Acids and Bases

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Acids and Bases

• Chapter 19 Section 1

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Properties Acids Bases

• Taste Sour
• Affect Indicators
• Turn Litmus Red
• Neutralize Bases
• pH below 7
• Strong or weak electrolyte in aqueous solution

• Taste Bitter
• Affect Indicators
• Turn Litmus Blue
• Neutralize Acids
• pH above 7
• Strong or weak electrolyte in aqueous solution
• Feel Slippery

3
Examples Acids Bases

• Citrus Fruits
• Apples
• Malic Acid
• Lime Away-H3PO4
• Phosphoric Acid
• Carbonated Soft Drinks
• Carbonic Acid-H2CO3
• Milk
• Lactic Acid

• Windex
• Ammonia
• Cleaning
• Sodium Hydroxide
• Lye soaps

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Arrhenius-Acids and Bases

• Acids are hydrogen containing compounds that
  ionize to yield hydrogen ions (H) in aqueous
  solutions.
• This results in hydronium ions in solution.
• H2O HCl ? H3O Cl-
• Bases are compounds that ionize to yield
  hydroxide ions (OH-) in aqueous solutions.
• Arrhenius focuses on end result the product
  formed

Svante Arrhenius 1887 Swedish chemist
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Arrhenius Acids

• Monoprotic Acids-contain one ionizable hydrogen
• nitric acid HNO3
• Diprotic Acids-contain two ionizable hydrogens
• sulfuric acid H2SO4
• Triprotic Acids-contain three ionizable hydrogens
• phosphoric acid H3PO4

Acetic acid
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Arrhenius Bases

• Bases are compounds that ionize to yield
  hydroxide ions (OH-) in aqueous solutions.
• NaOH(s) ? Na(aq) OH- (aq)
• KOH(s) ? K(aq) OH- (aq)

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Bronsted-Lowry Acids and Bases

• The Arrhenius definition was too narrow.
• It did not include some substances that act like
  acids or bases.
• Examples sodium carbonate (Na2CO3) and ammonia
  (NH3)

• 1923-Danish chemist Johannes Bronsted and English
  chemist Thomas Lowry

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Bronsted-Lowry Acids and Bases

• Acids Hydrogen-ion donor
• Bases Hydrogen-ion acceptor
• NH3 (aq) H2O(l) ? NH4(aq) OH- (aq)
• Ammonia is a hydrogen-ion acceptor therefore it
  is a Bronsted-Lowry base
• Water is a hydrogen-ion donor therefore it acts
  as a Bronsted-Lowry acid

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Conjugate Acids and Bases

• 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
• NH3 (aq) H2O(l) ? NH4(aq) OH- (aq)
• Base Acid Conjugate
  Conjugate
• acid
  base
• Conjugate acid-base pair two substances related
  by the loss or gain of a single hydrogen ion.

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Hydronium ion (H3O)
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Amphoteric

• Amphoteric A substance that can act as both an
  acid and a base.
• H2O is an example
• See book

12
Lewis Acids and Bases

• Lewis Acid electron pair acceptor
• Lewis Base electron pair donor
• Gilbert Lewis (1875-1946)

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The 3 Theories Compared
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Acids and Bases

• Chapter 19 Section 2

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Self-Ionization

• Self-ionization of water the reaction in which
  water molecules produce ions.
• H2O(l) ? H(aq) OH-(aq)
• In water or aqueous solution, hydrogen ions (H)
  are always joined to water molecules as hydronium
  ions (H3O).

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Self-Ionization
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Self-Ionization

• The self-ionization of water occurs to a very
  small extent.
• At 25C the equilibrium concentration of hydrogen
  ions (H) and hydroxide (OH-) ions are each
  only 1 x 10-7M.
• Neutral Solution any aqueous solution in which
  the H and OH- are equal.

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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.
• Kw H x OH- 1.0 x 10-14
• Ion-product constant for water (Kw) The product
  of the concentrations of hydrogen ions and
  hydroxide ions in water.

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Acidic and Basic Solutions

• H2O HCl ? H Cl-
• Acidic Solution is one in which the H is
  greater than the OH-.
• The H is greater then 1 x 10-7 M.
• NaOH(s) ? Na(aq) OH- (aq)
• Basic Solution (alkaline solution) is one in
  which the H is less than the OH-.
• The H is less then 1 x 10-7 M.

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pH Scale

• pH scale ranges from 0 to 14
• Neutral solutions have a pH of 7

0---------------------7--------------------14
Acidic ? Basic
(Neutral)

• Acidic pH below 7
• Basic pH above 7

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Calculations

• pH -log H
• pOH -log OH-
• pH pOH 14

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Acid-Base Indicators

• An indicator is a valuable tool for measuring pH
  because its acid form and base form have
  different colors in solution.

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Chapter 19

• Section 2 Section 3

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pH Scale
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pH Scale

• 0---------------------7--------------------14
• Acidic ? Basic
• (Neutral)

• Acidic pH below 7
• Basic pH above 7

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Calculations

• Kw HOH- 1.0 x 10-14
• pH -log H H10-pH
• pOH -log OH- OH-10-pOH
• pH pOH 14

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Measuring pH

• Most pH indicators change color within a narrow
  range of approximately two pH units.

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Strengths of Acids and Bases

• Acids are classified as strong or weak depending
  on the degree to which they ionize in water.
• Strong Acids Completely ionize in aqueous
  solution.
• Examples HCl and H2SO4
• H2O HCl ? H3O Cl- (100 ionized)
• Weak Acids Ionize only slightly in aqueous
  solution.
• Example Ethanoic Acid (Acetic Acid)
• CH3COOH(aq) H2O(l) ? H3O(aq) CH3COO-(aq)
  (about 1ionized)

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Acid Dissociation Constant

• Acid Dissociation Constant (Ka) the ratio of the
  concentration of the dissociated (or ionized)
  form of an acid to the concentration of the
  undissociated (nonionized) form.
• CH3COOH(aq) H2O(l) ? H3O(aq) CH3COO-(aq)
• Ka H3O x CH3COO-
• CH3COOH
• Ka reflects the fraction of an acid in the
  ionized form.

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Acid Dissociation Constant

• Weak acids have small Ka values.
• The stronger the acid is, the larger its Ka
  value.

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Base Dissociation Constant

• Strong Base dissociates completely into metal
  ions and hydroxide ions in aqueous solution.
• Ca(OH)2
• Weak Base react with water to form the hydroxide
  ion and the conjugate acid of the base
• NH3 H2O ? NH4 OH-
• Only about 1 the ammonia exists as NH4

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Base Dissociation Constant

• Base Dissociation Constant (Kb) the ratio of the
  concentration of the conjugate acid time the
  concentration of the hydroxide ion to the
  conjugate base.
• NH3 H2O ? NH4 OH-
• Kb NH4 x OH-
• NH3

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Chapter 19Acids and Bases

• Sections 3 4

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Dissociation Constants

• Acids
• Weak acids have small Ka values.
• The stronger the acid is, the larger its Ka value.

• Bases
• Weak bases have small Kb values.
• The stronger the base is, the larger its Kb value.

See Book Page 605
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Whats the difference?

• Concentrated or Dilute indicate how much acid
  or base is dissolved in solution.
• Moles per given volume (Molarity)
• Strong or Weak refer to the extent of
  ionization (dissociation) of an acid or base.
• How many particles dissociate into ions

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Calculating Dissociation Constants

• To find the Ka of a weak acid or the Kb of a weak
  base, substitute the measured concentrations of
  all the substances present at equilibrium into
  the expression for Ka or Kb.
• HA H2O ? H A-
• Ka H x A-
• HA
• Page 610 Work Sample Problem

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Neutralization Reactions

• Neutralization Reaction the reaction of an acid
  with a base produces water and a salt.
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)
• H2SO4 2KOH(aq) ? K2SO4(aq) 2H2O(l)

• Table salt isnt the only kind of salt.
• Salts are compounds consisting of an anion from
  an acid and a cation from a base.

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Chapter 19

• Section 4

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Neutralization

• Neutralization Reaction the reaction of an acid
  with a base produces water and a salt.
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)
• H2SO4 2KOH(aq) ? K2SO4(aq) 2H2O(l)
• Equivalence point when the number of moles of
  hydrogen ions equals the number of moles of
  hydroxide ions.

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Titration

• You can determine the concentration of an acid
  (or base) in a solution by performing a
  neutralization reaction.
• Titration the process of adding a known amount
  of a solution of known concentration to determine
  the concentration of another solution.

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Titration

• Standard Solution the solution of known
  concentration.
• Added from a buret
• End point the point at which the indicator
  changes color.

• The equivalence point needs to be very close to
  the end point of the titration. If not, you used
  the wrong indicator.

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Chapter 19

• Section 5

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Salts in Solution

• Salts that produce acidic solutions contain
  cations that release protons to water. (Cations
  from weak base)
• Salts that produce basic solutions contain anions
  that attract protons from water (Anions from weak
  acids)

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Buffers

• 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 is a solution in which the pH remains
  relatively constant when small amounts of acid or
  base is added
• Two buffer systems help keep human blood in a
  narrow pH range.

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Review Acid Base Definitions

• Arrhenius
• Acid H producer in solution
• Base OH- producer in solution
• Bronsted-Lowry
• Acid H donor
• Base H acceptor
• Lewis
• Acid Electron pair acceptor
• Base Electron pair donor

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Review Calculations

• pH -log H H10-pH
• pOH -log OH- OH-10-pOH
• pH pOH 14
• Kw HOH- 1.0 x 10-14

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Strong Weak Acids and Bases

• Acids
• Strong Acids Completely ionize in aqueous
  solution.
• Weak Acids Ionize only slightly in aqueous
  solution.
• Weak acids have small Ka values.
• The stronger the acid is, the larger its Ka value.

• Bases
• Strong Bases Completely ionize in aqueous
  solution.
• Weak Bases Ionize only slightly in aqueous
  solution.
• Weak bases have small Kb values.
• The stronger the base is, the larger its Kb value.