Transcript Topic B: Le Chatelier’s Principle and Optimum Conditions

Topic C – Part I:
Acid Base Equilibria and Ksp
Arrhenius Definition
• Acids produce hydrogen ions (H+) in aqueous solution.
• Bases produce hydroxide ions (OH-) when dissolved in
water.
• Only one kind of base.
• NH3 ammonia could not be an Arrhenius base.
Bronsted-Lowry Definitions
• And acid is an proton (H+) donor and a base is a proton
•
•
•
•
•
acceptor.
Acids and bases always come in pairs.
HCl is an acid.
When it dissolves in water it gives its proton to water.
HCl(g) + H2O(l)
H3O+ + ClWater is a base makes hydronium ion
Conjugate Pairs
 General equation
 HA(aq) + H2O(l)






H3O+(aq) + A-(aq)
Conj. acid + Conj. base
Acid + Base
This is an equilibrium.
Competition for H+ between H2O and AThe stronger base controls direction.
If H2O is a stronger base it takes the H+
Equilibrium moves to right.
Acid dissociation constant Ka
• The equilibrium constant for the general equation.
• HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
• Ka = [H3O+][A-]
•
[HA]
• H3O+ is often written H+ ignoring the water in equation (it
is implied).
Acid dissociation constant Ka
H+(aq) + A-(aq)
• HA(aq)
• Ka = [H+][A-]
•
[HA]
• We can write the expression for any acid.
• Strong acids dissociate completely.
• Equilibrium far to right.
• Conjugate base must be weak.
Sample Ex. 14.1
 Acid Dissociation (ionization) Reactions
 Write the simple dissociation reaction (omitting





water) for each of the following acids.
Hydrochloric Acid (HCl)
Acetic Acid (HC2H3O2)
The Ammonium Ion (NH4+)
The anilinium ion (C6H4NH3+)
The hydrated aluminum (III) ion [Al(H2O)6]3+
Describing Acid Strength
 Strong acids
 Weak acids
 Ka is large
 Ka is small
 [H+] <<< [HA]
 [H+] is equal to [HA]
 A- is a weaker
base than water
 A- is a stronger
base than water
Sample Ex. 14.2
• Using table 14.2 (p. 628) arrange the following species
according to their strength as bases: H2O, F-, Cl-, NO2 and
CN-
• The concentration of H3O+ and OH- are dependent upon
two, separate factors
• 1. The strength of the acid or base
• (the degree of ionization)
• 2.The amount of water present
• (the concentration of the solution)
• . It is possible to have a dilute, (but) strong acid
•
or to have a concentrated, (but) weak acid
• They might also have the same hydronium ion
concentration and pH
Examples of Acids and Bases
Types of Acids
• Polyprotic Acids- more than 1 acidic hydrogen (diprotic,
triprotic).
• Oxyacids - Proton is attached to the oxygen of an ion.
• Organic acids contain the Carboxyl group -COOH with the
H attached to O
• Generally very weak.
Amphoteric
 Behave as both an acid and a base.
 Water autoionizes
H3O+(aq) + OH-(aq)
KW= [H3O+][OH-]=[H+][OH-]
At 25ºC KW = 1.0 x10-14
In EVERY aqueous solution.
Neutral solution [H+] = [OH-]= 1.0 x10-7
Acidic solution [H+] > [OH-]
Basic solution [H+] < [OH-]
 2H2O(l)






Sample Ex. 14.3
• Calculate [H+] or [OH-] as required for each of the
following solutions at 25oC, and state whether the solution
is neutral, acidic, or basic.
• 1.0 x 10-5 M OH• 1.0 x 10-7 M OH• 10.0 M H+
pH
 pH= -log[H+]
 Used because [H+] is usually very small
 As pH decreases, [H+] increases exponentially
 Sig figs only the digits after the decimal place of a pH
are significant
 [H+] = 1.0 x 10-8 pH= 8.00 2 sig figs
 pOH= -log[OH-]
 pKa = -log K
Sample Ex. 14.5
• Calculating pH and pOH
• Calculate pH and pOH for each of the following solutions
at 25oC.
• a. 1.0 x 10-3 M OH• b. 1.0 M H+
Relationships
• KW = [H+][OH-]
• -log KW = -log([H+][OH-])
• -log KW = -log[H+]+ -log[OH-]
• pKW = pH + pOH
• ( KW = 1.0 x10-14 )
• 14.00 = pH + pOH
• [H+],[OH-],pH and pOH
• Given any one of these we can find the other three.
Sample Ex. 14.6
• The pH of a sample of human blood was measured to be
7.41 at 25oC. Calculate pOH, [H+], and [OH-] for the
sample.
[H+]
100 10-1 10-3 10-5 10-7 10-9 10-11 10-13 10-14
pH
0
1
Acidic
14 13
10-14 10-13
3
11
5
7 9
Neutral
9
7 5
11
3
13
14
Basic
1
0
pOH
10-11 10-9Basic
10-7 10-5 10-3 10-1 100
[OH-]
Calculating pH of Solutions
• Always write down the major ions in solution.
• Remember these are equilibria.
• Remember the chemistry.
• Don’t try to memorize there is no one way to do this.
Strong Acids
• HBr, HI, HCl, HNO3, H2SO4, HClO4
• ALWAYS WRITE THE MAJOR SPECIES
• Completely dissociated
• [H+] = [HA]
• [OH-] is going to be small because of equilibrium
• 10-14 = [H+][OH-]
• If [HA]< 10-7 water contributes H+
Sample Ex. 14.7
• a. Calculate the pH of 0.10 M HNO3
• b. Calculate the pH of 1.0 x 10-10 M HCl
Weak Acids
 Ka will be small.
 ALWAYS WRITE THE MAJOR SPECIES.
 It will be an equilibrium problem from the start.
 Determine whether most of the H
+ will come from the
acid or the water.
 Compare Ka or Kw
 The rest is just like previous equilibrium problems
Sample Ex. 14.8
 The pH of Weak Acids
 The hypochlorite ion (OCl-) is a strong oxidizing agent often found
in household bleaches and disinfectants. It is also the active
ingredient that forms when swimming pool water is treated with
clorine. In addition to its oxidizing abilities, the hypochorite ion has
a relatively high affinity for protons (it is a much stronger base than
Cl-, for example) and forms the weakly acidic hypochlorous acid
(HOCl, Ka = 3.5 x 10-8).
 Calculate the pH of a 0.100 M aqueous solution of
hypochlorous acid.
A mixture of Weak Acids
 The process is the same.
 Determine the major species.
 The stronger will predominate.
 Bigger Ka if concentrations are comparable
 Sample Ex. 14.9:
Calculate the pH of a
mixture 1.00 M HCN (Ka = 6.2 x 10-10) and 5.00
M HNO2 (Ka = 4.0 x 10-4)
 Calculate the concentration of cyanide ion (CN-)
Bases
-
• The OH is a strong base.
• Hydroxides of the alkali metals are strong bases because
they dissociate completely when dissolved.
• The hydroxides of alkaline earths Ca(OH)2 etc. are strong
dibasic bases, but they don’t dissolve well in water.
-
• Used as antacids because [OH ] can’t build up.
Bases without OH
 Bases are proton acceptors.
NH4+ + OHIt is the lone pair on nitrogen that accepts the proton.
Many weak bases contain N
B(aq) + H2O(l)
BH+(aq) + OH- (aq)
Kb = [BH+][OH- ]
 NH3 + H2O





[B]
Strength of Bases
• Hydroxides are strong.
• Others are weak.
• Smaller Kb weaker base.
• Calculate the pH of a solution of 4.0 M pyridine
(Kb = 1.7 x 10-9)
N:
Sample Ex. 14.13
The pH of Weak Bases
• Calculate the pH for a 15.0 M solution of NH3 (Kb =
1.8 x 10-5)
Polyprotic acids
• Always dissociate stepwise.
+ comes off much easier than the second.
• The first H
• Ka for the first step is much bigger than Ka for the second.
• Denoted Ka1, Ka2, Ka3
Polyprotic acid
H+ + HCO3Ka1= 4.3 x 10-7
• HCO3H+ + CO3-2
Ka2= 4.3 x 10-10
• Base in first step is acid in second.
• In calculations we can normally ignore the second
dissociation.
• H2CO3
Sulfuric acid is special
• In first step it is a strong acid.
• Ka2 = 1.2 x 10
-2
• Calculate the concentrations in a 1.0 M solution of
H2SO4 (Ex. 14.16)
• Calculate the concentrations in a 1.0 x 10
H2SO4 (Ex. 14.17)
-2 M solution of
Structure and Acid base Properties
• Any molecule with an H in it is a potential acid.
• The stronger the X-H bond the less acidic (compare
bond dissociation energies).
• The more polar the X-H bond the stronger the acid (use
electronegativities).
• The more polar H-O-X bond -stronger acid.
Strength of Carboxylic Acid
• The strength of a carboxylic acid depends upon the
stability of the anion that it forms when it loses its labile
(mobile proton).
• All carboxylic acids form this equilibrium, where ‘R’ can
vary.
•
RCOOH   RCOO- + H+
• If the anion (RCOO-) is relatively stable
• tends to donate H+ ions and the acid will be relatively strong
• The table below indicates that the methanoate ion is more
stable than the ethanoate ion, which is more stable than
the propanoate ion.
• (the smaller the pKa the stronger the acid)
• The ethanoate and propanoate ions have electron-
releasing alkyl groups (CH3 and C2H5) that ‘pump’
electrons into the COO- anion and make it reactive and
unstable.
• As a result, the H+ ions tend not to be released.
• This ‘pumping’ of electrons is called the inductive effect.
• Methanoic acid on the other hand does not have an alkyl
group so it does not push electrons into the COO–, the
anion remains relatively stable.
• A similar effect is observed when comparing the relative
strengths of halogen substituted carboxylic acids.
• It is found that as the number of chlorine atoms increases,
so does the strength of the acid.
Strength of oxyacids
• The more oxygen hooked to the central atom, the more
acidic the hydrogen.
• HClO4 > HClO3 > HClO2 > HClO
• Remember that the H is attached to an oxygen atom.
• The oxygens are electronegative
• Pull electrons away from hydrogen
Strength of oxyacids
Electron Density
Cl
O
H
Strength of oxyacids
Electron Density
O
Cl
O
H
Strength of oxyacids
Electron Density
O
Cl
O
O
H
Strength of oxyacids
Electron Density
O
O
O
Cl
O
H
Hydrated metals
• Highly charged
metal ions pull the
electrons of
surrounding water
molecules toward
them.
• Make it easier for
H+ to come off.
H
Al+3 O
H
Acid-Base Properties of Oxides
• Non-metal oxides dissolved in water can make acids.
(Covalent S-O bonds, polar O-H bonds)
•
SO3(g) + H2O(l)
H2SO4(aq)
• Ionic oxides dissolve in water to produce bases.
• (Ionic Ca-O bonds)
•
CaO(s) + H2O(l)
Ca(OH)2(aq)