chapter 14 acids and bases. homework assigned questions and problems (odd only) section 14.1...

Chapter 14

Acids and Bases

Homework• Assigned Questions and Problems (odd only)

• Section 14.1 (optional)• Section 14.2 Read this section• Section 14.3 Read this section• Section 14.4 33, 35, 37, 39• Section 14.5 Section on Neutralization Reactions, 43

• Section 14.6 51, 53, 55, 97, 103• Section 14.8 65, 67, 69, 109, 119• Section 14.9 71, 73, 75, 77, 79, 81, 83, 85, 87, 107, 121

14.2 Acids: Properties and Examples

• Among the most common and important compounds known

• Aqueous solutions are important in – biological systems– chemical industrial processes

• Early known characteristics– Causes a sour taste (lemons

and vinegar)– Causes litmus (dye) to change

from blue to red– Dissolves some metals

generating hydrogen gas (e.g., zinc and iron)

2Al(s)+ 6HCl(aq)→ 2AlCl3(aq)+3H2(g)

Acetic Acid HC2H3O2(aq)

14.2 Acids: Properties and Examples

• Common Acids (common laboratory reagents)– Sulfuric acid (H2SO4): most produced chemical in

U.S. – Used to produce phosphoric acid and phosphate

fertilizers, rust removal in iron and steel production, paper production

– Hydrochloric acid (HCl): steel production, organic compound production, pH control and neutralization, main component of stomach acid

– Nitric Acid (HNO3): mainly used in the production of fertilzers, explosives, and used to dissolve metals

– See table 14.1

Section 5.9 Naming Binary Acids

• Use the prefix hydro- before the root name of the element

• Add the suffix -ic and the word acid to the root name for the element

• Example: HCl– hydrochloric acid

• Example: HI– hydroiodic acid

Section 5.9 Naming Oxyacids

• Produce a hydrogen ion (H+) and a polyatomic ion when dissolved in water

• Composed of hydrogen, oxygen, and another nonmetal

• Use the root name of the polyatomic ion

• If it ends in -ate use the suffix -ic acid

• If it ends in -ite use the suffix -ous acid

• Example: H2SO4 (from SO42- , sulfate ion)

– sulfuric acid• Example: H2SO3 (from SO3

2- , sulfite ion)– sulfurous acid

14.3 Bases: Properties and Examples

• Early known characteristics– Causes a bitter taste (when

dissolved in water)– Causes litmus (dye) to change

from red to blue– Dissolves fats that are placed in

base solutions– Feel slippery like soap

• Also recognized as an important group of compounds

• Aqueous solutions are important in – biological systems– chemical industrial processes

Naming Bases• Most bases are usually ionic compounds • The hydroxide ion has a charge of (-1) and is

combined with a positively charged ion (group IA or IIA metal ion)

• Hydroxides (bases) are named by their cation first, then the word “hydroxide” is added to the metal cation

• Most common bases: – NaOH (sodium hydroxide)– KOH (potassium hydroxide) – Ca(OH)2 (calcium hydroxide)– NH4OH (ammonium hydroxide)

14.4 The Arrhenius Definition• Arrhenius (1884) first person to recognize the

essential nature of acids and bases• Defined them in terms of chemical composition• Arrhenius defined an acid as a substance that

produces hydrogen ions (H+) in water• Acids are covalent compounds that ionize in

water to produce H+

• Acids:(aq)Cl(aq)H(g)HCl OH2

hydrogen chloride gas

14.4 Molecular Definitions of Acids and Bases

• Arrhenius Acids• When dissolved in water will generate H+ and an

anion

– The H+ that is generated will give a sour taste– Vinegar (acetic acid)– Lemon (citric acid)

• Most acids are oxyacids (will also gen. H+)

(aq)Cl(aq)H(g)HCl OH2

(aq)NO(aq)H)(HNO 3OH

32 l

Types of Acids• Multiprotic Acids

–Can donate more than one proton–H2SO4

–H3PO4

H2SO4H2O HSO4

H

HSO4 H2O SO4

2 HStrong Acid

Weak Acid

Types of Acids• Binary Acids

• Molecular compounds in which hydrogen is attached to a second nonmetallic element

• Formed from the pure compound which has been dissolved in water

• Oxyacids• Molecular compounds composed of hydrogen, oxygen, and

a nonmetal (e.g. S, C, N, Cl, or P)• Molecular compounds which become acids when dissolved

in water

• Acid hydrogen is attached to an oxygen• Formulas are like polyatomic ions with hydrogen(s) attached


(aq)NO(aq)H)(HNO 3OH

32 l

Types of Acids• Organic Acids

– Organic compounds (carbon-containing compounds) with acidic properties. – Most common type are called carboxylic acids. The acidic portion of the

molecule is called a carboxyl group– For example, acetic acid (the active component of vinegar) contains a

carboxyl group

H3C CO

O HAcetate AnionAcetic Acid


• Arrhenius Bases• Arrhenius defined a base as a substance which

contains hydroxide and produces hydroxide ions (OH-) in water

• Bases are ionic compounds that produce hydroxide ions (OH-) in water

• Bases:NaOH(s) H2O Na(aq) OH (aq)


• Arrhenius Bases– When dissolved in water will generate OH- and a metal

ion. Two common examples:

• Bases are sometimes called alkalis– The OH- that is generated will give a bitter taste– Slippery feel (like soap)– Most bases that are generated are composed from

group 1A and 2A metals

(aq)OH(aq)Na(s)NaOH OH2

(aq)OH(aq)K(s)KOH OH2

14.4 The Brønsted-Lowry Definition of Acids and Bases

• Arrhenius definition is widely used, but model has limitations – Only for aqueous solutions– Does not explain how compounds such as

ammonia (does not contain OH-) produce a basic water solution

• New model (1923) by Brønsted and Lowry proposed a broader definition

• New model applied to both aqueous and non-aqueous solutions

• Explained how compounds without OH- could produce a basic solution when added to water

14.4 The Brønsted-Lowry Definition of Acids and Bases

• Brønsted-Lowry Acid– Any substance that can donate a proton (H+) to another

substance: a proton donor• Brønsted-Lowry Base

– Any substance that can accept a proton (H+) from another substance: a proton acceptor

– Proton donation does not occur unless a proton acceptor is present

• Arrhenius Acids/Bases vs. Br.-Lowry Acids/Bases– All acids in the Arrhenius definition are also acids according

to the Bronsted/Lowry model– The Br.-Lowry definition of bases is quite different from that

of Arrhenius– H+ ions produced by Br.-Lowry acids also react with water

Water as a Base• For example: HCl is a Bronsted/Lowry acid because

it produces H+ in water

• H+ does not exist in water due to the strong attraction to the polar water molecule

• In an aqueous solution the H+ ion generated reacts with water to form H3O+

H ClH O

H

H O

H

H Cl


Hydronium ionH+ acceptor

H+ donor

14.4 Identifying Brønsted-Lowry Bases and Their Conjugates

• Any chemical reaction that involves a Bronsted-Lowry acid must also involve a Br.-Lowry base

• The reaction involves a proton transfer• Acid can lose its proton to form a conjugate base

(something that could accept a proton back again)• Base accepts a proton to form a conjugate acid

(something that could donate a proton)

(aq)A(aq)OH)(OH(aq)HA 32

Acid Base Conjugate Acid

Conjugate Base

14.4 Identifying Brønsted-Lowry Bases and Their Conjugates

• HCl is the Brønsted-Lowry acid because it is donating a proton (H+) to the water molecule

• The water molecule is the Brønsted-Lowry base since it accepts a proton (H+)

(aq)Cl(aq)OH)(OH(g)HCl 32

Acid Base Conjugate Acid

Conjugate Base

14.4 Conjugate Acid/Base Pairs

• Acids/bases and conj acids/conj bases are related to each other by donating/accepting a single proton (H+)–The acid of the pair has the proton–The base of the pair does not–Every acid has a conjugate base–Every base has a conjugate acid–“Conjugate” is given to the part of the

pair that is produced in the reaction


• Each acid is related to a base on the opposite side–On the reactants side (left side)

substances are called “acid” or “base”–On the products side (right side)

substances are called “conjugate acid” or “conjugate base”

)(OH)(NH)(OH)(NH 423 aqaqlaq

Base Acid Conjugate Acid

Conjugate Base


• Which of the following represent conjugate acid-base pairs?• H2O, H3O+

• OH-, HNO3

• H2SO4, SO42-

• HC2H3O2, C2H3O2-


H2O, H3O+ H2O H H3O

OH-, HNO3Not Pairs

H2SO4, SO42- H2SO4 SO4

2- +2H+ Not Pairs

HC2H3O2, C2H3O2-

HC2H3O2 C2H3O2- + H+

+442 H+-SOHSOH

+33 H+NOHNO -

14.5 Reactions of Acids and BasesNeutralization Reactions

• When Arrhenius acids and bases are mixed, they will react with each other

• The acidic properties and the basic properties are destroyed and this means both substances are neutralized

• A neutralization is the reaction between an acid and a hydroxide base which forms a salt and water


• A neutralization is a double-displacement reaction

• Whenever an acid is completely reacted with a base, a neutralization occurs

• The net ionic equation is water formation

CBADCDAB

)(OH)(KCl)(KOH)(HCl 2 aqaqaqAcid Base Salt Water

)(OH)(OH)(H 2- aqaq


• Neutralization: The reaction between an acid and a base to form a salt and water

• H+ from the acid combines with the OH- from the base to form water

• The properties of both reactants are neutralized

• The salt contains the positive ion from the base and the negative ion from the acid


• The reaction hydrochloric acid and sodium hydroxide:• A double replacement reaction

• Ionic Equation

• Net Ionic Equation

)(OH )(NaCl)(NaOH )( HCl2

laqaqaq

acid base salt water

)(OH)(-Cl)(Na)(-OH)(Na)(-Cl )(H2

laqaqaqaqaqaq

)(OH)(-OH )(H2

laqaq

14.8 Water: Acid and Base in One (Ionization of Water)

• Substances that can act as acids or bases are termed amphoteric

• Water can act as an acid or a base– Amphoteric substance

H O

H

H O

H

H O

H

H O

H

Base Conj. BaseAcid Conj. Acid

H+ acceptor

H+ donor

)(OH)(OH)( OH)(OH -322 aqaq

Hydronium ion


• Experiments show that in a sample of pure water, a very small percentage has dissociated to produce ions

• It involves a proton transfer (a Br/L acid-base rxn)• The net result is an equal amount of H3O+ and OH-

)(OH)(OH)(OH)(OH 322 aqaq

H O

H

H O

H

H O

H

H O

H

Base Conj BaseAcid Conj Acid

Equal amounts produced


• The acid-base reaction of water with itself is called autoionization (self-ionization)– It is an equilibrium reaction– Ionization of water to form hydronium ion and

hydroxide ion is balanced by the recombination of ions to form water

– At 25 ºC, the concentration of each ion: 1.0×10-7 M– Square brackets around a compound denotes

“concentration” in moles per liter

H3O OH 110 7 M

OHOHOHOH 322

14.8 Ion-Product Constant for Water

• At any given temperature, the product of the concentration of H+ and OH- is always a constant

• This value can be calculated at 25 °C since the concentration (each ion) is known

• Kw (Ion-Product Constant for Water) is the product of the H3O+ and OH- ion molar concentrations in water

• Valid in pure water or water with solutes

H3O OH 110 14 = Kw M 10 1.00M) 10 (1.00 M) 10 (1.00 -14-7-7


• If the [H3O+] is increased by the addition of acidic solute, the [OH-] must decrease until the expression is 1.0 × 10-14 is satisfied

• Or, if [OH-] is increased by the addition of a basic solute to the water, the [H3O]+ must similarly decrease

14100.1][ -OHOH3

14100.1-OHOH ][ 3


• In an aqueous solution, neither [H3O+] or [OH-] is ever zero

• An acid is a substances that will increase the [H3O+] in solution

• All acidic solutions have a higher [H3O+] than [OH-]

• An base is a substance that will increase the [OH-] ions in solution

• All basic solutions have a higher [OH-] than [H3O+]

][ ][ - OH >OH3

][ ][ - OH OH3


• Neutral Solution

• Acidic Solution

• Basic Solution

• In all cases:

][ ][ - OH >OH3

][ ][ - OH OH3

][ ][ - OH OH3

H3O OH 110 14 = Kw

Using Kw in Calculations, Example 1

• Calculate [H3O+] in a solution in which [OH-] = 2.0x10-2 M. Is this solution acidic, basic or neutral?

= 5.010 13 M

basic][OH ]OH[ -3

H3O OH 110 14 = Kw

)([ M 10 1.010 2.0 14-2-]OH3

[

2-

14-

10 2.0

10 1.0OH ]3

14.9 The pH and pOH Scale

• H3O+ (from H+) concentrations range from very high values to extremely small valves

• Difficult and inconvenient to work with numbers over such a large range

• For example, [H3O+] of 10 M is 1000 trillion times greater than 10-14 M

• The pH scale of a solution was proposed as an easier and more practical way to handle such large numbers

M100.1]OH[M10]OH[ 1433

and

14.9 The pH Scale• The pH scale is defined as the

negative logarithm of the molar hydronium ion concentration

• Logarithms are exponents• The negative powers of 10 in the

concentrations are converted to positive numbers

OH logpH 3

14.9 The pH Scale

• The pH scale uses a common logarithm based on powers of 10

• Expressed mathematically:

• For a number expressed in scientific notation with a coefficient of 1, the logarithm of that number is equal to the value of the exponent

• Take the log of the number and then multiply by negative one (change the sign)

]OH[log 3pH

14.9 Calculating the pH from [H3O+]

• The negative log of the H3O+ concentration– To determine the number of significant figures

for logs: The number of decimal places for the log is equal to the number of sig. figs. in the concentration (original number)

M 101.0OH 43

)00.4(pH 00.4

Given:

M )101.0(log]OH[logpH 43

2 SF

2 D.P.

14.9 Calculating the pH from [H3O+]

• Calculate the pH for the following solutions

–A solution in which [H3O+]=1.0x10-3 M–A solution in which [OH-]= 5.0x10-5 M

Calculating the pH from [H3O+]Calculating the pH from [OH-]

70.9)70.9(pH

)100.1(log]OH[log 33

pH

00.3)00.3(pH

MOHGiven 33 100.1][:

1453 100.1)100.5(]OH[

5

14

3 100.5

100.1]OH[

103 100.2]OH[ )100.2(log]OH[log 10

3 pH

MGiven 5100.5]OH[: 143 100.1]OH[]OH[

Calculating the [H3O+] from pH

• It is often necessary to calculate the hydronium ion concentration for a solution from its pH value

• The logarithm must be undone

• To do this requires determining the antilog of the pH value

• The antilog can be obtained on a calculator using the antilog function

• Use inv log (-) pH value

Calculating the [H3O+] from pH

• The pH of rainwater in a polluted area was found to be 3.50. What is the [H3O+] for this rainwater?

]OH[log 3pH 50.3]OH[log 3

)50.3()]OH[(log 3 loginvloginv M43 1016.3]OH[

]OH[logpH 3

The pH Scale• pH is a log scale • A change in one unit on the pH scale

means a tenfold increase or decrease in [H3O+]–Every time an exponent changes by one,

the pH changes by one

• Lowering the pH increases the [H3O+] – low pH value= acidic solution–high pH value= basic solution

Measuring pH• Use a pH meter

– Electronic device that measures the pH of the solution• Use pH paper

– Paper has a chemical in it that changes to different colors depending on the pH of the solution

• Indicators– A compound that exhibits different colors depending on the pH of

its surroundings

The pOH Scale

• The negative log is also a way of expressing the magnitudes of other quantities

• The concentration of OH- can be expressed as pOH

• Same as pH scale, but associated with the [OH-]

– Low pOH value means high [OH-]

– High pOH value means low [OH-]

OH 1.0 10 4 M

pOH log OH log 1.0 10 4 )00.4(pOH 00.4

pH and pOH• In an aqueous solution, the sum

of the pH and pOH is always 14

• The value 14 corresponds to the of Kw equation

pH pOH 14.00

-log 1.0 × 10-14

pH and pOH

[H+] > [OH-]

[H+] < [OH-]

[H+] = [OH-]

Calculating pOH from pH

• A sample of rain in an area with severe air pollution has a pH of 3.5. What is the pOH of this rain water?

pH pOH 14.00

3.5 pOH 14.00

pOH 14.00 - 3.5 10.5

Calculating [OH-] from pOH

• The pOH of a liquid drain cleaner was found to be 10.50. What is the [OH-] for this cleaner?

pOH log OH 10.50

log OH 10.50

OH inv log( 10.50)OH 3.16 10 11 M

14.6 Acid-Base Titration• Determining the acid or base concentration

in a solution is a routing laboratory practice• The pH only determines the amount of

dissociated H+ in solution• Only dissociated molecules influence the

pH value

• Titration will give info about the total number of acid or base molecules present (concentration)

)(Cl)(H)(lCH aqaqaq

)(Cl)(H)(lCH aqaqaq

14.6 Acid-Base Titration• Titration involves the

gradual addition of one solution to another until the solute in the first solution has reacted completely with the solute in the second solution

• First measure a known volume of acid (unknown molarity) to the flask

•To determine the concentration of an acid solution, add a solution of base of known concentration to the flask by a buret

14.6 Acid-Base Titration• Base addition continues until all the acid has completely reacted with

the added base: endpoint or equivalence point• To determine endpoint, an indicator is used to detect when the acid-

base reaction is complete• If you know

– original volume of the acid– the volume and concentration of the base – the concentration of the acid can be calculated

At endpoint,molesacid = molesbase

Vacid

Mbase

endpoint

Vbase

14.6 Acid-Base Titration• As the base (NaOH) is added, OH- will react

with and neutralize HCl (and free H+), forming water

• At endpoint all of the acid has completely reacted with all of the base

At endpoint,molesacid = molesbase

At endpoint,Macid Vacid = Mbase Vbase

)(OH)(NaCl)(lCH )( NaOH 2 aqaqaq

• end

chapter 14 acids and bases. homework assigned questions and problems (odd only) section 14.1...

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