Chapter 19 Acids, Bases, and Salts 19.1 Acid-Base Theories

Chapter 19 Acids, Bases, and Salts 19.1 Acid-Base Theories
19.2 Hydrogen Ions and Acidity 19.3 Strengths of Acids and Bases 19.4 Neutralization Reactions 19.5 Salts in Solution Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 1

Why are high levels of ammonia harmful to you?
CHEMISTRY & YOU Why are high levels of ammonia harmful to you? Nitrogen compounds in bat urine can decompose and release ammonia into the air. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 2

Arrhenius Acids and Bases
How did Arrhenius define an acid and a base? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 3

Acids and bases have distinct properties.
Arrhenius Acids and Bases Acids and bases have distinct properties. Acids give foods a tart or sour taste. Aqueous solutions of acids are strong or weak electrolytes. Acids cause certain dyes, called indicators, to change color. Many metals, such as zinc and magnesium, react with aqueous solutions of acids to produce hydrogen gas. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 4

Acids and bases have distinct properties.
Arrhenius Acids and Bases Acids and bases have distinct properties. Soap is a familiar material that has the properties of a base. The bitter taste is a general property of bases. The slippery feel of soap is another property of bases. Bases will cause an indicator to change color. Bases also form aqueous solutions that are strong or weak electrolytes. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 5

In 1887, the Swedish chemist Svante Arrhenius proposed a new way of defining and thinking about acids and bases. According to Arrhenius, acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in solution. Bases are compounds that ionize to yield hydroxide ions (OH–) in aqueous solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 6

Acids vary in the number of hydrogens they contain that can form hydrogen ions. Some Common Acids Name Formula Hydrochloric acid HCl Nitric acid HNO3 Sulfuric acid H2SO4 Phosphoric acid H3PO4 Ethanoic acid CH3COOH Carbonic acid H2CO3 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 7

A hydrogen atom that can form a hydrogen ion is described as ionizable. Nitric acid (HNO3) has one ionizable hydrogen. Nitric acid is classified as a monoprotic acid. The prefix mono- means “one,” and the stem protic reflects the fact that a hydrogen ion is a proton. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 8

A hydrogen atom that can form a hydrogen ion is described as ionizable. Nitric acid (HNO3) has one ionizable hydrogen. Nitric acid is classified as a monoprotic acid. Acids that contain two ionizable hydrogens, such as sulfuric acid (H2SO4), are called diprotic acids. Acids that contain three ionizable hydrogens, such as phosphoric acid (H3PO4), are called triprotic acids. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 9

Not all compounds that contain hydrogen are acids.
Arrhenius Acids and Bases Arrhenius Acids Not all compounds that contain hydrogen are acids. Only a hydrogen that is bonded to a very electronegative element can be released as an ion. Such bonds are highly polar. When a compound that contains such bonds dissolves in water, it releases hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 10

In an aqueous solution, hydrogen ions are not present. Instead, the hydrogen ions are joined to water molecules as hydronium ions. A hydronium ion (H3O+) is the ion that forms when a water molecule gains a hydrogen ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 11

Methane (CH4) is an example of a hydrogen-containing compound that is not an acid. The four hydrogen atoms in methane are attached to the central carbon atom by weakly polar C—H bonds. Methane has no ionizable hydrogens and is not an acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 12

Ethanoic acid (CH3COOH), which is commonly called acetic acid, is an example of a molecule that contains both hydrogens that do not ionize and a hydrogen that does ionize. Although its molecules contain four hydrogens, ethanoic acid is a monoprotic acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 13

The three hydrogens attached to a carbon atom are in weakly polar bonds. They do not ionize. Only the hydrogen bonded to the highly electronegative oxygen can be ionized. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 14

The table below lists four common bases.
Arrhenius Acids and Bases Arrhenius Bases The table below lists four common bases. Some Common Bases Name Formula Solubility in Water Sodium hydroxide NaOH High Potassium hydroxide KOH Calcium hydroxide Ca(OH)2 Very low Magnesium hydroxide Mg(OH)2 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 15

The base sodium hydroxide (NaOH) is known as lye.
Arrhenius Acids and Bases Arrhenius Bases The base sodium hydroxide (NaOH) is known as lye. Sodium hydroxide is an ionic solid. It dissociates into sodium ions in aqueous solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 16

The base sodium hydroxide (NaOH) is known as lye.
Arrhenius Acids and Bases Arrhenius Bases The base sodium hydroxide (NaOH) is known as lye. Sodium hydroxide is extremely caustic. A caustic substance can burn or eat away materials with which it comes in contact. This property is the reason that sodium hydroxide is a major component of products that are used to clean clogged drains. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 17

Potassium hydroxide (KOH) is another ionic solid.
Arrhenius Acids and Bases Arrhenius Bases Potassium hydroxide (KOH) is another ionic solid. It dissociates to produce potassium ions and hydroxide ions in aqueous solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 18

Arrhenius Bases Sodium and potassium are Group 1A elements. Elements in Group 1A, the alkali metals, react violently with water. The products of these reactions are aqueous solutions of a hydroxide and a hydrogen gas. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 19

Sodium hydroxide and potassium hydroxide are very soluble in water.
Arrhenius Acids and Bases Arrhenius Bases Sodium hydroxide and potassium hydroxide are very soluble in water. The solutions would typically have the bitter taste and slippery feel of a base, but you would not want to test these properties. The solutions are extremely caustic to the skin. They can cause deep, painful, slow-healing wounds if not immediately washed off. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 20

CHEMISTRY & YOU Visitors to Bracken Cave wear protective gear to keep ammonia gas out of their eyes and respiratory tracts. Think about the properties of bases. Why are high levels of ammonia harmful? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 21

Ammonia is a base, and bases are caustic in high concentrations.
CHEMISTRY & YOU Visitors to Bracken Cave wear protective gear to keep ammonia gas out of their eyes and respiratory tracts. Think about the properties of bases. Why are high levels of ammonia harmful? Ammonia is a base, and bases are caustic in high concentrations. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 22

Arrhenius Bases Calcium hydroxide, Ca(OH)2, and magnesium hydroxide, Mg(OH)2, are compounds of Group 2A metals. These compounds are not very soluble in water. Their solutions are always very dilute, even when saturated. The low solubility of magnesium hydroxide makes the suspension safe to consume. Some people use this suspension as an antacid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 23

Can every hydrogen from every molecule form hydrogen ions, therefore acting as an Arrhenius acid?
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Can every hydrogen from every molecule form hydrogen ions, therefore acting as an Arrhenius acid?
No. Only hydrogens that are bonded to a very electronegative element can be released as ions. That means that only molecules containing hydrogens bonded to very electronegative elements are Arrhenius acids. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 25

Brønsted-Lowry Acids and Bases
What distinguishes an acid from a base in the Brønsted-Lowry theory? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 26

Sodium carbonate (Na2CO3) and ammonia (NH3) act as bases when they form aqueous solutions. Neither of these compounds is a hydroxide-containing compound, so neither would be classified as a base by the Arrhenius definition. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 27

In 1923, the Danish chemist Johannes Brønsted and the English chemist Thomas Lowry were working independently. Each chemist proposed the same definition of acids and bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 28

According to the Brønsted-Lowry theory, an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. This theory includes all the acids and bases that Arrhenius defined. It also includes some compounds that Arrhenius did not classify as bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 29

You can use the Brønsted-Lowry theory to understand why ammonia is a base. When ammonia dissolves in water, hydrogen ions are transferred from water to ammonia to form ammonium ions and hydroxide ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 30

You can use the Brønsted-Lowry theory to understand why ammonia is a base. When ammonia dissolves in water, hydrogen ions are transferred from water to ammonia to form ammonium ions and hydroxide ions. Ammonia is a Brønsted-Lowry base because it accepts hydrogen ions. Water is a Brønsted-Lowry acid because it donates hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 31

Conjugate Acids and Bases
Brønsted-Lowry Acids and Bases Conjugate Acids and Bases When the temperature of an aqueous solution of ammonia is increased, ammonia gas is released. HNH4+ reacts with OH– to form more NH3 and H2O. In the reverse reaction, ammonium ions donate hydrogen ions to hydroxide ions. NH4+ (the donor) acts as a Brønsted-Lowry acid, and OH− (the acceptor) acts as a Brønsted-Lowry base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 32

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases In essence, the reversible reaction of ammonia and water has two acids and two bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 33

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases A conjugate acid is the ion or molecule formed when a base gains a hydrogen ion. NH4+ is the conjugate acid of the base NH3. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 34

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases A conjugate base is the ion or molecule that remains after an acid loses a hydrogen ion. OH– is the conjugate base of the acid H2O. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 35

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases Conjugate acids are always paired with a base, and conjugate bases are always paired with an acid. A conjugate acid-base pair consists of two ions or molecules related by the loss or gain of one hydrogen ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 36

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases The ammonia molecule and the ammonium ion are a conjugate acid-base pair. The water molecule and the hydroxide ion are also a conjugate acid-base pair. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 37

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases In this reaction, hydrogen chloride is the hydrogen-ion donor and is by definition a Brønsted-Lowry acid. Water is the hydrogen-ion acceptor and a Brønsted-Lowry base. The chloride ion is the conjugate base of the acid HCl. The hydronium ion is the conjugate acid of the water base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 38

Brønsted-Lowry Acids and Bases Conjugate Acids and Bases The figure below shows the reaction that takes place when sulfuric acid dissolves in water. The products are hydronium ions and hydrogen sulfate ions. Use the figure to identify the two conjugate acid-base pairs. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 39

Some Conjugate Acid-Base Pairs
Interpret Data Some Conjugate Acid-Base Pairs Acid Base HCl Cl– H2SO4 HSO4– H3O+ H2O SO42– CH3COOH CH3COO– H2CO3 HCO3− HCO3– CO32– NH4+ NH3 OH– Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 40

Some Conjugate Acid-Base Pairs
Brønsted-Lowry Acids and Bases Some Conjugate Acid-Base Pairs Acid Base HCl Cl– H2SO4 HSO4– H3O+ H2O SO42– CH3COOH CH3COO– H2CO3 HCO3− HCO3– CO32– NH4+ NH3 OH– Amphoteric Substances Note that water appears in both the list of acids and the list of bases. Sometimes water accepts a hydrogen ion. At other times, it donates a hydrogen ion. How water behaves depends on the other reactant. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 41

Amphoteric Substances
Brønsted-Lowry Acids and Bases Amphoteric Substances A substance that can act as either an acid or a base is said to be amphoteric. Water is amphoteric. In the reaction with hydrochloric acid, water accepts a proton and is therefore a base. In the reaction with ammonia, water donates a proton and is therefore an acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 42

How can one substance, such as water, be both an acid and a base, according to the Brønsted-Lowry definition? Because water can act as both a hydrogen-ion donator and a hydrogen-ion acceptor, it can act as both an acid and a base according to the Brønsted-Lowry definition. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 44

How did Lewis define an acid and a base?
Lewis Acids and Bases Lewis Acids and Bases How did Lewis define an acid and a base? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 45

Lewis Acids and Bases According to Gilbert Lewis, an acid accepts a pair of electrons and a base donates a pair of electrons during a reaction. This definition is more general than those offered by Arrhenius or by Brønsted and Lowry. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 46

Lewis Acids and Bases A Lewis acid is a substance that can accept a pair of electrons to form a covalent bond. A Lewis base is a substance that can donate a pair of electrons to form a covalent bond. The Lewis definitions include all the Brønsted-Lowry acids and bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 47

Consider the reaction of H+ and OH–.
Lewis Acids and Bases Consider the reaction of H+ and OH–. The hydrogen ion donates itself to the hydroxide ion. H+ is a Brønsted-Lowry acid, and OH− is a Brønsted-Lowry base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 48

Consider the reaction of H+ and OH–.
Lewis Acids and Bases Consider the reaction of H+ and OH–. The hydroxide ion can bond to the hydrogen ion because it has an unshared pair of electrons. OH− is also a Lewis base, and H+, which accepts the pair of electrons, is a Lewis acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 49

Lewis Acids and Bases A second example of a reaction between a Lewis acid and a Lewis base is what happens when ammonia dissolves in water. Hydrogen ions from the dissociation of water are the electron-pair acceptor and the Lewis acid. Ammonia is the electron-pair donor and the Lewis base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 50

Acid-Base Definitions
Interpret Data This table compares the definitions of acids and bases. Acid-Base Definitions Type Acid Base Arrhenius H+ producer OH– producer Brønsted-Lowry H+ donor H+ acceptor Lewis electron-pair acceptor electron-pair donor The Lewis definition is the broadest. It extends to compounds that the Brønsted-Lowry theory does not classify as acids and bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 51

Identifying Lewis Acids and Bases
Sample Problem 19.1 Identifying Lewis Acids and Bases Identify the Lewis acid and the Lewis base in this reaction between ammonia and boron trifluoride. NH3 + BF3 → NH3BF3 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 52

Analyze Identify the relevant concepts.
Sample Problem 19.1 Analyze Identify the relevant concepts. 1 When a Lewis acid reacts with a Lewis base, the base donates a pair of electrons and the acid accepts the donated pair. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 53

Solve Apply concepts to this problem.
Sample Problem 19.1 Solve Apply concepts to this problem. 2 Identify the reactant with the unshared pair of electrons and the reactant that can accept the pair of electrons. Draw electron dot structures to identify which reactant has an unshared pair of electrons. Ammonia has an unshared pair of electrons to donate. The boron atom can accept the donated electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 54

Solve Apply concepts to this problem.
Sample Problem 19.1 Solve Apply concepts to this problem. 2 Classify the reactants based on their behavior. Lewis bases donate a pair of electrons, so ammonia is the Lewis base. Lewis acids accept a pair of electrons, so boron trifluoride is the Lewis acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 55

Are hydrogen-ion donors also electron-pair acceptors?

Are hydrogen-ion donors also electron-pair acceptors?
Yes. All substances defined as acids by the Brønsted-Lowry definition (an acid is a hydrogen-ion donor) are also defined as acids by the Lewis definition (an acid is an electron-pair acceptor). That means that these substances are both hydrogen-ion donors and electron-pair acceptors. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 57

Key Concepts According to Arrhenius, acids are hydrogen-containing compounds that ionize to yield hydrogen ions in aqueous solution. Bases are compounds that ionize to yield hydroxide ions in aqueous solution. According to Brønsted-Lowry theory, an acid is a hydrogen-ion donor and a base is a hydrogen-ion acceptor. According to Lewis, an acid accepts a pair of electrons and a base donates a pair of electrons. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 58

Glossary Terms hydronium ion (H3O+): the positive ion formed when a water molecule gains a hydrogen ion conjugate acid: the particle formed when a base gains a hydrogen ion; NH4+ is the conjugate acid of the base NH3 conjugate base: the particle that remains when an acid has donated a hydrogen ion; OH– is the conjugate base of the acid water Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 59

amphoteric: a substance that can act as both an acid and a base
Glossary Terms conjugate acid-base pair: two substances that are related by the loss or gain of a single hydrogen ion; ammonia (NH3) and the ammonium ion (NH4+) are a conjugate acid- base pair amphoteric: a substance that can act as both an acid and a base Lewis acid: any substance that can accept a pair of electrons to form a covalent bond Lewis base: any substance that can donate a pair of electrons to form a covalent bond Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 60

BIG IDEA Reactions Chemists define acids and bases according to the ions they yield in aqueous solution. Chemists also define acids and bases based on whether they accept or donate hydrogen ions, and whether they are electron-pair donors or acceptors. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 61

Chapter 19 Acids, Bases, and Salts 19.2 Hydrogen Ions and Acidity
19.1 Acid-Base Theories 19.2 Hydrogen Ions and Acidity 19.3 Strengths of Acids and Bases 19.4 Neutralization Reactions 19.5 Salts in Solution Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 62

CHEMISTRY & YOU What factors do you need to control so a fish has healthy water to live in? Goldfish can live for 20 years or more in an aquarium if the conditions are right. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 63

Hydrogen Ions from Water
How are [H+] and [OH−] related in an aqueous solution? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 64

Water molecules are highly polar and are in constant motion, even at room temperature. On occasion, the collisions between water molecules are energetic enough for a reaction to occur. A water molecule that loses a hydrogen ion becomes a hydronium ion (H3O+). A water molecule that loses a hydrogen ion becomes a hydroxide ion (OH−). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 65

Self-Ionization of Water The reaction in which water molecules produce ions is called the self-ionization of water. In an aqueous solution, hydrogen ions are always joined to water molecules as hydronium ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 66

Self-Ionization of Water
Hydrogen Ions from Water Self-Ionization of Water In pure water at 25°C, the concentration of hydrogen ions is only 1 × 10−7M. The concentration of OH− is also 1 × 10−7M because the numbers of H+ and OH− ions are equal in pure water. Any aqueous solution in which [H+] and [OH−] are equal is a neutral solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 67

Ion-Product Constant for Water
Hydrogen Ions from Water Ion-Product Constant for Water The ionization of water is a reversible reaction, so Le Châtelier’s principle applies. Adding either hydrogen ions or hydroxide ions to an aqueous solution is a stress on the system. In response, the equilibrium will shift toward the formation of water. The concentration of the other ion will decrease. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 68

Hydrogen Ions from Water Ion-Product Constant for Water The ionization of water is a reversible reaction, so Le Châtelier’s principle applies. In any aqueous solution, when [H+] increases, [OH−] decreases. Likewise, when [H+] decreases, [OH−] increases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 69

Hydrogen Ions from Water Ion-Product Constant for Water For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1.0 × 10−14. [H+] + [OH−] = 1.0 × 10−14 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 70

Hydrogen Ions from Water Ion-Product Constant for Water The product of the concentrations of the hydrogen ions and the hydroxide ions in water is called the ion-product constant for water (Kw). Kw = [H+] × [OH−] = 1.0 × 10−14 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 71

HCl(aq) → H+(aq) + Cl−(aq)
Hydrogen Ions from Water Ion-Product Constant for Water Acidic Solutions When some substances dissolve in water, they release hydrogen ions. HCl(aq) → H+(aq) + Cl−(aq) The hydrogen-ion concentration is greater than the hydroxide-ion concentration. A solution in which [H+] is greater than [OH−] is an acidic solution. The [H+] is greater than 1 × 10−7M. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 72

NaOH(aq) → Na+(aq) + OH−(aq)
Hydrogen Ions from Water Ion-Product Constant for Water Basic Solutions When sodium hydroxide dissolves in water, it forms hydroxide ions in solution. NaOH(aq) → Na+(aq) + OH−(aq) The hydrogen-ion concentration is less than the hydroxide-ion concentration. A basic solution is one in which [H+] is less than [OH−]. Basic solutions are also known as alkaline solutions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 73

Using the Ion-Product Constant for Water
Sample Problem 19.2 Using the Ion-Product Constant for Water If the [H+] in a solution is 1.0 × 10−5M, is the solution acidic, basic, or neutral? What is the [OH−] of this solution? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 74

Analyze List the knowns and the unknowns.
Sample Problem 19.2 Analyze List the knowns and the unknowns. 1 Use the expression for the ion-product constant for water and the known concentration of hydrogen ions to find the concentration of hydroxide ions. KNOWNS [H+] = 1.0 × 10−5M Kw= 1.0 × 10−14 UNKNOWNS Is the solution acidic, basic, or neutral? [OH−] = ?M Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 75

Calculate Solve for the unknowns.
Sample Problem 19.2 Calculate Solve for the unknowns. 2 Use [H+] to determine whether the solution is acidic, basic, or neutral. [H+] is 1.0 × 10−5M, which is greater than 1.0 × 10−7M. Thus, the solution is acidic. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 76

Sample Problem 19.2 Calculate Solve for the unknowns. 2 Rearrange the expression for the ion-product constant to solve for [OH−]. Kw = [H+] × [OH−] [OH−] = Kw [H+] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 77

Sample Problem 19.2 Calculate Solve for the unknowns. 2 Substitute the known values of [H+] and Kw. Then solve for [OH−]. = 1.0 × 10−9M [OH−] = 1.0 × 10−14 1.0 × 10−5 When you divide numbers written in scientific notation, subtract the exponent in the denominator from the exponent in the numerator. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 78

Evaluate Does the result make sense?
Sample Problem 19.2 Evaluate Does the result make sense? 3 If [H+] is greater than 1.0 × 10−7M, then [OH−] must be less than 1.0 × 10−7M. 1.0 × 10−9M is less than 1.0 × 10−7M. To check your calculation, multiply the values for [H+] and [OH−] to make sure that the result equals 1.0 × 10−14. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 79

A solution does not have an equal number of H+ and OH− ions
A solution does not have an equal number of H+ and OH− ions. What do you know about this solution? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 80

A solution does not have an equal number of H+ and OH− ions
A solution does not have an equal number of H+ and OH− ions. What do you know about this solution? You know that the solution is not neutral. Without knowing more information, you cannot say if it is acidic or basic (alkaline). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 81

How is pH used to classify a solution as neutral, acidic, or basic?
The pH Concept The pH Concept How is pH used to classify a solution as neutral, acidic, or basic? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 82

The pH Concept Expressing hydrogen-ion concentration in molarity is not practical. A more widely used system for expressing [H+] is the pH scale. The pH scale ranges from 0 to 14. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 83

The pH Concept Hydrogen Ions and pH The pH of a solution is the negative logarithm of the hydrogen-ion concentration. pH = −log[H+] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 84

The pH Concept Hydrogen Ions and pH In pure water or a neutral solution, the [H+] = 1 × 10−7M, and the pH is 7. pH = −log(1 × 10−7) = −(log 1 + log 10−7) = −(0.0 + (−7.0)) = 7.0 If the [H+] of a solution is greater than 1× 10−7M, the pH is less than 7.0. If the [H+] is less than 1× 10−7M, the pH is greater than 7.0. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 85

A solution with a pH less than 7.0 is acidic.
The pH Concept Hydrogen Ions and pH A solution with a pH less than 7.0 is acidic. A solution with a pH of 7.0 is neutral. A solution with a pH greater than 7.0 is basic. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 86

Hydrogen Ions and pH Interpret Data
When [H+] is given in the format 1 × 10–n, it’s easy to find the pH. It’s just the absolute value of the exponent n. Also, note that [H+] × [OH–] always equals 1 × 10–14. Hydrogen Ions and pH Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 87

CHEMISTRY & YOU In an aquarium, the pH of water is another factor that affects the ability of fish to survive. Most freshwater fish need a slightly acidic or neutral pH. For a saltwater tank, the ideal pH is slightly basic. What might explain this difference in the ideal pH range? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 88

CHEMISTRY & YOU In an aquarium, the pH of water is another factor that affects the ability of fish to survive. Most freshwater fish need a slightly acidic or neutral pH. For a saltwater tank, the ideal pH is slightly basic. What might explain this difference in the ideal pH range? Natural freshwater is slightly acidic, while natural saltwater typically contains compounds that make it slightly basic. Fish have adapted to these conditions in the wild and need them replicated in their tanks. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 89

Calculating pH from [H+]
The pH Concept Calculating pH from [H+] Expressing [H+] in scientific notation can make it easier to calculate pH. You would rewrite M as 1.0 × 10−3M. The coefficient 1.0 has two significant figures. The pH for this solution is 3.00. The two numbers to the right of the decimal point represent the two significant figures in the concentration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 90

The pH Concept Calculating pH from [H+] It is easy to find the pH for solutions when the coefficient is 1.0. The pH of the solution equals the exponent, with the sign changed from minus to plus. A solution with [H+] = 1 × 10−2M has a pH of 2.0. When the coefficient is a number other than 1, you will need to use a calculator with a log function key to calculate pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 91

Sample Problem 19.3 Calculating pH from [H+] What is the pH of a solution with a hydrogen-ion concentration of 4.2 × 10−10M? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 92

Analyze List the known and the unknown.
Sample Problem 19.3 Analyze List the known and the unknown. 1 To find the pH from the hydrogen-ion concentration, you use the equation pH = −log[H+]. KNOWN [H+] = 4.2 × 10−10M UNKNOWN pH = ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 93

Calculate Solve for the unknown.
Sample Problem 19.3 Calculate Solve for the unknown. 2 Start with the equation for finding pH from [H+]. pH = −log[H+] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 94

Sample Problem 19.3 Calculate Solve for the unknown. 2 Substitute the known [H+] and use the log function on your calculator to calculate the pH. Round the pH to two decimal places because the hydrogen-ion concentration has two significant figures. pH = −log(4.2 × 10−10) = −(− ) = = 9.38 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 95

Sample Problem 19.3 Evaluate Does the result make sense? 3 The value of the hydrogen-ion concentration is between 1 × 10−9M and 1 × 10−10M. So, the calculated pH should be between 9 and 10, which it is. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 96

Calculating [H+] from pH
The pH Concept Calculating [H+] from pH You can calculate the hydrogen-ion concentration of a solution if you know the pH. If the pH is an integer, it is easy to find [H+]. For a pH of 9.0, [H+] = 1 × 10−9M. Most pH values are not whole numbers. When the pH value is not a whole number, you will need a calculator with an antilog (10x) function to get an accurate value for the hydrogen-ion concentration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 97

Calculating [H+] from pH
Sample Problem 19.4 Calculating [H+] from pH The pH of an unknown solution is What is the hydrogen-ion concentration of the solution? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 98

Analyze List the known and the unknown.
Sample Problem 19.4 Analyze List the known and the unknown. 1 You will use the antilog function of your calculator to find the concentration. KNOWN pH = 6.35 UNKNOWN [H+] = ?M Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 99

Sample Problem 19.4 Calculate Solve for the unknown. 2 First, simply swap the sides of the equation for finding pH and substitute the known value. pH = −log[H+] −log[H+] = pH −log[H+] = 6.35 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 100

Sample Problem 19.4 Calculate Solve for the unknown. 2 Change the signs on both sides of the equation and then solve for the unknown. log[H+] = −6.35 [H+] = antilog(−6.35) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 101

Sample Problem 19.4 Calculate Solve for the unknown. 2 Use the antilog (10x) function on your calculator to find [H+]. Report the answer in scientific notation. [H+] = 4.5 × 10−7 On most calculators, use the 2nd or INV key followed by log to get the antilog. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 102

Sample Problem 19.4 Evaluate Does the result make sense? 3 The pH is between 6 and 7. So, the hydrogen-ion concentration must be between 1 × 10−6M and 1 × 10−7M. The answer is rounded to two significant figures because the pH was measured to two decimal places. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 103

Calculating pH from [OH−]
The pH Concept Calculating pH from [OH−] If you know the [OH−] of a solution, you can find its pH. You can use the ion-product constant to determine [H+] for a known [OH−]. Then you use [H+] to calculate the pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 104

Calculating pH from [OH−]
Sample Problem 19.5 Calculating pH from [OH−] What is the pH of a solution if [OH−] = 4.0 × 10−11M? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 105

Analyze List the knowns and the unknown.
Sample Problem 19.5 Analyze List the knowns and the unknown. 1 To find [H+], divide Kw by the known [OH−]. Then calculate pH as you did in Sample Problem 19.3. KNOWNS [OH−] = 4.0 × 10−11M Kw = 1.0 × 10−14 UNKNOWN pH= ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 106

Sample Problem 19.5 Calculate Solve for the unknown. 2 Start with the ion-product constant to find [H+]. Rearrange the equation to solve for [H+]. Kw = [OH−] × [H+] [H+] = Kw [OH−] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 107

Sample Problem 19.5 Calculate Solve for the unknown. 2 Substitute the values for Kw and [OH−] to find [H+]. [H+] = 1.0 × 10−14 4.0 × 10−11 = 0.25 × 10−3M = 2.5 × 10−4M Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 108

Sample Problem 19.5 Calculate Solve for the unknown. 2 Next use the equation for finding pH. Substitute the value for [H+]. pH = −log[H+] = −log(2.5 × 10−4) Use a calculator to find the log. Round the pH to two decimal places because the [OH−] has two significant figures. = −(− ) = 3.60 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 109

Sample Problem 19.5 Evaluate Does the result make sense? 3 A solution in which [OH−] is less than 1 × 10−7M is acidic because [H+] is greater than 1 × 10−7M. The hydrogen-ion concentration is between 1 × 10−3M and 1 × 10−4M. Thus, the pH should be between 3 and 4. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 110

Why do we use the pH scale to express hydrogen-ion concentration?

Why do we use the pH scale to express hydrogen-ion concentration?
It is more convenient and practical to use the pH scale. Expressing hydrogen-ion concentration in molarity takes up a lot of space and is not as easy to work with. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 112

What are two methods that are used to measure pH?
Measuring pH Measuring pH What are two methods that are used to measure pH? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 113

Either acid-base indicators or pH meters can be used to measure pH.
Measuring pH Either acid-base indicators or pH meters can be used to measure pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 114

Measuring pH Acid-Base Indicators An indicator (HIn) is an acid or a base that dissociates in a known pH range. Indicators work because their acid form and base form have different colors in solution. The acid form of the indicator (HIn) is dominant at low pH and high [H+]. The base form (In−) is dominant at high pH and high [OH−]. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 115

Measuring pH Acid-Base Indicators The change from dominating acid form to dominating base form occurs within a narrow range of about two pH units. At all pH values below the range, you would see only the color of the acid form. At all pH values above this range, you would see only the color of the base form. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 116

Measuring pH Acid-Base Indicators An indicator strip is a piece of paper or plastic that has been soaked in an indicator and then dried. The paper is dipped into an unknown solution. The color that results is compared with a color chart to measure the pH. Some indicator paper has absorbed multiple indicators. The colors that result will cover a wide range of pH values. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 118

Soil pH can affect how plants develop.
Measuring pH Acid-Base Indicators Soil pH can affect how plants develop. In acidic soils, hydrangeas produce blue flowers. In basic soils, hydrangeas produce pink flowers. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 119

A pH meter is used to make rapid, continuous measurements of pH.
Measuring pH pH Meters A pH meter is used to make rapid, continuous measurements of pH. The measurements of pH obtained with a pH meter are typically accurate to within 0.01 pH unit of the true pH. If the pH meter is connected to a computer or chart recorder, the user will have a record of the pH changes. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 120

Measuring pH pH Meters A pH meter can be easier to use than liquid indicators or indicator strips. The pH reading is visible in a display window on the meter. The color and cloudiness of the solution do not affect the accuracy of the pH value obtained. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 121

You’re planting a garden and want to know the approximate pH of your soil. What method should you use? While you could use a pH meter and have a very accurate reading of your soil’s pH, it is also OK to approximate the pH using a pH indicator strip. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 123

Either acid-base indicators or pH meters can be used to measure pH.
Key Concepts For aqueous solutions, the product of the hydrogen-ion concentration and the hydroxide-ion concentration equals 1 × 10–14. A solution with a pH less than 7.0 is acidic. A solution with a pH of 7 is neutral. A solution with a pH greater than 7.0 is basic. Either acid-base indicators or pH meters can be used to measure pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 124

Glossary Terms self-ionization: a term describing the reaction in which two water molecules react to produce ions neutral solution: an aqueous solution in which the concentrations of hydrogen and hydroxide ions are equal; it has a pH of 7.0 ion-product constant for water (Kw): the product of the concentrations of hydrogen ions and hydroxide ions in water; it is × 10–14 at 25°C Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 125

Glossary Terms acidic solution: any solution in which the hydrogen-ion concentration is greater than the hydroxide-ion concentration basic solution: any solution in which the hydroxide-ion concentration is greater than the hydrogen-ion concentration pH: a number used to denote the hydrogen-ion concentration, or acidity, of a solution; it is the negative logarithm of the hydrogen-ion concentration of a solution Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 126

The pH of a solution reflects the hydrogen-ion concentration.
BIG IDEA Reactions The pH of a solution reflects the hydrogen-ion concentration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 127

Chapter 19 Acids, Bases, and Salts 19.3 Strengths of Acids and Bases

What makes one acid safer than another?
CHEMISTRY & YOU What makes one acid safer than another? Lemon juice, which contains citric acid, has a pH of about 2.3. Yet, you consume lemon juice. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 129

Strong and Weak Acids and Bases
How are acids and bases classified as either strong or weak? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 130

In general, a strong acid is completely ionized in aqueous solution.
Strong and Weak Acids and Bases In general, a strong acid is completely ionized in aqueous solution. Hydrochloric and sulfuric acid are examples of strong acids. HCl(g) + H2O(l) → H3O+(aq) + Cl–(aq) % Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 132

A weak acid ionizes only slightly in aqueous solution.
Strong and Weak Acids and Bases A weak acid ionizes only slightly in aqueous solution. The ionization of ethanoic acid (CH3COOH), a typical weak acid, is not complete. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 133

Interpret Graphs Dissociation of an acid (HA) in water yields H3O+ and an anion, A–. The bar graphs compare the extent of the dissociation of a strong acid and a weak acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 135

Acid Dissociation Constant
Strong and Weak Acids and Bases Acid Dissociation Constant A strong acid, such as hydrochloric acid, completely dissociates in water. As a result, [H3O+] is high in an aqueous solution of strong acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 136

Strong and Weak Acids and Bases Acid Dissociation Constant By contrast, weak acids remain largely undissociated. In an aqueous solution of ethanoic acid, less than 1 percent of the molecules are ionized. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 137

Strong and Weak Acids and Bases Acid Dissociation Constant You can use a balanced equation to write the equilibrium-constant expression for a reaction. The equilibrium-constant expression shown below is for ethanoic acid. Keq= [H3O+] × [CH3COO–] [CH3COOH] × [H2O] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 138

Strong and Weak Acids and Bases Acid Dissociation Constant The acid dissociation constant (Ka) is the ratio of the concentration of the dissociated form of an acid to the concentration of the undissociated form. The dissociated form includes both the H3O+ and the anion. Keq × [H2O] = Ka = [H3O+] × [CH3COO–] [CH3COOH] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 139

Strong and Weak Acids and Bases Acid Dissociation Constant The acid dissociation constant (Ka) reflects the fraction of an acid that is ionized. For this reason, dissociation constants are sometimes called ionization constants. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 140

Strong and Weak Acids and Bases Acid Dissociation Constant If the degree of dissociation or ionization of the acid is small, the value of the dissociation constant will be small. Weak acids have small Ka values. If the degree of ionization of an acid is more complete, the value of Ka will be larger. The stronger an acid is, the larger its Ka value will be. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 141

Strong and Weak Acids and Bases Acid Dissociation Constant Nitrous acid (HNO2) has a Ka of 4.4 × 10−4, but ethanoic acid (CH3COOH) has a Ka of 1.8 × 10−5. This means that nitrous acid is more ionized in solution than ethanoic acid. Nitrous acid is a stronger acid than ethanoic acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 142

Strong and Weak Acids and Bases Acid Dissociation Constant Some acids have more than one dissociation constant because they have more than one ionizable hydrogen. Oxalic acid is a diprotic acid. It loses two hydrogens, one at a time. Therefore, it has two dissociation constants. Oxalic acid is found naturally in certain herbs and vegetables. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 144

Strong and Weak Acids and Bases Acid Dissociation Constant Observe what happens to the Ka with each ionization. The Ka decreases from first ionization to second. It decreases again from second ionization to third. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 145

Calculating Dissociation Constants
Strong and Weak Acids and Bases Calculating Dissociation Constants To calculate the acid dissociation constant (Ka) of a weak acid, you need to know the initial molar concentration of the acid and the [H+] (or alternatively, the pH) of the solution at equilibrium. You can use these data to find the equilibrium concentrations of the acid and the ions. These values are then substituted into the expression for Ka. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 146

Calculating Dissociation Constants
Strong and Weak Acids and Bases Calculating Dissociation Constants You can find the Ka of an acid in water by substituting the equilibrium concentrations of the acid, [HA], the anion from the dissociation of the acid, [A−], and the hydrogen ion, [H+], into the equation below. Fix equation so bracket is above the fraction-bar line Ka = [H+][A−] [HA] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 147

Calculating a Dissociation Constant
Sample Problem 19.6 Calculating a Dissociation Constant In a M solution of ethanoic acid, [H+] = 1.34 × 10−3M. Calculate the Ka of this acid. Refer to the table for the ionization equation for ethanoic acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 148

Sample Problem 19.6 Analyze List the knowns and the unknown. 1 KNOWNS [ethanoic acid] = M [H+] = 1.34 × 10−3M UNKNOWN Ka = ? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 149

Sample Problem 19.6 Calculate Solve for the unknown. 2 Start by determining the equilibrium concentration of the ions. [H+] = [CH3COO−] = 1.34 × 10−3M Each molecule of CH3COOH that ionizes gives an H+ ion and a CH3COO– ion. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 150

Sample Problem 19.6 Calculate Solve for the unknown. 2 Determine the equilibrium concentrations of each component. ( – )M = M Concentration [CH3COOH] [H+] [CH3COO−] Initial 0.1000 Change −1.34 × 10−3 1.34 × 10−3 Equilibrium 0.0987 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 151

Sample Problem 19.6 Calculate Solve for the unknown. 2 Substitute the equilibrium values into the expression for Ka. Ka = [H+] × [CH3COO–] [CH3COOH] = (1.34 × 10−3M) × (1.34 × 10−3M) 0.0987 = 1.82 × 10−5 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 152

Sample Problem 19.6 Evaluate Does the result make sense? 3 The calculated value of Ka is consistent with that of a weak acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 153

Base Dissociation Constant
Strong and Weak Acids and Bases Base Dissociation Constant Just as there are strong acids and weak acids, there are strong bases and weak bases. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 154

Strong and Weak Acids and Bases Base Dissociation Constant Just as there are strong acids and weak acids, there are strong bases and weak bases. A strong base dissociates completely into metal ions and hydroxide ions in aqueous solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 155

Strong and Weak Acids and Bases Base Dissociation Constant Just as there are strong acids and weak acids, there are strong bases and weak bases. A strong base dissociates completely into metal ions and hydroxide ions in aqueous solution. A weak base reacts with water to form the conjugate acid of the base and hydroxide ions. For a weak base, the amount of dissociation is relatively small. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 156

Strong and Weak Acids and Bases Base Dissociation Constant Ammonia is an example of a weak base. Window cleaners often use a solution of ammonia in water to clean glass. NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Ammonia Water Ammonium ion Hydroxide ion Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 157

Strong and Weak Acids and Bases Base Dissociation Constant When equilibrium is established, only about 1 percent of the ammonia is present as NH4+. This ion is the conjugate acid of NH3. The concentrations of NH4+ and OH− are low and equal. NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Ammonia Water Ammonium ion Hydroxide ion Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 158

Strong and Weak Acids and Bases Base Dissociation Constant The equilibrium-constant expression for the dissociation of ammonia in water is as follows: Keq = [NH4+] × [OH−] [NH3] × [H2O] NH3(aq) + H2O(l) NH4+(aq) + OH–(aq) Ammonia Water Ammonium ion Hydroxide ion Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 159

Strong and Weak Acids and Bases Base Dissociation Constant Recall that the concentration of water is constant in dilute solutions. This constant can be combined with the Keq for ammonia to give a base dissociation constant (Kb) for ammonia. Keq × [H2O] = Kb = [NH4+] × [OH−] [NH3] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 160

[conjugate acid] × [OH−]
Strong and Weak Acids and Bases Base Dissociation Constant The base dissociation constant (Kb) is the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base. Kb = [conjugate acid] × [OH−] [base] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 161

Strong and Weak Acids and Bases Base Dissociation Constant The magnitude of Kb indicates the ability of a weak base to compete with the very strong base OH– for hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 162

Strong and Weak Acids and Bases Base Dissociation Constant The magnitude of Kb indicates the ability of a weak base to compete with the very strong base OH– for hydrogen ions. Because bases such as ammonia are weak relative to the hydroxide ion, the Kb for such a base is usually small. The Kb for ammonia is 1.8 × 10−5. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 163

Strong and Weak Acids and Bases Base Dissociation Constant The magnitude of Kb indicates the ability of a weak base to compete with the very strong base OH– for hydrogen ions. Because bases such as ammonia are weak relative to the hydroxide ion, the Kb for such a base is usually small. The Kb for ammonia is 1.8 × 10−5. The smaller the value of Kb, the weaker the base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 164

Concentration Versus Strength
Strong and Weak Acids and Bases Concentration Versus Strength Sometimes people confuse the concepts of concentration and strength. The words concentrated and dilute indicate how much of an acid or base is dissolved in solution. These terms refer to the number of moles of the acid or base in a given volume. The words strong and weak refer to the extent of ionization or dissociation of an acid or base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 165

Comparing Concentration and Strength of Acids Quantitative (or Molar)
Interpret Data Concentration Versus Strength The table below shows four possible combinations of concentration and strength for acids. Comparing Concentration and Strength of Acids Acidic solution Concentration Strength Quantitative (or Molar) Relative Hydrochloric acid 12M HCl Concentrated Strong Gastric juice 0.8M HCl Dilute Ethanoic acid 17M CH3COOH Weak Vinegar 0.2M CH3COOH Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 166

Strong and Weak Acids and Bases Concentration Versus Strength The gastric juice in your stomach is a dilute solution of HCl. The relatively small number of HCl molecules in a given volume of gastric juice are all dissociated into ions. Even when concentrated hydrochloric acid is diluted with water, it is still a strong acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 167

Strong and Weak Acids and Bases Concentration Versus Strength Conversely, ethanoic acid (acetic acid) is a weak acid because it ionizes only slightly in solution. Vinegar is a dilute solution of ethanoic acid. Even at a high concentration, ethanoic acid is still a weak acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 168

CHEMISTRY & YOU Despite its relatively low pH, lemon juice is safe to consume because citric acid is a weak acid. Citric acid has three Ka values. What does this information tell you about citric acid? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 169

CHEMISTRY & YOU Despite its relatively low pH, lemon juice is safe to consume because citric acid is a weak acid. Citric acid has three Ka values. What does this information tell you about citric acid? This information tells you that citric acid has three ionizable hydrogen atoms. It is a triprotic acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 170

Concentration Versus Strength The same concepts apply to bases. A solution of ammonia can be either dilute or concentrated. However, in any solution of ammonia, the relative amount of ionization will be small. Thus, ammonia is a weak base at any concentration. Likewise, sodium hydroxide is a strong base at any concentration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 171

In strong acids, are all ionizable hydrogens completely ionized
In strong acids, are all ionizable hydrogens completely ionized? In weak acids? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 172

In strong acids, are all ionizable hydrogens completely ionized? In weak acids? In strong acids, all ionizable hydrogens are completely ionized. In weak acids, all ionizable hydrogens are partially ionized. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 173

Key Concept & Key Equation
Acids and bases are classified as strong or weak based on the degree to which they ionize in water. [H+][A−] Ka = [HA] Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 174

weak acid: an acid that is only slightly ionized in aqueous solution
Glossary Terms strong acid: an acid that is completely (or almost completely) ionized in aqueous solution weak acid: an acid that is only slightly ionized in aqueous solution acid dissociation constant (Ka): the ratio of the concentration of the dissociated form of an acid to the undissociated form; stronger acids have larger Ka values than weaker acids Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 175

Glossary Terms strong base: a base that completely dissociates into metal ions and hydroxide ions in aqueous solution weak base: a base that reacts with water to form the hydroxide ion and the conjugate acid of the base base dissociation constant (Kb): the ratio of the concentration of the conjugate acid times the concentration of the hydroxide ion to the concentration of the base Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 176

Chapter 19 Acids, Bases, and Salts 19.4 Neutralization Reactions

What could cause leaves to turn yellow during the growing season?
CHEMISTRY & YOU What could cause leaves to turn yellow during the growing season? This condition is called chlorosis because the plant lacks a pigment called chlorophyll. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 178

What products form when an acid and a base react?
Acid-Base Reactions Acid-Base Reactions What products form when an acid and a base react? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 179

HCl(g) + NaOH(aq) → NaCl(aq) + H2O(l)
Acid-Base Reactions Suppose you mix a solution of a strong acid, such as HCl, with a solution of a strong base, such as NaOH. The products are sodium chloride and water. HCl(g) + NaOH(aq) → NaCl(aq) + H2O(l) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 180

In general, acids and bases react to produce a salt and water.
Acid-Base Reactions In general, acids and bases react to produce a salt and water. The complete reaction of a strong acid and a strong base produces a neutral solution. Thus, this type of reaction is called a neutralization reaction. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 181

Acid-Base Reactions When you hear the word salt, you may think of the substance that is used to flavor food. Table salt (NaCl) is only one example of a salt. Salts are ionic compounds consisting of an anion from an acid and a cation from a base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 182

Acid-Base Reactions A reaction between an acid and a base will go to completion when the solutions contain equal numbers of hydrogen ions and hydroxide ions. The balanced equation provides the correct ratio of acid to base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 183

HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Acid-Base Reactions For hydrochloric acid and sodium hydroxide, the mole ratio is 1:1. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) 1 mol Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 184

For sulfuric acid and sodium hydroxide, the ratio is 1:2.
Acid-Base Reactions For sulfuric acid and sodium hydroxide, the ratio is 1:2. 1 mol 2 mol Two moles of the base are required to neutralize one mole of the acid. H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 185

2HCl(aq) + Ca(OH)2(aq) → CaCl2(aq) + 2H2O(l)
Acid-Base Reactions Similarly, hydrochloric acid and calcium hydroxide react in a 2:1 ratio. 2HCl(aq) + Ca(OH)2(aq) → CaCl2(aq) + 2H2O(l) 2 mol 1 mol Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 186

Finding the Moles Needed for Neutralization
Sample Problem 19.7 Finding the Moles Needed for Neutralization The term neutralization is used to describe both the reaction and the point at which a neutralization reaction is complete. How many moles of sulfuric acid are required to neutralize 0.50 mol of sodium hydroxide? The equation for the reaction is H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 187

Sample Problem 19.7 Analyze List the knowns and the unknown. 1 To determine the number of moles of acid, you need to know the number of moles of base and the mole ratio of acid to base. KNOWNS mol NaOH = 0.50 mol 1 mol H2SO4/2 mol NaOH (from balanced equation) UNKNOWN mol H2SO4 = ? mol Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 188

Sample Problem 19.7 Calculate Solve for the unknown. 2 Use the mole ratio of acid to base to determine the number of moles of acid. 0.50 mol NaOH × 1 mol H2SO4 2 mol NaOH = 0.25 mol H2SO4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 189

Sample Problem 19.7 Evaluate Does the result make sense? 3 Because the mole ratio of H2SO4 to NaOH is 1:2, the number of moles of H2SO4 should be half the number of the moles of NaOH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 190

Why are acid-base reactions called neutralization reactions?

Why are acid-base reactions called neutralization reactions?
The complete reaction of an acid with a base creates a solution of a salt in water. This solution has a neutral pH. It is neither acidic nor basic. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 192

At what point in a titration does neutralization occur?

Titration You can use a neutralization reaction to determine the concentration of an acid or base. The process of adding a measured amount of a solution of known concentration to a solution of unknown concentration is called a titration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 194

The steps in an acid-base titration are as follows:
A measured volume of an acid solution of unknown concentration is added to a flask. Several drops of an indicator are added to the solution while the flask is gently swirled. Measured volumes of a base of known concentrations are mixed into the acid until the indicator just barely changes color. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 195

Titration A flask with a known volume of acids (and an indicator) is placed beneath a buret that is filled with a base of known concentration. The base is slowly added from the buret to the acid. A change in the color of the solution is the signal that neutralization has occurred. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 196

The solution of known concentration is the standard solution.
Titration The solution of known concentration is the standard solution. You can use a similar procedure to find the concentration of a base using a standard acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 197

Titration Neutralization occurs when the number of moles of hydrogen ions is equal to the number of moles of hydroxide ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 198

Titration Neutralization occurs when the number of moles of hydrogen ions is equal to the number of moles of hydroxide ions. The point at which neutralization occurs is called the equivalence point. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 199

Titration The indicator that is chosen for a titration must change color at or near the pH of the equivalence point. The point at which the indicator changes color is the end point of the titration. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 200

Interpret Graphs This graph shows how the pH of a solution changes during the titration of a strong acid (HCl) with a strong base (NaOH). The initial acid solution has a low pH (about 1). As NaOH is added, the pH increases because some of the acid reacts with the base. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 201

Interpret Graphs This graph shows how the pH of a solution changes during the titration of a strong acid (HCl) with a strong base (NaOH). The equivalence point for this reaction occurs at a pH of 7. As the titration nears the equivalence point, the pH rises dramatically because hydrogen ions are being used up. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 202

Interpret Graphs This graph shows how the pH of a solution changes during the titration of a strong acid (HCl) with a strong base (NaOH). Extending the titration beyond the point of neutralization produces a further increase of pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 203

Interpret Graphs This graph shows how the pH of a solution changes during the titration of a strong acid (HCl) with a strong base (NaOH). If the titration of HCl and NaOH could be stopped right at the equivalence point, the solution in the beaker would consist of only H2O and NaCl, plus a small amount of indicator. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 204

CHEMISTRY & YOU Iron compounds need to dissociate before the iron can be absorbed by plants. However, these compounds become less soluble as the pH rises. For most plants, a pH between 5.0 and 6.5 will provide enough usable iron. How could you change the pH of soil? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 205

CHEMISTRY & YOU Iron compounds need to dissociate before the iron can be absorbed by plants. However, these compounds become less soluble as the pH rises. For most plants, a pH between 5.0 and 6.5 will provide enough usable iron. How could you change the pH of soil? You can change the pH of basic soil by adding an acidic compound. To acidic soil you can add a basic compound. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 206

H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l).
Sample Problem 19.8 Determining Concentration by Titration A 25-mL solution of H2SO4 is neutralized by 18 mL of 1.0M NaOH. What is the concentration of the H2SO4 solution? The equation for the reaction is H2SO4(aq) + 2NaOH(aq) → Na2SO4(aq) + 2H2O(l). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 207

Sample Problem 19.8 Analyze List the knowns and the unknown. 1 The conversion steps are as follows: L NaOH → mol NaOH → mol H2SO4 → M H2SO4. KNOWNS [NaOH] = 1.0M VNaOH = 18 mL = L VH2SO4 = 18 mL = L UNKNOWN [H2SO4] = ?M Convert volume to liters because molarity is in moles per liter. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 208

Sample Problem 19.8 Calculate Solve for the unknown. 2 Use the molarity to convert the volume of base to moles of base. 0.018 L NaOH × 1.0 mol NaOH 1 L NaOH = mol NaOH Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 209

Sample Problem 19.8 Calculate Solve for the unknown. 2 Use the mole ratio to find the moles of acid. 0.018 mol NaOH × 1.0 mol H2SO4 2 mol NaOH = mol H2SO4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 210

Sample Problem 19.8 Calculate Solve for the unknown. 2 Calculate the molarity by dividing moles of acid by liters of solution. molarity = mol of solute L of solution mol 0.025 L = = 0.36M H2SO4 Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 211

Sample Problem 19.8 Evaluate Does the result make sense? 3 If the acid had the same molarity as the base (1.0M), 50 mL of base would neutralize 25 mL of acid. Because the volume of the base is much less than 50 mL, the molarity of the acid must be much less than 1.0M. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 212

In strong acids, are all ionizable hydrogens completely ionized? In weak acids? In strong acids, all ionizable hydrogens are completely ionized. In weak acids, all ionizable hydrogens are partially ionized. Some hydrogens in these acids (those with larger Ka values) have a greater degree of ionization. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 214

In general, acids and bases react to produce a salt and water.
Key Concepts In general, acids and bases react to produce a salt and water. Neutralization occurs when the number of moles of hydrogen ions is equal to the number of moles of hydroxide ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 215

Glossary Terms neutralization reaction: a reaction in which an acid and a base react in an aqueous solution to produce a salt and water titration: process used to determine the concentration of a solution (often an acid or base) in which a solution of known concentration (the standard) is added to a measured amount of the solution of unknown concentration until an indicator signals the end point Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 216

Glossary Terms standard solution: a solution of known concentration used in carrying out a titration equivalence point: the point in a titration where the number of moles of hydrogen ions equals the number of moles of hydroxide ions end point: the point in a titration at which the indicator changes color Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 217

Chapter 19 Acids, Bases, and Salts 19.5 Salts in Solution

How is the pH of blood controlled in the human body?
CHEMISTRY & YOU How is the pH of blood controlled in the human body? The pH of human blood needs to be kept close to 7.4. A person cannot survive for more than a few minutes if the pH of blood drops below 6.8 or rises above 7.8. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 220

When is the solution of a salt acidic or basic?
Salt Hydrolysis Salt Hydrolysis When is the solution of a salt acidic or basic? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 221

A salt is one of the products of a neutralization reaction.
Salt Hydrolysis A salt is one of the products of a neutralization reaction. A salt consists of an anion from an acid and a cation from a base. The solutions of many salts are neutral. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 222

Some salts form acidic or basic solutions.
Salt Hydrolysis Some salts form acidic or basic solutions. Universal indicator was added to these 0.10M aqueous salt solutions. Based on the indicator color, the solutions can be classified as follows: Ammonium chloride, NH4Cl(aq), is acidic (pH of about 5.3). Sodium chloride, NaCl(aq), is neutral (pH of 7). Sodium ethanoate, CH3COONa(aq), is basic (pH of about 8.7). Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 223

Interpret Graphs One curve is for the addition of sodium hydroxide, a strong base, to ethanoic acid, a weak acid. An aqueous solution of sodium ethanoate exists at the equivalence point. CH3COOH(aq) + NaOH(aq) → CH3COONa(aq) + H2O(l) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 224

Interpret Graphs The second titration curve is for the reaction between hydrochloric acid, which is a strong acid, and sodium hydroxide. HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 225

Interpret Graphs The pH at the equivalence point for the weak acid-strong base titration is basic. For a strong acid-strong base titration, the pH at the equivalence point is neutral. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 226

In salt hydrolysis, the cations or anions of a
Interpret Graphs This difference in pH exists because hydrolysis occurs with some salts in solution. In salt hydrolysis, the cations or anions of a dissociated salt remove hydrogen ions from, or donate hydrogen ions to, water. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 227

Salt Hydrolysis Salts that produce acidic solutions have positive ions that release hydrogen ions to water. Salts that produce basic solutions have negative ions that attract hydrogen ions from water. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 228

CH3COONa(aq) → CH3COO−(aq) + Na+(aq)
Salt Hydrolysis Sodium ethanoate (CH3COONa) is the salt of a weak acid and a strong base. In solution, the salt is completely ionized. CH3COONa(aq) → CH3COO−(aq) + Na+(aq) Sodium ethanoate Ethanoate ion Sodium ion Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 229

(makes the solution basic)
Salt Hydrolysis The ethanoate ion is a Brønsted-Lowry base, which means it is a hydrogen acceptor. It reacts with water to form ethanoic acid and hydroxide ions. At equilibrium, the reactants are favored. CH3COO–(aq) H2O(l) CH3COOH(aq) + OH–(aq) H+ donor Brønsted-Lowry acid H+ acceptor Brønsted-Lowry base (makes the solution basic) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 230

(makes the solution basic)
Salt Hydrolysis CH3COO–(aq) H2O(l) CH3COOH(aq) + OH–(aq) H+ donor Brønsted-Lowry acid H+ acceptor Brønsted-Lowry base (makes the solution basic) This process is called hydrolysis because a hydrogen ion is split off a water molecule. The suffix -lysis comes from a Greek word meaning to “separate” or “loosen.” In the solution, the hydroxide-ion concentration is greater than the hydrogen-ion concentration. Thus, the solution is basic. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 231

NH4Cl(aq) → NH4+(aq) + Cl−(aq)
Salt Hydrolysis Ammonium chloride (NH4Cl) is the salt of the strong acid hydrochloric acid (HCl) and the weak base ammonia (NH3). It is completely ionized in solution. NH4Cl(aq) → NH4+(aq) + Cl−(aq) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 232

(makes the solution acidic)
Salt Hydrolysis The ammonium ion (NH4+) is a strong enough acid to donate a hydrogen ion to a water molecule. The products are ammonia molecules and hydronium ions. The reactants are favored at equilibrium, as shown by the relative sizes of the arrows. NH4+(aq) H2O(l) NH3(aq) + H3O+(aq) H+ donor Brønsted-Lowry acid H+ acceptor Brønsted-Lowry base (makes the solution acidic) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 233

(makes the solution acidic)
Salt Hydrolysis This process is another example of hydrolysis. At equilibrium the [H3O+] is greater than the [OH–]. Thus, a solution of ammonium chloride is acidic. NH4+(aq) H2O(l) NH3(aq) + H3O+(aq) H+ donor Brønsted-Lowry acid H+ acceptor Brønsted-Lowry base (makes the solution acidic) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 234

Salt Hydrolysis To determine if a salt will form an acidic or basic solution, remember the following rules: Strong acid + Strong base → Neutral solution Strong acid + Weak base → Acidic solution Weak acid + Strong base → Basic solution Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 235

Does the fact that a weak acid-strong base titration is basic mean that there is some base “left over” at the equivalence point? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 236

Does the fact that a weak acid-strong base titration is basic mean that there is some base “left over” at the equivalence point? No. All of the acid and base has been converted to a salt solution at the equivalence point. The solution is basic because the salt hydrolyzes. The salt has negative ions that attract hydrogen ions from water. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 237

Suppose you add 10 mL of 0.10M sodium hydroxide to 1 L of pure water.
Buffers Suppose you add 10 mL of 0.10M sodium hydroxide to 1 L of pure water. The pH will increase about 4 pH units—from 7.0 to about 11.0. This change is a relatively large increase in pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 239

Buffers Now consider a solution containing 0.20M each of ethanoic acid and sodium ethanoate. This solution has a pH of 4.76. If you add 10 mL of 0.10M sodium hydroxide to 1 L of this solution, the pH increases 0.01 pH units—from 4.76 to 4.77. This is a relatively small change in pH. If 10 mL of acid had been added instead of the base, the amount of change in pH would also have been small. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 240

Buffers The solution of ethanoic acid and sodium ethanoate is an example of a buffer. A buffer is a solution in which the pH remains fairly constant when small amounts of acid or base are added. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 241

Buffers The figure below compares what happens when 1.0 mL of 0.10M HCl solution is added to an unbuffered solution and to a solution with a buffer. HCl is added to each solution. buffered unbuffered buffered unbuffered The indicator shows that both solutions are basic (pH of about 8). The indicator shows no visible pH change in the buffered solution. The color change in the unbuffered solution indicates a change in pH from 8 to about 3. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 243

Buffers How Buffers Work A buffer solution is better able to resist drastic changes in pH than is pure water. A buffer solution contains one component that can react with hydrogen ions (hydrogen -ion acceptor) and one that can react with hydroxide ions (hydrogen-ion donor). These components act as reservoirs of neutralizing power that can be tapped when either hydrogen ions or hydroxide ions are added to the solution. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 244

Buffers How Buffers Work The ethanoic acid-ethanoate ion buffer can be used to show how a buffer works. When an acid is added to the buffer, the ethanoate ions (CH3COO–) act as a hydrogen-ion “sponge.” As the ethanoate ions react with the hydrogen ions, they form ethanoic acid. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 245

Buffers How Buffers Work Ethanoic acid is a weak acid and does not ionize extensively in water, so the change in pH is very slight. Hydrogen ion CH3COO–(aq) + H+(aq) CH3COOH(aq) Ethanoate ion Ethanoic acid Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 246

Buffers How Buffers Work When hydroxide ions are added to the buffer, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion. CH3COOH(aq) + OH–(aq) CH3COO–(aq) + H2O(l) Hydroxide ion Ethanoic acid Ethanoate ion Water Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 247

Buffers How Buffers Work When hydroxide ions are added to the buffer, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion. CH3COOH(aq) + OH–(aq) CH3COO–(aq) + H2O(l) Hydroxide ion Ethanoic acid Ethanoate ion Water The ethanoate ion is not a strong enough base to accept hydrogen ions from water to a great extent. Therefore, the reverse reaction is minimal and the change in pH is very slight. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 248

CHEMISTRY & YOU The equilibrium between carbonic acid (H2CO3) and hydrogen carbonate ions (HCO3–) helps keep the pH of blood within a narrow range (7.35–7.45). If the pH rises, molecules of carbonic acid donate hydrogen ions. What can happen if the pH drops, that is, if the [H+] increases? Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 249

CHEMISTRY & YOU The equilibrium between carbonic acid (H2CO3) and hydrogen carbonate ions (HCO3–) helps keep the pH of blood within a narrow range (7.35–7.45). If the pH rises, molecules of carbonic acid donate hydrogen ions. What can happen if the pH drops, that is, if the [H+] increases? Hydrogen carbonate ions can accept hydrogen ions when the pH drops. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 250

The Capacity of a Buffer
Buffers The Capacity of a Buffer Buffer solutions have their limits. As acid is added to an ethanoate buffer, eventually no more ethanoate ions will be present to accept the hydrogen ions. At that point, the buffer can no longer control the pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 251

Buffers The Capacity of a Buffer Buffer solutions have their limits. The ethanoate buffer also becomes ineffective when too much base is added. No more ethanoic acid molecules are present to donate hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 252

Buffers The Capacity of a Buffer Adding too much acid or base will exceed the buffer capacity of a solution. The buffer capacity is the amount of acid or base that can be added to a buffer solution before a significant change in pH occurs. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 253

Important Buffer Systems
Interpret Data The Capacity of a Buffer This table lists some common buffer systems. Important Buffer Systems Buffer name Formulas Buffer pH* Ethanoic acid-ethanoate ion CH3COOH/CH3COO− 4.76 Dihydrogen phosphate ion-hydrogen phosphate ion H2PO4−/HPO42− 7.20 Carbonic acid-hydrogen carbonate ion (solution saturated with CO2) H2CO3/HCO3− 6.46 Ammonium ion-ammonia NH4+/NH3 9.25 * Components have concentrations of 0.1M. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 254

Buffers The Capacity of a Buffer Two buffer systems help maintain optimal human blood pH. One is the carbonic acid-hydrogen carbonate buffer system. The other is the dihydrogen phosphate-hydrogen phosphate buffer system. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 255

Describing Buffer Systems
Sample Problem 19.9 Describing Buffer Systems Write balanced chemical equations to show how the carbonic acid-hydrogen carbonate buffer can “mop up” added hydroxide ions and hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 256

Analyze Identify the relevant concepts.
Sample Problem 19.9 Analyze Identify the relevant concepts. 1 A buffer contains two components: a hydrogen-ion acceptor (which can react with H+) a hydrogen-ion donor (which can react with OH−) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 257

Calculate Apply the concepts to this problem.
Sample Problem 19.9 Calculate Apply the concepts to this problem. 2 Identify the hydrogen-ion acceptor and the hydrogen-ion donor. H2CO3, a weak acid, can release hydrogen ions. HCO3– is the conjugate base, which can accept hydrogen ions. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 258

H2CO3(aq) + (OH−)(aq) HCO3−(aq) + H2O(l)
Sample Problem 19.9 Calculate Apply the concepts to this problem. 2 Write the equation for the reaction that occurs when a base is added to the buffer. When a base is added, the hydroxide ions react with H2CO3. H2CO3(aq) + (OH−)(aq) HCO3−(aq) + H2O(l) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 259

HCO3−(aq) + (H+)(aq) H2CO3(aq)
Sample Problem 19.9 Calculate Apply the concepts to this problem. 2 Write the equation for the reaction that occurs when an acid is added to the buffer. When an acid is added, the hydrogen ions react with HCO3−. HCO3−(aq) + (H+)(aq) H2CO3(aq) Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 260

How is the work of a buffer solution similar to a neutralization reaction?

How is the work of a buffer solution similar to a neutralization reaction?
A buffer solution contains compounds that are able to neutralize both acids and bases. It performs acid-base neutralization reactions without significant change in pH. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 262

Key Concepts Salts that produce acidic solutions have positive ions that release hydrogen ions to water. Salts that produce basic solutions have negative ions that attract hydrogen ions from water. A buffer is a solution of a weak acid and one of its salts or a weak base and one of its salts. Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 263

Glossary Terms salt hydrolysis: a process in which the cations or anions of a dissociated salt accept hydrogen ions from water or donate hydrogen ions to water buffer: a solution in which the pH remains relatively constant when small amounts of acid or base are added; a buffer can be either a solution of a weak acid and the salt of a weak acid or a solution of a weak base with the salt of a weak base buffer capacity: a measure of the amount of acid or base that may be added to a buffer solution before a significant change in pH occurs Copyright © Pearson Education, Inc., or its affiliates. All Rights Reserved. 264

Chapter 19 Acids, Bases, and Salts 19.1 Acid-Base Theories

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