CHM 120 CHAPTER 18 AcidBase Equilibria

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CHM 120CHAPTER 18 Acid-Base Equilibria

• Dr. Floyd Beckford
• Lyon College

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REVIEW

• Electrolyte Substances that dissolves in water
• to produce solutions that conduct electricity
• Nonelectrolytes Substances whose aqueous
• solutions do not conduct electricity
• Strong and weak relates to the degree of
• dissociation or ionization
• In aqueous solutions protons occasionally exist
• as hydrated molecules, H(H2O)n (n1)

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ACID-BASE CONCEPTS

• Acid-base reactions may be one of the most
• important class of reactions
• The most basic of the acid-base concepts is
• the Arrhenius theory
• Acids are substances that dissociate in water
• to produce hydronium ions, H3O and bases are
• substances that dissociate in water to produce
• hydroxide ions, OH-

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ACID-BASE STRENGTH

• A strong acid is almost completely dissociated
• or ionized

HCl(aq) ? H(aq) Cl-(aq)

• A weak acid is only partially dissociated or
• ionized

• So each reaction is an equilibrium with a K

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STRONG ACIDS AND BASES

• Common strong acids are either monoprotic or
• diprotic
• HA(aq) H2O(l) ? H3O(aq) A-(aq) 100
• Means HA H3O
• Similar situation for strong bases
• Typical strong bases Group 1A metal and Ca,
• Sr, and Ba hydroxides
• Strong acids HX, HNO3, HClO3, HClO4, H2SO4

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WEAK ACIDS

• Weak acids and bases are weak electrolytes
• For a weak acid, HA

• Ka the dissociation or ionization constant
• For acetic acid

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• Ka H3OCH3COO-/CH3COOH
• Acids with larger ionization constants ionize or
• dissociate to a greater extent than acids with
• smaller ionization constants
• The larger the value of Ka, the higher H3O
• and the stronger is the acid
• HIO3 Ka 1.6 x 10-1
• CH3COOH Ka 1.8 x 10-5
• HCN Ka 6.2 x 10-10

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AUTOIONIZATION OF WATER

• Reaction called autoionization of water
• K H3OOH-/H2O2
• KH2O2 H3OOH-
• Kw H3OOH-
• Kw is called ion product of water
• At 25 C, Kw 1.0 x 10-14
• Valid also for dilute aqueous solutions

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THE pH SCALE

• The hydronium ion concentration is a measure
• of a solutions acidity
• Usually small numbers
• The pH scale is used express acidity and
• basicity
• pH -logH3O so H3O 10-pH
• Note that as pH increases H3O decreases
• Value of Kw varies with temperature

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• Acidity depends on the concentration of
• H3O ions
• Acidic H3O gt OH-
• Basic H3O lt OH-
• Neutral H3O OH-
• Notice that neutral does NOT necessarily
• mean pH 7
• pH is usually quoted with the same number of
• significant digits as the concentration

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• pH can be measured using a pH meter or by an
• acid-base indicator
• An indicator is usually an organic acid that
  have
• different colors in solutions of different pH

• Indicators exist that cover the entire pH scale
• e.g. bromothymol blue 6.0 7.6
• phenolphthalein 8.0 10

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• The p-scale can also be applied to ionization
• constants
• pKa - log Ka
• The larger the value of Ka the smaller the value
  of pKa and the stronger the acid
• From the acid-dissociation constant we can
• calculate equilibrium concentrations as well as
• pH

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BRØNSTED-LOWRY THEORY

• An acid is a proton donor
• A base is a proton acceptor
• An acid-base reaction is the transfer of a
• proton from an acid to a base
• Reactions can be described in terms of what
• are called conjugate acid-base pairs
• - species which differ by a proton
• - a charge

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- B is a proton acceptor it is a base - BH is
a proton donor it is an acid So A- is the
conjugate base of HA and BH is the conjugate
acid of B

• HF/F- and H3O/H2O are conjugate pairs

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• Relationship of Ka and concentration
• In dilute solutions of weak acids, the
  assumption
• - is that all the H3O is coming from the acid
• - the concentration change of a species is small
• compared to the initial concentration of that
• species
• - at equilibrium, HA ? HAinit
• Always be cautious when using the assumptions
• given above

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• -dissociation
• Amount of acid that dissociates can be
• expressed as a percent

• In general -dissociation increases with the
• value of Ka
• For any weak acid, HA, -dissociation
• increases with dilution

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POLYPROTIC ACIDS

• Polyprotic acids Acids that provide more than
• one hydronium ion in solution
• e.g. H2SO4, H3PO4
• Polyprotic acids ionize in a stepwise manner
• Consider sulfuric acid, H2SO4

(1)
(2)

• In general, Ka1 gt Ka2

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WEAK BASES

• Weak bases undergo equilibria in water

• For a general weak base, B

• Kb is called the base-dissociation constant

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• For conjugate acid-base pairs the product of
• their equilibrium constants is the ionization
• constant for water
• Ka x Kb Kw

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FACTORS AFFECTING ACID STRENGTH

• The strength of an acid depends on the polarity
• of the H-E bond
• The polarity of the bond is related to the bond
• strength of H-E
• The weaker the H-E bond the stronger the
• acid
• Take the hydrohalic acids
• HX X F, Cl, Br, I

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• HF ltlt HCl lt HBr lt HI
• For binary acid in the same group H-A bond
• strength determines acid strength
• Same applies to other groups
• H2O lt H2S lt H2Se
• For binary acids in the same row polarity of the
  
• H-E bond determines acid strength
• - acid strength increase with the
• electronegativity of E

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• The oxoacids are also important
• e.g. HNO3, H2SO4, HClO4

• The acidity is dictated by factors that affect
• the O-H bond strength
• The electronegativity of the central element
• The oxidation number of the central element

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• Some small, highly charged metal ions are quite
• acidic
• - they are hydrated and transfers a proton
• to a water molecule

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SALTS ACID-BASE PROPERTIES

• Salts an ionic compound that is formed when
• an acid neutralizes a base
• NaOH(aq) HCl(aq) ? NaCl(aq) H2O
• Aqueous solutions of salts can be neutral,
  acidic
• or basic
• Strong acid strong base ? neutral solutions
• Strong acid weak base ? acidic solutions
• Strong base weak acid ? basic solutions

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• When a salt dissolves in water, its constituent
• ions may react with water reaction called
• hydrolysis
• NEUTRAL SOLUTIONS
• Salts of strong acids and strong bases
• e.g. NaCl
• Because the ions do not hydrolyze

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• Cl- is the conjugate base of HCl it is a weak
• base

• The same argument is made for Na
• Essentially H3O/OH- ratio does not
• change
• Group 1A metals, Ca2, Sr2, Ba2 and I-,
• Br-, Cl-, NO3-, ClO4- ? does not hydrolyze

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• Acidic Solutions
• NH3(aq) HCl(aq) ? NH4Cl(aq)
• Salts from weak bases and strong acids
• In aqueous solution NH4 undergo hydrolysis
• - the chloride ion does not

• The generation of H3O from the reaction
• makes these solutions acidic

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• Basic Solutions
• Salts from strong bases and weak acids give
• basic solutions
• This basic anions of weak acids hydrolyze to
• form hydroxide ions
• NaCH3COO ? CH3COO- H3O

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SALTS FROM WEAK ACID AND BASES

• e.g. NH4CH3COO, ammonium acetate
• Aqueous solution of the salts may be basic,
• acidic or neutral
• The pH depends on the relative Ka and Kb of
• the parent acid and base
• Consider the case when Ka Kb
• NH4CH3COO(aq) ? NH4(aq) CH3COO-(aq)

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• Both ions can undergo hydrolysis

• Hydrolysis constant for the acetate ion (Kbh) is
  
• equal to the hydrolysis constant for the
• ammonium ion (Kah)
• The same concentration of H3O as OH- is
• produced ? solution is NEUTRAL

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• If the parent Ka gt Kb solutions acidic
• e.g. NH4F

• Ka(HF) 7.2 x 10-4 gt Kb (NH3) 1.8 x 10-5
• Kb(F-) 1.4 x 10-11 lt Ka(NH4) 5.6 x 10-10
• So NH4 hydrolyzes to a greater extent than F-
• ? more H3O is produced than OH-

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• If the parent Ka lt Kb solutions basic
• e.g. NH4CN

• Ka(HCN) 4.9 x 10-10 lt Kb (NH3) 1.8 x 10-5
• Kb(CN-) 2.0 x 10-10 gt Ka(NH4) 5.6 x 10-10
• So NH4 hydrolyzes to a greater extent than F-
• ? more OH- is produced than H3O

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LEWIS THEORY

• The Lewis acid-base theory is the most
• complete that were are going to see
• In this theory
• - an acid is any species that can accept an
• electron pair
• - a base is any species that can donate an
• electron pair
• The Lewis acid typically have vacant orbitals

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• Common Lewis acids metal ions and electron
• deficient molecules such as AlCl3, BF3, etc.
• A typical Lewis base would be ammonia, NH3
• as well as the common anions
• Acid-base reactions
• BF3 NH3 ? F3B?NH3
• acid base acid-base adduct
• AlCl3 Cl- ?AlCl4-