Principles of Reactivity: Chemistry of Acids and Bases

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Principles of Reactivity Chemistry of Acids and
Bases

• Chapter 16

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Learning Objectives

• Students understand
• The similarities and differences between
  Brønsted-Lowry acid-base and Lewis acid-base
  theories.
• The influence of structure and bonding on
  acid-base properties.

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Learning Objectives

• Students will be able to
• Use the Brønsted-Lowry and Lewis theories of
  acids and bases
• Apply the principles of chemical equilibrium to
  acids and bases in aqueous solution
• Understand how Brønsted-Lowry theory is used to
  predict the outcome of reactions of acids and
  bases
• Calculate pH
• Given Ka and Kb, calculate concentration

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16.1 The Brønsted-Lowry Concept of Acids and Bases

• Acids are proton (H) donors
• Bases are proton acceptors
• Acids capable of donating one proton and so are
  called monoprotic acids.
• Other acids, called polyprotic acids, are capable
  of donating two or more protons.
• Polyprotic bases can accept more than one proton.

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The Brønsted-Lowry Concept of Acids and Bases

• Some molecules and ions can behave either as
  Brønsted acids or bases and are called
  amphiprotic.
• The reaction of a Brønsted acid and base produces
  a new acid and base.
• Conjugate acid-base pairs differ from each other
  by the presence of one hydrogen ion

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Practice Problem

• Write a balanced equation for the reaction that
  occurs when H3PO4, phosphoric acid, donates a
  proton to water to form the dihydrogen phosphate
  ion. Is the dihydrogen phosphate ion an acid, a
  base, or amphiprotic?
• Write a balanced equation for the reaction that
  occurs when the cyanide ion, CN-, accepts a
  proton from water to form HCN. Is CN- a Bronsted
  acid or base?

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Practice Problem

• In the following reaction, identify the acid on
  the left and its conjugate base on with right.
  Similarly, identify the base on the left and its
  conjugate acid on the right.
• HNO3 NH3 ?? NH4 NO3-

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16.2 Water and the pH Scale

• Two water molecules interact with each other to
  produce a hydronium ion and a hydroxide ion.
• This property is called autoionization. Water
  will still conduct a small amount of electricity
  since it contains low concentration of hydronium
  and hydroxide.

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Water Ionization Constant

• Kw H3OOH- 1.0 x 10-14 at 25 oC
• In pure water the two ion concentrations are
  equal and the water is said to be neutral.
• Adding acid increases the hydronium
  concentration, adding base increases the
  hydroxide. Both additions would disturb the
  equilibrium.

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Practice Problem

• A solution of the strong acid HCl has HCl 4.7
  x 10-3 M. What are the concentrations of H3O
  and OH- in this solution? (remember that HCl is
  100 ionized in water)

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pH Scale

• pH -logH3O
• pOH -logOH-
• pKw 14.00 pH pOH

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Indicators

• Approximate pH of solutions can be determined
  using an acid-base indicator.
• Substances that change color in a known pH range
  are Bronsted acids or bases for which the acid
  and its conjugate base have different colors.
• Modern pH meters are preferable!

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Practice Problem

• What is the pH of a 0.0012 M NaOH solution?
• The pH of a diet soda is 4.32 at 25oC. What are
  the hydronium and hydroxide ion concentrations in
  the soda?
• If the pH of a solution of the strong base
  Sr(OH)2 is 10.46, what is the concentration of
  the Sr(OH)2 in mol/L?

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Homework

• After reading sections 16.1-16.2, you should be
  able to do the following
• P. 629 (4-14)

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16.3 Equilibrium Constants

• The lower the pH the stronger the acid.
• For a strong base/acid, the OH-/H3O
  concentration is equal to the original base/acid
  concentration.
• For a weak base/acid, the ion concentration is
  much less than the original acid concentration.

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Equilibrium Constants

• For the general acid HA, we can write
• Ka H3OA-/HA
• where the Ka is the equilibrium constant for an
  acid in water.
• K is less than 1 for weak acids. Ka increases as
  acid strength increases.

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Equilibrium Constants

• For a weak base B in water
• Kb BHOH-/B
• Ionization Constants for Some Acids and their
  Conjugate Bases are on page 595.
• The weaker the acid, the stronger its conjugate
  base.

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Practice Problem

• Which is the stronger acid, H2SO4 or H2SO3?
• Is benzoic acid, C6H5CO2H, stronger or weaker
  than acetic acid?
• Which has the stronger conjugate base, acetic
  acid or boric acid?
• Which is the stronger base, ammonia or the
  acetate ion?
• Which has the stronger conjugate acid, ammonia or
  the acetate ion?

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Ka and Kb Values for Polyprotic Acids

• Ka values for each successive ionization step
  become smaller because it is more difficult to
  remove H from a negatively charged ion
• Successive Kb values for each ionization become
  smaller as well

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pKa

• The negative log of the Ka value (pKa) becomes
  smaller as the acid strength increases.
• pKa -logKa

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Practice Problem

• What is the pKa value for benzoic acid, C6H5CO2H?
• Is chloroacetic acid (ClCH2CO2H), pKa 2.87, a
  stronger or weaker acid than benzoic acid?
• What is the pKa for the conjugate acid of
  ammonia? Is this acid stronger or weaker than
  acetic acid?

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Relating Ionization Constants

• As Ka decreases, Kb increases. The product of
  the two is equal to the autoionization constant
  for water.
• Kw KaKb

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Practice Problem

• Ka for lactic acid, CH3CHOHCO2H, is 1.4 x 10-4.
  What is Kb for the conjugate base of this acid,
  CH3CHOHCO2-? Where does this base fit in Table
  16.2?

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16.4 Acid-Base Properties of Salts

• Anions that are conjugate bases of strong acids
  are such weak bases that they have no effect on
  solution pH. (Cl-, NO3-)
• There are numerous basic anions all are the
  conjugate bases of weak acids. (C2H3O2-)
• Acidic anions arise from polyprotic acids, and
  are amphiprotic as well. (HCO3-)
• Alkali metal and alkaline earth cations have no
  measurable effect on solution pH.
• All metal cations are hydrated in water. Only
  when the cation is a 2 or 3 does the ion act as
  an acid. (Al(H2O)63 is acidic)

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Practice Problem

• For each of the following salts in water, predict
  whether the pH will be greater than, less than,
  or equal to 7.
• KBr
• NH4NO3
• AlCl3
• Na2HPO4

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16.5 Predicting the Direction of Acid-Base
Reactions

• All proton transfer reactions proceed from the
  stronger acid and base to the weaker acid and
  base. (Equilibrium favors the weaker acid and
  base.)

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Practice Problem

• Which is the stronger Brønsted acid, HCO3- or
  NH4? Which has the stronger conjugate base?
• Is a reaction between HCO3- ions and NH3 product-
  or reactant-favored?
• HCO3-(aq) NH3(aq) ?? CO32-(aq) NH4(aq)

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Practice Problem

• Write the net ionic equation for the possible
  reaction between acetic acid and sodium hydrogen
  sulfate, NaHSO4. Does the equilibrium lie to the
  left or right?

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Homework

• After reading sections 16.3-16.5, you should be
  able to do the following
• P. 629b (19-37 odd)

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16.6 Types of Acid-Base Reactions

• Strong Acid with a Strong Base
• mixing equal quantities will produce a neutral
  solution
• Weak Acid with a Strong Base
• mixing equal quantities produces a basic
  solution pH depends on Kb for anion produced

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Types of Acid-Base Rxns

• Strong Acid with a Weak Base
• mixing equal quantities produces an acidic
  solution pH depends on Ka for the cation
  produced
• Weak Acid with a Weak Base
• mixing equal quantities produces a solution in
  which the pH depends on the Ka and Kb for the
  cations and anions produced

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Practice Problem

• Equal molar quantities of HCl and NaCN are mixed.
  Is the resulting solution acidic, basic, or
  neutral?
• Equal molar quantities of acetic acid and sodium
  sulfite, Na2SO3, are mixed. Is the resulting
  solution acidic, basic, or neutral?

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16.7 Calculations

• You can use an ICE table to calculate K from
  initial concentrations and measured pH.

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Strategy

• Write the K expression, set up an ICE table, and
  convert pH to H3O.
• Enter initial concentration.
• Assign x to represent changes, based on
  reaction stoichiometry and enter into ICE.
• Recognize that H3O x
• Enter expressions for equilibrium concentrations
  into ICE
• Solve for Ka

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Practice Problem

• A solution prepared from 0.055 mol of butanoic
  acid dissolved in sufficient water to give 1.0L
  of solution has a pH of 2.72. Determine Ka for
  butanoic acid. The acid ionizes according to the
  balanced equation
• CH3CH2CH2CO2H H2O ?? H3O CH3CH2CH2CO2-

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Equilibrium Constants

• Due to the fact that very little ionization
  occurs in a weak acid, we can assume that the
  acid concentration at equilibrium is basically
  the same as the initial acid concentration.
• This is valid whenever HA0 is greater than or
  equal to 100Ka.

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pH of Weak Acid or Base

• We can use ICE tables to calculate the pH of a
  solution of a weak acid or base and using known
  equilibrium constants.

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Practice Problem

• What are the equilibrium concentrations of acetic
  acid, acetate ion, and H3O for a 0.10M solution
  of acetic acid (Ka 1.8x10-5)? What is the pH
  of the solution?

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Practice Problem

• What are the equilibrium concentrations of HF,
  fluoride ion, and H3O when a 0.0015M solution of
  HF is allowed to come to equilibrium? What is
  the pH of the solution?

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Practice Problem

• Sodium hypochlorite, NaOCl, is used as a
  disinfectant in swimming pools and water
  treatment plants. What are the concentrations of
  HOCl and OH- and the pH of a 0.015M solution of
  NaOCl?

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pH After Acid-Base Reaction

• In order to calculate pH in a resulting solution,
  you must write a balanced equation and decide
  whether the products are acid or basic.
• You must then find initial concentrations and can
  calculate pH by solving an equilibrium problem.

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Practice Problem

• Calculate the pH after mixing 15mL of 0.12M
  acetic acid with 15mL of 0.12M NaOH. What are
  the major species in solution at equilibrium
  (besides water) and what are their concentrations?

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Homework

• After reading 16.6 and 16.7, you should be able
  to do the following
• P. 629 (41-42, 47-48, 55-58, 62-63)

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16.8 Polyprotic Acids and Bases

• Acids that can donate more than one proton are
  polyprotic.
• Each step has its own Ka, which becomes
  progressively smaller due to the increased energy
  required to remove a proton.
• The pH of many polyprotic acids depends primarily
  on the hydronium ion generated in the first
  ionization step, hydronium produced in the second
  step can be neglected.

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Practice Problem

• What is the pH of a 0.10M solution of oxalic
  acid, H2C2O4? What are the concentrations of
  H3O, HC2O4-, and the oxalate ion, C2O42-?

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16.9 Molecular Structure, Bonding, and Acids-Bases

• Stronger acids have weak H-X bonds (such as HCl)
  and weaker acids have strong H-X bonds (such as
  HF).
• Oxoacids, such as HNO3, contain an atom bonded to
  one or more oxygen atoms, some with hydrogen
  atoms attached.
• Inductive effect the attraction of electrons
  from adjacent bonds by more electronegative atoms

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Bonding and Acids/Bases

• Carboxylic acids, hydrated metal cations, and
  anions all act as Brønsted bases.
• Organic amines, such as ammonia compounds, act as
  Brønsted bases.

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Practice Problem

• Which is the stronger acid, H2SeO3 or H2SeO4?
• Which is the stronger acid, Fe(H2O)62 or
  Fe(H2O)63?
• Which is the stronger acid, HOCl or HOBr?

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16.10 Lewis Acids and Bases

• This concept is based on the sharing of electron
  pairs between acid and base.
• A Lewis acid is a substance that can accept a
  pair of electrons from another atom to form a new
  bond.
• A Lewis base is a substance that can donate a
  pair of electrons to another atom to form a new
  bond.

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Lewis Acids and Bases

• The product of an acid-base reaction in the Lewis
  sense is often called an acid-base adduct. This
  type of bond is called a coordinate covalent
  bond.
• The formation of hydronium and ammonium are both
  examples.

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Cationic Lewis Acids

• All metal cations form hydrated cations in which
  the metal ion is surrounded by water molecules,
  such a Fe(H2O)62.
• These structures are called complex ions, or
  coordination complexes.
• Hydroxide ion is a Lewis base and binds readily
  to metal cations to form metal hydroxides. Metal
  hydroxides are usually amphoteric. See table on
  p. 625.

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Molecular Lewis Acids

• Acidic oxides such as carbon dioxide and sulfur
  dioxide.
• Due to oxygens high electronegativity, electrons
  are polarized away from the other element which
  can therefore react with a Lewis base such as
  hydroxide ion.

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Practice Problem

• Describe each of the following as a Lewis acid or
  a Lewis base. (Draw the Lewis dot structure. Are
  there lone pairs on the central atom? If so, it
  may be a Lewis base. Does the central atom lack
  an electron pair? If so it can behave as a Lewis
  acid.
• PH3, BCl3, H2S, HS-

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Homework

• After reading sections 16.8-16.10, you should be
  able to do the following
• P. 629d-e (67-85 odd)