Chapter 10 Acids and Bases

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Chapter 10Acids and Bases
Chemistry, The Central Science, 10th
edition Theodore L. Brown H. Eugene LeMay, Jr.
and Bruce E. Bursten

• John D. Bookstaver
• St. Charles Community College
• St. Peters, MO
• ? 2006, Prentice Hall, Inc.

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Some Definitions

• Arrhenius
• Acid Substance that, when dissolved in water,
  increases the concentration of hydrogen ions.
• Base Substance that, when dissolved in water,
  increases the concentration of hydroxide ions.

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Polyprotic Acids
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Arrhenius Acids and Bases

• In 1884, Svante Arrhenius proposed these
  definitions
• acid a substance that produces H3O ions aqueous
  solution
• 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

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Arrhenius Acids and Bases

• when HCl, for example, dissolves in water, its
  reacts with water to give hydronium ion and
  chloride ion

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Arrhenius 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

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Acid and Base Strength

• Strong acid one that reacts completely or almost
  completely with water to form H3O ions
• 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

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Arrhenius Acids and Bases

• Acids produce H in aqueous solutions water
• HCl H(aq) Cl- (aq)
• Bases produce OH- in aqueous solutions
• water
• NaOH Na(aq) OH- (aq)

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Acids

• Produce H (as H3O) ions in water
• Produce a negative ion (-) too
• Taste sour
• Corrode metals
• React with bases to form salts and water

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Bases

• Produce OH- ions in water
• Taste bitter, chalky
• Are electrolytes
• Feel soapy, slippery
• React with acids to form salts and water

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Learning Check AB1
Describe the solution in each of the following
as 1) acid 2) base or 3)neutral. A.
___soda B. ___soap C. ___coffee D. ___
wine E. ___ water F. ___ grapefruit
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Solution AB1
Describe each solution as 1) acid 2) base
or 3) neutral. A. _1_ soda B. _2_
soap C. _1_ coffee D. _1_ wine E. _3_
water F. _1_ grapefruit
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Learning Check AB2
Identify each as characteristic of an A) acid
or B) base ____ 1. Sour taste ____ 2.
Produces OH- in aqueous solutions ____ 3.
Chalky taste ____ 4. Is an electrolyte ____ 5.
Produces H in aqueous solutions
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Solution AB2
Identify each as a characteristic of an A) acid
or B) base _A_ 1. Sour taste _B_ 2. Produces
OH- in aqueous solutions _B_ 3. Chalky
taste A, B 4. Is an electrolyte _A_ 5.
Produces H in aqueous solutions
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Acid and Base Strength

• Strong acids are completely dissociated in water.
• Weak acids only dissociate partially in water.

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Acid and Base Strength

• Substances with negligible acidity do not
  dissociate in water.

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Acid and Base Strength

• 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 a substance that dissociates only
  partially in water to produce OH- ions
• ammonia, for example, is a weak base

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Some Definitions

• BrønstedLowry
• Acid Proton donor
• Base Proton acceptor

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Strong Acids

• You will recall that the seven strong acids are
  HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.
• These are, by definition, strong electrolytes and
  exist totally as ions in aqueous solution.
• For the monoprotic strong acids,
• H3O acid.

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Strong Bases

• Strong bases are the soluble hydroxides, which
  are the alkali metal and heavier alkaline earth
  metal hydroxides (Ca2, Sr2, and Ba2).
• Again, these substances dissociate completely in
  aqueous solution.

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Strengths of Acids and Bases
Strong acids completely ionize (100) in
aqueous solutions HCl H2O
H3O Cl- (100 ions) Strong bases
completely (100) dissociate into ions in aqueous
solutions. NaOH Na (aq)
OH-(aq) (100 ions)
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Strong and Weak Acids and Bases
Strong acids HCl, HNO3 , H2SO4 Most other acids
are weak. Strong bases NaOH, KOH, and
Ca(OH)2 Most other bases are weak.
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What Happens When an Acid Dissolves in Water?

• Water acts as a BrønstedLowry base and abstracts
  a proton (H) from the acid.
• As a result, the conjugate base of the acid and a
  hydronium ion are formed.

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Conjugate Acids and Bases

• From the Latin word conjugare, meaning to join
  together.
• Reactions between acids and bases always yield
  their conjugate bases and acids.

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NH3, A Bronsted-Lowry Base
When NH3 reacts with water, most of the
reactants remain dissolved as molecules, but a
few NH3 reacts with water to form NH4 and
hydroxide ion. NH3 H2O
NH4(aq) OH- (aq) acceptor donor

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Bronsted-Lowry Acids
Acids are hydrogen ion (H) donors Bases are
hydrogen ion (H) acceptors HCl
H2O H3O Cl- donor
acceptor

-
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Brønsted-Lowry Acids Bases

• Acid a proton donor
• Base a proton acceptor
• Acid-base reaction a proton transfer reaction
• Conjugate acid-base pair any pair of molecules
  or ions that can be interconverted by transfer of
  a proton

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Brønsted-Lowry Acids Bases

• Brønsted-Lowry definitions do not require water
  as a reactant

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Brønsted-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 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

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If it can be either

• ...it is amphiprotic.
• HCO3-
• HSO4-
• H2O

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pH

• pH is defined as the negative base-10 logarithm
  of the hydronium ion concentration.
• pH -log H3O

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pH

• Therefore, in pure water,
• pH -log (1.0 ? 10-7) 7.00
• An acid has a higher H3O than pure water, so
  its pH is lt7
• A base has a lower H3O than pure water, so its
  pH is gt7.

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pH Range

• 0 1 2 3 4 5 6 7 8 9 10 11
  12 13 14
• Neutral
• HgtOH- H OH-
  OH-gtH

Acidic
Basic
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Other p Scales

• The p in pH tells us to take the negative log
  of the quantity (in this case, hydrogen ions).
• Some similar examples are
• pOH -log OH-
• pKw -log Kw

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pH of Some Common Acids

• gastric juice 1.0
• lemon juice 2.3
• vinegar 2.8
• orange juice 3.5
• coffee 5.0
• milk 6.6

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pH of Some Common Bases

• blood 7.4
• tears 7.4
• seawater 8.4
• milk of magnesia 10.6
• household ammonia 11.0

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pH

• These are the pH values for several common
  substances.

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How Do We Measure pH?

• For less accurate measurements, one can use
• Litmus paper
• Red paper turns blue above pH 8
• Blue paper turns red below pH 5
• An indicator

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How Do We Measure pH?

• For more accurate measurements, one uses a pH
  meter, which measures the voltage in the solution.

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Learning Check SW1
Identify each of the following as a 1) strong
acid or base 2) weak acid 3) weak
base A. ___ HCl (aq) B. ___ NH3(aq) C. ___
NaOH (aq) D. ___ H2CO3 (aq)
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Brønsted-Lowry Acids Bases

• the HCO3- ion, for example, can give up a proton
  to become CO32-, 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 amphiprotic
• 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 Bases

• 6. 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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Properties 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

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Properties 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

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Properties of Acids Bases

• Reaction with metal oxides
• strong acids react with metal oxides to give
  water plus a salt

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Properties 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

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Neutralization Reactions
When acid and bases with equal amounts of
hydrogen ion H and hydroxide ions OH- are mixed,
the resulting solution is neutral. NaOH (aq)
HCl(aq) NaCl H2O base
acid salt water Ca(OH)2 2
HCl CaCl2 2H2O base acid salt
water
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Neutralization
H3O and OH- combine to produce water H3O
OH- ?? 2 H2O
from acid from base
neutral Net ionic equation H
OH- ?? H2O
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Acid-Base Titrations

• 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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Titration

• A known concentration of base (or acid) is
  slowly added to a solution of acid (or base).

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Titration

• A pH meter or indicators are used to determine
  when the solution has reached the equivalence
  point, at which the stoichiometric amount of acid
  equals that of base.

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Acid-Base Titrations

• As an example, let us use 0.108 M H2SO4 to
  determine the concentration of a NaOH solution
• requirement 1 we know the balanced equation
• requirement 2 the reaction between H3O and OH-
  is rapid and complete
• 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 we use volumetric glassware

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Acid-Base Titrations

• experimental measurements
• doing the calculations

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Buffers

• Solutions of a weak conjugate acid-base pair.
• They are particularly resistant to pH changes,
  even when strong acid or base is added.

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Blood 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 H2CO3 and its conjugate base,
  HCO3-
• phosphate buffer H2PO4- and its conjugate base,
  HPO42-
• proteins discussed in Chapter 21