chapter 14 acids and bases table of contents · acids 1. aqueous solutions of acids have a sour...

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Table of Contents

Chapter 14 Acids and Bases

Section 1 Properties of Acids and Bases

Section 2 Acid-Base Theories

Section 3 Acid-Base Reactions



Lesson Starter

• The solutions in the beakers are different because

they have a different pH.

• One beaker contains a basic solution and the other

beaker contains an acidic solution

Chapter 14 Section 1 Properties of Acids and

Bases



Objectives

• List five general properties of aqueous acids and

bases.

• Name common binary acids and oxyacids, given

their chemical formulas.

• List five acids commonly used in industry and the

laboratory, and give two properties of each.

• Define acid and base according to Arrhenius’s theory

of ionization.

• Explain the differences between strong and weak

acids and bases.


Bases



Acids

1. Aqueous solutions of acids have a sour taste.

2. Acids change the color of acid-base indicators.

3. Some acids react with active metals and release

hydrogen gas, H2.

Ba(s) + H2SO4(aq) BaSO4(s) + H2(g)


Bases

4. Acids react with bases to produce salts and water.

5. Acids conduct electric current.



Visual Concepts

Properties of Acids

Chapter 14



Acids, continued Acid Nomenclature

• A binary acid is an acid that contains only two

different elements: hydrogen and one of the more

electronegative elements.

• HF, HCl, HBr, and HI

• Binary Acid Nomenclature

1. The name of a binary acid begins with the prefix

hydro-.

2. The root of the name of the second element

follows this prefix.

3. The name then ends with the suffix -ic.


Bases



Acids, continued Acid Nomenclature, continued


Bases




• An oxyacid is an acid that is a compound of

hydrogen, oxygen, and a third element, usually a

nonmetal.

• HNO3, H2SO4

• The names of oxyacids follow a pattern.

• The names of their anions are based on the names of

the acids.


Bases





Bases



Some Common Industrial Acids

• Sulfuric Acid • Sulfuric acid is the most commonly produced industrial

chemical in the world.

• Nitric Acid

• Phosphoric Acid

• Hydrochloric Acid

• Concentrated solutions of hydrochloric acid are commonly

referred to as muriatic acid.

• Acetic Acid

• Pure acetic acid is a clear, colorless, and pungent-smelling

liquid known as glacial acetic acid.


Bases



Bases

1. Aqueous solutions of bases taste bitter.

2. Bases change the color of acid-base indicators.

3. Dilute aqueous solutions of bases feel slippery.

4. Bases react with acids to produce salts and water.

5. Bases conduct electric current.


Bases



Visual Concepts

Properties of Bases

Chapter 14



Arrhenius Acids and Bases

• An Arrhenius acid is a chemical compound that

increases the concentration of hydrogen ions, H+, in

aqueous solution.

• An Arrhenius base is a substance that increases the

concentration of hydroxide ions, OH−, in aqueous

solution.


Bases



Arrhenius Acids and Bases, continued Aqueous Solutions of Acids

• Arrhenius acids are molecular compounds with

ionizable hydrogen atoms.

• Their water solutions are known as aqueous acids.

• All aqueous acids are electrolytes.


Bases



Arrhenius Acids and Bases, continued Aqueous Solutions of Acids, continued

• Common Aqueous Acids


Bases



Arrhenius Acids and Bases, continued Strength of Acids

• A strong acid is one that ionizes completely in

aqueous solution.

• a strong acid is a strong electrolyte

• HClO4, HCl, HNO3

• A weak acid releases few hydrogen ions in aqueous

solution.

• hydronium ions, anions, and dissolved acid

molecules in aqueous solution

• HCN

• Organic acids (—COOH), such as acetic acid


Bases



Arrhenius Acids and Bases, continued Aqueous Solutions of Bases

• Most bases are ionic compounds containing metal

cations and the hydroxide anion, OH−.

• dissociate in water

s aq + aq2H O –NaOH( ) Na ( ) OH ( )

aq + l aq + aq–

3 2 4NH ( ) H O( ) NH ( ) OH ( )


Bases

• Ammonia, NH3, is molecular

• Ammonia produces hydroxide ions when it reacts

with water molecules.



Arrhenius Acids and Bases, continued Strength of Bases

• The strength of a base depends on the extent to

which the base dissociates.

• Strong bases are strong electrolytes


Bases



Relationship of [H3O+] to [OH–]


Bases



Lesson Starter

• List three terms that describe the person in the

photo.

• The person has been described in many different

ways, but he or she is still the same person.

• Acids and bases also can be described differently

based on the circumstances.

Chapter 14 Section 2 Acid-Base Theories



Objectives

• Define and recognize Brønsted-Lowry acids and

bases.

• Define a Lewis acid and a Lewis base.

• Name compounds that are acids under the Lewis

definition but are not acids under the Brønsted-Lowry

definition.




Brønsted-Lowry Acids and Bases

• A Brønsted-Lowry acid is a molecule or ion that is a

proton donor.

• Hydrogen chloride acts as a Brønsted-Lowry acid

when it reacts with ammonia.

+ –

3 4HCl NH NH Cl

l + aq aq + aq–

2 3 4H O( ) NH ( ) NH ( ) OH ( )


• Water can act as a Brønsted-Lowry acid.



Brønsted-Lowry Acids and Bases, continued

• A Brønsted-Lowry base is a molecule or ion that is

a proton acceptor.

• Ammonia accepts a proton from the hydrochloric

acid. It acts as a Brønsted-Lowry base.

+ –

3 4HCl NH NH Cl


• The OH− ion produced in solution by Arrhenius

hydroxide bases (NaOH) is the Brønsted-Lowry

base.

• The OH− ion can accept a proton



Brønsted-Lowry Acids and Bases, continued

• In a Brønsted-Lowry acid-base reaction, protons

are transferred from one reactant (the acid) to

another (the base).

+ –

3 4HCl NH NH Cl


acid base



Monoprotic and Polyprotic Acids

• A monoprotic acid is an acid that can donate only

one proton (hydrogen ion) per molecule.

• HClO4, HCl, HNO3

• only one ionization step

g + l) aq + aq–

2 3HCl( ) H O( H O ( ) Cl ( )




Monoprotic and Diprotic Acids




Monoprotic and Polyprotic Acids, continued

• A polyprotic acid is an acid that can donate more

than one proton per molecule.

• H2SO4, H3PO4

• Multiple ionization steps

l + l aq + aq–

2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )

aq + l aq + l– 2–

4 2 3 4HSO ( ) H O( ) H O ( ) SO ( )

– 2–

4 4HSO and SO


(1)

(2)

• Sulfuric acid solutions contain H3O+,

ions



Monoprotic and Polyprotic Acids, continued

• A diprotic acid is the type of polyprotic acid that can

donate two protons per molecule

• H2SO4

• A triprotic acid is the type of polyprotic acid that can

donate three protons per molecule.

• H3PO4




Lewis Acids and Bases

• A Lewis acid is an atom, ion, or molecule that

accepts an electron pair to form a covalent bond.

• The Lewis definition is the broadest of the three

acid definitions.

• A bare proton (hydrogen ion) is a Lewis acid

aq + aq aq aq3 3 4H ( ) :NH ( ) [H—NH ] ( ) or [NH ] ( )




Lewis Acids and Bases, continued

• The formula for a Lewis acid need not include

hydrogen.

• The silver ion can be a Lewis acid

aq + aq aq 3 3 3 3 2Ag ( ) 2 :NH ( ) [H N— Ag —NH ] ( ) or [Ag(NH ) ]

aq aq aq– –

3 4BF ( ) F ( ) BF ( )


• Any compound in which the central atom has three

valence electrons and forms three covalent bonds

can react as a Lewis acid.



•A Lewis base is an atom, ion, or molecule that donates

an electron pair to form a covalent bond.

•Look for a negative charge on the molecule or ion



Lewis Acids and Bases, continued Acid Base Definitions




Lesson Starter

• What is the meaning of the word neutralization.

• How is the word used in everyday life?

• How is it likely to apply to acids and bases?

Chapter 14 Section 3 Acid-Base Reactions



Objectives

• Describe a conjugate acid, a conjugate base, and an

amphoteric compound.

• Explain the process of neutralization.

• Define acid rain, give examples of compounds that

can cause acid rain, and describe effects of acid rain.




Conjugate Acids and Bases

• The species that remains after a Brønsted-Lowry

acid has given up a proton is the conjugate base of

that acid.

aq + l aq + aq–

2 3HF( ) H O( ) F ( ) H O ( )


acid conjugate

base



Conjugate Acids and Bases, continued

• Brønsted-Lowry acid-base reactions involve two

acid-base pairs, known a conjugate acid-base pairs.

aq + l aq + aq–

2 3HF( ) H O( ) F ( ) H O ( )


acid1 base2 base1 acid2



Neutralization Reactions




Conjugate Acids and Bases, continued Strength of Conjugate Acids and Bases

• The stronger an acid is, the weaker its conjugate

base • The stronger a base is, the weaker its conjugate acid

g + l aq + aq–

2 3HCl( ) H O( ) H O ( ) Cl ( )


strong acid base acid weak base



Conjugate Acids and Bases, continued Strength of Conjugate Acids and Bases, continued

• Proton transfer reactions favor the production of the

weaker acid and the weaker base.

aq + l aq + aq–

4 2 3 4HClO ( ) H O( ) H O ( ) ClO ( )

aq + l aq + aq–

3 2 3 3CH COOH( ) H O( ) H O ( ) CH COO ( )


stronger acid stronger base weaker acid weaker base

• The reaction to the right is more favorable

weaker acid weaker base stronger acid stronger base

• The reaction to the left is more favorable



Relative Strengths of Acids and Bases


Relative Strengths of Acids and Bases



Amphoteric Compounds

• Any species that can react as either an acid or a

base is described as amphoteric.

• example: water

• water can act as a base

aq + l aq + aq–

2 4 2 3 4H SO ( ) H O( ) H O ( ) HSO ( )

g + l aq aq–

3 2 4NH ( ) H O( ) NH ( ) OH ( )


acid1 base2 acid2 base1

• water can act as an acid

base1 acid2 acid1 base2



Amphoteric Compounds, continued –OH in a Molecule

• The covalently bonded OH group in an acid is

referred to as a hydroxyl group.

• Molecular compounds containing —OH groups can

be acidic or amphoteric.

• The behavior of a compound is affected by the

number of oxygen atoms bonded to the atom

connected to the —OH group.




Oxyacids of Chlorine




Neutralization Reactions Strong Acid-Strong Base Neutralization

• In aqueous solutions, neutralization is the reaction

of hydronium ions and hydroxide ions to form water

molecules.

• A salt is an ionic compound composed of a cation

from a base and an anion from an acid.

2aq + aq aq lHCl( ) NaOH( ) NaCl( ) H O( )




Acid Rain

• NO, NO2, CO2, SO2, and SO3 gases from industrial

processes can dissolve in atmospheric water to

produce acidic solutions.

g + l aq3 2 2 4SO ( ) H O( ) H SO ( )


• example:

• Very acidic rain is known as acid rain.

• Acid rain can erode statues and affect ecosystems.



Visual Concepts

Acid Precipitation

Chapter 14



Objectives

• Describe the self-ionization of water.

• Define pH, and give the pH of a neutral solution at

25°C.

• Explain and use the pH scale.

• Given [H3O+] or [OH−], find pH.

• Given pH, find [H3O+] or [OH−].

Chapter 15 Section 1 Aqueous Solutions and

the Concept of pH



Hydronium Ions and Hydroxide Ions

Self-Ionization of Water

• In the self-ionization of water, two water molecules

produce a hydronium ion and a hydroxide ion by

transfer of a proton.

l + l aq + aq–

2 2 3H O( ) H O( ) H O ( ) OH ( )


the Concept of pH

• In water at 25°C, [H3O+] = 1.0 ×10−7 M and [OH−] =

1.0 × 10−7 M.

• The ionization constant of water, Kw, is expressed by

the following equation.

Kw = [H3O+][OH−]



Hydronium Ions and Hydroxide Ions, continued

Self-Ionization of Water, continued

• At 25°C,

Kw = [H3O+][OH−] = (1.0 × 10−7)(1.0 × 10−7) = 1.0 × 10−14

• Kw increases as temperature increases


the Concept of pH




Neutral, Acidic, and Basic Solutions

• Solutions in which [H3O+] = [OH−] is neutral.

• Solutions in which the [H3O+] > [OH−] are acidic.

• [H3O+] > 1.0 × 10−7 M

• Solutions in which the [OH−] > [H3O+] are basic.

• [OH−] > 1.0 × 10−7 M


the Concept of pH




Calculating [H3O+] and [OH–]

• Strong acids and bases are considered completely

ionized or dissociated in weak aqueous solutions.

s aq + aq2H O –NaOH( ) Na ( ) OH ( )

-14 -14

-12

3 – -2

1.0 10 1.0 10[H O ] 1.0 10 M

[OH ] 1.0 10


the Concept of pH

1 mol 1 mol 1 mol

• 1.0 × 10−2 M NaOH solution has an [OH−] of 1.0 × 10−2 M

• The [H3O+] of this solution is calculated using Kw.

Kw = [H3O+][OH−] = 1.0 × 10−14




Calculating [H3O+] and [OH–]

• If the [H3O+] of a solution is known, the [OH−] can be

calculated using Kw.

[HCl] = 2.0 × 10−4 M

[H3O+] = 2.0 × 10−4 M

Kw = [H3O+][OH−] = 1.0 × 10−14

-14 -14

– -10

-4

3

1.0 10 1.0 10[OH ] 5.0 10 M

[H O ] 2.0 10


the Concept of pH



Some Strong Acids and Some Weak Acids


the Concept of pH



Concentrations and Kw


the Concept of pH



Hydronium Ions and Hydroxide Ions,

continued Calculating [H3O

+] and [OH–]

Sample Problem A

A 1.0 10–4 M solution of HNO3 has been prepared for a

laboratory experiment.

a. Calculate the [H3O+] of this solution.

b. Calculate the [OH–].


the Concept of pH



Sample Problem A Solution Given: Concentration of the solution = 1.0 × 10−4 M HNO3

Unknown: a. [H3O+]

b. [OH−]

Solution:

• HNO3 is a strong acid

l + l aq + aq–

3 2 3 3HNO ( ) H O( ) H O ( ) NO ( )

3

3

mol HNOmolarity of HNO

1 L solution


the Concept of pH

a. 1 mol 1 mol 1 mol 1 mol



+] and [OH–], continued



Sample Problem A Solution, continued

3 3 33

3

mol HNO 1 mol H O mol H Omolarity of H O

L solution 1 mol HNO L solution

–14–

3

1.0 10[OH ]

[H O ]


the Concept of pH

a.

b. [H3O+][OH−] = 1.0 × 10−14






Sample Problem A Solution, continued

–4

3 3

3

–4–3 4

3

1.0 10 mol HNO 1 mol H O

1 L solution 1 mol HNO

1.0 10 mol H O

1 L solution1.0 10 M H O




-10

–14 –14–

-4

3

1.0 10 1.0 10[OH ]

[H O ] 1.0 101.0 10 M


the Concept of pH

a.

b.



The pH Scale

• The pH of a solution is defined as the negative of the

common logarithm of the hydronium ion concentration,

[H3O+].

pH = −log [H3O+]

• example: a neutral solution has a [H3O+] = 1×10−7

• The logarithm of 1×10−7 is −7.0.

pH = −log [H3O+] = −log(1 × 10−7) = −(−7.0) = 7.0


the Concept of pH



pH Values as Specified [H3O+]


the Concept of pH



The pH Scale

• The pOH of a solution is defined as the negative of the

common logarithm of the hydroxide ion concentration,

[OH−].

pOH = −log [OH–]

• example: a neutral solution has a [OH–] = 1×10−7

• The pH = 7.0.

• The negative logarithm of Kw at 25°C is 14.0.

pH + pOH = 14.0


the Concept of pH



Visual Concepts

pOH

Chapter 15



The pH Scale


the Concept of pH



Approximate pH Range of Common Materials


the Concept of pH



[H3O+], [OH–], pH and pOH of Solutions


the Concept of pH



Visual Concepts

Comparing pH and pOH

Chapter 15



Calculations Involving pH

• There must be as many significant figures to the right

of the decimal as there are in the number whose

logarithm was found.

• example: [H3O+] = 1 × 10−7

one significant figure

pH = 7.0


the Concept of pH



Using Logarithms in pH Calculations


the Concept of pH



Calculations Involving pH, continued

Calculating pH from [H3O+], continued

Sample Problem B

What is the pH of a 1.0 10–3 M NaOH solution?


the Concept of pH



–14 –14-11

3 – -3

1.0 10 1.0 10[H O ] 1.0 10 M

[OH ] 1.0 10


the Concept of pH

Sample Problem B Solution

Given: Identity and concentration of solution = 1.0 × 10−3 M NaOH

Unknown: pH of solution

Solution: concentration of base → concentration of OH−

→ concentration of H3O+ → pH

[H3O+][OH−] = 1.0 × 10−14

pH = −log [H3O+] = −log(1.0 × 10−11) = 11.00





• pH = −log [H3O+]

• log [H3O+] = −pH

• [H3O+] = antilog (−pH)

• [H3O+] = 10−pH

• The simplest cases are those in which pH values are

integers.


the Concept of pH






Calculating [H3O+] and [OH–] from pH, continued

Sample Problem D

Determine the hydronium ion concentration of an

aqueous solution that has a pH of 4.0.


the Concept of pH




Calculating [H3O+] and [OH–] from pH, continued

Sample Problem D Solution

Given: pH = 4.0

Unknown: [H3O+]

Solution:

[H3O+] = 10−pH

[H3O+] = 1 × 10−4 M


the Concept of pH




pH Calculations and the Strength of Acids and Bases

• The pH of solutions of weak acids and weak bases

must be measured experimentally.

• The [H3O+] and [OH−] can then be calculated from the

measured pH values.


the Concept of pH



pH of Strong and Weak Acids and Bases


the Concept of pH



pH Values of Some Common Materials


the Concept of pH



Objectives

• Describe how an acid-base indicator functions.

• Explain how to carry out an acid-base titration.

• Calculate the molarity of a solution from titration

data.

Chapter 15 Section 2 Determining pH and

Titrations



Indicators and pH Meters

• Acid-base indicators are compounds whose colors

are sensitive to pH.

• Indicators change colors because they are either

weak acids or weak bases.

– In + InH H


Titrations

• HIn and In− are different colors.

• In acidic solutions, most of the indicator is HIn

• In basic solutions, most of the indicator is In–




• The pH range over which an indicator changes color

is called its transition interval.

• Indicators that change color at pH lower than 7 are

stronger acids than the other types of indicators.

• They tend to ionize more than the others.

• Indicators that undergo transition in the higher pH

range are weaker acids.


Titrations




• A pH meter determines the pH of a solution by

measuring the voltage between the two electrodes

that are placed in the solution.

• The voltage changes as the hydronium ion

concentration in the solution changes.

• Measures pH more precisely than indicators


Titrations



Color Ranges of Indicators


Titrations




Titrations

Color Ranges of Indicators



Titration

• Neutralization occurs when hydronium ions and

hydroxide ions are supplied in equal numbers by

reactants.

H3O+(aq) + OH−(aq) 2H2O(l)


Titrations

• Titration is the controlled addition and measurement

of the amount of a solution of known concentration

required to react completely with a measured amount

of a solution of unknown concentration.



Titration, continued

Equivalence Point

• The point at which the two solutions used in a titration

are present in chemically equivalent amounts is the

equivalence point.

• The point in a titration at which an indicator changes

color is called the end point of the indicator.


Titrations




Equivalence Point, continued

• Indicators that undergo transition at about pH 7 are

used to determine the equivalence point of strong-

acid/strong base titrations.

• The neutralization of strong acids with strong bases

produces a salt solution with a pH of 7.


Titrations





• Indicators that change color at pH lower than 7 are

used to determine the equivalence point of strong-

acid/weak-base titrations.

• The equivalence point of a strong-acid/weak-base

titration is acidic.


Titrations





• Indicators that change color at pH higher than 7 are

used to determine the equivalence point of weak-

acid/strong-base titrations.

• The equivalence point of a weak-acid/strong-base

titration is basic.


Titrations



Titration Curve

for a Strong Acid

and a Strong

Base


Titrations



Titration Curve

for a Weak Acid

and a Strong

Base


Titrations



Molarity and Titration

• The solution that contains the precisely known

concentration of a solute is known as a standard

solution.

• A primary standard is a highly purified solid

compound used to check the concentration of the

known solution in a titration

• The standard solution can be used to determine the

molarity of another solution by titration.


Titrations



Performing a Titration, Part 1


Titrations



Performing a Titration, Part 2


Titrations



Molarity and Titration, continued

• To determine the molarity of an acidic solution, 10 mL

HCl, by titration

1. Titrate acid with a standard base solution

20.00 mL of 5.0 × 10−3 M NaOH was titrated

2. Write the balanced neutralization reaction

equation.

HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)


Titrations

1 mol 1 mol 1 mol 1 mol

3. Determine the chemically equivalent amounts

of HCl and NaOH.




4. Calculate the number of moles of NaOH used in

the titration.

• 20.0 mL of 5.0 × 10−3 M NaOH is needed to reach the

end point

-3-45.0 10 mol NaOH 1 L

20 mL 1.0 10 mol NaOH used1 L 1000 mL

-4-21.0 10 mol HCl 1000 mL

1.0 10 M HCl10.0 mL 1 L


Titrations

5. amount of HCl = mol NaOH = 1.0 × 10−4 mol

6. Calculate the molarity of the HCl solution




1. Start with the balanced equation for the

neutralization reaction, and determine the

chemically equivalent amounts of the acid and

base.

2. Determine the moles of acid (or base) from the

known solution used during the titration.

3. Determine the moles of solute of the unknown

solution used during the titration.

4. Determine the molarity of the unknown solution.


Titrations




Sample Problem F

In a titration, 27.4 mL of 0.0154 M Ba(OH)2 is added to

a 20.0 mL sample of HCl solution of unknown

concentration until the equivalence point is reached.

What is the molarity of the acid solution?


Titrations




Ba(OH)2 + 2HCl BaCl2 + 2H2O

1 mol 2 mol 1 mol 2 mol


Titrations

Sample Problem F Solution

Given: volume and concentration of known solution

= 27.4 mL of 0.0154 M Ba(OH)2

Unknown: molarity of acid solution

Solution:

1. balanced neutralization equation

chemically equivalent amounts




Sample Problem F Solution, continued

2. volume of known basic solution used (mL)

amount of base used (mol)

2

2 2

mol Ba(OH) 1 LmL of Ba(OH) solution mol Ba(OH)

1 L 1000 mL

2

2

2 mol HClmol of Ba(OH) in known solution mol HCl

mol Ba(OH)


Titrations

3. mole ratio, moles of base used

moles of acid used from unknown solution





4. volume of unknown, moles of solute in unknown

molarity of unknown

amount of solute in unknown solution (mol) 1000 mL

volume of unknown solution (mL) 1 L

molarity of unknown solution


Titrations





1. 1 mol Ba(OH)2 for every 2 mol HCl.

22

-4

2

0.0154 mol Ba(OH)24.7 mL of Ba(OH) solution

1 L

1 L4.22 10 mol Ba(OH)

1000 mL

–4

2

2

–4

2 mol HCl4.22 10 mol of Ba(OH)

1 mol Ba(OH)

8.44 10 mol HCl


Titrations

2.

3.





-2-48.44 10 mol HCl 1000 mL

20.0 m4.22 10

L 1M l

LHC


Titrations

4.

chapter 14 acids and bases table of contents · acids 1. aqueous solutions of acids have a sour...

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