format to make the following
conversions.
1. If you are not given it, write and balance the chemical equation for the
reaction.
2. Start your unit analysis in the usual way, setting the desired units of substance
2 equal to the given units of substance 1.
3. Convert from the units that you are given for substance 1 to moles of
substance 1.
• For pure solids and liquids, this means converting grams to moles using
the molar mass of the substance. (It might be necessary to insert one or
more additional conversion factors to convert from the given unit to
grams.)
• Molarity can be used to convert from volume of solution to moles of
solute. (It might be necessary to insert one or more additional conversion
factors to convert from the given unit to liters or milliliters.)
4. Convert from moles of substance 1 to moles of substance 2 using the
coefficients from the balanced equation.
5. Convert from moles of substance 2 to the desired units for substance 2.
• For pure solids and liquids, this means converting moles to mass using
the molar mass of substance 2.
• Molarity can be used to convert from moles of solute to volume of
solution.
6. Calculate your answer and report it with the correct significant figures (in
scientific notation, if necessary) and with the correct unit.
EXERCISE 5.9 - Stoichiometry and Molarity
How many milliliters of 6.00 M HNO3 are necessary to neutralize the carbonate in
75.0 mL of 0.250 M Na2CO3?
EXERCISE 5.10 - Molarity
What is the molarity of a AgClO4 solution made by dissolving 29.993 g of silver
perchlorate in water and diluting with water to 50.00 mL total?
99
EXERCISE 5.11 - General Stoichiometry
What is the maximum number of grams of silver chloride that will precipitate from
a solution made by mixing 25.00 mL of 0.050 M MgCl2 with an excess a AgNO3
solution?
�����������������������������������������
�����������������������������
�����������������������������
���������������������������
�����������������������������
�����������������������������
���������������������������
���������������
������������
�������������������
���������������
���������������
�����
�������������������
�������������������
�����
�������
�������
������������������
�������
����������������������������
����������������������
�������
����������������
����������������
����������������
����������������
�������������
��������
�����������������������
�������������
���������������������
�����������������������������
������������������
�������
����������������
�����������������
������������������
�������
����������������������
����������������������������
�����������������������
���������������������
�������������
�����������������������������
��������������������������������������������
Figure 5.15
General Equation Stoichiometry
Visit the following website to see a description of a procedure called titration that can
be used to determine the molarities of acidic and basic solutions:
www.mpcfaculty.net/mark_bishop/titration.htm
100
Chapter 5
Sample Study
Sheet 5.6:
Titration
Problems
Relating Quantities of Reactants and Products
TIP-OFF - For the problems that you will see in this text, you will be given the volume
of a solution of an acid or base (the titrant – solution 1) necessary to react completely
with a given volume of solution being titrated (solution 2). You will also be given
the molarity of the titrant (solution 1). You will be asked to calculate the molarity of
solution 2.
GENERAL PROCEDURE - You can use the following steps.
Use the unit analysis process, with the following general format.
(given) (volume unit) 1
? mol 2
=
(given) (volume unit) 2
1 L 2 soln
--- (voloume unit) 2
--- L
--- L (or mL)
--- (volume unit) 1
(molarity) mol 1
1 L (or 103 mL) 1 soln
(coef. 2) mol 2
(coef. 1) mol 1
∗ The first conversion factor is only necessary if you are not given liters of 2.
(Because you are usually given milliliters, this conversion factor often converts
from milliliters to liters.)
∗ The second conversion factor is only necessary if you are not given either
milliliters or liters of solution 1. (You are usually given milliliters, so if you use
the form of the molarity conversion factor that includes “103 mL 1 soln”, this
conversion factor is not necessary.)
∗ The coefficients in the final conversion factor come from the balanced
equation for the reaction.
∗ Complete the calculation in the usual way.
EXERCISE 5.12 - Titration Problem
When 34.2 mL of a 1.02 M NaOH solution is added from a burette to 25.00 mL of
a phosphoric acid solution that contains phenolphthalein, the solution changes from
colorless to red.
a. What is the titrant for this process?
b.
What is the molarity of the phosphoric acid?
101
TIP-OFF - You are asked to describe how you would make an aqueous solution from a Sample Study
pure solid, and you are given the volume and the molarity of solution desired.
Sheet 5.7:
GENERAL PROCEDURE – Follow these steps.
• Calculate the mass in grams of solid to be weighed.
? g X = (given) (volume unit) X soln
--- L (or --- mL)
--- (volume unit)
(molarity) mol X
1 L (or 103 mL) X soln
The first conversion factor is only necessary if you are not given liters or
milliliters of X solution.
• Answer the question with a sentence like the following.
Dissolve (calculated) grams of X in a minimum amount of water, and
dilute with water to the desired total volume.
EXERCISE 5.13 - Making Solution from Solid
An experiment calls for a total of 1.50 L of 0.200 M KMnO4 for a class of chemistry
students. How would this solution be made from pure, solid potassium permanganate
and water?
(molar mass) g X
1 mol X
Making
Solutions from
Pure, Solid
Substance
102
Chapter 5
Sample Study
Sheet 5.8:
Dilution
Problems
Relating Quantities of Reactants and Products
TIP-OFF - There are several tip-offs for this type of problem. Perhaps the most general
is the mention of two concentrations of the same substance. Other tip-offs include
mention of dilution or the specific reference to making a more dilute solution from a
more concentrated solution.
GENERAL PROCEDURE - There are two possible approaches.
UNIT ANALYSIS - You can use the general unit analysis thought-process, taking
care to clearly identify the units of concentrated solution, dilute solution, and
pure solute. The molarities given in the problem are used as conversion factors.
The following is one general form of this approach.
? mL conc soln = (given) mL dil soln soln
(dil molarity) mol solute
103 mL dil X soln
103 mL conc X
(conc molarity) mol solute
WITH DILUTION EQUATION - You can also use the following steps.
• Assign the variables MC, VC, MD, and VD to the values you are given and
the value that you want.
• Write the dilution equation.
MCVC = MDVD
• Solve the equation for the unknown variable.
• Plug in the values with their units for the other variables.
• Cancel your units.
• If your units do not cancel to yield the desired unit, make the necessary
unit conversions so the units do cancel correctly.
• Complete the calculation, round your answer to the correct significant
figures, and report the correct unit.
EXERCISE 5.14 - Dilution Problems
What is the molarity of a solution made by diluting 5.00 mL of a 14.8 M NH3 solution
to 75.0 mL?
103
Tip-off - You are asked to describe how to make a solution of a specific molarity.
General Procedure - The first step is to recognize whether you are starting with
pure solid or with a more concentrated solution.
• If you are starting with pure solid, use the procedure described on the study
sheet called Descriptions of Making Solutions From Solid and Water.
• If you are starting with a more concentrated solution, follow these steps. (In
this text, you will be given the volume and molarity desired for the more
dilute solution, and you will be given the molarity of the more concentrated
starting solution.)
∗ Calculate the volume of the more concentrated solution that must be
measured to start the process. This can be done with unit analysis or
with the shortcut for dilution problems. (See the sample study sheet for
Dilution Problems.)
? mL conc soln = (given) mL dil soln soln
MCVC = MDVD
VC =
(dil molarity) mol solute
103 mL dil X soln
Sample Study
Sheet 5.9:
Making
Solutions
103 mL conc X
(conc molarity) mol solute
MDVD
MC
∗ If the more concentrated solution is a pure or almostpure acid, such as
18 M H2SO4 or 17 M HC2H3O2, describe the process for making the
solution in the following way:
♦ Add the (calculated) volume of the concentrated acid solution to
enough water to dilute the acid significantly.
♦ Add water until the volume reaches the desired total.
∗ If the more concentrated solution is not a pure or almostpure acid,
describe the process for making the solution in the following way:
♦ Add the (calculated) volume of the more concentrated solution to a
volume-measuring instrument.
♦ Add water until the volume reaches the desired total.
EXERCISE 5.15 - Making Solution from Concentrated Acid
How would you make 250.0 milliliters of 2.00 M acetic acid from concentrated acetic
acid (called glacial acetic acid) that is 17.4 M HC2H3O2?
104
Chapter 5
Relating Quantities of Reactants and Products
EXERCISE 5.16 - Making Solution from More Concentrated
Solution
How would you make 50.0 milliliters of 0.250 M hydrochloric acid from
2.0 M HCl?
You can get more information about making solutions of a specific molarity at the
following Web address:
http://www.mpcfaculty.net/mark_bishop/making_solutions.htm