Chapter 15 Acid-Base Equilibria

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Chapter 15Acid-Base Equilibria
CHEMISTRY
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Acids and Bases

• Arrhenius Definition
• Acids - are substances that produce hydrogen ions
  (protons or H) in solution.
• Bases - are substances that produce hydroxide
  ions in solution.
• Strong Acids and Strong Bases totally ionize
  in solution
• Weak Acids and Weak Bases partially ionize in
  solution

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Acid Dissociation in Water

• General Rxn. when Acid dissolves in H2O
• HCl H2O H3O Cl-
  acid base conj. Acid conj. base

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Properties of H2O _at_ 25 oC

• H2O (l) D H (aq) OH- (aq)
• Neutral, can act as an acid and a base
• Kw HOH- 1.0 x 10-14 only _at_ 25 oC
• Kw water dissociation constant

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Acidity vs. Basicity

• If H gtOH- , solution is acidic
• If H ltOH- , solution is basic
• The term pX -log concentration of X
• So pH -log concentration of H
• pOH -log concentration of
  OH-
• pH power of hydrogen the power of H to which
  10 is raised

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pH

• Kw HOH- 1.0 x 10-14 _at_ 25 oC
• pKw -log H -log OH- - log 1.0 x 10-14
  14
• pH -log H
• pOH -log OH-

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Properties of H2O _at_ 25 oC

• pKw - log 1.0 x 10-14 14
• pH pOH 14
• pH 7 pOH 7
• pH pOH

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Things to Remember

• pKw - log 1.0 x 10-14 14 _at_ 25 oC
• pH pOH 14
• pH lt7 acidic
• pH gt 7 basic
• pH is between 0 - 14

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Broensted-Lowrys Definition

• Acid is a proton (H) donor.
• Base is a proton (H) acceptor.
• Broensted-Lowry Definition is more general
• It even applies to bases that have no OH such as
  NH3.

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Terminologies

• H proton
• OH- hydroxide ion
• H3O hydronium ion
• Conjugate base acid minus proton
• Conjugate acid base plus proton

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More Terminologies

• Conjugate acid-base pair
• Consists of 2 substances related to each other by
  the donation and acceptance of a single proton
  (H).
• Acid Dissociation Constant (Ka)

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Acid Equilibrium

• Equilibrium Expression for the reaction
• HCl H2O H3O Cl-
  acid base conj. Acid conj. Base
• Ka H3O Cl- HCl-
  HClH2O HCl

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Equations

• pH - log H
• pOH - log OH-
• H 10 pH
• OH- 10 - pOH
• Kw 10 - pKw
• pKw pH pOH
• pKw - log Kw
• Pw HOH-

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Sample Problem

• At 40 oC, a solution has Kw 2.916 x 10-14 pH
  7.51
• Calculate the following
• A. pOH of the solution
• B. hydrogen ion concentration H
• C. hydroxide ion concentration OH-
• D. pKw
• E. Is the solution acidic basic or neutral?

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Equilibrium

• K H3O Cl- HCl-
  HClH2O HCl
• H2O removed from top and bottom since H3O is
  simply H dissolved in water.
• Remember Keq products
  reactants

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Terminologies

• H proton
• OH- hydroxide ion
• H3O hydronium ion
• Conjugate base acid minus proton
• Conjugate acid base plus proton

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Problems on Acid Dissociation

• Write the simple dissociation reaction for each
  of the following acids. Omit water.
• A.) HNO3
• B.) CH3COOH (acetic acid)
• C.) NH4
• D.) Al(H2O)33

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Acid Strength

• Strength of acid is given by the equilibrium
  position of the dissociation reaction
• HA (aq) H2O (l) H3O A-
• Strong acid totally ionized and equilibrium
  lies far to the right
• Weak acid only partially ionized and
  equilibrium lies far to the left

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Strong Acid vs. Weak Acid

• Strong Acid yields a weak conjugate base (one
  that has weak affinity for proton weaker than
  H2O)
• Weak Acid yields a strong conjugate base (one
  that has strong affinity for proton stronger
  than H2O)

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Comparison
Property Strong Acid Weak Acid
Ka value Large Ka Small Ka
Equil. Position Far to the right Far to the Left
Equil. Concn H HA0 H ltlt HA0
Conj. Base Strength vs H2O A- much weaker base than H2O A- much stronger base than H2O
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Please Note!

• Tuesdays experiment is Experiment 29 Choice I.

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Sample Problems

• Given OH- 1.0 x 10-12 M, calculate pH. Is
  the solution basic, acidic or neutral?
• Given H 4.30 x 10-6 M, calculate pH. Is the
  solution basic, acidic or neutral?

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Strong Acids and Bases

• If the molarity of the acid or base is less than
  10-6 M then the autoionization of water needs to
  be taken into account. In other words, water is
  the primary source of H and OH-, so the pH would
  be neutral.

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Acids and Bases A Brief Review

• Acids taste sour and cause dyes to change color.
• Bases taste bitter and feel soapy.
• Arrhenius acids increase H bases increase
  OH- in solution.
• Arrhenius acid base ? salt water.
• Problem the definition confines us to aqueous
  solution.

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Brønsted-Lowry Acids and Bases

• Conjugate Acid-Base Pairs
• Whatever is left of the acid after the proton is
  donated is called its conjugate base.
• Similarly, whatever remains of the base after it
  accepts a proton is called a conjugate acid.
• Consider
• After HA (acid) loses its proton it is converted
  into A- (base). Therefore HA and A- are
  conjugate acid-base pairs.
• After H2O (base) gains a proton it is converted
  into H3O (acid). Therefore, H2O and H3O are
  conjugate acid-base pairs.
• Conjugate acid-base pairs differ by only one
  proton.

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Brønsted-Lowry Acids and Bases

• Relative Strengths of Acids and Bases
• The stronger the acid, the weaker the conjugate
  base.
• H is the strongest acid that can exist in
  equilibrium in aqueous solution.
• OH- is the strongest base that can exist in
  equilibrium in aqueous solution.

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Brønsted-Lowry Acids and Bases

• Relative Strengths of Acids and Bases
• The conjugate base of a strong acid (e.g. Cl-)
  has negligible acid-base properties.
• Similarly, the conjugate acid of a strong base
  has negligible acid-base properties.

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The Autoionization of Water

• The Ion Product of Water
• In pure water the following equilibrium is
  established
• at 25 ?C
• The above is called the autoionization of water.

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The pH Scale

• In most solutions H(aq) is quite small.
• We define
• In neutral water at 25 ?C, pH pOH 7.00.
• In acidic solutions, H gt 1.0 ? 10-7, so pH lt
  7.00.
• In basic solutions, H lt 1.0 ? 10-7, so pH gt
  7.00.
• The higher the pH, the lower the pOH, the more
  basic the solution.

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The pH Scale

• Most pH and pOH values fall between 0 and 14.
• There are no theoretical limits on the values of
  pH or pOH. (e.g. pH of 2.0 M HCl is -0.301.)

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The pH Scale

• Other p Scales
• In general for a number X,
• For example, pKw -log Kw.

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The pH Scale

• Measuring pH
• Most accurate method to measure pH is to use a pH
  meter.
• However, certain dyes change color as pH changes.
  These are indicators.
• Indicators are less precise than pH meters.
• Many indicators do not have a sharp color change
  as a function of pH.
• Most indicators tend to be red in more acidic
  solutions.

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The pH Scale
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Strong Acids and Bases

• Strong Acids
• The strongest common acids are HCl, HBr, HI,
  HNO3, HClO3, HClO4, and H2SO4.
• are strong electrolytes.
• All strong acids ionize completely in solution

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Strong Acids and Bases

• Strong Acids
• The strongest common acids are HCl, HBr, HI,
  HNO3, HClO3, HClO4, and H2SO4.
• Strong acids are strong electrolytes.
• All strong acids ionize completely in solution
• HNO3(aq) H2O(l) ? H3O(aq) NO3-(aq)
• Since H and H3O are used interchangeably, we
  write
• HNO3(aq) ? H(aq) NO3-(aq)

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Strong Acids and Bases

• Strong Acids
• In solutions the strong acid is usually the only
  source of H. (If the molarity of the acid is
  less than 10-6 M then the autoionization of water
  needs to be taken into account.)
• Therefore, the pH of the solution is the initial
  molarity of the acid.
• Strong Bases
• Most ionic hydroxides are strong bases (e.g.
  NaOH, KOH, and Ca(OH)2).

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Strong Acids and Bases

• If the molarity of the acid or base is less than
  10-6 M then the autoionization of water needs to
  be taken into account. In other words, water is
  the primary source of H and OH-, so the pH would
  be neutral.

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Strong Acids and Bases

• Strong Bases
• Strong bases are strong electrolytes and
  dissociate completely in solution.
• The pOH (and hence pH) of a strong base is given
  by the initial molarity of the base. Be careful
  of stoichiometry.
• In order for a hydroxide to be a base, it must be
  soluble.
• Bases do not have to contain the OH- ion
• O2-(aq) H2O(l) ? 2OH-(aq)
• H-(aq) H2O(l) ? H2(g) OH-(aq)
• N3-(aq) H2O(l) ? NH3(aq) 3OH-(aq)

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pH of Strong Acids and Bases

• The pH (and hence pOH) of a strong acid is given
  by the initial molarity of the acid.
• The pOH (and hence pH) of a strong base is given
  by the initial molarity of the base.
• Be careful of stoichiometric ratios!

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Please Note!

• Tuesdays experiment is Experiment 29 Choice I.

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Bronsted-Lowry Acids and Bases

• Bronsted-Lowry acids compounds that donate a
  proton (H)
• Bronsted-Lowry Bases compounds that accept a
  proton (H)
• Note that Bronsted-Lowry bases need not have the
  OH group on the formula

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Weak Acids

• Weak acids are only partially ionized in
  solution.
• There is a mixture of ions and unionized acid in
  solution.
• Therefore, weak acids are in equilibrium

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Weak Acids

• Ka is the acid dissociation constant.
• Note H2O is omitted from the Ka expression.
  (H2O is a pure liquid.)
• The larger the Ka the stronger the acid (i.e. the
  more ions are present at equilibrium relative to
  unionized molecules).
• If Ka gtgt 1, then the acid is completely ionized
  and the acid is a strong acid.

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NOTE

• For Weak Acids and Weak Bases
• USE ICE to determine H, OH-, pH and pOH.!

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Sample A Problem

• A solution of 0.10 M formic acid (HCOOH) has a pH
  of 2.38 at 25 oC.
• A. Calculate Ka for formic acid at this
  temperature.
• B. What percent of this solution is ionized?

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Important Reminder

• Please Take Note
• Kw 1.0 x 10-14 is only true at 25 oC
• Therefore, pH pOH 14 is also true
  ONLY at 25 oC
• If the temperature is not 25 oC, then Kw will be
  equal to something else and pKw will not be equal
  to 14.

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Sample Problem

• At the freezing point of water which is 0 oC,
  Kw 1.2 x 10-15.
• Calculate H and OH- for a neutral solution
  at this temperature.

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Sample Problem

• The Ka of acetic acid is 1.8 x 10-5.
• A. Calculate the pH of a 0.30 M solution of
  CH3COOH.
• B. Calculate OH- and pOH.
• C. Calculate Kb.
• Calculate ionization.

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A Simple Trick

• Use of approximation eliminates the difficulty
  of quadratic equations.
• Approximation is Valid if
• X_______ x 100 lt 5
• Initial Concn.

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Relationship between Ka and Kb

• Ka x Kb 1.0 x 10 -14 only at 25 oC.

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pH of polyprotic acids

• Treat polyprotic acids as separate steps!
• 1. H2A (aq) D H (aq) HA- (aq) Ka1
• 2. HA- (aq) D H (aq) A-2 (aq) Ka2
• Initial H in Step 2 is Equil. H from Step
  1.
• Total H SUM from Steps 1 2

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HOMEWORK

• What is the pH of a 1.00 M solution of tartaric
  acid, H2C4H4O6 (aq.) at 25.0 oC?
• Answer pH 1.49

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Sample Problem

• The Ka of acetic acid is 1.8 x 10-5. Calculate
  the Kb of of CH3COOH.

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Sample Problem

• The Ka of ammonia is 1.8 x 10-5. Calculate the
  pH of a 0.15 M solution of NH3.

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Sample Problem

• Calculate the concentration of an aqueous
  solution of NaOH that has a pH of 11.50.

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HOMEWORK

• What is the pH of a 1.00 M solution of tartaric
  acid, H2C4H4O6 (aq.) at 25.0 oC?
• Answer pH 1.49

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Weak Acids

• Calculating Ka from pH
• Weak acids are simply equilibrium calculations.
• The pH gives the equilibrium concentration of H.
• Using Ka, the concentration of H (and hence the
  pH) can be calculated.
• Write the balanced chemical equation clearly
  showing the equilibrium.
• Write the equilibrium expression. Find the value
  for Ka.
• Write down the initial and equilibrium
  concentrations for everything except pure water.
  We usually assume that the change in
  concentration of H is x.

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Weak Acids

• Calculating Ka from pH
• Substitute into the equilibrium constant
  expression and solve. Remember to turn x into pH
  if necessary.
• Using Ka to Calculate pH
• Percent ionization is another method to assess
  acid strength.
• For the reaction

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Weak Acids

• Using Ka to Calculate pH
• Percent ionization relates the equilibrium H
  concentration, Heqm, to the initial HA
  concentration, HA0.
• The higher percent ionization, the stronger the
  acid.
• Percent ionization of a weak acid decreases as
  the molarity of the solution increases.
• For acetic acid, 0.05 M solution is 2.0 ionized
  whereas a 0.15 M solution is 1.0 ionized.

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Weak Acids

• Polyprotic Acids
• Polyprotic acids have more than one ionizable
  proton.
• The protons are removed in steps not all at once
• It is always easier to remove the first proton in
  a polyprotic acid than the second.
• Therefore, Ka1 gt Ka2 gt Ka3 etc.

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Weak Acids
Polyprotic Acids
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Weak Bases

• Weak bases remove protons from substances.
• There is an equilibrium between the base and the
  resulting ions
• Example
• The base dissociation constant, Kb, is defined as

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Weak Bases

• Types of Weak Bases
• Bases generally have lone pairs or negative
  charges in order to attack protons.
• Most neutral weak bases contain nitrogen.
• Amines are related to ammonia and have one or
  more N-H bonds replaced with N-C bonds (e.g.,
  CH3NH2 is methylamine).
• Anions of weak acids are also weak bases.
  Example OCl- is the conjugate base of HOCl (weak
  acid)

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Relationship Between Ka and Kb

• We need to quantify the relationship between
  strength of acid and conjugate base.
• When two reactions are added to give a third, the
  equilibrium constant for the third reaction is
  the product of the equilibrium constants for the
  first two
• Reaction 1 reaction 2 reaction 3
• has

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Relationship Between Ka and Kb

• For a conjugate acid-base pair
• Therefore, the larger the Ka, the smaller the Kb.
  That is, the stronger the acid, the weaker the
  conjugate base.
• Taking negative logarithms

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Acid-Base Properties of Salt Solutions

• Nearly all salts are strong electrolytes.
• Therefore, salts exist entirely of ions in
  solution.
• Acid-base properties of salts are a consequence
  of the reaction of their ions in solution.
• The reaction in which ions produce H or OH- in
  water is called hydrolysis.
• Anions from weak acids are basic.
• Anions from strong acids are neutral.

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Acid-Base Properties of Salt Solutions

• An Anions Ability to React with Water
• Anions, X-, can be considered conjugate bases
  from acids, HX.
• If X- comes from a strong acid, then it is
  neutral.
• If X- comes from a weak acid, then
• The pH of the solution can be calculated using
  equilibrium!

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Acid-Base Properties of Salt Solutions

• An Cations Ability to React with Water
• Polyatomic cations with ionizable protons can be
  considered conjugate acids of weak bases.
• Some metal ions react in solution to lower pH.
• Combined Effect of Cation and Anion in Solution
• An anion from a strong acid has no acid-base
  properties.
• An anion that is the conjugate base of a weak
  acid will cause an increase in pH.

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Sample Problem

• A solution of NH3 in water has a pH of 10.50.
  What is the initial molarity of the solution?

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Other Weak Bases

• Amines ex. Methylamine (CH3NH2)
• carbonate ion (CO32-)
• hypochlorite ion (ClO-1)

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Weak Bases

• Also use ICE!
• Calculation is the same as for weak acids!
• Main difference is that you get OH- and pOH
  first.

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Effects of Salts on pH

• Conjugate bases of strong acids have no effect on
  pH.
• Conjugate acids of strong bases have no effect on
  pH.
• Conjugate bases of weak acids increase pH (more
  basic).
• Ex. F- (aq) H2O(l) D HF (aq)
  OH- (aq)
• Conjugate acids of weak bases decrease pH (more
  acidic).
• NH4(aq) H2O (l) D NH3 (aq)
  H3O (aq)

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Relationship Between Ka and Kb
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Acid-Base Properties of Salt Solutions

• Combined Effect of Cation and Anion in Solution
• A cation that is the conjugate acid of a weak
  base will cause a decrease in the pH of the
  solution.
• Metal ions will cause a decrease in pH except for
  the alkali metals (Grp. I) and alkaline earth
  metals.(Grp.II)
• When a solution contains both cations and anions
  from weak acids and bases, use Ka and Kb to
  determine the final pH of the solution.

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Sample Problem

• Determine whether the resulting solution in water
  will be acidic, basic or neutral.
• A. KClO3-
• B. NaCH3COO-
• C. Na2HPO4 Ka for HPO4- 4.2 x 10-13
• D. NH4Cl-

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Sample Problem

• Predict whether the potassium salt of citric acid
  (K2HC6H5O7-) will form an acidic, basic or
  neutral solution in water.

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Weak Acids

• Polyprotic Acids
• Polyprotic acids have more than one ionizable
  proton.
• The protons are removed in steps not all at once
• It is always easier to remove the first proton in
  a polyprotic acid than the second.
• Therefore, Ka1 gt Ka2 gt Ka3 etc.

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Weak Acids
Polyprotic Acids
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Sample Problem

• The solubility of CO2 in pure water at 25 oC and
  0.1 atm is 0.0037 M. The common practice is to
  assume that all of the dissolved CO2 is in the
  form of carbonic acid (H2CO3), which is produced
  by the reaction between the CO2 and H2O.
• What is the pH of a 0.0037 M solution of H2CO3?
• Ka1 4.3 x 10-7
• Ka2 5.6 x 10 -11

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Answer

• pH 4.4
• x1 4.0 x 10-5 M
• CO3- 5.6 x 10-11 M

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Sample Problem

• Calculate the pH and concentration of oxalate ion
  (C2O42-), in a 0.020 M solution of oxalic acid
  (H2C2O4)

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Answer

• pH 1.8
• oxalate 6.4 x 10-5 M

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Acid-Base Behavior and Chemical Structure

• Factors that Affect Acid Strength
• Consider H-X. For this substance to be an acid
  we need
• H-X bond to be polar with H? and X?- (if X is a
  metal then the bond polarity is H?-, X? and the
  substance is a base),
• the H-X bond must be weak enough to be broken,
• the conjugate base, X-, must be stable.

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Acid-Base Behavior and Chemical Structure

• Binary Acids
• Acid strength increases across a period and down
  a group.
• Conversely, base strength decreases across a
  period and down a group.
• HF is a weak acid because the bond energy is
  high.
• The electronegativity difference between C and H
  is so small that the C-H bond is non-polar and
  CH4 is neither an acid nor a base.

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Acid-Base Behavior and Chemical Structure

• Binary Acids

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Acid-Base Behavior and Chemical Structure

• Oxyacids
• Oxyacids contain O-H bonds.
• All oxyacids have the general structure Y-O-H.
• The strength of the acid depends on Y and the
  atoms attached to Y.
• If Y is a metal (low electronegativity), then the
  substances are bases.
• If Y has intermediate electronegativity (e.g. I,
  EN 2.5), the electrons are between Y and O and
  the substance is a weak oxyacid.

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Acid-Base Behavior and Chemical Structure

• Oxyacids
• If Y has a large electronegativity (e.g. Cl, EN
  3.0), the electrons are located closer to Y than
  O and the O-H bond is polarized to lose H.
• The number of O atoms attached to Y increase the
  O-H bond polarity and the strength of the acid
  increases (e.g. HOCl is a weaker acid than HClO2
  which is weaker than HClO3 which is weaker than
  HClO4 which is a strong acid).

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Acid-Base Behavior and Chemical Structure
Oxyacids
92
Acid-Base Behavior and Chemical Structure

• Carboxylic Acids
• Carboxylic acids all contain the COOH group.
• All carboxylic acids are weak acids.
• When the carboxylic acid loses a proton, it
  generate the carboxylate anion, COO-.

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Lewis Acids and Bases

• Brønsted-Lowry acid is a proton donor.
• Focusing on electrons a Brønsted-Lowry acid can
  be considered as an electron pair acceptor.
• Lewis acid electron pair acceptor.
• Lewis base electron pair donor.
• Note Lewis acids and bases do not need to
  contain protons.
• Therefore, the Lewis definition is the most
  general definition of acids and bases.

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Lewis Acids and Bases

• Lewis acids generally have an incomplete octet
  (e.g. BF3).
• Transition metal ions are generally Lewis acids.
• Lewis acids must have a vacant orbital (into
  which the electron pairs can be donated).
• Compounds with p-bonds can act as Lewis acids
• H2O(l) CO2(g) ? H2CO3(aq)

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Lewis Acids and Bases

• Hydrolysis of Metal Ions
• Metal ions are positively charged and attract
  water molecules (via the lone pairs on O).
• The higher the charge, the smaller the metal ion
  and the stronger the M-OH2 interaction.
• Hydrated metal ions act as acids
• The pH increases as the size of the ion increases
  (e.g. Ca2 vs. Zn2) and as the charge increases
  (Na vs. Ca2 and Zn2 vs. Al3).

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Lewis Acids and Bases
Hydrolysis of Metal Ions
97
End of Chapter 16Acid-Base Equilibria
98
Problem 1

• Give the conjugate base of the following
  Bronsted-Lowry acids
• H2SO3
• H2AsO4-
• NH4

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Problem 2

• By what factor does H change for a pH change
  of
• A. 2.00 units
• B. 0.50 units

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Problem 3

• Calculate OH- and pH for
• A.) 1.5 x 10-3 M Sr(OH)2. Sr(OH)2 is a strong
  base.
• B.) a solution formed by adding 10 mL of 0.100 M
  HBr to 20.0 mL of 0.200 M HCl

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Problem

• Calculate the pH of a solution made by adding
  15.00 grams of NaH in enough water to make 2.5 L
  of solution

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Problem

• Write the ionization and equilibrium expressions
  for HBrO2.

103
Problem

• A particular sample of vinegar has a pH of 2.9.
  Assuming acetic acid is the only acid in the
  vinegar, find the initial concentration of acetic
  acid in the vinegar.

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Problem

• The acid dissociation constant for benzoic acid
  (HC7H5O2) is 6.3 x 10-5. Calculate the
  equilibrium concentrations of H3O, C7H5O2- and
  HC7H5O2 if the initial concentration of HC7H5O2
  is 0.050 M.

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Problem

• Calculate the pH of 0.120 M pyridine (C5H5N).
  Kb for pyridine is 1.7 x 10-9.

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Problem

• A 0.200 M solution of a weak acid, HA is 9.4
  ionized. Using this information, calculate H,
  A-, HA and Ka for HA.

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Problem

• An unknown salt is either NaF, NaCl, or NaOCl.
  When 0.05 mole of the salt is dissolved in water
  to form 0.500 L of solution, the pH of the
  solution is 8.08. What is the identity of the
  salt?

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Problem

• Write the chemical equation and the Kb expression
  for the ionization of the following bases in
  aqueous solution
• A. Dimethylamine (CH3)2NH
• B. Formate ion (HCOO-)
• C. Carbonate ion (CO32-)

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Problem

• Calculate the molar concentration of OH- ions in
  a 0.075M solution of ethylamine.
• Kb of C2H5NH2 6.4 x 10-4.
• Calculate the pH of this solution.

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Problem

• Ka for acetic acid (CH3COOH) is 1.8 x 10-5 while
  Ka for hypochlorous (HClO) ion is 3.0 x 10-8.
• A. Which is the stronger acid?
• B. Which is the stronger conjugate base?
  Acetate ion (CH3COO-) or chlorous (ClO-) ion?
• C. Calculate kb values for CH3COO- and ClO-.

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Solubility vs. Ksp

• Solubility refers to the quantity that
  dissolves to form a saturated solution. Unit is
  gm/liter or moles/liter for molar solubility.
• - solubility if affected by temperature
• Solubility product constant is the equilibrium
  constant for the equilibrium that exists between
  the ionic solute and its saturated aqueous
  solution

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Ksp

• Solubility product constant the equilibrium
  constant indicating how soluble the product is in
  water.
• Example CaF2 (s) D Ca2 (aq) 2F- (aq)
• Ksp Ca2F-2

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Problem 1

• Give the ionization equation and Ksp expression
  for the reaction
• Ag2CrO4 (s) D ? ?

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Problem 2

• The Ksp for CaF2 is 3.9 x 10-11 at 25 oC.
  Assuming that CaF2 dissociates completely upon
  dissolving and that there are no other important
  equilibria affecting its solubility
• a.
  calculate the solubility of CaF2 in moles per
  liter.
• b. calculate the solubility of CaF2 in grams per
  liter.

115
Problem 3

• The Ksp for LaF3 is 2.0 x in 10-19. What is the
  solubility of LaF3 in water in moles per liter?
• What is the solubility of LaF3 in water in grams
  per liter?

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Answer

• A. 9.28 x 10-6 M

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Factors Affecting Solubility

• Common-Ion Effect
• Concentration

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Problem 4

• Calculate the molar solubility of CaF2 at 25 oC
  in a solution that is
• A. 0.010 M in Ca(NO3)2
• B. 0.025 M in NaF

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Precipitation of Ions

• Remember Q, the reaction quotient?
• If Q gt Ksp, prepitations occurs until Q
  Ksp Q Ksp, equilibrium exists (saturated
  solution) Q lt Ksp, solid dissolves until Q Ksp.

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Problem 1

• A solution contains 1.0 x 10-12 M Ag and 2.0 x
  10-2 M Pb2. When Cl- is added, both AgCl and
  Ksp precipitate from the solution.
• What concentration of Cl- is necessary to begin
  the precipitation of each salt?
• Which salt precipitates first?

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Insoluble Chlorides

• Of the common metals ions, only Ag, Hg2 2, Pb
  2 form insoluble chlorides.

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Qualitative Analysis

• Order of separation of ions Cl-
  ? S2- ? (OH)- ? (PO4) 3- ? NH4
• 1st step add 6M HCl
• 2nd step add H2S and 0.20 M HCl
• 3rd step add (NH4)2S
• 4th step add (NH4)2HPO4 and NH3