Chapter 14 Acids and Bases

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Chapter 14Acids and Bases
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Types of Electrolytes

• salts water soluble ionic compounds
• all strong electrolytes
• acids form H1 ions in water solution
• bases combine with H1 ions in water solution
• increases the OH-1 concentration
• may either directly release OH-1 or pull H1 off
  H2O

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Properties of Acids

• Sour taste
• react with active metals
• i.e. Al, Zn, Fe, but not Cu, Ag or Au
• 2 Al 6 HCl 2 AlCl3 3 H2
• corrosive
• react with carbonates, producing CO2
• marble, baking soda, chalk, limestone
• CaCO3 2 HCl CaCl2 CO2 H2O
• change color of vegetable dyes
• blue litmus turns red
• react with bases to form ionic salts

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Common Acids
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Structures of Acids

• binary acids have acid hydrogens attached to a
  nonmetal atom
• HCl, HF

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Structure of Acids

• oxy acids have acid hydrogens attached to an
  oxygen atom
• H2SO4, HNO3

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Structure of Acids

• carboxylic acids have COOH group
• HC2H3O2, H3C6H5O3
• only the first H in the formula is acidic
• the H is on the COOH

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Properties of Bases

• also known as alkalis
• taste bitter
• alkaloids plant product that is alkaline
• often poisonous
• solutions feel slippery
• change color of vegetable dyes
• different color than acid
• red litmus turns blue
• react with acids to form ionic salts
• neutralization

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Common Bases
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Structure of Bases

• most ionic bases contain OH ions
• NaOH, Ca(OH)2
• some contain CO32- ions
• CaCO3 NaHCO3
• molecular bases contain structures that react
  with H
• mostly amine groups

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Arrhenius Theory

• bases dissociate in water to produce OH- ions and
  cations
• ionic substances dissociate in water
• NaOH(aq) ? Na(aq) OH(aq)
• acids ionize in water to produce H ions and
  anions
• because molecular acids are not made of ions,
  they cannot dissociate
• they must be pulled apart, or ionized, by the
  water
• HCl(aq) ? H(aq) Cl(aq)
• in formula, ionizable H written in front
• HC2H3O2(aq) ? H(aq) C2H3O2(aq)

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Arrhenius Acid-Base Reactions

• the H from the acid combines with the OH- from
  the base to make a molecule of H2O
• it is often helpful to think of H2O as H-OH
• the cation from the base combines with the anion
  from the acid to make a salt
• acid base ? salt water
• HCl(aq) NaOH(aq) ? NaCl(aq) H2O(l)

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Problems with Arrhenius Theory

• does not explain why molecular substances, like
  NH3, dissolve in water to form basic solutions
  even though they do not contain OH ions
• does not explain acid-base reactions that do not
  take place in aqueous solution
• H ions do not exist in water. Acid solutions
  contain H3O ions
• H a proton!
• H3O hydronium ions

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Brønsted-Lowery Theory

• in a Brønsted-Lowery Acid-Base reaction, an H is
  transferred
• does not have to take place in aqueous solution
• broader definition than Arrhenius
• acid is H donor, base is H acceptor
• base structure must contain an atom with an
  unshared pair of electrons
• in the reaction, the acid molecule gives an H to
  the base molecule
• HA B ? A HB

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Amphoteric Substances

• amphoteric substances can act as either an acid
  or a base
• have both transferable H and atom with lone pair
• HCl(aq) is acidic because HCl transfers an H to
  H2O, forming H3O ions
• water acts as base, accepting H
• HCl(aq) H2O(l) ? Cl(aq) H3O(aq)
• NH3(aq) is basic because NH3 accepts an H from
  H2O, forming OH(aq)
• water acts as acid, donating H
• NH3(aq) H2O(l) ? NH4(aq) OH(aq)

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Brønsted-Lowery Acid-Base Reactions

• one of the advantages of Brønsted-Lowery theory
  is that it allows reactions to be reversible
• HA B ? A HB
• the original base has an extra H after the
  reaction so it could act as an acid in the
  reverse process
• and the original acid has a lone pair of
  electrons after the reaction so it could act
  as a base in the reverse process
• A HB ? HA B
• a double arrow, ?, is usually used to indicate a
  process that is reversible

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Conjugate Pairs

• In a Brønsted-Lowery Acid-Base reaction, the
  original base becomes an acid in the reverse
  reaction, and the original acid becomes a base in
  the reverse process
• each reactant and the product it becomes is
  called a conjugate pair
• the original base becomes the conjugate acid and
  the original acid becomes the conjugate base

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Brønsted-Lowery Acid-Base Reactions
HA B ? A HB
acid base conjugate conjugate
base acid
HCHO2 H2O ? CHO2 H3O acid
base conjugate conjugate base
acid
H2O NH3 ? HO NH4 acid
base conjugate conjugate base
acid
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Conjugate Pairs
In the reaction H2O NH3 ? HO NH4
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Practice Identify the Brønsted-Lowery Acids and
Bases and their Conjugates in each Reaction
H2SO4 H2O ? HSO4 H3O
HCO3 H2O ? H2CO3 HO
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Neutralization Reactions

• H OH- ??H2O
• acid base ??salt water
• double displacement reactions
• salt cation from base anion from acid
• cation and anion charges stay constant
• H2SO4 Ca(OH)2 ? CaSO4 2 H2O
• some neutralization reactions are gas evolving
  where H2CO3 decomposes into CO2 and H2O

H2SO4 2 NaHCO3 ? Na2SO4 2 H2O 2 CO2
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Nonmetal Oxides are Acidic

• nonmetal oxides react with water to form acids
• causes acid rain
• CO2 (g) H2O(l) ? H2CO3(aq)
• 2 SO2(g) O2(g) 2 H2O(l) ? 2 H2SO4(aq)
• 4 NO2(g) O2(g) 2 H2O(l) ? 4 HNO3(aq)

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Acid ReactionsAcids React with Metals

• acids react with many metals
• but not all!!
• when acids react with metals, they produce a salt
  and hydrogen gas

3 H2SO4(aq) 2 Al(s) ? Al2(SO4)3(aq) 3 H2(g)
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Acid ReactionsAcids React with Metal Oxides

• when acids react with metal oxides, they produce
  a salt and water
• 3 H2SO4 Al2O3 ? Al2(SO4)3 3 H2O

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Base Reactions

• the reaction all bases have is common is
  neutralization of acids
• strong bases will react with Al metal to form
  sodium aluminate and hydrogen gas
• 2 NaOH 2 Al 6 H2O ? 2 NaAl(OH)4 3 H2

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Titration

• using reaction stoichiometry to determine the
  concentration of an unknown solution
• Titrant (unknown solution) added from a buret
• indicators are chemicals added to help determine
  when a reaction is complete
• the endpoint of the titration occurs when the
  reaction is complete

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Titration
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Titration
The base solution is the titrant in the buret.
As the base is added to the acid, the H reacts
with the OH to form water. But there is still
excess acid present so the color does not change.
At the titrations endpoint, just enough base has
been added to neutralize all the acid. At this
point the indicator changes color.
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Example 14.4Acid-Base Titration
The titration of 10.00 mL of HCl solution of
unknown concentration requires 12.54 mL of 0.100
M NaOH solution to reach the end point. What is
the concentration of the unknown HCl solution?
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Strong or Weak

• a strong acid is a strong electrolyte
• practically all the acid molecules ionize, ?
• a strong base is a strong electrolyte
• practically all the base molecules form OH ions,
  either through dissociation or reaction with
  water, ?
• a weak acid is a weak electrolyte
• only a small percentage of the molecules ionize,
  ?
• a weak base is a weak electrolyte
• only a small percentage of the base molecules
  form OH ions, either through dissociation or
  reaction with water, ?

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Strong Acids

• The stronger the acid, the more willing it is to
  donate H
• use water as the standard base
• strong acids donate practically all their Hs
• 100 ionized in water
• strong electrolyte
• H3O strong acid
• molarity

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Strong Acids
Pure Water
HCl solution
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Weak Acids

• weak acids donate a small fraction of their Hs
• most of the weak acid molecules do not donate H
  to water
• much less than 1 ionized in water
• H3O ltlt weak acid

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Weak Acids
Pure Water
HF solution
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Strong Bases

• The stronger the base, the more willing it is to
  accept H
• use water as the standard acid
• strong bases, practically all molecules are
  dissociated into OH or accept Hs
• strong electrolyte
• multi-OH bases completely dissociated
• HO strong base x ( OH)

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Weak Bases

• in weak bases, only a small fraction of molecules
  accept Hs
• weak electrolyte
• most of the weak base molecules do not take H
  from water
• much less than 1 ionization in water
• HO ltlt strong base

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Relationship between Strengths of Acids and their
Conjugate Bases

• the stronger an acid is, the weaker the
  attraction of the ionizable H for the rest of the
  molecule is
• the better the acid is at donating H, the worse
  its conjugate base will be at accepting a H
• strong acid HCl H2O ? Cl H3O weak conj.
  base
• weak acid HF H2O ? F H3O strong
  conj. base

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Autoionization of Water

• Water is actually an extremely weak electrolyte
• therefore there must be a few ions present
• about 1 out of every 10 million water molecules
  form ions through a process called autoionization
• H2O Û H OH
• H2O H2O Û H3O OH
• all aqueous solutions contain both H and OH
• the concentration of H and OH are equal in
  water
• H OH 10-7M _at_ 25C

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Ion Product of Water

• the product of the H and OH concentrations is
  always the same number
• the number is called the ion product of water and
  has the symbol Kw
• H x OH 1 x 10-14 Kw
• as H increases the OH must decrease so the
  product stays constant
• inversely proportional

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Acidic and Basic Solutions

• neutral solutions have equal H and OH
• H OH 1 x 10-7
• acidic solutions have a larger H than OH
• H gt 1 x 10-7 OH lt 1 x 10-7
• basic solutions have a larger OH than H
• H lt 1 x 10-7 OH gt 1 x 10-7

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Example - Determine the H1 for a 0.00020 M
Ba(OH)2 and determine whether the solution is
acidic, basic or neutral
Ba(OH)2 Ba2 2 OH therefore OH 2 x
0.00020 0.00040 4.0 x 10-4 M
H 2.5 x 10-11 M
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Practice - Determine the H1 concentration and
whether the solution is acidic, basic or neutral
for the following

• OH 0.000250 M
• OH 3.50 x 10-8 M
• Ca(OH)2 0.20 M

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Complete the TableH vs. OH-
H 100 10-1 10-3 10-5 10-7
10-9 10-11 10-13 10-14
OH-
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Complete the TableH vs. OH-
Acid
Base
H 100 10-1 10-3 10-5 10-7
10-9 10-11 10-13 10-14
OH-10-14 10-13 10-11 10-9 10-7
10-5 10-3 10-1 100
even though it may look like it, neither H of
OH- will ever be 0
the sizes of the H and OH- are not to scale
because the divisions are powers of 10 rather
than units
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pH

• the acidity/basicity of a solution is often
  expressed as pH
• pH -logH, H 10-pH
• exponent on 10 with a positive sign
• pHwater -log10-7 7
• need to know the H concentration to find pH
• pH lt 7 is acidic pH gt 7 is basic, pH 7 is
  neutral

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pH

• the lower the pH, the more acidic the solution
  the higher the pH, the more basic the solution
• 1 pH unit corresponds to a factor of 10
  difference in acidity
• normal range 0 to 14
• pH 0 is H 1 M, pH 14 is OH 1 M
• pH can be negative (very acidic) or larger than
  14 (very alkaline)

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pH of Common Substances
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Example - Calculate the pH of a 0.0010 M Ba(OH)2
solution determine if is acidic, basic or
neutral
Ba(OH)2 Ba2 2 OH- therefore OH- 2 x
0.0010 0.0020 2.0 x 10-3 M
pH -log H -log (5.0 x 10-12) pH 11.3
pH gt 7 therefore basic
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Practice - Calculate the pH of the following
strong acid or base solutions

• 0.0020 M HCl
• 0.0050 M Ca(OH)2
• 0.25 M HNO3

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Complete the TablepH
pH
H 100 10-1 10-3 10-5 10-7
10-9 10-11 10-13 10-14
OH-10-14 10-13 10-11 10-9 10-7
10-5 10-3 10-1 100
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Complete the TablepH
Acid
Base
pH 0 1 3 5 7
9 11 13 14
H 100 10-1 10-3 10-5 10-7
10-9 10-11 10-13 10-14
OH-10-14 10-13 10-11 10-9 10-7
10-5 10-3 10-1 100
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Sample - Calculate the concentration of H for
a solution with pH 3.7
H 10-pH
H 10-3.7 means 0.0001 lt H1 lt 0.001
H 2 x 10-4 M 0.0002 M
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Practice - Determine the H for each of the
following

• pH 2.7
• pH 12
• pH 0.60

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Buffers

• buffers are solutions that resist changing pH
  when small amounts of acid or base are added
• they resist changing pH by neutralizing added
  acid or base
• buffers are made by mixing together a weak acid
  and its conjugate base
• or weak base and it conjugate acid

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How Buffers Work

• the weak acid present in the buffer mixture can
  neutralize added base
• the conjugate base present in the buffer mixture
  can neutralize added acid
• the net result is little to no change in the
  solution pH

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A Buffer made from Acetic acid and Sodium Acetate

• a buffer solution with a pH of 4.75 can be made
  by mixing equal volumes of 1 M HC2H3O2 and 1 M
  NaC2H3O2
• adding 10 mL of 0.1 M HCl to 1 L of this solution
  will give a solution with a pH of 4.75
• adding 10 mL of 0.1 M HCl to 1 L of distilled
  water will give a solution with pH of 3.0
• adding 10 mL of 0.1 M NaOH to 1 L of this
  solution will give a solution with a pH of 4.75
• adding 10 mL of 0.1 M NaOH to 1 L of distilled
  water will give a solution with pH of 11.0

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Acetic Acid/Acetate Buffer
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What is Acid Rain?

• natural rain water has a pH of 5.6
• naturally slightly acidic due mainly to CO2
• rain water with a pH lower than 5.6 is called
  acid rain
• acid rain is linked to damage in ecosystems and
  structures

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What Causes Acid Rain?

• many natural and pollutant gases dissolved in the
  air are nonmetal oxides
• CO2, SO2, NO2
• nonmetal oxides are acidic
• CO2 H2O ? H2CO3
• 2 SO2 O2 2 H2O ? 2 H2SO4
• processes that produce nonmetal oxide gases as
  waste increase the acidity of the rain
• natural volcanoes and some bacterial action
• man-made combustion of fuel
• weather patterns may cause rain to be acidic in
  regions other than where the nonmetal oxide is
  produced

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pH of Rain in Different Regions
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Sources of SO2 from Utilities
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Damage from Acid Rain

• acids react with metals, and materials that
  contain carbonates
• acid rain damages bridges, cars and other
  metallic structures
• acid rain damages buildings and other structures
  made of limestone or cement

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Damage from Acid Rain