AP Chemistry

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AP Chemistry

• Unit 15 Acids, Bases, and Acid-Base Equilibrium

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Definitions of Acids Bases

• Arrhenius Definition of Acids and Bases
• a) Acids produce H ions in solutions
• b) Bases produce (OH)- ions in solutions

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Problems with Arrhenius

• The Arrhenius definitions of acids and bases did
  not properly explain why other substances are
  acids or bases
• Example) NH3

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Problems with Arrhenius

• It was once thought that NH3 existed as NH4OH in
  solution
• However a Lewis Structure cannot be drawn for
  NH4OH
• This called for a new definition to be developed

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Bronsted-Lowry Definition

• These two scientists devise better definitions
  for acids and bases which would encompass more
  substances
• Bronsted-Lowry Acid A proton donor
• Bronsted-Lowry Acid A proton acceptor

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More About Bronsted-Lowry

• A proton is simply a H ion
• The hydrogen in substances is described as being
  ionizable
• This ionizable hydrogen is attracted to a center
  of negative charge (lone pairs of electrons)

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More About Bronsted-Lowry

• HCl H2O ? H3O Cl-

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More About Bronsted-Lowry

• 1) Show the reaction between HCl and NH3 and
  clearly identify the acid and the base
• 2) Identify the acid and the base when NH3 is
  placed in water

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Conjugate Acids and Bases

• When an acid-base reactions occurs, a conjugate
  acid and base is formed
• Conjugate acids are the original base plus a
  hydrogen ion (ie it is the acid on the product
  side of the equation)

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Conjugate Acids and Bases

• Conjugate bases are the original acid minus a
  hydrogen ion (ie it is the base on the product
  side of the equation)

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Identify the Bronsted-Lowry Acids and Bases
their conjugates

• 1) H2S NH3 ? NH4 HS-
• 2) OH- H2PO4- ? H2O HPO42-
• What is special about these reactions?

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AP Chemistry February 1, 2002

• 1) Review Homework
• 2) Notes on Strengths of Acids/Bases
• 3) Notes on Ka Kb Values
• 4) Notes on pH Scale
• 5) HW Pg. 687 3-9

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Ka Kb Values

• Ka values are written for weak acids in water
• Kb values are written for weak bases in water
• NH3 H2O ? NH4 OH-
• HC2H3O2 H2O ? H3O (C2H3O2)-

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Significance of Acid/Base Strength

• When equilibrium is established, the side in
  which there are stronger acids/bases will shift
  toward the weaker sides
• Thus the concentration of substances will favor
  the weaker members

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Significance of Conjugate Acid-Bases

• The stronger an acid, the weaker is its conjugate
  base
• The stronger a base, the weaker is its conjugate
  acid
• Acid-Base reactions favor the direction of the
  stronger member to the weaker member of each pair

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Strong Acids

• HI - hydroiodic
• HBr-hydrobromic
• HCl - hydrochloric
• HClO4 - perchloric
• H2SO4 - sulfuric
• HClO3 - chloric
• HNO3 - nitric

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Which conjugate acid-base pair is favored?

• 1) HBr, Br-
• 2) Cl- , HCl
• 3) NH4, NH3

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Factors Affecting Acid Strength

• A) Binary Acids The lower the bond dissociation
  energy, the easier the bond is broken.
• Thus the more likely that acid will donate a H
  ion. (Low BE Stronger Acid (Weaker Conjugate
  Base)
• Low BE Stronger Acid

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Factors Affecting Acid Strength

• The larger the anion, the stronger that acid is
  (the easier the bond is broken)
• Acid strength increases going across the table
  while increasing going down the table the
  greater the distribution in charge (polarity) the
  more likely the substance will lose H ion

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Factors Affecting Acid Strength

• Nonpolar covalent acids are weaker than polar
  covalent acids which are weaker than ionic acids
• Essentially, the greater the dipole in the acid,
  the more likely the acid is strong!

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Determine the Stronger Acid

• H2O H2Se
• HF HI

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AP Chemistry Daily Notes Quiz

• 1. Explain how to determine between a
  Bronsted-Lowry Acid and Base.
• 2. Which substance is favored in
• a) HCl and Cl-
• b) HC2H3O2 and C2H3O2-
• Explain why.
• 3. Why is HBr is a stronger acid than H2S?

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Oxoacids

• B) Oxoacids Contain hydrogen, oxygen, and some
  other element (nonmetal). At least one H bonded
  to an O.
• The other elements tendency to attract other
  electrons assists in determining the strength of
  the acid

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Oxoacids

• If the other element attracts electrons very
  strongly, electrons are withdrawn from Oxygen
  Hydrogen bond. This weakens the O-H bond and
  results in stronger acids.

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Oxoacids

• In oxoacids, the more electronegative the
  nonmetallic element is, the stronger the acid
  will be (dissociate completely).

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• Adding more oxygen atoms that are added (H2SO4
  and H2SO3) is the same as adding more
  electronegative elements.
• Since oxygen has a high electronegativity, this
  withdraws electrons from the O-H bond and
  results in stronger acids

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Rank the following based on strength

• HClO
• HClO2
• HClO3
• HClO4

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Question

• Why is H3PO4 a weak acid?
• Which is lower, H3PO4 or H3PO3,in terms of acidic
  strength?

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Carboxylic Acids Strengths

• C) The strength of carboxylic acids depends on
  how easily electrons can be withdrawn from O-H
  bond
• Since all carboxylic acids contain COOH, the
  rest of the molecule (R-) is important

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Carboxylic Acids Strengths

• If R is simply hydrocarbons, then electrons are
  not withdrawn and the acid is usually weak
• CH3COOH Ka 1.8 x 10-5
• CH3(CH2)3COOH Ka 1.5 x 10-5

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Carboxylic Acids Strengths

• If R contain elements of high electronegativity,
  then electrons are withdrawn and the acid is
  strong
• The closer the electronegative elements are to
  COOH, the more easily electrons are withdrawn
  (ie strong acid)

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Examples

• I-CH2CH2COOH 8.3 x 10-5
• Cl-CH2CH2HCOOH 1 x 10-4
• CH2CHClCOOH 1.4 x 10-4
• CH3CCl2COOH 8.7 x 10-3

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Base Strengths

• Any atom or groups which withdraws electrons
  makes the base weaker
• Ex) NH3 Kb 1.8 x10-5
• BrNH2 Kb 2.5 x 10-8
• Aromatic amines are weaker than nonaromiatic
  amines

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pH Scale

• The pH Scale runs 1-14 (some scales start at 0)
• pH values lt 7 are acidic
• pH values gt 7 are basic (alkaline)
• pH values 7, Neutral

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AP Chemistry Daily Notes Quiz

• 1. Explain how to determine between a
  Bronsted-Lowry Acid and Base.
• 2. Which substance is favored in
• a) HCl and Cl-
• b) HC2H3O2 and C2H3O2-
• Explain why.
• 3. Why is HBr is a stronger acid than H2S?

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Mathematical Determination of pH

• pH - log H OR
• - log H3O
• Determine the pH of a 0.500 M HCl solution

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Mathematical Determination of pOH Other
Relationships

• pOH - log OH-
• pH pOH 14
• H OH- 1 x 10-14 Kw
• Determine the pH of a 0.500 M NaOH solution

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pKa and pKb Values

• pKa - log Ka
• pKb - log Kb
• Low values for pKa pKb correspond to large
  values for Ka and Kb

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Example Problem

• 1) Calculate the pH of a 0.500 M Benzoic Acid
  (HC7H5O2) Solution. The Ka for benzoic acid is
  6.3 x 10-5.
• Set up ICE Chart

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Example Problem 2

• 2) If you have a 0.0010 M formic acid, what is
  its pH? The Ka is 1.8x 10-4
• Ans) 3.47

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Example Problem 3

• 3) Calculate the pH of a 0.010 M solution of
  pyridine. The Kb is 1.5 x 10-9
• Ans) 8.59

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Questions

• 4) Explain which solution you would expect to
  have a higher pH
• 0.5 M HBr
• 0.5 M HC2H3O2
• 5) What is the pH of a 0.345 M Ammonia solution
  (Kb 1.8 x 10-5)

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Polyprotic Acids

• Polyprotic acids are acids that have the ability
  to ionize more than once
• Ex) H3PO4, H2SO4, H2CO3
• The Ka for the second and third ionization is
  much smaller than the first

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Polyprotic Acids

• H3PO4 ? H3O (H2PO4)-
• (H2PO4)- ? H3O (HPO4)2-
• H(PO4)2- ? H3O (PO4)3-
• Ka1 7.1 x 10-3, Ka2 6.3 x 10-8, Ka3 4.3 x
  10-13
• 6) Use this information to calculate the pH of a
  0.500 M H3PO4 solution

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Polyprotic Acids

• Carbonic acid (H2CO3) Ka1 4.2 x 10-7 Ka2
  4.7 x 10-11

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Question

• 7) Explain which solution will have the lower
  pH
• a) 0.33 M NaOH
• b) 0.33 M NH3

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Polyprotic Acids

• Sulfuric acid (H2SO4) first ionization is
  strong
• Ka2 1.1 x 10-2
• In concentrated solutions, the first ionization
  produces all the H3O
• In dilute solutions (less than 0.0010 M), the
  second dissociation goes to completion

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Polyprotic Acids

• In dilute solutions (less than 0.0010 M), the
  second dissociation goes to completion assume
  complete dissociation of 2 H ions
• In intermediate concentrations (0.0010 to 0.50
  M), consider both ionizable H ions ICE table
  for 2nd Ionization

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Questions

• 8) Determine pH of
• a.) 0.95 M H2SO4
• b.) 0.35 M H2SO4
• c.) 8.5 x 10-4 M H2SO4

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Ions as Acids and Bases

• Na2CO3 is basic
• Why??
• Consider solubility of compound
• Na salts are soluble
• Carbonates are insoluble

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• Na2CO3 ? 2 Na (CO3)2-
• Carbonate sets up an equilibrium with H2O
• (CO3)2- H2O ? (HCO3)- (OH)-
• This reaction represents hydrolysis
• Hydrolysis is the breaking up of water
• (CO3)2- hydrolyzes Na does not

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Rules for Acid/Base Hydrolysis

• 1) Salts of strong acids and strong bases
  produce neutral solutions
• Ex) HCl and NaOH
• 2) Salts of strong acids and weak bases produce
  acidic solutions
• Ex) HCl and NH3

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Rules for Acid/Base Hydrolysis

• 3) Salts of weak acids and strong bases produce
  alkaline solutions
• Ex) HC2H3O2 and NaOH
• 4) Salts of weak acids and weak bases produce
  either neutral, acidic, or alkaline solutions

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Essential Rule

• Only ions that are the conjugates of weak acids
  or weak bases hydrolyze appreciably

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Practice

• Indicate whether the following solutions are
  acidic, basic, or neutral
• 1) NH4I (aq)
• 2) CH3COONH4 (aq)

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Quantitative Analysis of Hydrolysis

• Ka Kb Kw
• Kw 1 x 10-14 (Waters Values)
• 9) Calculate the pH of a 0.25 M CH3COONa(aq)
  solution.

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Quantitative Analysis of Hydrolysis

• Calculate the pH of a 0.25 M CH3COONa(aq)
  solution.
• Hint Na is soluble in water
• CH3COO- is conjugate base of weak acid
• CH3COO- H2O ? CH3COOH OH-

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Solution

• Construct ICE Chart
• CH3COO- H2O ? CH3COOH OH-
• Since this is hydrolysis (splitting of water),
  use Ka Kb Kw to determine the Keq of the
  reaction (in this case, we solve for Kb) Ka is
  in the chart, 1.8 x 10-5

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Problem

• 10) What is the pH of 0.034 M NH4Cl solution?
• Remember solubility rules, predict pH as acidic,
  basic, or neutral
• Kb of NH3 1.8 x 10-5
• pH 5.46

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Problem 11

• 11) What is the molarity of a NH4NO3 solution
  that has a pH 4.80?
• (Is it reasonable to expect that the pH is
  acidic?)
• 0.46 M

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The Common Ion Effect

• Suppose we add some of the ions on the products
  side of an equilibrium before that equilibrium is
  established
• 12) Calculate the pH of a solution that contains
  1.00 M CH3COOH and 1.00 M CH3COONa.

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The Common Ion Effect

• Calculate the pH of a solution that contains 1.00
  M CH3COOH and 1.00 M CH3COONa.

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The Common Ion Effect

• CH3COOH ? H CH3COO-
• 1.00 0 1.00
• -X X X
• 1.00-X X 1.00 X
• Notice (-X and X are insignificant since Ka is
  small)

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