ACID – BASE EQUILIBRIA
Acids and Bases
 An Arrhenius acid is a substance that when dissolved in water
increases the H+ ion concentration
 An Arrhenius base is a substance that when dissolved in water
increases the OH- concentration
 Narrow definition that is limited to aqueous solutions
Bronsted-Lowry Acids and Bases
 In the Bronsted-Lowry system, an acid is a chemical species
that donates H+ and a base is a species that accepts H+
 A Bronsted-Lowry base does not need to contain OH
o NH3 is a Bronsted base but not an Arrhenius base
 Consider NH3 + H2O  NH4+ + OHo Water donates a proton therefore it is an acid
o Ammonia accepts the proton therefore it is a base
 Amphoteric substances can behave as an acid or a base. Water
is an example
Conjugate Acid-Base Pairs
 Whatever is left of an acid after it donates a H+ is called its
conjugate base.
 A conjugate acid is formed by adding a H+ to a base
 Consider HA + H2O  H3O+ + A HA and A- differ only in the presence of a proton
o They are a conjugate acid-base pair
o A- is the conjugate base
o When H2O(base) gains a proton it is converted toH3O+(acid)
 H3O+ is the conjugate acid
Relative Strength of Acids and Bases
 The stronger an acid, the weaker its conjugate base will be
 Acids and bases can be categorized by their behavior in water
o Strong Acids completely transfer their protons to water
 No undissociated molecules remain in solution
 Their conjugate bases do not become protonated ie
HCl
o Weak acids only partly dissociate
 They exist as a mixture molecules and component
ions
 Their conjugate bases show a tendency to abstract
protons from water
 The conjugate bases are weak bases
 Acetic acid is a weak acid; acetate ion is a
weak base
o In every acid-base reaction, the position of the
equilibrium favors the transfer of a proton from the
stronger acid to the stronger base
 H+ is the strongest acid in aqueous equilibrium
 OH- is the strongest base
Autoionization of Water
 In pure water the following equilibrium is established
o 2H2O(l)  H3O+ + OH Ion Product of Water
o The equilibrium expression can be written as
Kc =[H3O+] [OH-] / [H2O]2 since H2O is a pure liquid we
can simplify to:
[H2O] 2 Kc = [H3O+] [OH-] = Kw
 Kw is called the ion product constant of water
o At 25C , the ion product constant is 1.0 X 10-14= Kw =
[H3O] [OH]
The pH Scale
 We express [H+] in terms of pH: pH = -log[H+]= log [H3O+]
 We use a similar system to describe [OH] : pOH = -log[OH-]
 The value of Kw is 1 X 10-14 , we can describe the relation
between pH and pOH as pH + pOH = -logKw = 14.00
Strong Acids
 Strong acids are strong electrolytes
 Ionize completely in water
 In solution, strong acids are usually the only source of H+
 The ph of a strong monoprotic acid may be calculated from the
initial molarity
Strong Bases
 Ionic hydroxides of alkali metals and heavier alkaline earth
metals are strong bases
 Strong bases are strong electrolytes and ionize completely
 Not all bases contain OHo Ionic metal oxides, hydrides and nitrides are strong bases
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o
o
o
They are able to abstract a proton from water
O-2 + H2O  2OHH- + H2O  H2 + OHN-3+ H2ONH3 + 3OH-
Weak Acids
 Weak acids only partly ionize so there is a mixture of ions and
unionized acid in solution
 An equilibrium expression can be written for the reaction
HA  H+ + A- Ka = [H+] [A-] / [HA]
 Ka is called the acid dissociation constant
 The larger the Ka the stronger the acid
Calculating pH from Ka
 To calculate Ka we need to know the equilibrium
concentrations of all species.
 pH gives the concentration of H+
 We then use stoichiometric coefficients and an ICE box to
calculate the other concentrations
 Substitute and calculate
 See Ex 16.10 page 608
Using Ka to calculate pH
 Write the balanced equation
 Write the equilibrium expression. Look up value for Ka
 Use an ICE box to determine equilibrium concentrations. It is
customary to use X for the change in concentration of H+
 Then substitute in the equilibrium expression. This will yield a
quadratic equation. This can be simplified.
 If Ka is very small, assume that X is negligible compared to the
initial concentration of HA.
 Calculate X. If it is less than 5% of the initial concentration the
assumption is OK. If greater than 5% use the quadratic formula
to solve for X
 See Ex. 16.11 page 610-611
Polyprotic Acids
 Polyprotic acids have more than one ionizable proton. Ie.
Sulfuric or phosphoric acid
 The protons are removed in successive steps.
 It is much easier to ionize the first proton. Ka1 Ka2Ka3
 The majority of H+ comes from the first ionization. If
successive Ka values differ by 103 or more , pH can be
calculated from Ka1
Weak Bases
 Weak bases remove protons from water
 There is an equilibrium between the base and the resulting ions
 Weak base + H2O  conjugate acid + OH Example : NH3 + H2O  NH4+ + OH Kb = [NH4+] [OH-] / [NH3] the larger Kb , the stronger the base
Relationship between Ka and Kb
 Consider the following : NH4+  NH3+ + H+ and NH3 + H2O
 NH4+ + OH Ka = [NH3+] [H+] / [ NH4+]
Kb= [NH4+] [OH-] / [NH3]
 If we add the 2 equations we get H2O  H+ + OH- the
autoionization of water. Recall that Kw = [H+] [OH- ] = Ka X K b
 Or pKa + pKb = pKw = 14 at 25C
Acid-Base Properties of Salt Solutions
 Nearly all salts are strong electrolytes. Therefore they exist in
solution as ions.
 Acid-base properties are a function of the reaction of their ions
in solution
 Many salt ions can react with water to form H+ or OH- . This is
termed hydrolysis.
 Anions from weak acids are basic
 Anions from strong acids are neutral
 Anions with ionizable protons( HSO4-) are amphoteric
 All cations , except those of alkaline metals and heavier
alkaline earth metals, are weak acids
 pH of a solution may be predicted
o Salts derived from a strong acid and strong base are
neutral ie. NaCl, Ca(NO3)2
o Salts derived from a strong base and a weak acid are
basic ie Ba(C2H3O2)2
o Salts derived from a weak base and a strong acid are
acidic ie NH4Cl
o Salts from a weak acid and weak base can be either.
 Compare Ka and Kb to determine
Acid-Base Behavior and Chemical Structure
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Factors affecting acid strength
Consider H X
For this substance to be an acid the bond must be polar
The bond must be weak enough to be broken
The greater the stability of the conjugate base (X-) the more
acidic
Binary Acids
 The HX bond strength determines relative acid strength in any
group in the periodic table
 HX bond strength tends to decrease as X increases in size
 Relative acid strength increases down a group, base strength
decreases
 Bond polarity determines acid strength in a period
 Relative acid strength increases left to right across a period as
the electronegativity of X increases
 HF is a weak acid because the bond strength is extremely high
 CH4 is neither an acid or a base because the bond is non-polar
Oxyacids
 Many acids contain one or more OH bonds
 These are oxyacids of the form YOH
 The strength of the acid depends on the electronegativity of Y(
as it increases, more acid)
 A general trend is :
o If Y is a metal( low electronegativity) the substance is a
base
o If Y has intermediate electronegativity ( ie I) the
substance is a weak oxyacid
o If Y has high electronegativity(Cl) then it is a strong
oxyacid.
o As the number of O atoms attached to Y increase so does
the bond polarity and the strength of the acid
o HClO is weaker than HClO2 is weaker than HClO3
Carboxylic Acids
 There are a large class of acids that contain COOH
groups(carboxyl group)
 Why are they acidic?
o The additional O increases the polarity
Lewis Acids and Bases
 A Lewis acid is an electron pair acceptor; A base is an electron
pair donor
 Types of compounds that act as Lewis Acids
o Lewis acids have an incomplete octet(BF3)
o Transition metal ions
o Compounds with multiple bonds( CO2)
Hydrolysis of Metal Ions
 Hydrated metal ions act as acids
 Fe(H2O)6+3  Fe(H2O)5(OH)+2 + H+
 In general
o The higher the charge, more acidic
o The smaller the metal ion, the more acidic