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Acids and Bases
Bettelheim, Brown, Campbell and Farrell
Chapter 9
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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
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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
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Solutions
Acidic Solution: H3O+ > OH-
(low pH)
Basic (Alkaline)
Solution: H3O+ < OH-
(high pH)
Neutral Solution: H3O+ = OH-
(pH ~ 7)
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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
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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
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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)
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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
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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)
conjugate acid-base pair
conjugate acid-base pair
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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)
conjugate acid-base pair
+ +Acetic acid Ammonia
(acid)
conjugate acid-base pair
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
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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
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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
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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-
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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-
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Fig. 9.2Strong
Weak
Weak
Strong
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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
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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
..
..
..
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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
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Acid-Base Equilibria
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.
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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
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Acid-Base Equilibria
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+ +
?
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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
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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] =
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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
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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
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
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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
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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
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Self-Ionization of Water
– 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
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Self-Ionization of 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
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pH Scale
By definition:
pH = - log [H3O+]
p = “- log”
Brackets used to show concentration (M)
Scale ranges from 0 to 14
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pH scale
Range from 0 to 14
pH = -log [H3O+]
pH < 7 acidic
pH = 7 neutral
pH > 7 basic
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Solutions
Acidic Solution: H3O+ > OH-
Basic (Alkaline)
Solution: H3O+ < OH-
Neutral Solution: H3O+ = OH-
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pH ScaleTypical values range from 0 to 14
pH = 7 – neutral
pH > 7 – basic
pH < 7 – acidic
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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
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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
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pH of Salt Solutions
• NH4Cl Ammonium chloride
• NH4+ + OH- → NH4OH (weak base)
• H+ + Cl- → HCl (strong acid)
• Strong acid + weak base = acidic salt
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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
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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
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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
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Equivalence point:
Amount of acid = Amount of base
End point:
Indicator color changes
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Fig. 9.6
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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
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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
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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
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pH indicator (acid-base indicator): substance that turns color when the H3O+ (acid) concentration changes
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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
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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
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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+
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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
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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
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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
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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