SEE IMMEDIATELY::: CHANG 16.4

Electrolytes:  compounds that ionize (or dissociate) to
    produce aqueous solutions that conduct an electric
    current.  Ex.:  NaCl

Strong Electrolytes:  ionized or dissociated, completely,
    or very nearly completely, in dilute aqueous solutions.
    These are:  strong acids, strong bases, and most soluble
    salts.  Ex.:  HCl, NaOH, KBr

Weak Electrolytes:  not ionized completely in dilute aqueous
    solutions.  These are weak acids, weak bases, salts of
    weak acids and bases.
    Ex.:  acetic acid, formic acid, ammonia,
            methylamine chloride, etc.

Non-electrolytes:  exist as molecules in aqueous solutions
    (do NOT make ions), and therefore, do NOT conduct
    electric current.  Ex.:  sugars, glycols, alcohols

ACID AND BASE EQUILIBRIA:

Acids are sour; they change colors of dyes:  litmus from
    blue to red.
  a) Arrhenius definition:  acids increase [H⁺] in water.
            Ex.:  HNO₃(aq) + H₂O ---->  H₃O⁺(aq) + NO₃^-(aq)

  b) React with Zn(s) or Mg(s) to form H₂ gas:
            2HCl(aq) + Zn(s)----> ZnCl₂(aq) + H₂(g)

  c) React with CO₃^2- to form CO₂ gas:
   2HCl(aq) + MgCO₃(s) ----> MgCl₂(aq) + CO₂(g) + H₂O(l)
 

Bases taste bitter and feel soapy; change litmus from
    red to blue.
  Arrhenius definition:  bases increase [OH^-] in water.
            Ex.:  NaOH(s) + H₂O ---->  Na⁺(aq) + OH^-(aq)

These definitions are ONLY for aqueous solutions!
 

16.1 Brønsted Acids and Bases
Focus on the  H⁺(aq)
Brønsted-Lowry:  acid donates H⁺ and base accepts H⁺.
Therefore, a Brønsted-Lowry base does NOT need to
    contain OH^- (e.g., H₂O can be a Brønsted-Lowry base
    but cannot be an Arrhenius base); water can also be a
    Brønsted-Lowry acid!!

                NH₃ + H₂O -----> NH₄⁺+ + OH^-
    Here, water is donating a proton to ammonia, and is
            an acid. 

Water can act as a base or an acid.
 

Conjugate Acid-Base Pairs
After the proton is donated by an acid, what is left is
        called its conjugate base.
Similarly, whatever remains of the base after it accepts a
       proton is called its conjugate acid.

Consider:       HA(aq) + H₂O(l) ----> H₃O⁺(aq) + A^-(aq)
                                                   <----

After HA (acid) loses its proton, it is converted into
            A^- (base).
Therefore, HA and A^- are conjugate acid-base pairs.
After H₂O (base) gains a proton, it is converted into
        H₃O⁺ (acid).
Therefore, H₂O and H₃O⁺ are conjugate acid-base pairs.

Conjugate acid-base pairs differ by only one proton.

The stronger the acid, the weaker the conjugate base,
    and vice-versa.

H⁺ is the strongest acid that can exist in equilibrium in
            aqueous solution.
OH^- is the strongest base that can exist in equilibrium
        in aqueous solution.
Any acid or base stronger than H⁺ or OH^- simply reacts
        stoichiometrically to produce H⁺ and OH^-.

The conjugate base of a strong acid (e.g., Cl^-) has
    negligible acid-base properties.
Similarly, the conjugate acid of a strong base has
    negligible acid-base properties.
 

At 25°C:
               K_w = K[H₂O] = [H⁺][OH^-] = 1.0 x 10^-14

The above is called the autoionization of water.

The H⁺(aq) is simply a proton with no electrons.  (The H
    atom has one proton, one electron and no neutrons.)

In water, the H⁺(aq) forms clusters:
            2 H₂O(l) ----> H₃O⁺(aq) + OH^-(aq)
                           <----

The simplest cluster is H₃O⁺(aq), THE HYDRONIUM ION.
    Larger clusters are  H₅O₂⁺ and H₉O₄⁺.
Mostly, we use H⁺(aq) and H₃O⁺(aq) interchangeably.
 

16.3 pH - A Measure of Acidity
The pH scale:
 In most solutions, [H⁺](aq) or [H₃O⁺] is quite small,
    from about 10^-1 to 10^-7 M.
 We define pH = -log[H⁺] or -log[H₃O⁺].
 Similarly, pOH = -log[OH^-].
 

 In neutral water at 25°C, pH = pOH = 7.00.
    Therefore, [H⁺] = [OH^-] = 1.0 x 10^-7M 

 In acidic solutions, [H⁺] > 1.0 x 10^-7M, so pH < 7.00.
 In basic solutions, [H⁺] < 1.0 x 10^-7M, so pH > 7.00.

                    pH + pOH = 14.00

The higher the pH, the lower the pOH, and the more
            basic the solution.
Most pH and pOH values fall between 0 and 14. 
    However, there are no theoretical limits on the values
    of pH or pOH.  (e.g., pH of 0.10 M HCl is 1.00; pH of
    1.00 M HCl = 0.00, pH of 2.00 M HCl is -0.301,  
    pH of 10.0 M HCl = -1.00).  There are some
    concentrated acid mixtures whose pH is
    approximately   -35.0!!

Other p scales:
In general, pX = - logX
           so     pK = - logK,  pK_w = - logK_w
            and  pK_w = pH + pOH = 14.00

Measuring pH:
 The most accurate method to measure pH is to use a
        pH meter.
 However, certain dyes change color as pH changes.
    These are indicators.
 Indicators are less precise than pH meters.
 Some natural products can be used as indicators (tea
    is colorless in acid and brown in base; red cabbage
    extract and red rose petals are other natural
    indicators).
 

7 STRONG ACIDS:  HCl, HBr, HI, HNO₃, HClO₃, HClO₄, H₂SO₄
 

Strong Bases are also strong electrolytes and dissociate
    completely in water.
In strong bases, the [OH^-] is due solely to the base, and
    [OH^-] = molarity of the base, and pOH can be obtained
    directly from the molarity of the hydroxide ion.

    You have to multiply the number of hydroxide ions per
    formula unit of base times the molarity to get [OH^-]:

    E.g., for CsOH, x = 1; for Ca(OH)₂, x = 2.
    Note: the base MUST be soluble; if insoluble, get NO
                    [OH^-]!
    Also, the formula for the base does not necesarily
        contain the hydroxide ion:
                        e.g., O^2- + H₂O -----> 2OH^-

7 STRONG BASES:  LiOH, NaOH, KOH, CsOH, RbOH,
                                  Ba(OH)₂, Sr(OH)₂

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