Chapter 17: Acids and Bases

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Chapter 17 Acids and Bases

• Acid-base reactions involve proton (hydrogen ion,
  H) transfer
• The generalization of the Arrhenius definition of
  acids and bases is called the Brønsted-Lowry
  definitions
• An acid is a proton donor
• A base is a proton acceptor
• This allows for gas phase acid-base reactions

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• Species that differ by a proton, like H2O and
  H3O, are called conjugate acid-base pairs

The reaction of HCl and H2O. HCl is the acid
because it donates a proton. Water is the base
because it accepts a proton.
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(a) Formic acid transfers a proton to a water
molecule. HCHO2 is the acid and H2O is the base.
(b) When a hydronium ion transfers a proton to
the CHO2- ion, H3O is the acid and formate ion
is the base.
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• An amphoteric substances can act as either an
  acid or base
• These are also called amphiprotic, and can be
  either molecules or ions
• For example, the hydrogen carbonate ion

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• The strength of an acid is a measure of its
  ability to transfer a proton
• Acids that react completely with water (like HCl
  and HNO3) are classified as strong
• Acids that are less than completely ionized are
  called weak acids
• Bases can be classified in a similar fashion
• Strong bases, like the oxide ion, react
  completely
• Weak bases, like NH3, undergo incomplete reactions

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• The strongest acid in water is the hydronium ion
• If a more powerful proton donor is added to
  water, it quantitatively reacts with water to
  produce H3O
• Similarly, the strongest base that can be found
  in water is the hydroxide ion, because more
  powerful proton acceptors react quantitatively
  with water to produce OH-

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• Acetic acid (HC2H3O2) is a weak acid
• It ionizes only slightly in water
• The hydronium ion is a better proton donor than
  acetic acid (it is a stronger acid)
• The acetate ion is a better proton acceptor than
  water (it is a stronger base)
• The position of an acid-base equilibrium favors
  the weaker acid and base

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• This can be generalized
• Stronger acids and bases tend to react with each
  other to produce their weaker conjugates
• The stronger a Brønsted acid is, the weaker is
  its conjugate base
• The weaker a Brønsted acid is, the stronger is
  its conjugate base
• These ideas can be applied to the binary acids
  (acids made from hydrogen and one other element)

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• The strengths of the binary acids increases from
  left to right within the same period
• For example, HCl is stronger acid than H2S which
  is a stronger acid than PH3
• The strengths of the binary acids increase from
  top to bottom within a group
• For example, HI is a stronger acid than HBr which
  is a stronger acid than HCl
• Trends are also present in the oxoacids (acids of
  hydrogen, oxygen, and one other element)

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• When the central atom holds the same number of
  oxygen atoms, the acid strength increases from
  the bottom to top within a group and from left to
  right within a period
• Acid strength HClO4 gt HBrO4 gt HIO4
• Acid strength HClO4 gt H2SO4 gt H3PO4
• For a given central atom, the acid strength of an
  oxoacid increases with the number of oxygens held
  by the central atom
• Acid strength H2SO4 gt H2SO3

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• The strength of an acid can be analyzed in terms
  the the basicity of the anion formed during the
  ionization
• The basicity is the willingness of the anion to
  accept a proton from the hydronium ion
• Consider H2SO4 and H3PO4

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• The anions are
• In oxoanions, the lone oxygens carry most of the
  negative charge, making the hydrogen sulfate ion
  a weaker base than the hydrogen phosphate ion

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• In terms of the percentage of molecules that are
  ionized, sulfuric acid is a stronger acid than
  phosphoric acid
• There is a third definition for acid and bases
• It is a further generalization, or broadening, of
  the species that can be classified as either an
  acid or base
• The definitions are based on electron pairs and
  are called Lewis acids and bases

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• A Lewis acid is any ionic or molecular species
  that can accept a pair of electrons in the
  formation of a coordinate covalent bond
• A Lewis base is any ionic or molecular species
  that can donate a pair of electrons in the
  formation of a coordinate covalent bond
• Neutralization is the formation of a coordinate
  covalent bond between the donor (base) and
  acceptor (acid)

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NH3 (a Lewis base) forms a coordinate covalent
bond with BF3 (a Lewis acid) during
neutralization. NH3BF3 is called an addition
compound because it was made by joining two
smaller molecules.
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Carbon dioxide (a Lewis acid) reacts with
hydroxide ion (a Lewis base) in solution to form
the bicarbonate ion. The electrons in the
coordinate covalent bond come from the oxygen
atom in the hydroxide ion.
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• Lewis acids
• Molecules or ions with incomplete valence shells
  (for example BF3 or H)
• Molecules or ions with complete valence shells,
  but with multiple bonds that can be shifted to
  make room for more electrons (for example CO2)
• Molecules or ions that have central atoms capable
  of holding additional electrons (usually, atoms
  of elements in Period 3 and below, for example
  SO2)

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• Lewis bases
• Molecules or ions that have unshared pairs of
  electrons and that have complete shells (for
  example O2- or NH3)
• All Brønsted acids and bases are Lewis acids and
  bases, just like all Arrhenius acids and bases
  are Brønsted acids and bases
• Consider a proton transfer from the Lewis
  perspective

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• For example, the proton transfer between the
  hydronium ion and ammonia

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• The elements most likely to form acids are the
  nonmetals in the upper right-hand corner of the
  periodic table
• The elements most likely to form basic hydroxides
  are the IA and IIA metals along the left of the
  periodic table
• The elements themselves can be classified
  according to the ability of their oxides to form
  acids or bases

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• In general, most metal oxides react with water to
  form bases, and nonmetal oxides react with water
  to form acids
• In Section 5.5 metal oxides were called base
  anhydrides and nonmetal oxides were called acid
  anhydrides
• When cations dissolve in water, they form species
  called hydrated ions
• Hydrated metal ions tend to be Brønsted acids

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• For the monohydrate of the metal ion Mn the
  equilibrium can be represented as

The metal ion makes the hydrogens on the water
more acidic.
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• The charge density of a cation is its charge
  divided by its volume
• The higher the charge density, the better a
  cation is at drawing electron density from a O-H
  bond and the more acidic it is
• Within a given period, the cation size increases,
  and the charge density decreases, from top to
  bottom
• As a result, the most acidic hydrated cations are
  found at the top of a group
• As the cation charge increases, it becomes more
  acidic

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• When the charge (oxidation number) is small, its
  oxide tends to be basic
• When the cation ion charge is 3, the oxide
  begins to become acidic
• An amphoteric species is capable of acting as
  both an acid and base
• Aluminum oxide is an example of an amphoteric
  compound

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• Many nonmetal oxides are acid anhydrides
• For example

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• Water undergoes self-ionization or autoionization
  making it a weak electrolyte
• This equilibrium is described by the ion product
  of water

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• At 25oC in pure water it has been found that
• In any aqueous solution, the product of H and
  OH- equals Kw
• This provides an alternate a way to define the
  acidity or basicity of a solution

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• Neutral solutions H3O OH- or H
  OH-
• Acidic solutions H3O gt OH- or H gt
  OH-
• Basic solutions H3O lt OH- or H lt
  OH-
• To make the comparison of small values of H
  easier, the pH was defined
• In terms of the pH
• Neutral solutions pH 7.00
• Acidic solutions pH lt 7.00
• Basic solutions pH gt 7.00

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The pH of some common solutions. H decreases,
while OH- increases, from top to bottom.
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• The pH of a solution can be measured with a pH
  meter or estimated using a visual acid-base
  indicator (see Table 17.3, page 762)
• An acid-base indicator is a species that changes
  color based on the pH
• Calculating the pH of a strong acid or base is
  easy because they are 100 dissociated in
  aqueous
• For example, the pH of 0.10 M HCl is 1.00 and the
  pH of 0.10 M NaOH is 13.00

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• In the last example it was assumed that the total
  concentration of H was due to the strong acid
  (HCl) and OH- was due to the strong base (NaOH)
  
• This assumption is valid because the
  autoionization of water is suppressed in strongly
  acidic or strongly basic solutions
• This assumption fails for very dilute solutions
  of acids or bases (less than 10-6 M)