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Acids and bases

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					Theories of acids and bases
 Example:
HF + NH3 ↔ NH4+ + F-

 The HF transfers a H+ to the NH3, so it acts as an
 acid; the NH3 accepts the H+ so it acts as a base.
   In Bronsted-Lowry theory, an acid can only behave as a proton
    donor if there is also a base present to act as a proton acceptor.
    Consider the acid-base reaction between a generic acid HA and
    base B:
    HA + B ↔ A- + BH+
    We can see that HA acts as an acid, while B acts as a base. But, if we
    look at the reverse reaction, BH+ acts as an acid (donates a proton)
    and A- acts as a base (accepts a proton).
    So, acids react to form bases and bases react to form acids. Acid-
    base pairs that are related to each other in this way are called
    conjugate acid-base pairs.
    Conjugate acid-base pairs differ by one proton.
   One example of a conjugate pair is H2O and H3O+, which
    is found in all acid-base reactions in aqueous solution.
   H2O + H+ ↔ H3O+ -- this reaction takes place when an
    acid dissolves in water. H3O+ is called the hydroxonium
    ion, the oxonium ion, or the hydronium ion.
   For most reactions it is convenient to simply write H+
    (aq).
   So, H3O+ is the conjugate acid and H2O is the conjugate
    base
   Water can act as a base or as an acid, depending what it
    is reacting with.
     Substances that can act as acids or bases
     are said to be amphoteric or amphiprotic.
   To act as a Bronsted-Lowry acid, substances must be
    able to dissociate and release H+.
   To act as a Bronsted-Lowry base, substances must be
    able to accept H+, which means, they must have a lone
    pair of electrons.
   Example: HCO3- acts as acid and base
   HCO3-(aq) + H2O(l) ↔ CO32-(aq) + H3O+(aq)
   HCO3-(aq) + H2O(l) ↔ H2CO3(aq) + OH-(aq)
 Lewis acids and Bases
   The Lewis definition of acids is broader than the Bronsted-Lowry
    theory.
   By the Lewis definition, an acid is any species capable of accepting a
    lone pair of electrons (no longer restricted to just H+).
   Lewis acid-base reactions result in the formation of a covalent bond
    which will always be a dative bond, because both electrons come from
    the base.
   For example:




   BF3 has an incomplete octet, so it is able to act as a Lewis acid by
    accepting a pair of electrons; NH3 acts as a Lewis base by donating
    its lone pair of electrons. The arrow shows the bond is a coordinate
    covalent bond (dative bond)
   Other good examples of Lewis acid-base reactions are found
    in the chemistry of transition elements.
   Transition metals often form ions with vacant orbitals. Thus
    they are able to act as Lewis acids and accept lone pairs of
    electrons.
   The Lewis base, which donates the lone pair, is called a ligand,
    and usually surrounds the ion in a fixed number ratio.
   Dative bonds form between the metal ion and the ligands,
    resulting in a complex ion that has a characteristic color.
   Typical ligands found in complex ions include H2O, CN- and
    NH3. Note that these all possess lone pairs of electrons.
Theory                  Definition of acid       Definition of base
Bronsted-Lowry          Proton donor             Proton acceptor
Lewis                   Electron pair acceptor   Electron pair donor

 Although all Bronsted-Lowry acids are Lewis acids, not all Lewis
 acids are Bronsted-Lowry acids, so the term Lewis acids are usually
 reserved for those species which can only be described by Lewis
 theory (that is those that do not release H+).

				
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posted:1/17/2012
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