15
CHAPTER
Acids and
Bases React
Chapter Preview
Sections
15.1 Acid and Base Reactions
MiniLab15.1 Acidic, Basic, or
Neutral?
15.2 Applications of Acid-Base
Reactions
MiniLab 15.2 What Does a Buffer Do?
ChemLab
Titration of Vinegar
514
What Happened Here?
W
hen industrial plants burn fossil fuels, air
pollutants are released into the atmosphere. Sulfur and nitrogen oxides often
combine with moisture in the air, forming acid
rain. This precipitation, in turn, may have a negative impact on otherwise healthy vegetation.
Start-up Activities
Physical or Chemical Change?
Determine if a physcial or chemical change took place.
Safety Precautions
Materials
• 500-mL Erlenmeyer flask
• 1,000-mL graduated cylinder
• one-hole stopper with 15-cm length of glass tube
• 1,000-mL beaker
• 45-cm length of rubber (or plastic) tubing
• stopwatch or clock with second hand
• weighing dish
• balance
• baking soda
• vinegar
it a physical or chemical change? How do you
know? Was this process endothermic or
exothermic? Calculate the average reaction rate
based on the number of bubbles per second.
What I Already Know
Review the following concepts
before studying this chapter.
Chapter 12: the mole concept; using
the factor label method
Chapter 13: solution concentration
using molarity
Chapter 14: properties and definitions of acids and bases
Procedure
1. Measure 300 mL of water. Pour water into the
2.
3.
4.
5.
6.
7.
8.
500-mL Erlenmeyer flask.
Weight 15 g of baking soda. Carefully pour the baking soda into the flask. Swirl the flask until the solution is clear.
Insert the rubber stopper with the glass tubing into
the flask.
Measure 600 mL of water and pour it into the
1,000-mL beaker.
Attach one end of the rubber tubing to the top of the
glass tubing. Place the other end of the rubber tubing in the beaker. Be sure the rubber tubing remains
under water.
Remove the stopper from the flask. Carefully add
250 mL of vinegar to the flask. Replace the stopper.
Count the number of bubbles coming into the beaker
for 20 s. Repeat this two more times.
Record your data in a data table.
Reading Chemistry
Quickly skim the chapter, writing
down the titles of the section heads.
As you read, locate the main idea of
each section and write it, in your
own words, beneath each title.
Preview this chapter’s content and
activities at chemistryca.com
Conclude and Apply
1. Describe what you observed in the flask after the
acid was added to the baking soda solution. Was
515
SECTION
15.1
Acid and Base
Reactions
A
s you learned in Chapter 14, acids
and bases are opposites in most of
their properties. But you also
learned that there is a big difference
between strong acids and weak acids
and between strong bases and weak
bases. Do all acids react with all bases?
Do all acid-base reactions produce a
neutral solution?
SECTION PREVIEW
Objectives
✓ Distinguish the
overall, ionic, and net
ionic equations for an
acid-base reaction.
✓ Classify acids and
bases using the hydrogen transfer definition.
Types of Acid-Base Reactions
✓ Predict and explain
the final results of an
acid-base reaction.
The reaction of an acid and a base is called a neutralization reaction
because the properties of both the acid and base are diminished or neutralized when they react.
In most cases, the reaction of an acid with a base produces water and a
salt. Salt is a general term used in chemistry to describe the ionic compound formed from the negative part of the acid and the positive part of
the base. In the language of chemistry, sodium chloride, common table
salt, is just one of a large number of ionic compounds that are called salts.
KCl, NH4NO3, and Fe3(PO4)2 are other examples of salts.
Consider the following neutralization reaction. Hydrochloric acid, HCl,
is a common household and laboratory acid. Muriatic acid is the common
household name of hydrochloric acid. It is often sold in hardware stores
to be used in masonry work to remove excess mortar from brick. Sodium
hydroxide, NaOH, is a common household and laboratory base. The common name of sodium hydroxide is lye. It is the primary component of
many drain cleaners. Figure 15.1 shows litmus tests before and after mixing these substances together.
Review Vocabulary
pH: mathematical
scale in which the concentration of hydronium ions in a solution
is expressed as a number from 0 to 14.
New Vocabulary
neutralization
reaction
salt
ionic equation
spectator ion
net ionic equation
Figure 15.1
Neutralization Reactions
A solution of hydrochloric acid,
HCl, is added to exactly the
amount of a solution of basic
sodium hydroxide, NaOH, that
will react with it. Litmus papers
show that the resulting salt solution is neither acidic nor basic.
NaOH(aq) HCl(aq) ˇ
NaCl(aq) H2O(l)
516
Chapter 15
Acids and Bases React
Basic
Acidic
No color
changes
After the reaction, the mixture contains only the salt
Table 15.1 Types of Acid-Base
sodium chloride, NaCl, dissolved in water. The litmus test
Reactions
shows no acid or base present in the reaction products.
Because both acids and bases may be either strong or
weak, four possible combinations of acid-base reactions
Acid
Base
may occur. Table 15.1 summarizes the possibilities. As
Strong
Strong
you will find out later in this section, only three of the
Strong
Weak
types are significant in everyday chemistry.
Weak
Strong
As long as one of the reactants is strong, the acidWeak
Weak
base reaction goes to completion. As you learned in
Chapter 6, a reaction goes to completion when the limiting reactant is completely consumed.
Although all of the reactions in Table 15.1 are acid-base reactions, the
submicroscopic interactions in each are different. Examine each possible
type of acid-base reaction and see how they compare.
Strong Acid Strong Base
A typical type of acid-base reaction is one in which both the acid and
base are strong. The reaction of aqueous solutions of hydrochloric acid
and sodium hydroxide shown in Figure 15.1 is a good example of this
type of reaction.
A Macroscopic View
It is easy to write and balance equations for strong acid-strong base
reactions. In the previous example, HCl is the acid; NaOH is the base. The
products are NaCl, which is a salt, and water. Now take a closer look at
these reactants and their products.
The Submicroscopic View: Ionic Reactions
Recall from Chapter 14 that HCl, when dissolved in water, completely
ionizes into hydronium ions and chloride ions because HCl is a strong
acid. As you learned in Chapter 14, the hydronium ion is more conveniently written in shorthand as H.
HCl(aq) ˇ H(aq) Cl(aq)
You also know that sodium hydroxide in water completely dissociates
into sodium ions and hydroxide ions because NaOH is a strong base.
NaOH(aq) ˇ Na(aq) OH(aq)
An overall equation for the reaction between NaOH and HCl shows
each substance involved in the reaction. An overall equation does not
indicate whether these substances exist as ions. The best way for you to
model the submicroscopic behavior of an acid-base reaction is to show
reactants and products as they actually exist in solution. Instead of an
overall equation, an ionic equation, in which substances that primarily
exist as ions in solution are shown as ions, can be written.
H(aq) Cl(aq) Na(aq) OH(aq) ˇ
Na(aq) Cl(aq) H2O(l)
15.1
Acid and Base Reactions
517
Acidic, Basic, or Neutral?
1
The salt potassium bromide forms in the acid-base neutralization
reaction between hydrobromic acid and potassium hydroxide.
HBr(aq) KOH(aq) ˇ H2O(l) KBr(aq)
Hydrobromic acid and potassium hydroxide are referred to as the parent
acid and the parent base of potassium bromide. Test several aqueous salt
solutions with bromothymol blue indicator to determine whether the
solutions are acidic (yellow), basic (blue), or neutral (green).
Procedure
1. Wear laboratory aprons and
safety goggles.
2. Use labeled microtip pipets to
put six drops of sodium acetate
solution in A1, potassium nitrate
solution in A2, ammonium
chloride solution in A3, sodium
carbonate solution in A4, sodium chloride solution in A5, and
aluminum sulfate solution in A6
of a 24-well microplate.
3. Add two drops of bromothymol blue indicator solution to
each of the salt solutions. Stir
each with a separate toothpick.
4. Set the microplate on a piece of
white paper and look down
through each well to determine
the color of the solution.
Record your results.
Analysis
1. According to the color of each
solution that has bromothymol
blue indicator added, is each of
the solutions acidic, basic, or
neutral?
2. Relate the relative strengths of
the parent acids and bases to
the results of using the indicator with the salt solutions.
3. What type of salt might be
added to a product, such as
shampoo, in order to make it
slightly acidic so that it will not
be harmful to skin and hair?
Check the label of several
brands of shampoo for the
presence of such a salt.
Notice in the previous equation that in addition to showing the acid as
completely ionized and the base as completely dissociated, the ionic compound NaCl is also dissociated. Water does not ionize much, so it is indicated as a molecule rather than H and OH ions. The ionic aspects of this
ionic equation are confirmed in Figure 15.2.
Figure 15.2
What ions are present?
When a conductivity apparatus is placed in
solutions of each reactant and product of the
reaction between a strong acid and a strong
base, you can see that the acid, base, and salt
exist as ions in solution.
518
Chapter 15
Acids and Bases React
Taste
Although the total flavor of a food comes from the complex combination of taste, smell, touch, texture or consistency, and temperature sensations, taste is a major factor. Three
of the four fundamental tastes are directly linked to acids and
bases. Your tongue has four different types of taste buds—
sweet, salty, bitter, and sour—that are located at different
places on your tongue. Only certain molecules and ions can
react with these specific buds to produce a
signal that is sent to a certain region of
your brain. When these signals are
received, your brain processes
them, and you sense taste.
1. The taste buds that sense
a bitter taste are located
at the base of the
tongue. Bases taste bitter. Many medications
are basic, and pharmaceutical companies spend
a lot of time and research
trying to mask the bitter
taste with other tastes.
Bitter
2. The taste buds that detect a sour
taste are along the sides of the
tongue. The sour taste comes from
acids in your food. Sour-tasting foods
include vinegar and citrus fruits.
Sour
Salty
Salty and sweet
Sweetness of Some Compounds
Compound
Lactose
Glucose
Sucrose
Fructose
Aspartame
Saccharin
Relative Sweetness
16
74
100
173
16 000
50 000
As the table shows, aspartame, the popular artificial sweetener in soft drinks, is 160
times sweeter than sucrose, common table
sugar. The sweetening ability of aspartame
comes from a molecular structure that creates an impact 160 times that of sucrose on
your taste buds that detect the sweet taste.
3. The taste buds that detect the salty and sweet tastes
are located at the tip of your tongue. A salt is the product of an acid-base reaction. The sweet taste seems to
depend a great deal on the properties of both acids and
bases that are combined on a single molecule, but this
taste is not as clear-cut as the other tastes.
Thinking Critically
1. How are three of the
four fundamental
tastes linked to the
properties of acids
and bases?
2. What types of companies do you think
15.1
would support
research on taste?
3. What benefits might
artificial sweeteners,
such as aspartame,
have over natural
sweeteners?
Acid and Base Reactions
519
Spectator Ions and the Net Ionic Reaction
Note that the ionic equation gives more information about how a
strong acid-strong base reaction occurs. When you examine the two sides
of the ionic equation on page 517, you see that Na and Cl are present
both as reactants and as products. Although they are important components of an overall equation, they do not directly participate in the chemical reaction. They are called spectator ions because they are present in the
solution but do not participate in the reaction, Figure 15.3.
Figure 15.3
Spectators
The presence of spectators at
a sporting event is important,
but spectators do not actually
participate in the game and
do not determine the final
outcome.
Why show the spectator ions in an equation if they aren’t really
involved in the reaction? An ionic equation can be simplified to take care
of that problem. Just as in a mathematical equation, items common to
both sides of the equation can be subtracted. This process simplifies the
equation so that the reactants and products that actually change can be
seen more clearly.
H (aq) Cl (aq) Na (aq) OH (aq) ˇ Na (aq) Cl (aq) H2O(l)
+
+
+
When ions common to both sides of the equation are removed from
the equation, the result is called the net ionic equation for the reaction of
HCl with NaOH.
H(aq) OH(aq) ˇ H2O(l)
The net ionic equation describes what is really happening at the submicroscopic level. Although solutions of HCl and NaOH are mixed, the
net ionic equation is hydrogen ions reacting with hydroxide ions to form
water.
Even though the strong acid and strong base in Sample Problem 1 are
different from those in the HCl reaction with NaOH, the net ionic equation is the same. Hydrogen ions from the acid react with hydroxide ions
from the base to form water. This equation is always the net ionic equation for a strong acid-strong base reaction.
520
Chapter 15
Acids and Bases React
1
SAMPLE PROBLEM
Equations—Strong Acid, Strong Base
Write the overall, ionic, and net ionic equations for the reaction of
sulfuric acid with potassium hydroxide.
Analyze
Set Up
whether the acid is a strong acid or a weak acid and whether
• Decide
the base is strong or weak. A list of strong acids and bases can be
•
found in Chapter 14, Table 14.2. By looking at this table, you can see
that sulfuric acid, H2SO4, is a strong acid. Potassium hydroxide, KOH,
is a strong base.
Write an equation for the overall reaction. Because sulfuric acid is a
diprotic acid, you need two moles of KOH for every one mole of
H2SO4. Two moles of water and one mole of K2SO4 will be produced.
H2SO4(aq) 2KOH(aq) ˇ K2SO4(aq) 2H2O(l)
Write the ionic equation by showing H SO , KOH, and K SO as ions.
• You
must keep track of the coefficients from the overall equation and
2
4
2
4
the formulas of the substances when writing the coefficients of the ions.
2H(aq) SO42(aq) 2K(aq) 2OH(aq) ˇ
2K(aq) SO42(aq) 2H2O(l)
Solve
Look for spectator ions. In this reaction, K and SO are spectator
• ions.
Subtract them from both sides of the equation to get the net
4
2
ionic equation.
2H(aq) SO42(aq) 2K(aq) 2OH(aq) ˇ
2K(aq) SO42(aq) 2H2O(l)
2H(aq) 2OH(aq) ˇ 2H2O(l)
the balanced net reaction by
• Simplify
dividing coefficients on both sides of the
equation by the common factor of 2.
Problem-Solving
H I N T
Coefficients should be in the smallest
whole-number ratio possible.
2H (aq) 2OH (aq) ˇ 2H2O(l)
+
results in H (aq) OH (aq) ˇ H2O(l)
2
+
Check
a final look at the net ionic equation to make sure no ions are
• Take
common to both sides of the equation.
PRACTICE PROBLEMS
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
Write overall, ionic, and net ionic equations for each of the following
reactions.
1. hydroiodic acid, HI, and calcium hydroxide, Ca(OH)2
2. hydrobromic acid, HBr, and lithium hydroxide, LiOH
3. sulfuric acid, H2SO4, and strontium hydroxide, Sr(OH)2
4. perchloric acid, HClO4, and barium hydroxide, Ba(OH)2
15.1
Acid and Base Reactions
521
The pH Perspective
The net ionic equation also shows why this reaction is called a neutralization reaction. The hydrogen ion from the acid reacts with the hydroxide ion from the base to form water, which is neutral.
Figure 15.4 shows the reaction of 50.0 mL of 0.100M HCl with 50.0 mL
of 0.100M NaOH, which is exactly the right amount of base to react with
all of the acid. As the NaOH solution is added to the HCl solution, the pH
of the solution increases. After all of the NaOH solution is added, the pH
of the final solution is 7.
What happens if one of the reactants is strong and the other is weak? A
similar approach can be used.
Figure 15.4
Final pH
The reaction of a strong acid and a
strong base is definitely a neutralization. The pH of a 0.100M HCl solution
(top) is 1. The pH of a 0.100M NaOH
solution (bottom left) is close to 13.
When the reaction is complete, the
pH is 7, which is the pH of a neutral
solution (bottom right).
1.00
7.08
12.52
Strong Acid Weak Base
Do the reactions change when the strength of an acid or base changes?
Look at an example of what happens when a strong acid and a weak base
are mixed together. Consider the reaction of hydrobromic acid and aluminum hydroxide. The overall equation shows the reactants and products.
3HBr(aq) Al(OH)3(s) ˇ AlBr3(aq) 3H2O(l)
Hydrobromic acid, a strong acid, completely ionizes in water. All of the
Al(OH)3 that dissolves dissociates, so it is technically a strong base. However, because it is so insoluble, few OH ions are produced, and Al(OH)3
acts as a weak base. Therefore, the ionic equation shows little dissociation
of the base. The dissociated salt, AlBr3, is also shown as ions.
3H(aq) 3Br(aq) Al(OH)3(s) ˇ Al3(aq) 3Br 3H2O(l)
522
Chapter 15
Acids and Bases React
The spectator ions in this equation are bromide ions. They are removed
from both sides of the equation to produce the net ionic equation.
3H(aq) Al(OH)3(s) ˇ Al3(aq) 3H2O(l)
Compare this equation to the net ionic equation for a strong acidstrong base reaction.
A Strong Acid NH3
Recall from Chapter 14 that the most common weak base does not
contain the hydroxide ion. Consider an equation for the reaction between
hydrochloric acid and ammonia.
HCl(aq) NH3(aq) ˇ NH4Cl(aq)
Notice that although the product is a salt, NH4Cl, no water is produced
in this overall reaction. As before, use the net ionic equation to understand the submicroscopic processes for this reaction of a weak base with a
strong acid.
In medieval times, natural deposits of ammonium chloride, a compound derived from
ammonia, were first
mined in Egypt near a
temple of the Egyptian
god Ammon.
The Ionic Reaction: What’s in Solution?
Recall that when ammonia dissolves in water, some of the ammonia
molecules react with water to form ammonium ions and hydroxide ions.
However, most of the ammonia molecules remain as molecules. Other
than water, the major particle present in an aqueous solution of ammonia
is ammonia molecules, NH3.
A solution of ammonia is best represented by NH3(aq). A solution of
HCl, as in the last case, is best represented as H(aq) and Cl(aq). The
ionic reaction is written, as in Sample Problem 1, by representing what is
actually in the reactant and product solutions.
H(aq) Cl(aq) NH3(aq) ˇ NH4(aq) Cl(aq)
The ionic salt NH4Cl is written as dissociated ions. Figure 15.5 shows
which of these particles are actually involved in the reaction.
Figure 15.5
What actually reacts?
A quick look at the ionic reaction shows that the chloride ion is a
spectator ion because it appears on both sides of the reaction.
You get the net ionic equation by subtracting the spectator Cl
from both sides of the ionic equation.
H (aq) Cl (aq) NH3(aq) ˇ NH4 (aq) Cl (aq)
+
+
H (aq) NH3(aq) ˇ
+
Cl
+
NH4 (aq)
H
Cl
NH4
NH3
15.1
Acid and Base Reactions
523
EARTH SCIENCE
CONNECTION
Cave Formation
When you think of a cave, what images
come to mind? A massive, damp, cool,
underground chamber running deep into
Earth or a fantasy underground world with
stone icicles rising from the floor and ornate
columns seeming to support the ceiling?
Whatever visual pictures you may imagine,
caves are one of nature’s wonders.
How caves are formed Caves form in limestone regions throughout the world. Limestone is calcium carbonate, which is only
slightly soluble in water. The caves that form
within these rocks are called solution caves.
What causes natural water to be acidic?
Most rainwater is slightly acidic because it
contains carbon dioxide from the atmosphere. A small amount of the CO2 dissolves
in the water, but some of it reacts with the
water to form carbonic acid.
CO2(g) H2O(l) ˇ H2CO3(aq)
Carbonic acid forms a hydronium ion and a
hydrogen carbonate ion.
H2CO3(aq) H2O(l) ˇ
H3O(aq) HCO3(aq)
How does this acid form caves? The hydronium ions react with limestone to produce
soluble ions.
H3O(aq) CaCO3(s) ˇ
Ca2(aq) HCO3(aq) H2O(l)
The acidic water dissolves the limestone rocks,
producing open spaces that contain water.
Stalactites and stalagmites During the second phase, clay, silt, sand, or gravel moves into
the spaces. In the third phase, streams partially
remove these materials and modify and
enlarge the spaces. Stalactites and stalagmites
now form by the reverse of the chemical and
physical processes that formed the cave. The
524
Chapter 15
Acids and Bases React
water containing dissolved CO2 and H2CO3 is
saturated with Ca(HCO3)2. As it seeps through
the roof of the cave, the water in each droplet
slowly evaporates.
Some of the carbonic acid changes back
into carbon dioxide and water. The pH of the
water increases, and the solubility of
Ca(HCO3)2 decreases. The CaCO3 precipitates out slowly, forming stalactites over
thousands of years. As the saturated water
drops hit the floor, the same processes slowly
form stalagmites. Sometimes, the two formations grow together, forming pillars.
Connecting to Chemistry
1. Applying Identify
two physical
changes and two
chemical changes
that occur in cave
formation.
2. Thinking Critically
Write the equations
for cave formation
from the MgCO3
part of dolomite,
CaCO3 MgCO3.
2
SAMPLE PROBLEM
Equations—Strong Acid, Weak Base
Write the overall, ionic, and net ionic equations for the reaction of
nitric acid with ammonia.
Analyze
whether the acid is strong or weak and whether the base is
• Decide
strong or weak. Nitric acid, HNO , is a strong acid. Ammonia is a weak
3
Set Up
Solve
•
•
base.
Write an equation for the overall reaction.
HNO3(aq) NH3(aq) ˇ NH4NO3(aq)
Write the ionic equation. Because HNO3 is a strong acid, you write it
as completely ionized. NH3 is a weak base, so you write it as NH3. You
dissociate the salt, ammonium nitrate, NH4NO3, into its component
ions because it is an ionic compound.
H(aq) NO3(aq) NH3(aq) ˇ NH4(aq) NO3(aq)
for spectator ions. In this reaction, only the nitrate ion, NO , is
• Look
a spectator ion. Subtract NO from both sides of the equation to get
3
3
Check
•
•
the net ionic equation.
H(aq) NO3(aq) NH3(aq) ˇ NH4(aq) NO3(aq)
H(aq) NH3(aq) ˇ NH4(aq)
Note that this is the same net equation as in the HCl and NH3 example
shown in Figure 15.5.
Take a final look at the equation to make sure no ions are common to
both sides of the equation.
PRACTICE PROBLEMS
Write overall, ionic, and net ionic equations for each of the following
reactions.
5. perchloric acid, HClO4, and ammonia, NH3
6. hydrochloric acid, HCl, and aluminum hydroxide, Al(OH)3
7. sulfuric acid, H2SO4, and iron(III) hydroxide, Fe(OH)3
For all these examples of strong acid-weak base reactions, the net ionic
equation differs from that for a strong acid and a strong base. The submicroscopic interactions in these strong acid-weak base reactions are
between hydrogen ions and the bases.
Strong Acid-Weak Base and pH
Figure 15.6 shows that a solution of 0.100M ammonia is definitely a
base. It has pH greater than 7. If you compare the pH of 0.100M NaOH
with the pH of 0.100M NH3, you see that ammonia is a weaker base
because it has a lower pH.
15.1
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
Acid and Base Reactions
525
Figure 15.6
The Product Is Acidic
The reaction of a strong acid and weak base is
not quite a neutralization. The pH of a 0.100M
HCl solution is 1, and the pH of a 0.100M NH3
solution is approximately 11. When equal volumes of the solutions are mixed and the reaction is complete, the pH is approximately 5.
1.00
10.97
5.18
Figure 15.6 also shows that when equal volumes of solutions of ammonia and hydrochloric acid of equal molarity are mixed, the pH of the final
mixture is less than 7. A pH less than 7 means that the final reaction mixture is acidic. Therefore, more hydronium than hydroxide ions must be
present in the final reaction mixture.
How can equal moles of base and acid react to produce a neutral solution in one reaction and an acidic solution in the second? The result must
have something to do with the relative strengths of the acid and the base.
A Broader Definition of Acids and Bases
The reaction of a strong acid with a weak base demonstrates the need
for a slightly broader definition of acids and bases. As you learned in the
last chapter, much of the behavior of acids and bases in water can be
explained by a model that focuses on the hydrogen ion transfer from the
acid to the base. This model will also help explain why every acid-base
reaction does not result in a neutral solution.
Hydrogen-Ion Donor or Acceptor
You can use the ability to exchange a hydrogen ion as the basis of a
broader definition of an acid or a base. In this definition, called the
Brønsted-Lowry definition of acids and bases, an acid is defined as a
substance that donates, or gives up, a hydrogen ion in a chemical reaction.
A base, not surprisingly, is just the opposite. A base is a substance that
accepts a hydrogen ion in a chemical reaction.
526
Chapter 15
Acids and Bases React
Figure 15.7
Defining Acids and Bases by H Transfer
In a reaction between aqueous HCl and aqueous NH3, several H transfers occur.
ˇ
A. HCl(aq) H2O(l) ˇ H3O (aq) Cl (aq)
A In the transfer of H from
B In the transfer of H from
acid
base
ˇ
+
+
HCl to a water molecule,
a hydronium ion to an
B. H3O (aq) NH3(aq) ˇ NH4 (aq) H2O(l)
HCl acts as an acid, and
ammonia molecule, the
acid
base
water acts as a base.
hydronium ion acts as an
ˇ
+
C. H2O(l) NH3(aq) ˇ NH4 (aq) OH (aq)
acid, and the ammonia
molecule acts as a base.
acid
base
+
C Water also reacts with ammonia
molecules. Water acts as an acid,
and ammonia acts as a base.
Take another look at the net ionic equation of HCl with NH3 on page
523. If you adhere strictly to the definition of a base as a hydroxide-ion
producer in water, none of these equations define ammonia as a base.
Remember that hydroxide ions are produced, but the amount is so small
that it is not shown in the equations.
H(aq) NH3(aq) ˇ NH4(aq)
Using the new definition, you can definitely say that ammonia is acting
as a base. It is accepting a hydrogen ion. Figure 15.7 shows this reaction
written in its most complete form and clearly shows the hydrogen ion
transfer. The hydronium ion is acting as the acid because it donates the
hydrogen ion.
Lab
See page 870 in
Appendix F for
Testing for Ammonia
It Takes Two to Transfer
Notice from Figure 15.7 that the definitions of an acid as a hydroniumion producer and a base as a hydroxide-ion producer are included in this
H-transfer definition. When HCl reacts with water, it acts as an H
donor, so it is an acid. Water acts as an H acceptor, so it is a base. When
ammonia reacts with water, the ammonia molecule accepts H from the
water. Ammonia is the base and water is the acid. Remember that it takes
two to transfer, so for every acid (an H donor), there must be a base (an
H acceptor).
Although you generally think of water as neutral, a unique property of
water can now be observed. Water can act as either an acid or a base,
depending upon what else is in solution.
Water Not Required
Although most of the reactions that you will study occur in water, the
H -transfer definition does not require water to be present. For example, the
reaction between HCl(g) and NH3(g) is shown in Figure 15.8. This reaction
occurs in the gas phase. It involves the transfer of a hydrogen ion from a
gaseous HCl molecule to a gaseous ammonia molecule to form a solid product. This gas reaction can now be classified as an acid-base reaction.
15.1
Acid and Base Reactions
527
Figure 15.8
A Gas Phase Acid-Base Reaction
HCl(g) and NH3(g) react to form NH4Cl(s). Gases
from the concentrated aqueous solutions react to
form a smoke of solid ammonium chloride.
Weak Acid Strong Base
vinegar:
vinaigre (Fr) sour
wine
When wine
becomes sour, it
has turned to
vinegar.
Considering this new H-transfer acid-base definition, take a look at
the type of acid-base reaction in which the acid is weak and the base is
strong. An example is the reaction of acetic acid, HC2H3O2, the weak acid
present in vinegar, with sodium hydroxide. The equation of the overall
reaction is similar to that of a strong acid-strong base reaction.
HC2H3O2(aq) NaOH(aq) ˇ NaC2H3O2(aq) H2O(l)
The Ionic Reaction: What’s in Solution?
Vinegar is a dilute
solution of acetic acid.
Pure acetic acid is
often called glacial
acetic acid. It was first
purified in 1700 by the
distillation of vinegar.
At room temperature,
pure acetic acid is a
liquid, but it freezes at
17°C. The term glacial
means “icelike.” In
poorly heated chemistry labs, pure acetic
acid freezes. Acetic
acid is 34th on the
1994 list of top industrial chemicals.
528
Chapter 15
As you know, acetic acid is a weak acid. In a solution of acetic acid, only
a small fraction of the acetic acid molecules ionize. Other than water, the
major particle present in an aqueous solution of acetic acid is HC2H3O2.
NaOH, as seen before, completely dissociates. The ionic equation shows
what is present when the acid and base react. As in the previous cases, the
salt is also written as dissociated ions.
HC2H3O2(aq) Na(aq) OH(aq) ˇ
Na(aq) C2H3O2(aq) H2O(l)
Net Ionic Reaction: H Transfer
The ionic equation contains the sodium ion as a spectator ion. Subtracting Na from each side of the ionic equation gives the net ionic
equation.
HC2H3O2(aq) Na (aq) OH (aq) ˇNa (aq) C2H3O2 (aq) H2O(l)
HC2H3O2(aq) OH (aq) ˇ C2H3O2 (aq) H2O(l)
+
+
The net ionic equation shows that a weak acid and a strong base react
by hydrogen ion transfer from the weak acid to the hydroxide ion. The
acetic acid is the H donor and serves as the acid. The hydroxide ion is
the H acceptor and serves as the base. Although this is an acid-base reaction, notice that H is not involved as a reactant or product of the reaction. However, using the H-transfer definition, it is easy to include this
reaction as an acid-base reaction.
Acids and Bases React
2.95
12.52
9.05
Figure 15.9
The photos in Figure 15.9 show what happens when 0.100M solutions
of sodium hydroxide and acetic acid are mixed. Notice that the pH of the
0.100M HC2H3O2 is greater than the pH of the 0.100M HCl solution in
Figure 15.6.
As the sodium hydroxide solution is added to the acetic acid solution,
the pH increases as the acidic hydrogens from acetic acid molecules react
with the hydroxide ions. When equal volumes of the two are mixed, the
final pH is greater than 7. The final reaction mixture is basic.
As in the case of the weak base-strong acid reaction, mixing equal
moles of acid and base does not produce a neutral solution. Because the
final pH in this case is basic, more hydroxide than hydronium ions must
be present in the final reaction mixture.
3
SAMPLE PROBLEM
Weak Acid, Strong Base
The reaction of a weak acid
and strong base does not
result in a neutral solution.
The pH of a 0.100M HC2H3O2
solution is approximately 3,
and the pH of a 0.100M NaOH
solution is approximately 13.
When the reaction is complete, the pH of the resulting
solution is approximately 9.
Equations—Weak Acid, Strong Base
Write the overall, ionic, and net ionic equations for the reaction of
phosphoric acid with lithium hydroxide.
Analyze
Set Up
the acid and base as strong or weak. Phosphoric acid, H PO ,
• isIdentify
a weak acid. Lithium hydroxide, LiOH, is a strong base.
Write
a balanced equation for the overall reaction. The salt is lithium
• phosphate,
Li PO . Water is also a product. You’ll need three moles of
3
3
Solve
•
•
4
4
LiOH for each mole of H3PO4. Three moles of water will be produced.
H3PO4(aq) 3LiOH(aq) ˇ Li3PO4(aq) 3H2O(l)
Now, write the ionic equation. Because H3PO4 is a weak acid, it is only
partially ionized, and you write it in the ionic equation as H3PO4.
LiOH is completely dissociated, as is the salt lithium phosphate. Be
careful to keep the ionic equation balanced.
H3PO4(aq) 3Li(aq) 3OH(aq) ˇ
3Li(aq) PO43(aq) 3H2O(l)
Check for spectator ions. Li is a spectator ion. Subtract Li from both
sides of the equation to get the net ionic equation.
H3PO4(aq) 3Li (aq) 3OH (aq) ˇ 3Li (aq) PO4 (aq) 3H2O(l)
3–
H3PO4(aq) 3OH (aq) ˇ PO4 (aq) 3H2O(l)
+
Check
+
3–
to see that the reaction occurs through hydrogen ion transfer
• Check
from phosphoric acid molecules to hydroxide ions. Every weak acidstrong base reaction occurs by this type of H transfer.
15.1
Acid and Base Reactions
529
PRACTICE PROBLEMS
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
Write overall, ionic, and net ionic equations for the following
reactions.
18. carbonic acid, H2CO3, and sodium hydroxide, NaOH
19. boric acid, H3BO3, and potassium hydroxide, KOH
10. acetic acid, HC2H3O2, and calcium hydroxide, Ca(OH)2
Weak and Weak: It’s Uncertain
Figure 15.10
Weak Acid-Weak Base
The weak acetic acid does
not react with the weak
base aluminum hydroxide.
The strong-strong reaction plus the two types of weak-strong reactions
are the favorable acid-base reactions. Looking at an acid-base reaction as
occurring by H transfer helps you to understand why the weak-weak
reaction is not considered a favorable reaction.
Because neither a weak acid nor
a weak base has a strong tendency
to transfer a hydrogen ion, transfer between the two may occur,
but it is uncommon. Reactions
between a weak acid and a weak
base generally do not play an
important role in acid-base chemistry, as shown in Figure 15.10.
SECTION REVIEW
Understanding Concepts
Thinking Critically
1. Write the overall, ionic, and net ionic equations
for the following reactions.
a) perchloric acid and sodium hydroxide
b) sulfuric acid and ammonia
c) citric acid, H3C6H5O7, and potassium
hydroxide
2. For each of the reactions in question 1, predict
whether the pH of the product solution is
acidic, basic, or neutral. Explain.
3. Identify the acid and the base in each of the following reactions.
a) HBr(aq) H2O(l) ˇ H3O(aq) Br(aq)
b) NH3(aq) H3PO4(aq) ˇ
NH4(aq) H2PO4(aq)
c) HS(aq) H2O(l) ˇ H2S(aq) OH(aq)
4. Applying Concepts Think about what acid and
what base would react to form each of the following salts. When each of the salts is dissolved in
water, will its solution be acidic, basic, or neutral?
If not neutral, use an equation to explain why.
c) LiC2H3O2
a) NH4Cl
b) NaCl
d) NH4C2H3O2
530
Chapter 15
Acids and Bases React
Applying Chemistry
5. Lactic Acid Reaction Lactic acid is produced in
muscles during exercise. It is also produced in
sour milk due to the action of lactic acid bacteria. The formula for lactic acid is HC3H5O3.
Write the overall, ionic, and net ionic equations
for the reaction of lactic acid with sodium
hydroxide. Will the pH of the product solution
be greater than 7, exactly 7, or less than 7?
chemistryca.com/self_check_quiz
SECTION
Applications of
Acid-Base Reactions
15.2
A
SECTION PREVIEW
cid-base reactions play an
important role in many
chemical systems, whether
you’re interested in blood chemistry, the chemistry of acid rain
and lakes, or the chemistry of
your favorite shampoo. How is
the acid-base balance maintained in each of these situations? Look at several applications of acid-base reactions and
how balance is maintained.
Objectives
✓ Evaluate the importance of a buffer in
controlling pH.
✓ Design strategies
for doing acid-base
titrations, and
calculate results from
titration data.
Review Vocabulary
Buffers to Regulate pH
Much blood chemistry depends on the acid-base balance in blood. The
pH of your blood is slightly basic, about 7.4. If you’re healthy, the pH of
your blood does not vary by more than one-tenth of a pH unit. If you
stop and think about all of the materials that go into and out of your
blood, it is amazing that the pH remains so constant. Many of those
materials are acidic or basic, which is why the constancy of the pH of
blood is fascinating. Blood is an example of an effective buffer.
Spectator ions: ions
that are present in a
solution but do not
participate in a
reaction.
New Vocabulary
buffer
titration
standard solution
Buffers Defined
A buffer is a solution that resists changes in pH when moderate
amounts of acids or bases are added. It contains ions or molecules that
react with OH or H if one of these ions is introduced into the solution.
Buffer solutions are prepared by using a weak acid with one of its salts
or a weak base with one of its salts. For example, a buffer solution can be
prepared by using the weak base ammonia, NH3, and an ammonium salt,
such as NH4Cl. If an acid is added, NH3 reacts with the H.
NH3(aq) H(aq) ˇ NH4(aq)
If a base is added, the NH4 ion from the salt reacts with the OH.
NH4(aq) OH(aq) ˇ NH3(aq) H2O(l)
Look at another system that contains the weak acid acetic acid,
HC2H3O2, and the salt sodium acetate, NaC2H3O2. If a strong base, OH, is
added to the buffer system, the weak acid reacts to neutralize the addition.
HC2H3O2(aq) OH(aq) ˇ C2H3O2(aq) H2O(l)
15.2
Applications of Acid-Base Reactions
531
What does a buffer do?
2
Compare the amounts of acidic and basic solutions required to cause
similar pH changes in two solutions: one a sodium chloride solution at
a pH of 7 and the other a buffered solution with a pH of 7.
Procedure
7. Add two drops of phenolphthalein indicator to tubes 3
and 4. The phenolphthalein is
1. Wear aprons and safety goggles.
colorless if the pH is lower than
2. Obtain and number four test
8.2 and will change to pink or
tubes 1-4.
magenta as the pH is raised
3. Use a small, graduated cylinder
from 8.2 to 10.0.
to measure 5.0 mL of NaCl
8. Add 0.1M NaOH solution drop
solution into tubes 1 and 3.
by drop to tubes 3 and 4, stir4. Use another small, graduated
ring after each drop, just until
cylinder to measure 5.0 mL of
the solution color changes to
pH 7 buffer solution into tubes
pink or magenta. Record the
2 and 4.
number of drops required for
5. Add two drops of methyl
each solution.
orange indicator to tubes 1 and Analysis
2. The methyl orange is yellow 1. Compare the number of drops
if the pH is greater than 4.4 and
of acid required to lower the
will change to orange-red as the
pH of the two solutions suffipH is lowered from 4.4 to 3.2.
ciently that the methyl orange
6. Add 0.100M HCl solution drop
changed color.
by drop to tubes 1 and 2, stir2. Compare the number of drops
ring after each drop, just until
of base required to raise the pH
the solution color changes to
of the two solutions sufficiently
orange-red. Record the number
that the phenolphthalein
of drops required for each
changed color.
solution.
3. How would you describe the
effect of the buffer on the pH
of the solution?
This reaction takes care of the added OH. If H is added, the acetate ion
from the NaC2H3O2 is available to neutralize the added H.
C2H3O2(aq) + H(aq) ˇ HC2H3O2(aq)
This system is shown in Figure 15.11.
Notice in Figure 15.11 that the pH does not remain constant in the
buffer solution. It changes slightly in the direction of the pH of the added
acid or base. These pH changes are insignificant when you compare them
to the changes that occur in the unbuffered solution.
These two buffer systems are common ones used in many laboratories.
Take a closer look at the specific buffer chemistry of your blood, which
operates just like these laboratory buffers.
532
Chapter 15
Acids and Bases React
Figure 15.11
Buffered
4.90
Buffers
A buffer maintains the pH of a
solution at a fairly constant
value. Compare what happens
when acid and base are added
to an acetic acid/sodium
acetate buffer system at pH 5 (top) to what happens
when the same amount of
acid and base are added to an
unbuffered solution of pH 5
(bottom).
5.10
Unbuffered
2.00
12.00
Blood Buffer: Dissolved CO2
The ability of blood to maintain a constant pH of 7.4 is due to several
buffer systems. Dissolved carbon dioxide makes up one of the systems.
Remember that when carbon dioxide dissolves in water, it produces carbonic acid, H2CO3.
CO2(g) H2O(l) ˇ H2CO3(aq)
The other part of the blood buffer is the hydrogen carbonate ion,
HCO3. If something happens to increase OH in your blood, H2CO3
reacts to lower the OH concentration and keep the pH from increasing.
If H enters the blood, HCO3 reacts to keep the pH from decreasing.
added OH: H2CO3(aq) OH(aq) ˇ HCO3(aq) H2O(l)
added H: HCO3(aq) H(aq) ˇ H2CO3(aq)
Keep in mind that the level of CO2 in the blood, and therefore the level
of carbonic acid, is ultimately controlled by the lungs. If the amount of
H in the blood increases, a large amount of H2CO3 is produced, which
lowers the H concentration. In order to reduce the carbonic acid concentration produced, the lungs work to remove CO2. Yawning is a mechanism that your body uses to get rid of extra CO2.
On the other hand, rapid and deep breathing can cause a deficiency of
carbon dioxide in the blood. This problem is called hyperventilation. It
often occurs when a person is nervous or frightened. In this case, the air
in the lungs is exchanged so rapidly that too much CO2 is released. CO2
blood levels drop, which causes the amount of carbonic acid in the blood
to decrease. This causes the blood pH to increase and can be fatal if steps
are not taken to stop this type of breathing.
When a person hyperventilates, he or she needs to become calm and
breathe regularly. If the person breathes with a paper bag covering his or
her nose and mouth, the concentration of CO2 is increased in the air
breathed. More CO2 is forced into the blood through the lungs, and the
pH of the blood drops to its normal level. Figure 15.12 shows the narrow
pH range of blood and summarizes the behavior of the buffer that controls the pH.
15.2
7.8
Death
Serious
illness
pH
7.45
7.35
Normal
Serious
illness
6.8
Death
Figure 15.12
Blood pH
The pH of human blood is
maintained within a narrow
range by a mixture of buffers.
The H2CO3 /HCO3 system is
one of the important parts of
the blood buffer.
Applications of Acid-Base Reactions
533
Chemistry
Hiccups
Normal Breathing and Hiccuping
Have you ever wondered what causes hiccups? Hiccups afflict almost
everyone, but remain a scientific puzzle. They serve no useful purpose.
What are hiccups? Most of the time,
hiccups are a harmless annoyance lasting a few minutes to several hours.
They are repeated, involuntary spasms
of your diaphragm, which is the domeshaped breathing muscle separating the
chest area from the abdomen. When
hiccups start, the diaphragm jerks, and
the air coming in is stopped when a
small flap (the epiglottis) suddenly
closes the opening to the windpipe (the
glottis), resulting in the familiar “hic.”
No one is sure why this happens. It
might start in a hiccup center in the
brain stem by abnormal stimulation of
nerves that control the diaphragm and
the glottis.
A When you
breathe in, the
diaphragm
contracts and
flattens. The
glottis and
epiglottis
open, and air
rushes in.
Position of
ribs when
exhaling
Acids and Bases React
Lung
when
inhaling
Position of
diaphragm
when
exhaling
Position of
diaphragm
when
inhaling
C When you hiccup, your diaphragm twitches. Air
is forced past the glottis and epiglottis, which
snap shut and cause a hiccup.
Possible cures A number of common cures stop
the rhythmic reflex. These methods include massaging the area with a cotton swab, gargling with water,
sipping ice water, eating a spoonful of dry sugar, and
biting on a lemon. Interruptions of normal breathing such as sneezing, coughing, and sudden pain or
fright may stop hiccups.
Most cures that seem to work are related to an
increase of CO2 in the blood, which slightly lowers
the pH. The body has a buffer system to maintain a
blood pH range of 7.35 to 7.45. Slight decreases in
pH might turn off certain nervous controls that
cause hiccups.
Chapter 15
Lung
when
exhaling
Position of
ribs when
inhaling
How are they caused? Hiccups occur
when your stomach is distended due
to gas, overeating, or drinking carbonated beverages. Other causes are sudden temperature changes, such as
drinking hot or cold beverages or taking a cold shower, and excitement or stress.
534
Epiglottis
Glottis
Larynx
B When you
breathe out,
the diaphragm
relaxes and
resumes its
dome shape,
pushing the
air out.
Exploring Further
1. Comparing and Contrasting Compare and
contrast possible cures for hiccups.
2. Hypothesizing Hypothesize as to why breathing in and out of a paper bag or holding your
breath might stop hiccups.
For more information about the mystery of hiccups,
visit the Chemistry Web site at chemistryca.com
Acid Rain Versus Acid Lakes
In Chapter 14, you learned about the sources of acid rain and about its
impact on plant and animal life, as well as on human-made objects such
as monuments and buildings. When acid rain falls on lakes and streams, it
might be expected that the acidity of the water increases and the pH
decreases.
This effect is true in some cases. Many lakes in the northeastern United
States, southern Canada, northern Europe, and the Scandinavian countries have pHs as low as 4.0. The pH of a healthy lake is about 6.5.
Other lakes, many of which are in the midwestern United States, appear
to get the same amount of acid rain as these low-pH lakes, but they do
not show a dramatic lowering of pH. The key to the ability of these lakes,
one of which is shown in Figure 15.13, to resist the pH-lowering effects of
acid rain is their local geology.
Figure 15.13
Buffered Lakes
If the rock and soil that compose and surround a lake bed
are rich in limestone, the lake
can neutralize acid rain by
acid-base reactions. These
lakes have a capacity to
absorb acid rain without an
appreciable change in pH. The
water in these lakes behaves
as a buffer.
Limestone is primarily calcium carbonate, CaCO3. Calcium carbonate
reacts with carbon dioxide and water to form calcium hydrogen carbonate, Ca(HCO3)2, a water-soluble compound. Lakes in areas rich in limestone have significant concentrations of hydrogen carbonate ions.
CaCO3(s) CO2(g) H2O(l) ˇ Ca(HCO3)2(aq)
Just as in the regulation of pH in the blood, hydrogen carbonate ions
produced from calcium hydrogen carbonate form a base that can neutralize acid in lakes.
HCO3(aq) H(aq) ˇ CO2(g) H2O(l)
The Acid-Base Chemistry of Antacids
The acid-base chemistry of stomach upset is big business. Even if you
never use commercial antacids, a lot of other people do. The terms acid
indigestion or acid stomach are a bit misleading. You need an acid stomach
in order to be healthy. Remember that the pH of stomach acid, which is
mostly hydrochloric acid, is about 2.5.
15.2
Applications of Acid-Base Reactions
535
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
Some Stomach Chemistry
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
@@@@@@@@@
@@@@@@@@@@@@@
Mucous
membrane
Stomach
An alternative
approach to controlling
acid indigestion is now
available in over-thecounter medication.
These drugs work by
decreasing the secretion of acid in the
stomach. These drugs
were previously available only by prescription for persons who
suffer severe acid indigestion or who have a
tendency for gastric
ulcers.
The combination of an acidic environment and enzymes in the stomach works together to break down the complex molecules that you eat
into smaller molecules that can be transported through your blood and
delivered to every cell in your body. These smaller molecules are the
source of energy and structural material for the cells.
Your stomach is made of protein that is not much different from the
protein in a hamburger. The inner walls of your stomach are coated by a
basic mucous membrane that protects the stomach from the digestive
power of the acid and enzymes. If the stomach contents become too
acidic, this basic membrane breaks down by acid-base neutralization reactions. At the point where the barrier is neutralized, the gastric juices can
begin to digest the protein that makes up the stomach wall, as shown in
Figure 15.14. This causes the discomfort of acid indigestion and may
cause more serious problems.
In most individuals, this breakdown of the mucous membrane corrects
itself in a short time, and there is no long-term damage to the stomach.
However, in some cases, the damage may be more long-term, and if not
treated, a stomach ulcer may result. The correction process can be speeded
up by using an antacid.
Antacids Anti-Acids Bases
Although a variety of over-the-counter and prescription antacids is
available, they all share common acid-base chemistry. Table 15.2 gives a
list of the bases that are most widely used in antacids. The compounds
can be divided into two categories: the hydroxide-containing bases and
the carbonate-containing bases.
Table 15.2 Compounds Used in Antacids
Insoluble Hydroxides
Aluminum hydroxide, Al(OH)3
Magnesium hydroxide, Mg(OH)2
536
Chapter 15
Figure 15.14
Stomach Ulcer
A gastric ulcer forms when
the membrane that protects
the stomach lining breaks
down and stomach acid
attacks the stomach wall.
Acids and Bases React
Carbonate-Based
Calcium carbonate, CaCO3
Magnesium carbonate, MgCO3
Sodium hydrogen carbonate, NaHCO3
Potassium hydrogen carbonate, KHCO3
Chemistry and
The Development of
Artificial Blood
You have a serious accident and need a blood
transfusion immediately. There is no time to type
your blood. Would you rather have someone else’s
blood or artificial blood? Right now, this choice is
not available, but it may be in the near future.
Under the described situation, artificial blood
might be the better choice.
Why is artificial blood needed? The fear of
blood transfusions is on the rise. Even though
many safeguards are used, people are afraid of
contracting diseases such as AIDS and hepatitis.
Shortages of certain types of blood occur frequently, especially if a catastrophic disaster happens. Some people think the solution to these
problems is artificial blood. Researchers have
invested hundreds of millions of dollars to develop
such a substance.
Types of artificial blood In real blood, hemoglobin in red blood cells carries oxygen. Many types
of artificial blood focus on the hemoglobin aspect
of the blood. One group is working on using the
hemoglobin from outdated human blood from
blood banks. Other scientists obtain it from genetically engineered bacteria, cattle, or pigs. Hemoglobin that has been removed or that was never in
red blood cells causes problems. Its molecule falls
apart. It does not carry oxygen well, and it sometimes clogs blood vessels. Scientists from different
companies are working on these problems. Some
are doing animal testing, some are already testing
humans, and others have plans to do so in the near
future.
Another approach to this problem is to use an
oxygen-carrying perfluorocarbon emulsion that
has been around since 1966. At that time, its oxygen-holding power was demonstrated by the fact
that mice could breathe oxygen and survive when
submerged in it. It is not considered a blood substitute by its manufacturer, but rather a drugdelivery system.
Artificial success? It will probably be several
years before most oxygen-carrying blood substitutes will be approved by the FDA. However, one
artificial blood component is currently in use. In
the past, hemophiliacs have run a much higher
risk of getting HIV because many units of blood
were needed to get enough blood-clotting factor
VIII. Now, the FDA has approved two genetically
engineered factor VIIIs—Recombinate and
Kogenate—so people who need factor VIII have a
safe alternative.
Artificial
hemoglobin
Analyzing the Issue
1. Acquiring Information Research the structure
of hemoglobin, and make a simplified model of
hemoglobin.
2. Inferring Analyze why scientists working to
develop artificial blood need an extensive
knowledge of chemistry.
3. Thinking Critically What problems might arise
with government control of artificial blood
research? What advantages might there be?
15.2
Applications of Acid-Base Reactions
537
Hydroxide Antacids
If you examine the ingredient list on an antacid, you probably recognize any hydroxide-containing compounds as bases. The hydroxides used
in antacids have low solubility in water. Because the pH of saliva is neutral
to slightly basic, these insoluble hydroxides do not dissolve and react until
they get past your mouth and upper digestive tract and into the highly
acidic environment of the stomach.
Milk of magnesia, a suspension of magnesium hydroxide in water, is a
good example of this type of antacid. If you have ever used unflavored
milk of magnesia, you may have noticed that it has a bitter taste, which is
typical of a base.
Carbonate Antacids
Figure 15.15
Antacids
Antacids are hydroxidecontaining bases or carbonate-containing compounds.
You are familiar with the acid-neutralizing abilities of the carbonates
and hydrogen carbonates from the discussion about blood buffering and
acid rain. Carbonate and hydrogen carbonate antacids react with HCl to
form carbonic acid, which decomposes into carbon dioxide gas and water.
CaCO3(s) 2HCl(aq) ˇ CaCl2(aq) H2CO3(aq)
NaHCO3(aq) HCl(aq) ˇ NaCl(aq) H2CO3(aq)
H2CO3(aq) ˇ H2O(l) CO2(g)
Many carbonates and hydrogen carbonates are insoluble in water and
have great neutralizing power. They are the primary components of many
over-the-counter antacid tablets. See an example of such an antacid in
Figure 15.15.
As with any over-the-counter medication, antacids are designed for
occasional use. Anyone who needs to use them frequently for attacks of
indigestion may have a more serious health problem and should consult a
physician.
Some antacids contain a carbonate or hydrogen
carbonate and a weak acid, such as citric acid.
When added to water, these compounds react,
producing a basic salt and carbonic acid, which
decomposes into water and carbon dioxide. 䊳
538
Chapter 15
Acids and Bases React
Stoichiometry Revisited: Acid-Base Titrations
Although acid-base reactions do not have to happen in aqueous solutions, many are water-based. In these reactions, it is frequently important
to know the concentrations of the solutions involved.
The general process of determining the molarity of an acid or a base
through the use of an acid-base reaction is called an acid-base titration.
In a titration, the molarity of one of the reactants, acid or base, is known,
but the other is unknown. The known reactant molarity is used to find
the unknown molarity of the other solution. Solutions of known molarity
that are used in this fashion are called standard solutions.
Titration of an Acid with a Base
Suppose you are in charge of cleaning up and doing inventory in the
chemistry stockroom. Your first job is to catalog the variety of solutions
that are on the shelf. One of the solutions you find is a tightly stoppered,
well-labeled bottle of 0.100M NaOH.
On the next shelf, you find a tightly stoppered bottle that is not as welllabeled. The label has been damaged, and the only thing left of the label is
M HCl. This tells that the solution contains HCl, but it does not tell the
molarity. For the inventory form, you need to know the molarity of the
solution. What can you do?
Do a Titration
You know that NaOH and HCl react completely.
HCl(aq) NaOH(aq) ˇ NaCl(aq) H2O(l)
You know the concentration of the NaOH solution, so it is your standard solution. You can use the reaction, the volumes of acid and base
used, plus the molarity of the base to determine the molarity of the unlabeled HCl. Follow the first part of this process in Figure 15.16.
Figure 15.16
Set Up the Titration
A titration requires mixing measured volumes of
the standard solution and the unknown solution.
To add the NaOH to the HCl and
to measure the amount of NaOH
solution needed, a burette is used.
A burette is a long, calibrated tube
with a valve at the bottom. 䊳
䊴 You need to carefully measure a volume
of the unknown HCl solution into a
flask. For careful volume measurements,
you can use a pipet. For this titration,
add 20.0 mL of the unknown acid to the
flask using a 20.0-mL pipet.
15.2
Applications of Acid-Base Reactions
539
Using Indicators in a Titration
When the solution is neutral, you know that you have added exactly
enough base to react with the amount of acid present, and you are at what
is known as the endpoint of the titration. But how do you know when the
endpoint is reached? Probably the best way to indicate the endpoint of a
titration is to use an acid-base indicator. Different indicators change color
at different pH values. Look at the indicators in Figure 15.17. For a titration of NaOH and HCl, the endpoint is reached when the solution reaches a pH of 7. Therefore, you need an indicator that changes color as close
to pH 7 as possible. The best indicator for this titration is bromothymol blue, which changes from yellow to blue close to pH 7.
Figure 15.17
Indicators and Titration
The graphs show some of the
more popular acid-base indicators used in the chemistry
laboratory and the pH at
which they change color.
Titration Process
Assume you have a burette containing an NaOH standard solution and
a flask containing 20.0 mL of HCl of unknown concentration, as
described in Figure 15.16. You know what indicator to use, according to
Figure 15.17. To actually perform the NaOH-HCl titration, follow the
process outlined in Figure 15.18, and record the data collected.
The reaction between a strong acid and a strong
base results in a completely neutral solution. Bromothymol blue is an effective indicator for such
reactions because it changes color at a pH of 7. 䊳
14
12
pH
10
14
6
12
4
pH
2
Phenolphthalein
pH = 8.80 at endpoint
10
8
0
Bromothymol blue
10 20 30 40 50 60 70
Volume of strong base added (mL)
6
4
Phenolphthalein
Bromothymol blue
pH = 7.00
at endpoint
Methyl red
8
Methyl red
2
0
10 20 30 40 50 60 70
Volume of strong base added (mL)
14
12
10
pH
䊱 Because the reaction between a weak acid and a
strong base results in a slightly basic solution, the
endpoint pH for a weak acid-strong base titration
is greater than 7. For such a titration, phenolphthalein changes color at the endpoint.
8
6
4
The titration of a weak base with a
strong acid has an endpoint pH that is
less than 7. For this titration, methyl red
is a good indicator because its pH range
matches closely the endpoint pH. 䊳
540
Chapter 15
Acids and Bases React
Phenolphthalein
Bromothymol blue
pH = 5.27
at endpoint
Methyl red
2
0
10 20 30 40 50 60 70
Volume of strong acid added (mL)
Figure 15.18
Titration Process
The following process shows how to perform
and obtain data for an HCl-NaOH titration.
A Fill the burette with the
NaOH solution. Be sure to
record the exact reading
on the burette. 䊳
䊴 B Add a few drops of bromothymol
blue indicator to the HCl solution in
the flask. The solution turns yellow.
C Titrate the HCl with the NaOH by slowly
adding NaOH from the burette to the flask as
the solution is constantly stirred. Continue to
add the NaOH slowly until the color changes
from yellow to blue. 䊲
D Measure the final volume
of NaOH solution, which
is 19.9 mL in this example. This reading means
that the volume of NaOH
used in this titration is
19.9 mL. 䊲
Determining Concentration: Using Stoichiometry
You now have all your experimental data. How do you combine these
experimental data into an experimental result—the molarity of the HCl
solution?
First, summarize what you know. You know that 20 mL of the HCl
solution reacts with 19.9 mL of 0.100M NaOH solution to reach the endpoint. From the balanced equation for the reaction, you know that one
mole of HCl reacts with one mole of NaOH. Therefore, the number of
moles of HCl in 20.0 mL of the HCl solution equals the number of moles
of NaOH in 19.9 mL of 0.100M NaOH solution.
Now, use the factor label method to solve this solution stoichiometry
problem, just as you used it to solve other stoichiometry problems.
Because you know the concentration of the NaOH solution, first find the
number of moles of NaOH involved in the reaction.
19.9 mL soln
–3
1 L 0.100 mol NaOH
1.99 10 mol NaOH
103 mL 1.00 L soln
Next, examine the balanced equation for the reaction and determine
that, because their coefficients are the same, equal numbers of moles of
NaOH and HCl react.
15.2
Applications of Acid-Base Reactions
541
SMALL
SCALE
Titration of Vinegar
Vinegar is a solution of mostly acetic acid in
water. It varies in concentration from about three
percent to five percent by volume. The acetic acid
in vinegar may be neutralized by adding sodium
hydroxide.
HC2H3O2(aq) NaOH(aq) ˇ
H2O(l) NaC2H3O2(aq)
You will titrate several brands of commercial
vinegar with sodium hydroxide solution of a
known concentration, and use your data to calculate the molarities and volume percentages of
acetic acid in the vinegars.
Problem
What are the volume percentages of acetic acid
in several brands of vinegar?
Objectives
Observe acid-base titrations of several vinegars
with a standard sodium hydroxide solution.
Calculate the volume percentages of acetic acid
in the vinegars.
Compare the acetic acid concentrations of various brands of vinegar.
•
•
•
542
Chapter 15
Acids and Bases React
Materials
24-well microplate
several brands of vinegar
labeled microtip pipets
standard NaOH solution
phenolphthalein solution
toothpicks
distilled water in a wash bottle
Safety Precautions
Sodium hydroxide is caustic and can damage
skin and eyes. If you come into contact with
any of this solution, rinse the affected area with
a large volume of water and notify the teacher.
Wash hands thoroughly when you complete
the lab.
1. Make a data table like the one shown.
2. Use a microtip pipet to add ten drops of the
first type of vinegar to each of the wells—A1,
B1, and C1—of the microplate. Record the
brand of the first vinegar in your data table.
3. Use a clean microtip pipet to add one drop of
phenolphthalein indicator solution to each of
the three wells.
4. Set the microplate on a piece of white paper.
5. Use a clean pipet to carefully add a drop of the
standard NaOH solution to the solution in well
A1, and stir with a toothpick. Pause for about
30 seconds and look down through the well for
evidence of a persistent, light pink, phenolphthalein color that indicates the endpoint of the
titration. Repeat this process with each drop
until the endpoint is reached. Record the num-
ber of drops of sodium hydroxide solution
3. Comparing and Contrasting Which of the
required to titrate the vinegar to the endpoint.
brands of vinegar you tested contained the
highest volume percent of acetic acid?
6. Repeat procedure 5 with the second sample of
this vinegar, which is in well B1. If the result
differs by more than one drop from that of the
first titration, repeat again with the sample in
well C1.
1. How could you have changed the experimental
7. Repeat procedures 2 through 6 with the other
procedure in order to more accurately deterbrands of vinegar, using other columns of
mine the concentrations of the vinegars?
wells of the microplate. Record your data after
2. If the cost and volume of
each titration.
each brand of vinegar are
available, calculate the
cost per percent of
acetic acid per unit
volume for each.
1. Interpreting Data From the two closest trials,
Which is the
find the average number of drops of NaOH
best buy based
required to titrate each vinegar. Use this averupon this
age number of drops and the given molarity of
criterion?
the NaOH to calculate the molarity of acetic
acid in each of the brands of vinegar. Assuming identical volumes for drops of vinegar and
drops of NaOH solution, the ratio of reacting
volumes in liters is the same as the ratio of
reacting volumes in drops.
2. Interpreting Data Use your results to calculate
the volume percentage of acetic acid in each
brand of vinegar according to the formula:
Macetic acid (1.00 percent acetic acid/0.175M
acetic acid) percent acetic acid by volume.
Type of Vinegar
Trial 1—Drops
of NaOH
Trial 2—Drops
of NaOH
Trial 3—Drops
of NaOH
Brand A
15.2
Applications of Acid-Base Reactions
543
NaOH(aq) HCl(aq) ˇ NaCl(aq) H2O(l)
Because 1.99 103 mol NaOH react, 1.99 103 mol HCl present in
solution also react.
Finally, use the volume to find the molarity of the acid.
–3
1.99 10 mol HCl 10 3 mL
0.0995 mol HCl
0.0995M HCl
20.0 mL soln
1L
1 L soln
Based on your single titration, the molarity of the HCl solution is
0.0995M. However, before you put this value on the label, you probably
would repeat the titration for several additional trials in order to verify
your analysis and be more confident of the value on the label. Study the
following Sample Problem that gives another example of a titration.
4
SAMPLE PROBLEM
Finding Molarity
A 15.0-mL sample of a solution of H2SO4
with an unknown molarity is titrated with
32.4 mL of 0.145M NaOH to the bromothymol
blue endpoint. Based upon this titration, what
is the molarity of the sulfuric acid solution?
Analyze
the molarity of the NaOH solution
• Because
is known, the number of moles of NaOH
Set Up
•
Solve
•
involved in the titration can be calculated.
The corresponding number of moles of
H2SO4 can then be determined, and this
figure can be used to calculate the molarity of
the acid.
Write the balanced equation for the reaction. Remember that sulfuric
acid is a diprotic acid.
H2SO4(aq) 2NaOH(aq) ˇ 2H2O(l) Na2SO4(aq)
Because the concentration of the NaOH solution is known, find the
number of moles of NaOH used in the titration.
32.4 mL soln 1 L
0.145 mol NaOH
–3
4.70 10 mol NaOH
3
10 mL 1.00 L soln
the balanced equation, find the number of moles of H SO that
• Using
react with 4.70 10 mol NaOH.
2
3
4
–3
4.70 10 mol NaOH 1 mol H2SO4
–3
2.35 10 mol H2SO4
2 mol NaOH
• Find the concentration of the H SO .
2
–3
4
3
2.35 10 mol H2SO4 10 mL
0.157 mol H2SO4
0.157M H2SO4
15.0 mL soln
1L
1 L soln
Check
544
Chapter 15
to make sure that the final units are what they are supposed
• Check
to be.
Acids and Bases React
Indicators
Indicators are acids and bases that have complicated structures and change color as they lose or gain hydrogen ions.
The molecules must have acidic hydrogens, and they must
have a relatively large number of carbon-carbon double
bonds. The structure of the indicator thymol blue,
H2C27H28O5S, satisfies both of these criteria.
1. At low pH, thymol blue
contains two acidic
hydrogens. In this form,
the molecule is red.
3. At pH of 8.9, the
thymol blue molecule
loses its second acidic
hydrogen. With no
acidic hydrogens, the
molecule is blue.
2. When the pH rises above 2, the concentration of H drops sufficiently
that the molecule loses one of its
acidic hydrogens. The monoprotic
form of thymol blue is yellow.
However, a single indicator can give only a
general indication of pH. For example, if a
solution is blue with thymol blue, it does not
reveal the actual pH. It indicates only that the
pH is above 8.9. By combining indicators that
span the range of pH from 1 to 14, it is possible
to create color changes that allow more precise
measurement of pH. pH paper is such a mixture. When the color of the paper is compared
to the color chart, the pH can be determined to
within about one pH unit.
Thinking Critically
1. What are two characteristics that molecules of indicators
must have?
2. What do you think is
the relationship
between the number
of acidic hydrogens
15.2
in an indicator molecule and the number
of color changes?
3. What might be some
advantages of using
indicator papers
instead of a pH
meter?
Applications of Acid-Base Reactions
545
PRACTICE PROBLEMS
For more practice with solving
problems, see Supplemental
Practice Problems,
Appendix B.
11. A 0.100M LiOH solution was used to titrate an HBr solution of
unknown concentration. At the endpoint, 21.0 mL of LiOH solution had neutralized 10.0 mL of HBr. What is the molarity of the
HBr solution?
12. A 0.150M KOH solution fills a burette to the 0 mark. The solution
was used to titrate 25.0 mL of an HNO3 solution of unknown
concentration. At the endpoint, the burette reading was 34.6 mL.
What was the molarity of the HNO3 solution?
13. A Ca(OH)2 solution of unknown concentration was used to
titrate 15.0 mL of a 0.125M H3PO4 solution. If 12.4 mL of
Ca(OH)2 are used to reach the endpoint, what is the concentration of the Ca(OH)2 solution?
Connecting Ideas
Acid-base reactions are usually double-displacement reactions. If you
examine the oxidation number of each element involved in a doubledisplacement reaction, you can see that the oxidation numbers of the elements do not change. Are there types of reactions in which the oxidation
numbers do change?
Many important reactions, such as the rusting of an old car or the
burning of a fuel, involve chemical reactions in which oxidation numbers
do change. It’s important to find out what causes these changes and investigate this important type of reaction.
SECTION REVIEW
Understanding Concepts
1. How does a solution that contains dissolved
ammonia and ammonium chloride act as a
buffer? Use net ionic equations to show how
this buffer responds to added H and OH.
2. Use chemical equations to show that the buffer
system in blood works like the buffer in question 1.
3. Sequence the steps in an acid-base titration.
Thinking Critically
4. Applying Concepts The following are the endpoint pHs for three titrations. From the endpoint pH, indicate whether the titration
involves a weak acid-strong base, a weak base546
Chapter 15
Acids and Bases React
strong acid, or a strong acid-strong base reaction. Use Figure 15.17 to select the best indicator for the reaction. What color will you see at
the endpoint?
a) pH 4.65 b) pH 8.43 c) pH 7.00
Applying Chemistry
5. Antacids You wish to compare two different
antacid tablets, brand X and brand Y. You crush
each tablet and add each to 100 mL of 1.00M
HCl. After stirring, you titrate the leftover HCl
in the resulting solution with 1.00M NaOH.
The brand X tablet requires less NaOH than the
brand Y tablet requires. Which antacid neutralizes more acid?
chemistryca.com/self_check_quiz
CHAPTER 15 ASSESSMENT
REVIEWING MAIN IDEAS
15.1 Acid and Base Reactions
■
■
■
■
Acid-base reactions are classified by the
strength of the acid and base. Three reactions
are of interest: strong acid-strong base, weak
acid-strong base, and weak base-strong acid.
Representing acid-base reactions by ionic and
net ionic equations shows what is happening
submicroscopically.
Acids and bases in reactions can be identified
using a hydrogen-ion transfer definition. An
acid is an H donor; a base is an H acceptor.
When acids and bases react, the pH of the
final solution is dependent upon the nature
of the reactants.
15.2 Applications of Acid-Base Reactions
■
A buffer is a solution that maintains a relatively constant pH when H ions or OH
ions are added.
■
■
■
The pH of blood is controlled in part by a
buffer composed of carbonic acid, H2CO3,
and the hydrogen carbonate ion, HCO3.
Antacids are bases that react with stomach acid.
An acid-base titration uses an acid-base reaction to determine the molarity of an
unknown acid or base.
Vocabulary
For each of the following terms, write a sentence that shows
your understanding of its meaning.
buffer
ionic equation
net ionic equation
neutralization reaction
salt
spectator ion
standard solution
titration
UNDERSTANDING CONCEPTS
1. What are the three types of acid-base reactions
that always go to completion? Give an example
of each by writing an overall equation.
2. Describe what it means for a reaction to go to
completion.
3. Write and balance the overall equation for each
of the following reactions. Identify the type of
acid-base reaction represented by the equation.
a) potassium hydroxide phosphoric acid
b) formic acid, HCHO2 calcium hydroxide
c) barium hydroxide sulfuric acid
4. For each of the reactions in question 3, write
the ionic and net ionic equations.
5. For each of the reactions in question 4, what
ions are spectators? Will the final reaction mixture be acidic, basic, or neutral? Explain.
chemistryca.com/vocabulary_puzzlemaker
6. In words, not equations, explain the use and
differences in the overall, ionic, and net ionic
equations.
7. Define a buffer solution.
8. Write the pH control reactions for the carbon
dioxide-based buffer in blood.
9. Why is it incorrect to define a buffer as a solution that maintains a constant pH?
Chapter 15
Assessment
547
CHAPTER 15 ASSESSMENT
10. If the H concentration in blood increases,
what happens to the concentrations of H2CO3,
HCO3, and H?
11. What role do the lungs play in regulating
blood pH?
12. Write two reactions that explain why lakes in
limestone areas are capable of resisting pH
decreases due to acid rain.
13. Antacids are classified into two types. What are
they? Give an example of each type.
14. Consider an antacid that contains aluminum
hydroxide. Write the overall equation that
shows how this antacid reduces the acidity of
stomach acid.
15. A student found that 53.2 mL of a 0.232M
solution of NaOH was required to titrate
25.0 mL of an acetic acid solution of unknown
molarity to the endpoint. What is the molarity
of the acetic acid solution?
16. A student neutralizes 30.0 mL of a sample of
sodium hydroxide with 28.9 mL of 0.150M
HCl. What is the molarity of the sodium
hydroxide?
17. A student finds that 23.1 mL of 0.200M potassium hydroxide are required to react completely with 25.0 mL of a phosphoric acid solution.
What is the molarity of the H3PO4?
18. How does the endpoint pH of a strong acidstrong base titration compare with that of a
weak acid-strong base titration?
19. How does the endpoint pH of a strong basestrong acid titration compare with that of a
weak base-strong acid titration?
20. A 50.0-mL sample of an unknown monoprotic
acid is titrated to the endpoint with 45.5 mL of
0.200M Ca(OH)2. What is the molarity of the
acid solution?
21. A 50.0-mL sample of aqueous ammonia is
titrated to the endpoint with 36.3 mL of
0.100M H2SO4. What is the molarity of the
ammonia solution?
22. What is hyperventilation? How does it change
the pH of blood?
548
Chapter 15
Acids and Bases React
23. When 25.0 g of baking soda, NaHCO3, and
25.0 mL of 1.00M HCl are mixed, is the final
solution acidic, basic, or neutral?
APPLYING CONCEPTS
24. Methylamine, CH3NH2, is a weak base, as
ammonia is. When methylamine completely
reacts with hydrochloric acid, the final solution has a pH less than 7. Why are the products of this “neutralization” reaction not neutral? Use the net ionic equation to help in
your explanation.
25. How many milliliters of 0.200M HCl are
required to react with 25.0 mL of 0.100M
methylamine, CH3NH2?
26. Vitamin C is also known as ascorbic acid,
HC6H7O6. A solution made from a vitamin C
tablet is titrated to the endpoint with 12.3 mL of
0.225M NaOH. Assuming that vitamin C is the
only acid present in the tablet, how many moles
of vitamin C are in the tablet?
27. Why are magnesium hydroxide and aluminum
hydroxide effective antacids, but sodium
hydroxide is not?
28. How many milliliters of 0.100M NaOH are
required to neutralize 25.0 mL of 0.150M HCl?
29. Concentrated HCl solutions are prepared by
dissolving HCl(g) in water. Concentrated HCl
is usually sold as a 12 M HCl solution. How
many liters of HCl(g) at 25°C and 1 atmosphere pressure are required to make 10.0 liters
of 12 M HCl?
30. Complete and balance the following overall
equations.
a) KOH(aq) HNO3(aq) ˇ
b) Ba(OH)2(aq) HCl(aq) ˇ
c) NaOH(aq) H3PO4(aq) ˇ
d) Ca(OH)2(aq) H3PO4(aq) ˇ
31. Dihydrogen phosphate and monohydrogen
phosphate ions play an important role in
maintaining the pH in intracellular fluid. Write
equations that show how these ions maintain
the pH.
chemistryca.com/chapter_test
CHAPTER 15 ASSESSMENT
32. The concentration of H2CO3 in blood is 1⁄20
of the concentration of HCO3, yet the blood
buffer is capable of buffering the pH against
bases, as well as against acid. Explain.
33. Tartaric acid is often added to artificial fruit
drinks to increase tartness. A sample of a certain beverage contains 1.00 g of tartaric acid,
H2C4H4O6. The beverage is titrated with
0.100M NaOH. Assuming no other acids are
present, how many milliliters of base are
required to neutralize the tartaric acid?
34. How many grams of tartaric acid, H2C4H4O6,
must be added to 150 mL of 0.245M NaOH to
completely react?
35. Stomach acid is approximately 0.0200M HCl.
What volume of stomach acid does an antacid
tablet that contains 45.5 percent Mg(OH)2 and
weighs 355 mg neutralize?
36. Suppose the tablet in question 35 is composed
of calcium carbonate. Is this tablet more effective than the one composed of Mg(OH)2?
Chemistry and Society
37. What function of blood is most important
when developing artificial blood?
Earth Science Connection
38. What acid most likely causes groundwater to
be acidic? How does groundwater become
acidic?
Everyday Chemistry
39. How is the CO2 concentration in the blood
related to hiccups?
How It Works
40. Explain how molecules and ions are related to
taste.
How It Works
41. Can an indicator provide an exact pH?
Explain.
THINKING CRITICALLY
Observing and Inferring
42. A sample of rainwater turns blue litmus red.
Fresh portions of the rainwater turn thymol
blue indicator yellow, bromophenol blue indicator green, and methyl red indicator red. Estimate the pH of the rainwater.
Interpreting Data
43. When formic acid, HCHO2, reacts completely
with NaOH, the resulting solution has a pH
greater than 7. Why are the products of this
neutralization reaction not neutral? Use the
net ionic equation to help in your explanation.
Applying Concepts
44. Write an overall equation for the acid-base
reaction that would be required to produce
each of the following salts.
a) NaCl
c) MgCl2
e) KBr
b) CaSO4
d) (NH4)2SO4
Observing and Inferring
45. ChemLab Explain why different bottles of the
same brand of vinegar might contain solutions
that have different pHs.
Making Predictions
46. MiniLab 1 Would a solution of iron(III) bromide, FeBr3, be acidic, basic, or neutral?
Relating Cause and Effect
47. MiniLab 2 Explain why phenolphthalein and
methyl orange are used as indicators in MiniLab 2.
CUMULATIVE REVIEW
48. Give the name of the compound represented
by the formula Mn(NO3)2 4H2O, and determine how many atoms of each element are
present in three formula units of the compound. (Chapter 5)
Chapter 15
Assessment
549
CHAPTER 15 ASSESSMENT
49. Terephthalic acid is an organic compound used
in the formation of polyesters. It contains 57.8
percent C, 3.64 percent H, and 38.5 percent O.
The molar mass is known to be approximately
166 g/mol. What is the molecular formula of
terephthalic acid? (Chapter 12)
50. Write equations for the dissociation of the following ionic compounds when they dissolve in
water. (Chapter 13)
a) CuSO4
b) Ca(NO3)2
c) Na2CO3
51. What is a monoprotic acid? A triprotic acid?
Give an example of each. (Chapter 14)
SKILL REVIEW
52. Data Table Solutions of five different monoprotic acids are all 0.100M. The pH of each
solution is given. Rank the acids in the following table from weakest to strongest. For each
solution, use the indicator table in Figure
14.22 to predict the color that each solution
would produce with the given indicator.
Solution
pH
A
5.45
B
1.00
C
3.45
D
4.50
E
2.36
Solution (weakest to
strongest acid)
550
Chapter 15
pH
Acids and Bases React
WRITING IN CHEMISTRY
53. Write an article about the effect of acid rain on
a specific aspect of a local environment such as
a lake or a forest. Give some history of the
problem and indicate when local residents first
realized a problem exists. What, if any, corrective measures have been taken to correct the
problem? Is the environmental damage
reversible?
PROBLEM SOLVING
54. Because antacids are frequently insoluble in
water, they are often analyzed by dissolving
them in a known volume of HCl with a known
molarity. After the antacid has completely
reacted, there is still HCl left in the solution.
This excess HCl is then titrated with a standard
NaOH solution. A 165-mg sample of an
antacid tablet containing calcium carbonate is
dissolved in 50.0 mL of 0.100M HCl. After
complete reaction, the excess HCl is titrated
with 15.8 mL of 0.150M NaOH. Sketch a flowchart that shows the steps in the analysis.
Color in
Bromphenol Blue
Color in
Methyl Red
Color in
Thymol Blue
Standardized Test Practice
1. What is the symbol for a hydronium ion?
a) H
c) OH
b) H
d) OH
Use the chemical equation to answer questions 2
and 3.
Mg(OH)2(aq) 2HCl → MgCl2(aq) 2H2O(l)
2. Which of the compounds in the equation is
considered a base?
c) MgCl2
a) Mg(OH)2
b) HCl
d) H2O
3. Which of the compounds in the equation is
considered a salt?
c) MgCl2
a) Mg(OH)2
b) HCl
d) H2O
4. Which ions are excluded from a net ionic
equation?
a) weak acids or bases c) positive ions
b) negative ions
d) spectator ions
5. Which of the following is true about a solution
with a pH lower than 7?
a) The solution is a strong acid.
b) The quantity of hydronium ions is greater
than the quantity of hydroxide ions.
c) The quantity of hydroxide ions is greater
than the quantity of hydronium ions.
d) The solution is an equal mixture of moles
of acid and moles of base.
6. In the equation, HBr is
HBr(aq) H2O(l) → H3O (aq) Br(aq)
a) considered a strong base because it completely ionizes in water.
b) considered a weak base because it loses a
positive charge to a water molecule.
c) considered a strong acid because it completely ionizes in water.
d) considered a weak acid because it loses a
positive charge to a water molecule.
chemistryca.com/standardized_test
Weak Acid Ionization Constants
Weak Acid
Ionization Constant
Hydrofluoric acid
6.3 ⴛ 10ⴚ4
Methanoic acid
1.8 ⴛ 10ⴚ4
Ethanoic acid
1.8 ⴛ 10ⴚ5
Hypochlorous acid
4.0 ⴛ 10ⴚ8
Use the table above to answer question 7.
7. The ionization constant of a weak acid is a calculation of the number of ionized molecules (products) in a dilute aqueous solution divided by the
number of un-ionized molecules (reactants) in
the solution. Which acid is the weakest?
a) hydrofluoric acid
b) methanoic acid
c) ethanoic acid
d) hypochlorous acid
8. A Brønsted-Lowry acid is defined as a(n)
a) acid that donates a hydrogen ion during a
chemical reaction.
b) acid that accepts a hydrogen ion during a
chemical reaction.
c) strong acid that donates the maximum
number of hydrogen ions.
d) weak acid that donates a limited number of
hydrogen ions.
9. A solution of 0.600M HCl is used to titrate 15.00
mL of KOH solution. The endpoint of the titration is reached after the addition of 27.13 mL of
HCl. What is the concentration of the KOH
solution?
a) 9.000M
c) 0.332M
b) 1.09M
d) 0.0163M
Test Taking Tip
Tables If a test question involves a table, skim
the table before reading the question. Read the
title, column heads, and row heads. Then read the
question and interpret the information in the table.
Standardized Test Practice
551