Chapter 19 Acids, Bases, and Salts

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Chapter 19Acids, Bases, and Salts
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Acids and Bases

• Acids
• vinegar ? citrus fruits
• carbonated drinks ? car battery
• lemon juice ? tea

• Bases
• calcium hydroxide in mortar ? antacids
• household cleaning agents

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Properties of Acids

• Give foods a tart or sour taste
• lemon vinegar for example
• Aqueous solutions of acids are electrolytes
  (conduct electricity)
• Acids cause certain chemical indicators to change
  color.
• Acid Base Salt water

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Properties of Bases

• Bases have a bitter taste
• soap
• Bases have a slippery feel
• Aqueous solutions of bases are electrolytes
  (conduct electricity)
• Bases cause certain chemical indicators to change
  color.
• Acid Base Salt water

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Arrhenius Acids Bases
Chemists recognized the properties of acids and
bases, but were unable to propose a theory to
explain their behavior.
In 1887, Swedish chemist Svante Arrhenius
proposed a revolutionary way of defining and
thinking about acids and bases

• Acids are hydrogen-containing compounds that
  ionize to yield hydrogen ions (H) in aqueous
  solution.
• Bases are compounds that ionize to yield
  hydroxide ions (OH-) in aqueous solution

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Arrhenius Acids
Monoprotic acids acids that contain one
ionizable hydrogen HNO3 nitric acid Diprotic
acids acids that contain two ionizable
hydrogens H2SO4 sulfuric acid Triprotic
acids acids that contain three ionizable
hydrogens H3PO4 phosphoric acid
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Arrhenius Acids

• Not all compounds that contain hydrogen are acids
• Ex. CH4 methane has weak polar C H bonds
  and no ionizable hydrogens. Not an acid.
• Not all hydrogens in an acid may be released as
  hydrogen ions.
• Only hydrogens in very polar bonds are ionizable.
  In the case where hydrogen is joined to a very
  electronegative element.
• Ex. HCl hydrogen chloride very polar covalent
  molecule

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Arrhenius Acids

• When HCL dissolves in water, it releases hydrogen
  ions because the hydrogen ions are stabilized by
  solvation.

H2O H Cl (g) H (aq)
Cl- (aq) Hydrogen
Hydrogen Chloride chloride
ion ion Ionizes
to form an aqueous solution of hydronium ions and
chloride ions HCl H2O H3O
Cl-
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Arrhenius Acids

• Ethanoic acid CH3COOH is a monoprotic acid due to
  its structure
• H O
• H C C O H
• H
• The three H attached to the carbon are in weak
  polar bonds. They do not ionize.
• Only the H bonded to the highly electronegative O
  can be ionized

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Arrhenius Bases
Sodium hydroxide dissociates into sodium ions and
hydroxide ions in aqueous solution.
H2O NaOH (s) Na (aq) OH- (aq) Sodium
Sodium Hydroxide Hydroxide
Ion ion
Potassium hydroxide dissociates into sodium ions
and hydroxide ions in aqueous solution.
H2O KOH (s) K (aq) OH- (aq) Potassium
Potassium Hydroxide Hydroxide
Ion ion
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Arrhenius Bases
Group IA, the alkali metals, react with water to
produce solutions that are basic. Group IA
metals are very soluble in water and can produce
concentrated solutions. Group 2A metals are not
very soluble in water. Their solutions are always
very dilute.
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Bronsted-Lowry Acids and Bases
Arrhenius definition of acids and bases is not a
very comprehensive one. If defines acids and
bases narrowly and does not include certain
substances that have acidic or basic
properties. Na2CO3 (aq) is basic
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Bronsted-Lowry Acids and Bases
The Bronste-Lowry theory defines acid a
hydrogen-ion donor base a hydrogen-ion
acceptor
All acids and bases included in the Arrhenius
theory are also acids and bases according to the
Bronsted-Lowry theory.
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Ammonia as a BaseBronsted-Lowry Theory

• NH3 (aq) H2O (l) NH4 (aq) OH- (aq)
• ammonia is the hydrogen-ion acceptor and
  therefore a BL base
• water is the hydrogen-ion donor and therefore a
  BL acid.
• Hydrogen ions are transferred from water to
  ammonia, which causes the hydroxide-ion
  concentration to be greater than it is in pure
  water.

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Conjugate Acids and Bases

• NH3 (aq) H2O (l) NH4 (aq) OH-
  (aq)
• base acid
  conjugate acid conjugate base
• When ammonia dissolves and reacts with water,
  NH4 is the conjugate acid of the base NH3.
• OH- is the conjugate base of acid H2O

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Conjugate Acids and Bases

• HCl (g) H2O (l) Ý H3O (aq) Cl- (aq)
• acid base
  conjugate acid conjugate base
• HCl is the hydrogen-ion donor thus a BL acid.
• Water is the hydrogen-ion acceptor thus BL base

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Conjugate Acid-Base Pair

• Conjugate acid the particle formed when a base
  gains a hydrogen ion
• Conjugate base the particle that remains when
  an acid has donated a hydrogen ion..
• Conjugate acids and bases are always paired with
  a base or an acid, respectively.
• Conjugate acid-base pairs consists of two
  substances related by the loss or gain of a
  single hydrogen ion.

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Common Conjugate Acid-Base Pairs
Acid Base
HCl Cl-
H2SO4 HSO4-
H3O H2O
HSO4- SO42-
CH3COOH CH3COO-
H2CO3 HCO3-
HCO3- CO32-
NH4 NH3
H2O OH-
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Bronsted-Lowry Acids and Bases
A water molecule that gains a hydrogen ion
becomes a positively charged hydronium ion
(H3O) Amphoteric a substance that can act as
both an acid and a base Ex water H2SO4
H2O H3O HSO4- NH3 H2O
NH4 OH-
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Lewis Acids and Bases
Gilbert Lewis proposed a third Acid Base
theory Acid accepts a pair of electrons
during a reaction Base donates a pair of
electrons during a reaction Concept is more
general than either the Arrhenius theory or the
Bronsted-Lowry theory.
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Lewis Acids and Bases
Lewis Acid a substance that can accept a pair
of electrons to form a covalent bond. Lewis Base
a substance that can donate a pair of electrons
to form a covalent bond.
.. H -O H
O
.. H
H Lewis Lewis Acid
Base
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Acid Base Definitions
Type Acid Base
Arrhenius H producer OH- producer
Bronsted Lowry H H acceptor
Lewis Electron-pair acceptor Electron-pair donor
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End of Section 19.1
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Hydrogen Ions From Water
Water molecules are highly polar and are in
continuous motion. Occasionally, the collisions
between water molecules are energetic enough to
transfer a hydrogen ion from one water molecule
to another. Self ionization of water the
reaction in which water molecules produce ions

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Hydrogen Ions From Water
A water molecule that loses a hydrogen ion
becomes a negatively charged hydroxide ion A
water molecule that gains a hydrogen ion becomes
a positively charged hydronium ion
H2O (l) OH- (aq) H (aq)

Hydroxide ion Hydroxide ion

Self ionization of water the reaction in which
water molecules produce ions
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Self Ionization of Water

• Hydrogen ions in aqueous solution have several
  names.
• Some chemists call them protons
• Some chemists call them hydrogen ions or
  hydronium ions.
• For our purposes, either H or H3O will
  represent hydrogen ions in aqueous solution.
• H2O H2O H3O
  OH-

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Self Ionization of Water

• The self-ionization of water occurs to a very
  small extent.
• In pure water at 25C, the equilibrium
  concentration of hydrogen ions and hydroxide ions
  are each only 1 x 10-7.
• In other words the concentration of OH- and H
  are equal in pure water

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Neutral Solutions
Any aqueous solution in which H and OH- are
equal is a neutral solution.
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Ion Product Constant for Water
When H increases OH- decreases When H
decreases OH- increases
LeChateliers principle when a stress is
applied to a system in dynamic equilibrium, the
system changes in a way that relieves the stress
If additional ions (either H or OH-) are added
to a solution, the equilibrium shifts. The
concentration of the other type of ion decreases.
More water molecules are formed in the
process. H (aq) OH- (aq)
H2O (l)
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Ion Product Constant for Water
For aqueous solutions, the product of the
hydrogen ion concentration and the hydroxide ion
concentration equals 1.0 x 10-14 H x OH-
1.0 x 10-14 This equation is true for all
dilute aqueous solutions at 25C. Ion-Product
Constant for Water (Kw) the product of the
concentrations of the hydrogen ions and hydroxide
ions in water Kw H x OH- 1.0 x
10-14
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Ion Product Constant for Water
Not all solutions are neutral When some
substances dissolve in water, they release
hydrogen ions. When hydrogen chloride dissolves
in water, it forms hydrochloric acid.
H2O HCl (g) H
(aq) Cl- (aq)
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Ion Product Constant for Water
In the previous HCl solution, the hydrogen-ion
concentration is greater than the hydroxide-ion
concentration. Acidic Solution one in which
H is greater than OH-. The H of an
acidic solution is greater than 1 x 10-7
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Ion Product Constant for Water
When sodium hydroxide dissolves in water, it
forms hydroxide ions in solution.
H20 NaOH(s) Na(aq)
OH-(aq) In the above solution, the hydrogen-ion
concentration is less than the hydroxide-ion
concentration. Basic Solution one in which
H is less than OH- The H of a basic
solution is less than 1 x 10-7 Basic solutions
are also known as alkaline solutions.
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The pH Concept
The pH scale was proposed by Danish Scientist
Soren Sorensen in 1909. The pH scale is used to
express H 1 2 3 4 5 6 7 8 9
10 11 12 13 14 Strongly
Neutral

Strongly Acidic

Basic
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Calculating pH
The pH of a solution is the negative logarithm of
the hydrogen-ion concentration. pH -logH
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Calculating pH
In neutral solution, the H 1 x 10-7M. The pH
is 7 pH -logH pH -log(1 x 10-7) pH
-(log 1 log 10-7) pH -(0.0 -7.0) pH 7.0
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Classifying Solutions
A solution in which H is greater than 1 x 10-7
has a pH less than 7.0 and is acidic. A solution
in which H is less than 1 x 10-7 has a pH
greater than 7.0 and is basic. The pH of pure
water or a neutral aqueous solution is
7.0 Acidic solution pH lt 7.0 H gt 1 x
10-7M Neutral solution pH 7.0 H equals 1 x
10-7M Basic solution pH gt 7.0 H lt 1 x 10-7
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Calculating pH
pH can be read from the value of H if it is
written in scientific notation and has a
coefficient of 1. Then the pH of the solution
equals the exponent, with the sign changed from
minus to plus H 1 x 10-2 has a pH of
2.0 H 1 x 10-13 has a pH of 13.0
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Calculating pH
If the pH is an integer, it is also possible to
directly write the value of H. pH 9.0 then
H of 1 x 10-9M pH 4 then H 1 x 10-4M
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Calculating pOH
The pOH of a solution equals the negative
logarithm of the hydroxide-ion concentration
pOH -log OH- A neutral solution has a pOH
of 7
Acidic solution pOH gt 7.0 OH- lt 1 x
10-7M Neutral solution pOH 7.0 OH- equals 1
x 10-7M Basic solution pOH lt 7.0 OH- gt 1 x
10-7
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pH and pOH Relationship
pOH pH 14 pH 14 pOH pOH 14 - pH
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pH Significant Figures
For pH calculation, you should express the
hydrogen-ion concentration in scientific
notation H 0.0010M should be written 1.0 x
10-3 0.0010M has two sig figs Write pH 3.00
with 2 zeros to the right of the decimal place
representing the 2 sig figs
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Problem Example
Colas are slightly acidic. If the H in a
solution is 1.0 X 10-5 M , is the solution
acidic, basic or neutral. What is the OH- of
this solution? H 1.0 X 10-5 M which is
greater than 1.0 X 10-7 M so solution is
acidic Kw OH- x H 1.0 X 10-14 OH-
1.0 X 10-14 / H OH- 1.0 X 10-14 / 1.0 X
10-5 OH- 1.0 X 10-9
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Problem Example
What is the pH of a solution with a hydrogen-ion
concentration of 4.2 x 10-10 M? pH -log
H pH -log (4.2 x 10-10) pH -(9.3765) pH
9.38
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Problem Example
pH of an unknown solution is 6.35. What is its
hydrogen-ion concentration? pH -log H 6.35
-log H -6.35 log H Using calculator
find the antilog of -6.35 4.5 x 10-7 M H
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Problem Example
What is the pH of a solution if the OH-
4.0X10-11M? Kw H x OH- 1 x 10-14 H
1 x 10-14 / OH- H 1 x 10-14 / 4.0 x
10-11 H 0.25 x 10-3 M H 2.5 x 10-4 M
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Problem Example (cont)
What is the pH of a solution if the OH-
4.0X10-11M? pH -log H pH -log (2.5 x
10-4) pH - (-3.60205) pH 3.60
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Acid-Base Indicators
Indicator - (HIn) is an acid or a base tht
undergoes dissociation in a know pH range An
indicator is a valuable tool for measuring pH
because its acid form and base form have
different color in solution.
OH- HIn (aq) H (aq) In- (aq)
acid
H base
form
form The acid form dominates the
dissociation equilibrium at low pH (high H),
and the base form dominates the equilibrium at
high pH (high OH-)
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Acid-Base Indicators
For each indicator, the change from dominating
acid from to dominating base form occurs in a
narrow range of approximately two pH units.
Within this range, the color of the solution is
a mixture of the colors of the acid and the base
forms. Knowing the pH range over which this
color change occurs, can give you a rough
estimate of the pH of the solution.

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Acid-Base Indicators

• Many different indicators are needed to span the
  entire pH spectrum.
• Indicator characteristics that limit their
  usefulness.
• Listed pH values of indicators are usually given
  for 25ºC. At other temperatures, an indicator
  may change color at a different pH.
• If the solution being tested is not colorless,
  the color of the indicator may be distorted.
• Dissolved salts in a solution may also affect the
  indicators dissociation.
• Using indicator strips can help overcome these
  problems.

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pH Meters

• A pH meter makes rapid, accurate pH measurements.
  
• often easier to use than liquid indicators or
  indicator strips.
• Measurements of pH obtained with a pH meter are
  typically accurate to within 0.01 pH unit of the
  true pH.
• Color and cloudiness of the unknown solution do
  not affect the accuracy of the pH value
• If the solution being tested is not colorless,
  the color of the indicator may be distorted.
  

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End of section 19.2
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Strong Acids

• Acids are classified as strong or weak depending
  on the degree to which they ionize in water.
• In general, strong acids are completely ionized
  in aqueous solution.
• HNO3 - nitric acid HCl - hydrochloric
  acid H2SO4 - sulfuric acid HClO4 -
  perchloric acid HBr - hydrobromic acid HI
  - hydroiodic acid
• HCl(g) H2O(l) H3O(aq) Cl-(aq)

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Weak Acids

• Weak acids ionize only slightly in aqueous
  solution.
• Some Weak Acids
• Acetic Acid H3COOH
• Boric Acid H3BO3 (all three are weak)
• Phosphoric Acid H3PO4 (all three are weak)
• Sulfuric Acid HSO4- (first ionization is
  strong)
• CH3COOH(aq) H2O(l) H3O(aq)
  CH3COO-(aq)
• ethanoic acid water
  hydronium ethanoate
• ion ion

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Acid Strength
A strong acid completely dissociates in water
(H3O is high). A weak acid remains largely
undissociated. (H3O is low).
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Equilibrium Constant (Keq)
Write the equilibrium-constant expression from
the balanced chemical equation. CH3COOH(aq)
H2O(l) H3O(aq) CH3COO-(aq) Keq
H3O x CH3COO- H3COOH x
H2O H2O constant in dilute solutions
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Acid Dissociation Constant (Ka)
Ka Ratio of the concentration of the
dissociated form of an acid to the concentration
of the undissociated form. H3COOH(aq)
H2O(l) H3O(aq) CH3COO-(aq) Acid
Dissociation Constant Ka H3O x CH3COO-
CH3COOH
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Acid Dissociation Constant (Ka)
Acid dissociation constant reflects the fraction
of an acid in the ionized form. (Ka sometimes
called ionization constant) If the value of
the Ka is small, then the degree of dissociation
or ionization of the acid in the solution is
small. Weak acids small Ka values Stronger the
acid larger the Ka
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Acid Dissociation Constant (Ka)
Nitrous acid (HNO2) has a Ka of 4.4 x
10-4 Acetic acid (CH3COOH) has a Ka of 1.8 x
10-5 Nitrous acid is more ionized in solution
and a stronger acid
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Acids

• Strong Acids
• Have high H3O
• Large dissociation constant
• Weak Acids
• Have low H3O
• Small dissociation constant

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Acids
Diprotic and triprotic acids lose their hydrogens
one at a time. Each ionization reaction has a
separate dissociation constant. H3PO4 3
separate dissociation constants.
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Base Dissociation Constant (Kb)

• Strong bases dissociate completely into metal
  ions and hydroxide ions in aqueous solution.
• Some strong bases are not very soluble in water
  (calcium hydroxide and magnesium hydroxide)
• Small amounts that do not dissolve dissociate
  completely
• Weak bases react with water to form the hydroxide
  ion and the conjugate acid of the base.
• NH3(aq) H2O(l) NH4(aq) OH-(aq)
• Ammonia Water
  Ammonium Ion Hydroxide ion

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Base Dissociation Constant (Kb)
NH3(aq) H2O(l) NH4(aq)
OH-(aq) Ammonia Water
Ammonium Ion Hydroxide ion Only
about 1 of ammonia is present as NH4
Equilibrium Constant Keq NH4 x
OH- NH3 x H2O H2O
constant in dilute solutions Base Dissociation
Constant Kb NH4 x OH- NH3
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Concentration and Strength

• The words concentrated and dilute indicate how
  much of an acid or base is dissolved in solution.
• Number of moles of the acid or base in a given
  volume
• The words strong and weak refer to the extent of
  ionization or dissociation of an acid or base
• How many of the particles ionize or dissociate
  into ions
• A sample of HCl added to a large volume of water
  becomes more dilute, but it is still a strong
  acid.
• Vinegar is a dilute solution of a weak acid.

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End of section 19.3
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Acid-Base Reactions

• If you mix a solution of a strong acid containing
  hydronium ions with a solution of a strong base
  that has an equal number of hydroxide ions, a
  neutral solution results.
• Final solution has properties that are
  characteristic of neither an acidic nor a basic
  solution.
• HCl(aq) NaOH(aq) NaCl(aq)
  H2O(l)
• H2SO4(aq) 2KOH(aq) K2SO4(aq)
  H2O(l)

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Neutralization Reactions
Reactions of weak acids and weak bases do not
usually produce a neutral solution. In general,
reactions with which an acid and a base react in
an aqueous solution to produce a salt and water
are called neutralization reactions.
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Making Salts
Prepare potassium chloride by mixing equal molar
quantities of hydrochloric acid and potassium
hydroxide. HCl KOH KCl
H20 Heating the solution to evaporate the water
will leave the salt potassium chloride. In
general, the reaction of an acid with a base
produced water and salt
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Titration
The number of moles of hydrogen ions provided by
the acid are equivalent to the number of
hydroxide ions provided by the base. HCl(aq)
NaOH(aq) NaCl (aq) H20 (l)
1 mole 1 mole 1
mole 1 mole H2SO4(aq)
2NaOH(aq) Na2SO4(aq) 2H20 (l)
1 mole 2 mole
1 mole 2 mole When and acid
base are mixed, the Equivalence point is when
the number of moles of hydrogen ions equals the
number of moles of hydroxide ions.
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Sample Problem

• How many moles of sulfuric acid are required to
  neutralize 0.50 mol of sodium hydroxide?
• H2SO4(aq) 2NaOH(aq) Na2SO4(aq)
  2H20 (l)
• Mole ratio of H2SO4 to NaOH is 12
• 0.50 mol NaOH 1 mol H2SO4 0.25 mol
  H2SO4
• 2
  mol NaOH

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Practice Problem
How many moles of potassium hydroxide are needed
to completely neutralize 1.56 mol of phosphoric
acid? H3PO4(aq) 3KOH(aq)
K3PO4(aq) 3H2O(l) 1.56 mol H3PO4
3 mol KOH 4.68 mol H3PO4
1 mol H3PO4
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Titration
You can determine the concentration of acid or
base in a solution by performing a neutralization
reaction. You must use an appropriate
acid-base indicator to show when neutralization
has occurred. In the lab, typically
phenolphthalein for acid base neutralization
reactions. Solutions that contain
phenolphthalein turn from colorless to deep pink
as the pH of the solution changes from acidic to
basic.
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Titration

1. Measured volume of an acid solution of unknown
   concentration is added to a flask

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Titration

1. Several drops of the indicator are added to the
   solution while the flask is swirled

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Titration

1. Measured volumes of the base of known
   concentration are mixed into the acid until the
   indicator just barely changes color.

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Titration
Titration the process of adding a known amount
of solution of known concentration to determine
the concentration of another solution. Standard
solution the solution of known
concentration End point the point at which the
indicator changes color You can also use
titration to find the concentration of a base
using a standard acid.
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Titration
Titration the process of adding a known amount
of solution of known concentration to determine
the concentration of another solution. Standard
solution the solution of known
concentration End point the point at which the
indicator changes color. The point of
neutralization Equivalence point the point in
a titration where the number of moles of hydrogen
ions number of moles of hydroxide ions..
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Sample Problem
A 25ml solution of H2SO4 is completely
neutralized by 18ml of 1.0M NaOH. What is the
concentration of the H2SO4 solution?
H2SO4(aq) 2NaOH(aq) Na2SO4(aq)
2H20 (l) 0.018 L NaOH 1.0 mol NaOH 1
mol H2SO4
1L NaOH 2 mol
NaOH 0.025L 0.36M H2SO4
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Practice Problem
How many milliliters of 0.45M HCl will neutralize
25.0ml of 1.00M KOH? HCl KOH
H2O KCl 0.025 L KOH 1.0 mol KOH
1 mol HCl
1L KOH 1 mol KOH 1 L HCl
1000 ml HCl 56 ml HCl
0.45 mol HCl 1 L HCl
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Practice Problem
What is the molarity of H3PO4 if 15.0 ml is
completely neutralized by 38.5 ml of 0.150 M?
H3PO4 3NaOH 3H2O Na3PO4
0.0385 L NaOH 0.150 mol NaOH 1 mol H3PO4
1L
NaOH 3 mol NaOH
0.129M H3PO4 0.015L H3PO4
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End of section 19.4
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Salt Hydrolysis
A salt consists of an anion from an acid and a
cation from a base. The salt forms as a result
of a neutralization reaction Although solutions
of many salts are neutral, some are acidic and
others are basic. ..
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Salt Hydrolysis
Salt Hydrolysis the cations or anions of a
dissociated salt remove hydrogen ions from or
donate hydrogen ion to water. Hydrolyzing
salts are usually derived from a strong acid and
weak base or from a weak acid and a strong base.
In general, salts that produce acidic solutions
contain positive ions that release protons to
water. Salts that produce basic solutions
contain negative ions that attract protons from
water.
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Salt Hydrolysis
CH3COONa (aq) CH3COO- (aq)
Na (aq) Sodium ethanoate
ethanoate ion sodium
ion CH3COONa is the salt from a weak acid
CH3COOH and a strong base NaOH In solution the
salt is completely ionized.
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Salt Hydrolysis
Salt Hydrolysis the cations or anions of a
dissociated salt remove hydrogen ions from or
donate hydrogen ion to water. CH3COO-(aq)
H2O(l) CH3COOH (aq) OH- (aq)
BL base BL acid
makes
hydrogen-ion hydrogen-ion
solution
acceptor donor

basic This process is called hydrolysis because
it splits a hydrogen ion off a water molecule.
Resulting solution contains a hydroxide-ion
concentration greater than the hydrogen-ion
concentration. Thus the solution is basic
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Salt Hydrolysis
NH4Cl (aq) NH4 (aq) Cl- (aq)
Ammonium
Ammonium ion Chloride ion
chloride NH4Cl is the salt from a strong acid
(hydrochloric acid, HCl) and a weak base
(ammonia, NH3) In solution the salt is
completely ionized.
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Salt Hydrolysis
NH4(aq) H2O(l) NH3(aq)
H3O(aq) BL acid BL base
makes
hydrogen-ion hydrogen-ion
solution donor
acceptor
acidic This process is also called
hydrolysis because it splits a hydrogen ion off a
water molecule. Resulting solution contains a
hydrogen-ion concentration greater than the
hydroxide-ion concentration. Thus the solution is
acidic
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Salt Hydrolysis
Equivalence Point Equivalence Point Equivalence Point
Strong Acid Strong Base pH 7 neutral
Weak Acid Strong Base pH gt 7 basic
Strong Acid Weak Base pH lt 7 acidic
Equivalence point the point in a titration
where the number of moles of hydrogen ions
number of moles of hydroxide ions
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Buffers
Buffer a solution in which the pH remains
relatively constant when small amounts of acid or
base are added. A buffer is a solution of a
weak acid and one of its salts, or a solution of
a weak base and one of its salts. A buffer
solution is better able to resist drastic changes
in pH than is pure water.
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Buffers
A solution of ethanoic acid (CH3COOH) and
sodium ethanoate (CH3COONa) is an example of a
typical buffer. CH3COOH and CH3COO- (source
is the completely ionized CH3COONa) act as
reservoirs of neutralizing power.
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Buffers
CH3COO-(aq) H(aq) CH3COOH
(aq) ethanoate ion
hydrogen ion ethanoic acid When
an acid is added to the solution, the ethanoate
ions act as a hydrogen-ion sponge. CH3COOH
(aq) OH-(aq) CH3COO-(aq) H2O (l)
Ethanoic acid hydroxide ion
ethanoate ion water When a base
is added to the solution, the ethanoic acid and
the hydroxide ions react to produce water and
the ethanoate ion.
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Buffers
The ethanoate ion is not strong enough base to
accept hydrogen ions from water extensively.
The buffer solution cannot control the pH when
too much acid is added, because no more ethanoate
ions are present to accept hydrogen ions.
Buffer also become ineffective when too much
base is added. No more ethanoic acid molecules
are present to donate hydrogen ions.
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Buffers
When too much acid or base is added, the buffer
capacity is exceeded. Buffer capacity the
amount of acid or base than can be added to a
buffer solution before a significant change in pH
occurs.
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Buffers
When a base is added to a buffered solution,
the acidic form removes hydroxide ions from the
solution. When an acid is added to a buffered
solution, the basic form removes hydrogen ions
from the solution.
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Buffers Your Blood
Your body function properly only when the pH of
your blood lies between 7.35 and 7.45. Your
blood contains buffers (hydrogen carbonate ions
and carbonic acid) HCO3- (aq) H
(aq) H2CO3 (aq) hydrogen
hydrogen ions
carbonic acid carbonate ion As long as
there are hydrogen carbonate ions available, the
excess hydrogen ions are removed, and the pH of
the blood changes very little.
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End of Section Chapter 19