AP Chemistry
Chapter 16 Outline
16.1 Acids and Bases
16.1.1 Operational Definitions (characteristic properties)
16.1.1.1
Acids taste sour, turn litmus paper red and react with certain metals to release
hydrogen gas.
16.1.1.2
Bases taste bitter, feel slippery, turn litmus paper blue and phenolphthalein pink.
16.1.2 Arrhenius acid-base theory (aqueous solutions)
16.1.2.1
Acids produce H+ ions in water
16.1.2.1.1 Ex. HCl
16.1.2.2
Bases produce OH- ions in water
16.1.2.2.1 Example, NaOH
16.2 Brønsted-Lowry Acids and Bases
16.2.1 H+ ion simply a bare proton, with no electrons
16.2.1.1
Tends to form hydronium ions, H3O+
16.2.2 Bronsted-Lowry theory emphasizes proton transfer reactions
16.2.2.1
B-L acid: a substance that can donate a proton to another substance
16.2.2.2
B-L base: a substance that can accept a proton
16.2.2.3
Amphiprotic: can act as either an acid OR a base
16.2.3 Conjugate acid-base pairs: HX and X16.2.3.1
One is a reactant, the other is a product
16.2.3.2
Formulas differ only by a H+
16.2.3.3
Two sets of conjugate pairs in any B-L reaction
16.2.4 Not all acids and bases are equally good at donating or accepting protons
16.2.4.1
The stronger the acid, the weaker its conjugate base; the stronger a base, the
weaker its conjugate acid.
16.2.4.1.1 The conjugate bases of strong acids have practically no tendency to accept protons.
16.2.4.1.2 The conjugate bases of weak acids are weak bases.
16.2.4.1.3 The equilibrium position favors transfer of the proton to the stronger base. OR, the
stronger acid and the stronger base will react to form the weaker acid and the weaker
base.
16.3 The autoionization of water H2O + H2O  H3O+ + OH16.3.1 No individual molecule remains ionized for very long!
16.3.2 Kw = [H3O+][ OH-] = 1.0 x 10-14 at 25oC
Memorize this!
16.3.2.1
This is true in ANY (dilute) aqueous solution
16.3.2.1.1 If [H3O+] = [ OH-] the solution is “neutral”
16.3.2.1.2 If [H3O+] > [ OH-] the solution is acidic
16.3.2.1.3 If [H3O+] < [ OH-] the solution is basic
16.4 The pH scale
Know this entire section!
16.4.1 [H3O+] in a solution is usually very small
16.4.1.1
pH = -log[H3O+] or pH = -log[H+]
16.4.1.2
the pH of a neutral solution is 7.00 at 25oC
16.4.1.3
the pH of an acidic solution is <7
16.4.1.4
the pH of a basic solution is >7
16.4.1.5
for mental math: if [H3O+] is 1 x 10-x, pH = x (you can use this to estimate pH)
16.5.2 pOH = -log[OH-]
Be able to do calculations with these equations!
16.5.3 pH + pOH = 14
16.5.4 Measuring pH
16.5.4.1
Use a pH meter
16.5.4.2
Use indicators (less precise)
16.5 Strong acids and bases
16.5.1 Strong electrolytes that exist in aqueous solutions entirely as ions
16.5.1.1
Memorize the strong acids: HCl, HBr, HI, HNO3, HClO3, HClO4, H2SO4
16.5.1.2
In aqueous solutions of strong acids, the acid is the only significant source of H+
ions
16.5.1.2.1 Typically ignore the H+ concentration from the autoionization of water
16.5.1.2.2
[H+] = the original concentration of the acid for monoprotic acids
16.5.1.3
Memorize the strong bases: hydroxides of the alkali metals and the heavier
alkaline earth metals
16.5.1.3.1 Basic anhydrides: metal oxides will form hydroxides when dissolved in water
16.6 Weak acids
16.6.1 Most acidic substances are weak acids, only partially ionized in solution
16.6.1.1
HA(aq)  H+(aq) + A-(aq)
or HA(aq) + H2O(l)  H3O+(aq) + A-(aq)
[ H O  ][ A  ]
[ H  ][ A  ]
16.6.1.1.1
or K a  3
Ka 
[ HA]
[ HA]
16.6.1.1.2 note that only H atoms attached to oxygen atoms show acidic behavior
16.6.1.1.3 the larger the value of Ka, the stronger the acid
16.6.2 Calculating Ka from pH
16.6.2.1
Write out balanced equation & Ka expression
16.6.2.2
Set up an ICE table to find concentrations of the species involved
16.6.2.2.1
Assume that (initial acid concentration – change) (initial acid concentration);
16.6.2.2.2 OK if % ionization is < 5% or if Ka is 10-4 (this will avoid needing to use the
quadratic formula)
16.6.2.3
Substitute into Ka expression and evaluate
16.6.2.4
The pH of a weak acid solution is higher than that of a strong acid of the same
molarity
change in [ HA]
% ionization 
 100
16.6.2.5
initial [ HA]
16.6.3 Polyprotic acids have more than one ionizable hydrogen atom.
16.6.3.1
H2A  H+ + HAKa1
+
216.6.3.1.1 HA  H + A
Ka2
16.6.3.2
It is always easier to remove the first proton than following protons.
16.6.3.2.1 Ka1 > Ka2
16.6.3.2.2 As long as successive Ka values differ by a factor of 103 or more, consider only Ka1
16.7 Weak bases
16.7.1 Weak bases may either …
16.7.1.1
Have a pair of nonbonding electrons that can accept a proton OR
16.7.1.2
be the anions of weak acids
16.7.2 B(aq) + H2O  HB+(aq) + OH-(aq)
[ HB  ][OH  ]
16.7.2.1
be able to do ICE-style problems with this
Kb 
[ B]
16.8 Relationship between Ka and Kb
16.8.1 Ka x Kb = Kw
for a conjugate acid-base pair! MEMORIZE THIS
16.8.1.1
If you know Ka for a weak acid, you can use this relationship to solve for Kb of
its conjugate base
16.8.2 Often, Ka values are reported as pKa (-log Ka)
16.8.2.1
pKa + pKb = pKw = 14.00 at 25oC
16.9 Acid-base properties of salt solutions
16.9.1 Ions can exhibit acidic or basic behavior
16.9.2 Hydrolysis reaction: salt + water  acid + base
16.9.3 Predicting the pH of salt solutions:
16.9.3.1
An anion that is the conjugate base of a strong acid will not affect the pH.
16.9.3.2
An anion that is the conjugate base of a weak acid will produce a basic solution.
16.9.3.3
A cation that is the conjugate acid of a weak base will produce an acidic solution.
16.9.3.4
The cations of the strong Arrhenius bases will not affect the pH.
16.9.3.5
Other metal ions will produce an acidic solution.
16.9.3.6
If both the conjugate base of a weak acid and the conjugate acid of a weak base
are present, the ion with the larger Ka or Kb value will determine the pH of the solution.
16.10Acid-Base behavior and chemical structure
16.10.1Three main factors:
16.10.1.1
A molecule containing hydrogen will transfer a proton only if the H-X bond is
polarized so that the H atom has a + charge and the X atom has a - charge
16.10.1.2
Very strong bonds are less easily dissociated than weaker ones.
16.10.1.3
The greater the stability of the conjugate base, the stronger the acid.
16.10.2Binary Acids
16.10.2.1
Acidity decreases down a column
16.10.2.1.1 H-X bond strength decreases as X increases in atomic radius
16.10.2.2
Acidity increases L-R in a period
16.10.2.3
H-X bonds become more polar going L-R
16.10.3Oxyacids
16.10.3.1
As central atom becomes more electronegative, acid strength increases
16.10.3.2
For the same central atom, acid strength increases as the number of oxygen atoms
increases
KNOW THIS
16.10.3.2.1
Because electron density is pulled away from the OH bond, making it more polar
16.10.4Carboxylic acids
16.10.4.1
Resonance structures for the conjugate base help to stabilize it
16.11 Lewis acids and bases—broadest theory of the 3
Know this whole section!
16.11.1
Lewis acid = electron pair acceptor
16.11.1.1
Lewis acids have an empty valence orbital
16.11.1.2
Many cations can act as Lewis acids
16.11.2
Lewis base = electron pair donor
16.11.2.1
Lewis bases have a lone pair of electrons
16.11.3
Hydrolysis of metal ions
16.11.3.1
hydration as a Lewis acid-base reaction
16.11.3.1.1
as charge increases, metal ion is more acidic
16.11.3.1.2
as radius decreases, metal ion is more acidic
16.11.3.1.3
Note: Lewis acid-base theory also relevant for complex ion reactions