acids and bases bettelheim, brown, campbell and farrell chapter 9

Acids and Bases

Bettelheim, Brown, Campbell and Farrell

Chapter 9

Arrhenius Acids and Bases

– acid:acid: a substance that produces H3O+ ions aqueous solution

– When HCl dissolves in water, its reacts with water to give hydronium ion and chloride ion

H+(aq) + H2O(l) H3O+(aq)Hydronium ion

HCl(aq)+H2O(l) H3O+(aq) + Cl-(aq)

H O

:

+ H Cl

:: : H O H

:

+H

+Cl -

::: ::

H

Arrhenius Acids and Bases

base:base: a substance that produces OH- ions in aqueous solution

– other bases are not hydroxides; these bases produce OH- by reacting with water molecules

NaOH(s) H2O Na+(aq) +OH-(aq)

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

HO H:

:+ H N HH

H+ + O

:::-

H NH

H: H

Solutions

Acidic Solution: H3O+ > OH-

(low pH)

Basic (Alkaline)

Solution: H3O+ < OH-

(high pH)

Neutral Solution: H3O+ = OH-

(pH ~ 7)

Acid and Base Strength

– Strong acid:Strong acid: one that reacts completely or almost completely with water to form H3O+ ions

– Strong base:Strong base: one that reacts completely or almost completely with water to form OH- ions

Strong Acids & Bases

All others are weak acids or bases

HClHBrHIHNO3

H2SO4

HClO4

LiOHNaOHKOH

Ba(OH)2

Hydrochloric acidHydrobromic acidHydroiodic acidNitric acidSulfuric acidPerchloric acid

Lithium hydroxideSodium hydroxidePotassium hydroxideBarium hydroxide

Formula Name Formula Name

Acid and Base Strength

• Weak acid:Weak acid: a substance that dissociates only partially in water to produce H3O+ ions

• Weak base:Weak base: a substance that dissociates only partially in water to produce OH- ions

CH3COOH(aq) + H2O(l) CH3COO-(aq) + H3O+(aq)

Acetic acid Acetate ion

NH3(aq) + H2O(l) NH4+(aq) + OH-(aq)

Brønsted-Lowry Acids & Bases– Acid:Acid: a proton donor– Base:Base: a proton acceptor– Acid-base reaction:Acid-base reaction: a proton transfer reaction– Conjugate acid-base pair:Conjugate acid-base pair: any pair of

molecules or ions that can be interconverted by transfer of a proton

HCl(aq) + H2O(l) H3O+(aq)+Cl-(aq)

WaterHydrogenchloride

Hydroniumion

Chlorideion

(base)(acid) (conjugateacid of water)

(conjugatebase of HCl)

conjugate acid-base pair


Conjugate acid base pair differ only by a proton H+

HCl(aq) + H2O(l) H3O+(aq)+Cl-(aq)

WaterHydrogenchloride

Hydroniumion

Chlorideion

(base)(acid) (conjugateacid of water)

(conjugatebase of HCl)



Brønsted-Lowry Acids & Bases

• Brønsted-Lowry definitions do not require water as a reactant

NH4+CH3COOH CH3COO-

NH3

(base) (conjugate baseacetic acid)

(conjugate acidof ammonia)


+ +Acetic acid Ammonia

(acid)


Acetate ion

Ammoniumion

CH3-C-OO

H N HH

H CH3-C-O -

OH N H

H

H+ +

Acetic acid(proton donor)

Acetate ion

+:

:: ::

:: :: :

Ammonia(proton acceptor)

Ammoniumion

Conjugate Acids and Bases

HA + B ↔ BH+ + A-

acid base conjugate conjugate acid of B base of HA

An acid will react to form its conjugate base.A base will react to form its conjugate acid.

Conjugate acid-base pairs differ only by a hydrogen

Conjugate Acid-Base Pairs

If acid is very strong, its corresponding conjugate base is very weak.

Stronger acid ionizes more completely, so the base will not attract the H3O+ well

Example:

HCl Cl-

Conjugate Acid-Base Pairs

If base is stronger, its corresponding conjugate acid is weaker.

Stronger base accepts H+ more easily, so the acid will not donate H+ well

Examples: OH- HOH

PO33- HPO4

2-

Fig. 9.2Strong

Weak

Weak

Strong

C2H5OH C2H5O-H2O OH-HPO4

2- PO43-

HCO3- CO3

2-

C6H5OH C6H5O-HCN CN-

NH3NH4+

H2PO4- HPO4

2-

H2S HS-H2CO3 HCO3

-CH3COOH CH3COO-H3PO4 H2PO4

-HSO4

- SO42-

H2OH3O+HNO3 NO3

-H2SO4 HSO4

-HCl Cl-HI I-Hydroiodic acid

Hydrochloric acidSulfuric acid

Dihydrogen phosphateAcetateBicarbonate

Hydrogen phosphateAmmonia

Phenoxide

Carbonate

PhosphateHydroxideEthoxide

Hydrogen sulfide

Nitric acidHydronium ion

Hydrogen sulfate ion

Name of acid Name of ion

Phosphoric acidAcetic acidCarbonic acid

Dihydrogen phosphateAmmonium ion

Phenol

Bicarbonate ion

Hydrogen phosphate ionWaterEthanol

Hydrogen sulfide

AcidConjugate Base

IodideChlorideHydrogen sulfateNitrateWater

Sulfate

StrongAcids

Weak Acids

Weak Bases

StrongBases

Hydrocyanic acid Cyanide

Water as an Acid and a Base

• Water Water is is amphoteric:amphoteric:

HC2H3O2 + H2O H3O+ + C2H3O2-

acid base acid base

H2O + NH3 NH4+ + OH-

acid base acid base

• Water Water is is amphoteric:amphoteric:

HC2H3O2 + H2O H3O+ + C2H3O2-

acid base acid base

H2O + NH3 NH4+ + OH-

acid base acid base

..

..

..

Acid-Base Equilibria

– For weak acids, significant amounts of both the acid and its conjugate base will be present and form an equilibrium

HCl + H2O H3O++Cl-

H3O+CH3COO-H2OCH3COOH + +

Acetic acid Acetate ionWeak Acid

Strong Acid


What if the base is not water? How can we determine which are the major species present?

CH3COOH NH3 CH3COO-NH4

++ +

Acetic acid Acetate ion

?

Ammonia Ammonium ion(conjugate baseof CH3COOH

(conjugate acidof NH3

(acid) (base)

Equilibrium lies on the side of the

weaker acid and weaker base.

C2H5OH C2H5O-H2O OH-HPO4

2- PO43-

HCO3- CO3

2-

C6H5OH C6H5O-HCN CN-

NH3NH4+

H2PO4- HPO4

2-

H2S HS-H2CO3 HCO3

-CH3COOH CH3COO-H3PO4 H2PO4

-HSO4

- SO42-

H2OH3O+HNO3 NO3

-H2SO4 HSO4

-HCl Cl-HI I-Hydroiodic acid

Hydrochloric acidSulfuric acid

Dihydrogen phosphateAcetateBicarbonate

Hydrogen phosphateAmmonia

Phenoxide

Carbonate

PhosphateHydroxideEthoxide

Hydrogen sulfide

Nitric acidHydronium ion

Hydrogen sulfate ion

Name of acid Name of ion

Phosphoric acidAcetic acidCarbonic acid

Dihydrogen phosphateAmmonium ion

Phenol

Bicarbonate ion

Hydrogen phosphate ionWaterEthanol

Hydrogen sulfide

AcidConjugate Base

IodideChlorideHydrogen sulfateNitrateWater

Sulfate

StrongAcids

Weak Acids

Weak Bases

StrongBases

Hydrocyanic acid Cyanide


CH3COOH NH3 CH3COO-NH4

++ +

Acetic acid(stronger acid)

Acetate ion(weaker base)

Ammonia(stronger base)

Ammonium ion(weaker acid)

?

The position of this equilibrium lies to the right—

Formation of weaker acid and weaker base is favoredExample:Example: Predict the position of equilibrium in this

acid-base reaction

H2CO3 OH- HCO3- H2O+ +

?

H2CO3 OH- HCO3- H2O+ +

?

Stronger

acid

Stronger

baseWeaker base

Weaker acid

Can see from table:

Bases: OH- stronger than HCO32-

Acids: H2CO3 stronger than H2O

Right side favored—Equilibrium to the right

Acid Ionization Constants

– The equilibrium constant, Keq, is

– Treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L

– combine the two constants to give a new constant, which we call an acid ionization constant, Ka

HA H2O A- H3O++ +

[HA][H2O]

[A-][H3O+]Keq =

[HA]

[A-][H3O+]Ka = Keq[H2O] =

Acid Ionization Constants: Ka

– Acetic Acid: Ka for acetic acid is 1.8 x 10-5

– pKa for acetic acid is 4.75

pKa = -log Ka

H3PO4

HCOOH

CH3CH(OH)COOH

CH3COOH

H2CO3

H2PO4-

H3BO3

NH4+

C6H5OH

HPO42-

HCO3-

HCN

Phosphoric acid

Formic acid

Lactic acid

Acetic acid

Carbonic acid

Dihydrogen phosphate ion

Name

7.21

pKa

9.14

9.25

9.89

12.66

10.25

Boric acid

Ammonium ion

Phenol

Hydrogen phosphate ion

Bicarbonate ion

Acid

7.5 x 10-3

1.8 x 10-4

8.4 x 10-4

1.8 x 10-5

4.3 x 10-7

6.2 x 10-8

Ka

7.3 x 10-10

5.6 x 10-10

1.3 x 10-10

2.2 x 10-13

5.6 x 10-11

2.12

3.75

3.08

4.75

6.37

Hydrocyanic acid 4.9 x 10-10 9.31

p = -log

Acid Ionization Constants: Ka

– Weak acid has the small Ka, but large pKa

pKa = -log Ka

H3PO4

HCOOH

CH3CH(OH)COOH

CH3COOH

H2CO3

H2PO4-

H3BO3

NH4+

C6H5OH

HPO42-

HCO3-

HCN

Phosphoric acid

Formic acid

Lactic acid

Acetic acid

Carbonic acid


Name

7.21

pKa

9.14

9.25

9.89

12.66

10.25

Boric acid

Ammonium ion

Phenol


Bicarbonate ion

Acid

7.5 x 10-3

1.8 x 10-4

8.4 x 10-4

1.8 x 10-5

4.3 x 10-7

6.2 x 10-8

Ka

7.3 x 10-10

5.6 x 10-10

1.3 x 10-10

2.2 x 10-13

5.6 x 10-11

2.12

3.75

3.08

4.75

6.37


p = -log

H3PO4

HCOOH

CH3CH(OH)COOH

CH3COOH

H2CO3

H2PO4-

H3BO3

NH4+

C6H5OH

HPO42-

HCO3-

HCN

Phosphoric acid

Formic acid

Lactic acid

Acetic acid

Carbonic acid


Name

7.21

pKa

9.14

9.25

9.89

12.66

10.25

Boric acid

Ammonium ion

Phenol


Bicarbonate ion

Acid

7.5 x 10-3

1.8 x 10-4

8.4 x 10-4

1.8 x 10-5

4.3 x 10-7

6.2 x 10-8

Ka

7.3 x 10-10

5.6 x 10-10

1.3 x 10-10

2.2 x 10-13

5.6 x 10-11

2.12

3.75

3.08

4.75

6.37


Self-Ionization of Water

[H2O] as a constant = 55.5 mol/L

H2O+H2O H3O++OH-

BaseAcid Conjugateacid of H2O

Conjugatebase of H2O

[H2O]2

[H3O+][HO-]Keq = waterionization.mov


– Ion product of water, KIon product of water, Kww = = 1.0 x 10-14

– Product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14

[H3O+][OH-]Kw = Keq[H2O]2 =

Kw = 1.0 x 10-14

[H3O+]

[OH-]

= 1.0 x 10-7 mol/L

= 1.0 x 10-7 mol/Lin pure water


– Example: add 0.010 mole of HCl to 1 liter of pure water, in this solution, [H3O+] is 0.010 or 1.0 x 10-2. What is hydroxide ion concentration?

pH = -log [H3O+] p = -log

pOH = -log [OH-]

[OH-] = 1.0 x 10-14

1.0 x 10-2= 1.0 x 10-12

pH Scale

By definition:

pH = - log [H3O+]

p = “- log”

Brackets used to show concentration (M)

Scale ranges from 0 to 14

pH scale

Range from 0 to 14

pH = -log [H3O+]

pH < 7 acidic

pH = 7 neutral

pH > 7 basic

Solutions

Acidic Solution: H3O+ > OH-

Basic (Alkaline)

Solution: H3O+ < OH-

Neutral Solution: H3O+ = OH-

pH ScaleTypical values range from 0 to 14

pH = 7 – neutral

pH > 7 – basic

pH < 7 – acidic

pH• pH of some common materials

pH

Battery acidGastric juiceLemon juiceVinegarTomato juiceCarbonated beveragesBlack coffee

UrineRain (unpolluted)

Milk

SalivaPure waterBloodBilePancreatic fluidSeawaterSoap

Milk of magnesiaHousehold ammonia

Lye (1.0 M NaOH)

0.51.0-3.02.2-2.42.4-3.44.0-4.44.0-5.05.0-5.1

5.5-7.56.2

6.3-6.6

6.5-7.57.0

7.35-7.456.8-7.07.8-8.08.0-9.08.0-10.0

10.511.7

14.0

Material pHMaterial

pH of Salt Solutions

NaCH3COO = Sodium Acetate

In water:

NaCH3COO → Na+ + CH3COO-

Na+ + OH- → NaOH (strong base)

CH3COO- + H+ → CH3COOH (weak acid)

Strong base + weak acid = basic salt

pH of Salt Solutions

• NH4Cl Ammonium chloride

• NH4+ + OH- → NH4OH (weak base)

• H+ + Cl- → HCl (strong acid)

• Strong acid + weak base = acidic salt

Neutralization ReactionsAlso known as Acid-Base Reactions

H3O+ + OH- → 2 H2O

Acid + Base → Water

Neutralization is Special Case of Double Replacement Reaction--- Water is a product

HCl + NaOH → NaCl + H2O

Acid + Base → Salt + Water

Neutralization Reaction

Acid and Base react with each other to form water and a salt

Frequently use a pH indicator to show when end point has been reached

End point is pH at which [H3O+] = [OH-]

and color chosen indicator changes

Titration

Titration:

Volume of a solution of knownconcentration is added to a solution

of unknown concentration.

Measure amount of known solution needed to react exactly with original material - at endpoint

Equivalence point:

Amount of acid = Amount of base

End point:

Indicator color changes

Fig. 9.6

Acid-Base Titrations

• Use 0.108 M H2SO4 to determine the concentration of a NaOH solution – HOW?HOW?

2NaOH(aq)+H2SO4(aq) Na2SO4(aq) + 2H2O(l)(concentration

known)(concentrationnot known)

titration.mov

Example:

Add NaOH (a base) to a solution of acid that contains the pH indicator phenolphthalein

Phenolphthalein is colorless in acidic solution and is pink in basic solution

When enough NaOH is added so that it neutralizes all of the acid, any additional NaOH makes the solution basic and it turns pink

What is the concentration of an acid solution if 12.54 mL of 0.1000 M NaOH

neutralizes 10.00 mL of acid?At endpoint: mol Acid = mol Base

MaVa = MbVb

Ma = MbVb = (0.1000 M) ( 12.54 mL)

Va 10.00 mL

= 0.1254 M

pH indicator (acid-base indicator): substance that turns color when the H3O+ (acid) concentration changes

pH indicators may be extracted from many natural products

lichens, apple skins, blueberries, red cabbage, etc.

pH indicators may be in many forms: embedded in paper (pH or litmus

paper)liquid

Buffer Solutions

• Contain a weak acid and its conjugate base

• Contain a weak base and its conjugate acid

• Two species differ only by an H+

• Add extra acid—the conjugate base will react to remove the added H+

• Add extra base—the conjugate acid will react to remove the added OH-

• Extra acid or base is removed, so pH will remain relatively constant

pH Buffers

– add a strong acid, such as HCl, adds H3O+ ions react with acetate ions and are removed from solution

– add a strong base, such as NaOH, adds OH- ions react with acetic acid and are removed from solution

CH3COO- H3O+ CH3COOH H2O+ +

CH3COOH OH- CH3COO- H2O+ +

CH3COOH H2O CH3COO- H3O++ +

Added asCH3COOH

Added asCH3COO-Na+

pH Buffers

– Consider a phosphate buffer of 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base)

waterpH

0.10 M phosphate buffer7.07.21

2.0 12.07.12 7.30

pH afteraddition of

0.010 mole HCl

pH afteraddition of

0.010 mole NaOH

Illustration of Buffer Effects• What is the pH of a solution 0.1 M in formic acid (HCOOH) and 0.10 M in HCOONa when Ka=1.8 x 10 -4?

What is the pH of the solution after 0.03 mol of NaOH is added to 1.0 L of the buffer?

[H3O+]= Ka [HA]/[A-] [H3O+] = 1.8 x 10-4 (0.10)/(0.10)

pH= 3.8

[H3O+] = 1.8 x 10-4 (0.07)/(0.10)pH = 3.9

Blood Buffers

• The average pH of human blood is 7.4

• The body uses three buffer systems– carbonate buffer:carbonate buffer: H2CO3 and its conjugate

base, HCO3-

– phosphate buffer:phosphate buffer: H2PO4- and its conjugate

base, HPO42-

– proteins:proteins: discussed in Chapter 21

Henderson-Hasselbalch Eqn

• Henderson-Hasselbalch equation:Henderson-Hasselbalch equation: shows mathematical relationship between – pH, – pKa of the weak acid, HA – concentrations HA, and its conjugate base, A-

[HA]

[A-]+ logpH = pKa Henderson-Hasselbalch Equation

acids and bases bettelheim, brown, campbell and farrell chapter 9

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