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Chemistry Chapter 19.

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slide2
Figure 19.1 Many items contain acids or bases, or produce acids and bases when dissolved in water. a. Citrus fruit contain citric acid. b. Tea contains tannic acid. c. Antacids use bases to neutralize excess stomach acid. d. The base calcium hydroxide is a component of mortar.
slide3
Acids taste sour, will change the color of an acid-base indicator, and can be strong or weak electrolytes in aqueous solution.
slide4
Bases taste bitter, feel slippery, will change the color of an acid-base indicator, and can be strong or weak electrolytes in aqueous solution.
slide5
Arrhenius said that acids are hydrogen-containing compounds that ionize to yield hydrogen ions (H+) in aqueous solution. He also said that bases are compounds that ionize to yield hydroxide ions (OH−) in aqueous solution.
slide6
monoprotic acids.
  • any acid that contains one ionizable proton (hydrogen ion); nitric acid (HNO3) is a monoprotic acid
slide7
diprotic acids.
  • any acid that contains two ionizable protons (hydrogen ions); sulfuric acid (H2SO4) is a diprotic acid
slide8
triprotic acids.
  • any acid that contains three ionizable protons (hydrogen ions); phosphoric acid (H3PO4) is a triprotic acid
slide13
The Brønsted-Lowry theory defines an acid as a hydrogen-ion donor, and a base as a hydrogen-ion acceptor.
slide15
In this reaction, ammonia is the hydrogen-ion acceptor and therefore is a Brønsted-Lowry base. Water, the hydrogen-ion donor, is a Brønsted-Lowry acid. Hydrogen ions are transferred from water to ammonia
slide16
conjugate acid
  • the particle formed when a base gains a hydrogen ion; NH4+ is the conjugate acid of the base NH3
slide17
conjugate base
  • the particle that remains when an acid has donated a hydrogen ion; OH- is the conjugate base of the acid water
slide18
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 p
slide23
amphoteric.
  • a substance that can act as both an acid and a base
slide24
). Lewis proposed that an acid accepts a pair of electrons during a reaction, while a base donates a pair of electrons.
slide25
Lewis acid
  • any substance that can accept a pair of electrons to form a covalent bond
slide26
Lewis base
  • any substance that can donate a pair of electrons to form a covalent bond
slide32
(3) Key Concept What are the properties of acids and bases? Hint
  • (4) Key Concept How did Arrhenius define an acid and a base?Hint
  • (5) Key Concept How are acids and bases defined by the Brønsted-Lowry theory?Hint
  • (6) Key Concept What is the Lewis theory of acids and bases?Hint
slide33
(7)
  • What is a conjugate acid-base pair?Hint
  • Write equations for the ionization of HNO3 in water and the reaction of CO32− with water. For each equation, identify the hydrogen-ion donor and hydrogen-ion acceptor. Then label the conjugate acid-base pairs in each equation.
slide34
)Identify the following acids as monoprotic, di-protic, or triprotic. Explain your reasoning.
  • H2CO3
  • H3PO4
  • HCl
  • H2SO4
slide35
hydronium ion (H3O+).
  • the positive ion formed when a water molecule gains a hydrogen ion
slide36
Hydrogen Ions from Water

self-ionization

a term describing the reaction in which two water molecules react to produce ions

slide38
In water or aqueous solution, hydrogen ions (H+) are always joined to water molecules as hydronium ions (H3O+). Hydrogen ions in aqueous solution have several names. Some chemists call them protons. Others prefer to call them hydrogen ions or hydronium ions. In this textbook, either H+ or H3O+ is used to represent hydrogen ions in aqueous solution.
slide39
The self-ionization of water occurs to a very small extent. In pure water at 25°C, the equilibrium concentration of hydrogen ions ([H+]) and the equilibrium concentration of hydroxide ions ([OH−]) are each only 1 × 10−7M. This means that the concentrations of H+ and OH− are equal in pure water.
slide40
neutral solution.
  • an aqueous solution in which the concentrations of hydrogen and hydroxide ions are equal; it has a pH of 7.0
slide42
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
slide43
ion-product constant for water (Kw).
  • the product of the concentrations of hydrogen ions and hydroxide ions in water; it is 1 X 10-14 at 25°C
slide44
acidic solution
  • any solution in which the hydrogen-ion concentration is greater than the hydroxide-ion concentration
slide45
basic solution
  • any solution in which the hydroxide-ion concentration is greater than the hydrogen-ion concentration
slide46
alkaline solutions.
  • a basic solution
slide52
The pH Concept
  • Expressing hydrogen-ion concentration in molarity is cumbersome. A more widely used system for expressing [H+] is the pH scale, proposed in 1909 by the Danish scientist Søren Sørensen (1868–1939). On the pH scale, which ranges from 0 to 14, neutral solutions have a pH of 7. A pH of 0 is strongly acidic. A solution with a pH of 14 is strongly basic.
slide53
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
slide55
A solution in which [H+] is greater than 1 × 10−7M has a pH less than 7.0 and is acidic. The pH of pure water or a neutral aqueous solution is 7.0. A solution with a pH greater than 7 is basic and has a [H+] of less than 1 × 10−7M.
slide56
Acidic solution: pH < 7.0 [H+] greater than 1 × 10−7M
  • Neutral solution: pH = 7.0 [H+] equals 1 × 10−7M
  • Basic solution: pH > 7.0 [H+] less than 1 × 10−7M
slide59
Calculating pOH
  • The pOH of a solution equals the negative logarithm of the hydroxide-ion concentration.
  • pOH = −log[OH−]
slide76
An indicator is a valuable tool for measuring pH because its acid form and base form have different colors in solution.
slide78
pH Meters
  • A pH meter makes rapid, accurate pH measurements. Most chemistry laboratories have a pH meter
slide79
(17) Key Concept What is the relationship between [H+] and [OH−] in an aqueous solution?Hint
  • (18) Key Concept What is true about the relative concentrations of hydrogen ions and hydroxide ions in each kind of solution?Hint
  • basic
  • acidic
  • neutral
slide80
19) Key Concept What is true about the colors of a pH indicator?Hint
  • (20)Determine the pH of each solution.
  • [H+] = 1 × 10−6M
  • [H+] = 0.00010M
  • [OH−] = 1 × 10−2M
  • [OH−] = 1 × 10−11M
slide81
21)What are the hydroxide-ion concentrations for solutions with the following pH values?
  • 6.00
  • 9.00
  • 12.00
slide82
strong acids
  • an acid that is completely (or almost completely) ionized in aqueous solution
slide83
Weak acids
  • an acid that is only slightly ionized in aqueous solution
slide85
Acid Dissociation Constant
  • In an aqueous solution of ethanoic acid, fewer than 1% of ethanoic acid molecules are ionized at any instant. Therefore, ethanoic acid is considered a weak acid
slide89
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
slide92
Strong bases
  • a base that completely dissociates into metal ions and hydroxide ions in aqueous solution
slide93
Weak bases
  • a base that reacts with water to form the hydroxide ion and the conjugate acid of the base
slide94
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 conjugate base
slide95
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. They indicate how many of the particles ionize or dissociate into ions.
slide97
To find the Ka of a weak acid or the Kb of a weak base, substitute the measured concentrations of all the substances present at equilibrium into the expression for Ka or Kb
slide104
(24) Key Concept Compare a strong acid and a weak acid in terms of the acid dissociation constant.Hint
  • (25) Key Concept How do you determine the Ka of a weak acid or the Kb of a weak base? Hint
  • (26)Which acid in Table 19.6 would you expect to have the lowest ionization constant?
slide105
(27)Acid HX has a very small value of Ka. How do the relative amounts of H+ and HX compare at equilibrium?
slide106
(28)Write the equations for the ionization or dissociation of the following acids and bases in water.
  • nitric acid Hint
  • ethanoic acid Hint
  • ammonia Hint
  • magnesium hydroxide Hint
slide109
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
  • H2SO4(aq) + 2KOH(aq) → K2SO4(aq) + 2H2O(l)
slide110
neutralization reactions.
  • a reaction in which an acid and a base react in an aqueous solution to produce a salt and water
slide113
In general, the reaction of an acid with a base produces water and one of a class of compounds called salts.
slide114
equivalence point
  • the point in a titration where the number of moles of hydrogen ions equals the number of moles of hydroxide ions
slide120
The steps in a neutralization reaction are as follows.
  • A measured volume of an acid solution of unknown concentration is added to a flask.
  • Several drops of the indicator are added to the solution while the flask is gently swirled.
  • Measured volumes of a base of known concentration are mixed into the acid until the indicator just barely changes color.
slide121
.
  • 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
slide122
standard solution.
  • a solution of known concentration used in carrying out a titration
slide123
end point
  • the point in a titration at which the indicator changes color
slide132
(34) Key Concept What are the products of a reaction between an acid and a base? Hint
  • (35) Key Concept What occurs at the endpoint of a titration? Hint
slide133
36)How many moles of HCl are required to neutralize aqueous solutions of these bases?
  • 2 mol NH3
  • 0.1 mol Ca(OH)2
slide134
(37)Write complete balanced equations for the following acid-base reactions?
  • H2SO4(aq) + KOH(aq) →
  • H3PO4(aq) + Ca(OH)2(aq) →
  • HNO3(aq) + Mg(OH)2(aq) →
slide135
A salt consists of an anion from an acid and a cation from a base. It forms as a result of a neutralization reaction. Although solutions of many salts are neutral, some are acidic and others are basic
slide138
The pH at the equivalence point is 8.7—basic.
  • For a strong acid–strong base titration, the pH at the equivalence point is 7, or neutral. This difference exists because some salts promote hydrolysis
slide139
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
slide140
Depending on the direction of the hydrogen-ion transfer, solutions containing hydrolyzing salts may be either acidic or basic. Hydrolyzing salts are usually derived from a strong acid and a weak base, or from a weak acid and a strong base.
slide141
In general, salts that produce acidic solutions contain positive ions that release protons to water. Salts that produce basic solutions contain negative ions that attract protons from water.
slide143
The ethanoate ion is a Brønsted-Lowry base, which means it is a hydrogen-ion acceptor. It establishes an equilibrium with water, forming electrically neutral ethanoic acid and negative hydroxide ions.
slide144
This process is called hydrolysis because it splits a hydrogen ion off a water molecule. The resulting solution contains a hydroxide-ion concentration greater than the hydrogen-ion concentration. Thus the solution is basic.
slide145
Ammonium chloride (NH4Cl) is the salt of a strong acid (hydrochloric acid, HCl) and a weak base (ammonia, NH3). It is completely ionized in solution.
slide147
This process is also called hydrolysis. It results in the formation of unionized ammonia and hydronium (hydrogen) ions. The [H3O+] is greater than the [OH−]. Thus, a solution of ammonium chloride is acidic.
slide148
Strong acid + Strong base → Neutral solution
  • Strong acid + Weak base → Acidic solution
  • Weak acid + Strong base → Basic solution
slide149
A buffer is a solution in which the pH remains relatively constant when small amounts of acid or base are added. A buffer is a solution of a weak acid and one of its salts, or a solution of a weak base and one of its salts.
slide150
A buffer solution is better able to resist drastic changes in pH than is pure water. Figure 19.28 illustrates how a buffer works. Ethanoic acid (CH3COOH) and its anion (CH3COO−) act as reservoirs of neutralizing power. They react with any hydroxide ions or hydrogen ions added to the solution
slide154
When an acid is added to the solution, the ethanoate ions (CH3COO−) act as a hydrogen-ion “sponge.” This creates ethanoic acid, which does not ionize extensively in water; thus the pH does not change appreciably
slide156
A base is a source of hydroxide ions. When a base is added to the solution, the ethanoic acid and the hydroxide ions react to produce water and the ethanoate ion.
slide158
For example, consider the buffer solution in which the sodium ethanoate (CH3COONa) is completely ionized.
slide159
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
slide166
Key Concept What type of salt produces an acidic solution? A basic solution?Hint
  • (41) Key Concept What substances are combined to make a buffer? Hint
slide167
(42)Which of these salts would form an acidic aqueous solution?
  • KC2H3O2
  • LiCl
  • NaHCO3
  • (NH4)2SO4
slide168
43)Using equations, show what happens when acid is added to an ammonium ion–ammonia buffer. What happens when base is added