general, organic, and biochemistry, 7e
DESCRIPTION
General, Organic, and Biochemistry, 7e. Bettelheim, Brown, and March. Chapter 8. Acids and Bases. Arrhenius Acids and Bases. In 1884, Svante Arrhenius proposed these definitions acid: a substance that produces H 3 O + ions aqueous solution - PowerPoint PPT PresentationTRANSCRIPT
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General, Organic, and General, Organic, and Biochemistry, 7eBiochemistry, 7e
Bettelheim,Bettelheim,
Brown, and MarchBrown, and March
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Chapter 8Chapter 8
Acids and BasesAcids and Bases
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Arrhenius Acids and BasesArrhenius Acids and Bases• In 1884, Svante Arrhenius proposed these
definitions • acid:acid: a substance that produces H3O+ ions aqueous
solution• base:base: a substance that produces OH- ions in aqueous
solution• this definition of an acid is a slight modification of the
original Arrhenius definition, which was that an acid produces H+ in aqueous solution
• today we know that H+ reacts immediately with a water molecule to give a hydronium ion
H+(aq) + H2O(l) H3O+(aq)Hydronium ion
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Arrhenius Acids and BasesArrhenius Acids and Bases• when HCl, for example, dissolves in water, its reacts
with water to give hydronium ion and chloride ion
• we use curved arrows to show the change in position of electron pairs during this reaction
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 BasesArrhenius Acids and Bases• With bases, the situation is slightly different
• many bases are metal hydroxides such as KOH, NaOH, Mg(OH)2, and Ca(OH)2
• these compounds are ionic solids and when they dissolve in water, their ions merely separate
• 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)
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Arrhenius Acids and BasesArrhenius Acids and Bases• we use curved arrows to show the transfer of a proton
from water to ammonia
HO H
::+ H N H
H
H+ + O
:::
-H N
H
H: H
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Acid and Base StrengthAcid 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
• here are the six most common strong acids and the four most common strong 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 StrengthAcid and Base Strength• Weak acid:Weak acid: a substance that dissociates only
partially in water to produce H3O+ ions• acetic acid, for example, is a weak acid; in water, only 4
out every 1000 molecules are converted to acetate ions
• Weak base:Weak base: a substance that dissociates only partially in water to produce OH- ions• ammonia, for example, is a weak base
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 & BasesBrønsted-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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Brønsted-Lowry Acids & BasesBrønsted-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
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Brønsted-Lowry Acids & BasesBrønsted-Lowry Acids & Bases• we can use curved arrows to show the transfer of a
proton from acetic acid to ammonia
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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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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Brønsted-Lowry Acids & BasesBrønsted-Lowry Acids & Bases• Note the following about the conjugate acid-base
pairs in the table
1. an acid can be positively charged, neutral, or negatively charged; examples of each type are H3O+, H2CO3, and H2PO4
-
2. a base can be negatively charged or neutral; examples are OH-, Cl-, and NH3
3. acids are classified a monoprotic, diprotic, or triprotic depending on the number of protons each may give up; examples are HCl, H2CO3, and H3PO4
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Brønsted-Lowry Acids & BasesBrønsted-Lowry Acids & Bases• carbonic acid, for example can give up one proton to
become bicarbonate ion, and then the second proton to become carbonate ion
4. several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base
Carbonic acid
Bicarbonateion
Bicarbonateion
Carbonateion
H2CO3 H2O
HCO3- H2O
HCO3-
CO32-
H3O+
H3O+
+ +
+ +
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Brønsted-Lowry Acids & BasesBrønsted-Lowry Acids & Bases• the HCO3
- ion, for example, can give up a proton to become CO3
2-, or it can accept a proton to become H2CO3
• a substance that can act as either an acid or a base is said to be amphiproticamphiprotic
• the most important amphiprotic substance in Table 8.2 is H2O; it can accept a proton to become H3O+, or lose a proton to become OH-
5. a substance cannot be a Brønsted-Lowry acid unless it contains a hydrogen atom, but not all hydrogen atoms in most compounds can be given up• acetic acid, for example, gives up only one proton
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Brønsted-Lowry Acids & BasesBrønsted-Lowry Acids & Bases6. there is an inverse relationship between the
strength of an acid and the strength of its conjugate base• the stronger the acid, the weaker its conjugate base• HI, for example, is the strongest acid in Table 8.2, and
its conjugate base, I-, is the weakest base in the table
• CH3COOH (acetic acid) is a stronger acid that H2CO3 (carbonic acid); conversely, CH3COO- (acetate ion) is a weaker base that HCO3
- (bicarbonate ion)
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Acid-Base EquilibriaAcid-Base Equilibria• we know that HCl is a strong acid, which means that the position
of this equilibrium lies very far to the right
• in contrast, acetic acid is a weak acid, and the position of its equilibrium lies very far to the left
• but what if the base is not water? How can we determine which are the major species present?
HCl + H2O H3O++Cl-
H3O+CH3COO-H2OCH3COOH + +
Acetic acid Acetate ion
CH3COOH NH3 CH3COO-NH4
++ +
Acetic acid Acetate ion
?
Ammonia Ammonium ion(conjugate baseof CH3COOH
(conjugate acidof NH3
(acid) (base)
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Acid-Base EquilibriaAcid-Base Equilibria• To predict the position of an acid-base
equilibrium such as this, we do the following• identify the two acids in the equilibrium; one on the left
and one on the right• using the information in Table 8.2, determine which is
the stronger acid and which is the weaker acid• also determine which is the stronger base and which is
the weaker base; remember that the stronger acid gives the weaker conjugate base, and the weaker acid gives the stronger conjugate base
• the stronger acid reacts with the stronger base to give the weaker acid and weaker base; equilibrium lies on the side of the weaker acid and weaker base
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Acid-Base EquilibriaAcid-Base Equilibria• identify the two acids and bases, and their relative
strengths
• the position of this equilibrium lies to the right
CH3COOH NH3 CH3COO-NH4
++ +
Acetic acid(stronger acid)
Acetate ion(weaker base)
Ammonia(stronger base)
Ammonium ion(weaker acid)
?
CH3COOH NH3 CH3COO- NH4++ +
Acetic acid(stronger acid)
Acetate ion(weaker base)
Ammonia(stronger base)
Ammonium ion(weaker acid)
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Acid-Base EquilibriaAcid-Base Equilibria• Example:Example: predict the position of equilibrium in
this acid-base reaction
H2CO3 OH- HCO3- H2O+ +
?
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Acid-Base EquilibriaAcid-Base Equilibria• Example:Example: predict the position of equilibrium in
this acid-base reaction
• Solution:Solution: the position of this equilibrium lies to the right
H2CO3 OH- HCO3- H2O+ +
?
H2CO3 OH- HCO3- H2O+ +
Strongeracid
Strongerbase
Weakerbase
Weakeracid
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Acid Ionization ConstantsAcid Ionization Constants• when a weak acid, HA, dissolves in water
• the equilibrium constant, Keq, for this ionization is
• because water is the solvent and its concentration changes very little when we add HA to it, we treat [H2O] as a constant equal to 1000 g/L or 55.5 mol/L
• we 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 ConstantsAcid Ionization Constants• Ka for acetic acid, for example is 1.8 x 10-5
• because the acid ionization constants for weak acids are numbers with negative exponents, we commonly express acid strengths as pKa where
• the value of pKa for acetic acid is 4.75
• values of Ka and pKa for some weak acids are given in Table 8.3
• as you study the entries in this table, note the inverse relationship between values of Ka and pKa
• the weaker the acid, the smaller its Ka, but the larger its pKa
pKa = -log Ka
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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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Properties of Acids & BasesProperties of Acids & Bases• Neutralization
• acids and bases react with each other in a process called neutralization; these reactions are discussed in Section 8.10
• Reaction with metals• strong acids react with certain metals (called active
metals) to produce a salt and hydrogen gas, H2
• reaction of a strong acid with a metal is a redox reaction; the metal is oxidized to a metal ion and H+ is reduced to H2
Mg(s) + 2HCl(aq) MgCl2(aq) + H2(g)Magnesium Hydrochloric
acidMagnesium
chlorideHydrogen
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Properties of Acids & BasesProperties of Acids & Bases• Reaction with metal hydroxides
• reaction of an acid with a metal hydroxide gives a salt plus water
• the reaction is more accurately written as
• omitting spectator ions gives this net ionic equation
HCl(aq)Hydrochloric
acid
+ KOH(aq)
Water
+KCl(aq)Potassiumchloride
Potassiumhydroxide
H2O
H3O+ Cl- K+ OH- 2H2O Cl- K++ + + + +
H3O+ OH- 2H2O+
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Properties of Acids & BasesProperties of Acids & Bases
• Reaction with metal oxides• strong acids react with metal oxides to give water plus
a salt
2H3O+(aq) + CaO(s) 3H2O(l) + Ca2+(aq)Calciumoxide
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Properties of Acids & BasesProperties of Acids & Bases• Reaction with carbonates and bicarbonates
• strong acids react with carbonates to give carbonic acid, which rapidly decomposes to CO2 and H2O
• strong acids react similarly with bicarbonates
2H3O+(aq) + CO32-(aq) H2CO3(aq) + 2H2O(l)
H2CO3(aq) CO2(g) + H2O(l)
2H3O+(aq) + CO32-(aq) CO2(g) + 3H2O(l)
H3O+(aq) + HCO3-(aq) H2CO3(aq) + H2O(l)
H2CO3(aq) CO2(g) + H2O(l)
H3O+(aq) + HCO3-(aq) CO2(g) + 2H2O(l)
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Properties of Acids & BasesProperties of Acids & Bases• Reaction with ammonia and amines
• any acid stronger than NH4+ is strong enough to react
with NH3 to give a salt• in the following reaction, the salt formed is ammonium
chloride, which is shown as it would be ionized in aqueous solution
• in Ch 16 we study amines, compounds in which one or more hydrogens of NH3 are replaced by carbon groups
HCl(aq) + NH3(aq) NH4+(aq) + Cl-(aq)
HCl(aq) + CH3NH2(aq) CH3NH3+(aq) +Cl-(aq)
Methylamine Methylammoniumion
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Self-Ionization of WaterSelf-Ionization of Water• pure water contains a very small number of H3O+ ions
and OH- ions formed by proton transfer from one water molecule to another
• the equilibrium expression for this reaction is
• we can treat [H2O] as a constant = 55.5 mol/L
H2O+H2O H3O++OH-
BaseAcid Conjugateacid of H2O
Conjugatebase of H2O
[H2O]2
[H3O+][HO-]Keq =
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Self-Ionization of WaterSelf-Ionization of Water• combining these constants gives a new constant called
the ion product of water, Kion product of water, Kww
• in pure water, the value of Kw is 1.0 x 10-14
• in pure water, H3O+ and OH- are formed in equal amounts (remember the balanced equation for the self-ionization of water)
• this means that in pure water
[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 WaterSelf-Ionization of Water• the equation for the ionization of water applies not only
to pure water but also to any aqueous solution
• the product of [H3O+] and [OH-] in any aqueous solution is equal to 1.0 x 10-14
• for example, if we add 0.010 mole of HCl to 1 liter of pure water, it reacts completely with water to give 0.010 mole of H3O+
• in this solution, [H3O+] is 0.010 or 1.0 x 10-2
• this means that the concentration of hydroxide ion is[OH-] = 1.0 x 10-14
1.0 x 10-2= 1.0 x 10-12
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pH and pOHpH and pOH• because hydronium ion concentrations for most
solutions are numbers with negative exponents, we commonly express these concentrations as pH, where
pH = -log [H3O+]
• we can now state the definitions of acidic and basic solutions in terms of pH
• acidic solution:acidic solution: one whose pH is less than 7.0• basic solution:basic solution: one whose pH is greater than 7.0• neutral solution:neutral solution: one whose pH is equal to 7.0
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pH and pOHpH and pOH• just as pH is a convenient way to designate the
concentration of H3O+, pOH is a convenient way to designate the concentration of OH-
pOH = -log[OH-]
• the ion product of water, Kw, is 1.0 x 10-14
• taking the logarithm of this equation gives
pH + pOH = 14• thus, if we know the pH of an aqueous solution, we can
easily calculate its pOH
Kw = [H3O+][OH-] = 1.0 x 10-14
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pH and pOHpH and pOH• 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 SolutionspH of Salt Solutions• When some salts dissolve in pure water, there is
no change in pH from that of pure water• Many salts, however, are acidic or basic and
cause a change the pH when they dissolve• We are concerned in this section with basic salts
and acidic salts
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pH of Salt SolutionspH of Salt Solutions• Basic salt: Basic salt: the salt of a strong base and a weak
acid; when dissolved in water, it raises the pH• as an example of a basic salt is sodium acetate• when this salt dissolves in water, it ionizes; Na+ ions do
not react with water, but CH3COO- ions do
• the position of equilibrium lies to the left• nevertheless, there are enough OH- ions present in 0.10
M sodium acetate to raise the pH to 8.88
OH-CH3COOHH2OCH3COO- + +Acetic acid
(stronger acid)Acetate ion
(weaker baseHydroxide ion(stronger base)
Water(weaker base)
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pH of Salt SolutionspH of Salt Solutions• Acidic salt:Acidic salt: the salt of a strong acid and a weak
base; when dissolved in water, it lowers the pH• an example of an acidic salt is ammonium chloride• chloride ion does not react with water, but the
ammonium ion does
• although the position of this equilibrium lies to the left, there are enough H3O+ ions present to make the solution acidic
NH4+ + H2O NH3 + H3O+
Ammonia(stronger base)
Ammonium ion(weaker acid)
Hydronium ion(stronger acid
Water(weaker base)
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Acid-Base TitrationsAcid-Base Titrations• Titration:Titration: an analytical procedure in which a
solute in a solution of known concentration reacts with a known stoichiometry with a substance whose concentration is to be determined• in this chapter, we are concerned with titrations in
which we use an acid (or base) of known concentration to determine the concentration of a base (or acid) in another solution
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Acid-Base TitrationsAcid-Base Titrations• An acid-base titration must meet these
requirement1. we must know the equation for the reaction so that we
can determine the stoichiometric ratio of reactants to use in our calculations
2. the reaction must be rapid and complete
3. there must be a clear-cut change in a measurable property at the end pointend point (when the reagents have combined exactly)
4. we must have accurate measurements of the amount of each reactant
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Acid-Base TitrationsAcid-Base Titrations• As an example, let us use 0.108 M H2SO4 to
determine the concentration of a NaOH solution• requirement 1:requirement 1: we know the balanced equation
• requirement 2:requirement 2: the reaction between H3O+ and OH- is rapid and complete
• requirement 3:requirement 3: we can use either an acid-base indicator or a pH meter to observe the sudden change in pH that occurs at the end point of the titration
• requirement 4:requirement 4: we use volumetric glassware
2NaOH(aq)+H2SO4(aq) Na2SO4(aq) + 2H2O(l)(concentration
known)(concentrationnot known)
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Acid-Base TitrationsAcid-Base Titrations• experimental measurements
• doing the calculations
Trial I
Volume of 0.108 M H2SO4
Volumeof NaOH
25.0 mL 33.48 mLTrial II 25.0 mL 33.46 mL
Trial III 25.0 mL 33.50 mL
average = 33.48 mL
2 mol NaOH1 mol H2SO4
= 0.161 mol NaOHL NaOH
= 0.161 M
mol NaOHL NaOH = 0.108 mol H2SO4
1 L H2SO4x x0.0250 L H2SO4
0.03348 L NaOH
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pH BufferspH Buffers• pH buffer:pH buffer: a solution that resists change in pH
when limited amounts of acid or base are added to it• a pH buffer as an acid or base “shock absorber”• a pH buffer is common called simply a buffer• the most common buffers consist of approximately
equal molar amounts of a weak acid and a salt of the weak acid; that is, approximately equal molar amounts of a weak acid and a salt of its conjugate base
• for example, if we dissolve 1.0 mole of acetic acid and 1.0 mole of its conjugate base (in the form of sodium acetate) in water, we have an acetate buffer
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pH BufferspH Buffers
• How an acetate buffer resists changes in pH• if we add a strong acid, such as HCl, added H3O+ ions
react with acetate ions and are removed from solution
• if we add a strong base, such as NaOH, added 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 BufferspH Buffers• The effect of a buffer can be quite dramatic
• consider a phosphate buffer prepared by dissolving 0.10 mole of NaH2PO4 (a weak acid) and 0.10 mole of Na2HPO4 (the salt of its conjugate base) in enough water to make 1 liter of solution
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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pH BufferspH Buffers• Buffer pHBuffer pH
• if we mix equal molar amounts of a weak acid and a salt of its conjugate base, the pH of the solution will be equal to the pKa of the weak acid
• if we want a buffer of pH 9.14, for example, we can mix equal molar amounts of boric acid (H3BO3), pKa 9.14, and sodium dihydrogen borate (NaH2BO3), the salt of its conjugate base
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pH BufferspH Buffers• Buffer capacity depends both its pH and its
concentration
pH The closer the pH of the buffer is to the pKaof the weak acid, the greater the buffer capacity
Concentration The greater the concentration of the weak acid and its conjugate base, the greater the buffer capacity
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Blood BuffersBlood Buffers• The average pH of human blood is 7.4
• any change larger than 0.10 pH unit in either direction can cause illness
• To maintain this pH, 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, HPO4
2-
• proteins:proteins: discussed in Chapter 21
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Henderson-Hasselbalch Eg.Henderson-Hasselbalch Eg.• Henderson-Hasselbalch equation:Henderson-Hasselbalch equation: a mathematical
relationship between • pH, • pKa of the weak acid, HA • concentrations HA, and its conjugate base, A-
• It is derived in the following way
• taking the logarithm of this equation gives
HA H2O A- H3O++ +
[HA]
[A-][H3O+]Ka =
[HA]log [H3O+] + log
[A-]log Ka =
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Henderson-Hasselbalch Eg.Henderson-Hasselbalch Eg.• multiplying through by -1 gives
• -log Ka is by definition pKa, and -log [H3O+] is by definition pH; making these substitutions gives
• rearranging terms gives
[HA]-log [H3O+] - log
[A-]-log Ka =
[HA]
[A-]+ logpH = pKa Henderson-Hasselbalch Equation
[HA][A-]
pKa = pH - log
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Henderson-Hasselbalch Eg.Henderson-Hasselbalch Eg.• Example:Example: what is the pH of a phosphate buffer
solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 dissolved in enough water to make 1.0 liter of solution
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8-52© 2003 Thomson Learning, Inc.All rights reserved
Henderson-Hasselbalch Eg.Henderson-Hasselbalch Eg.• Example:Example: what is the pH of a phosphate buffer
solution containing 1.0 mole of NaH2PO4 and 0.50 mole of Na2HPO4 in enough water to make one liter of solution
• SolutionSolution• the equilibrium we are dealing with and its pKa are
• substituting these values in the H-H equation gives
H2PO4- H2O HPO4
2- H3O++ + pKa = 7.21
1.0 mol/L 0.50 mol/L
= 7.21 - 0.30 = 6.91
+ logpH = 7.21 0.501.0
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End End Chapter 8Chapter 8
Chapter 8 Acids and BasesChapter 8 Acids and Bases