Transcript No Slide Title

Acid Base Equilibria

Acids and bases are found in many
common substances and are important
in life processes.
 Group
Work: Make a list of some
common acids and bases. How do we
know which is which?
 There are several models for what
constitutes an acid or a base -- three
models to be discussed.
14.1 Acids and Bases: A Brief
Review

Acid:
tastes sour
stings skin
corrosive to metals
releases CO2 from carbonates
turns litmus red
turns phenolphthalein colorless

Base:
tastes bitter
feels slippery
turns litmus blue
turns ph. pink
React together to form a salt with loss of the
characteristic acid/base properties
Arrhenius Definition
Acids produce hydrogen ions (H+ ) in
aqueous solution.
 Bases produce hydroxide ions (OH-)
when dissolved in water.
 Limits to aqueous solutions.
 Only one kind of base.
 NH3 ammonia could not be an
Arrhenius base.

Arrhenius Definition

Not realistic: H+ has a radius of 10-13
cm, which gives a very concentrated
charge, so it associates with H2O as
H(H2O)4+, which we usually simplify to
H3O+ or H+ (aq)

OH– is also associated with H2O as
OH(H2O)3– which we usually write as
OH–(aq).
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.

H3O+ + Cl-

HCl(g) + H2O(l)

Water is a base makes hydronium ion
Acid/Base Pairs
General equation
 HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
 Acid + Base
Conjugate acid +
Conjugate base
 This is an equilibrium.
+
 Competition for H between H2O and A
 The stronger base controls direction.
+
 If H2O is a stronger base it takes the H
 Equilibrium moves to right.

Acid/Base Pairs – Group Work

Write a balanced equation showing how the
following substances behave as acids in
water and identify the conjugate acid-base
pairs.
HNO3





HCO3-
H3PO4
H2PO4-
HNO3(aq) + H2O(l) ⇌ H3O+ (aq) + NO3-(aq)
HCO3-(aq) + H2O(l) ⇌ H3O+ (aq) + CO32-(aq)
H3PO4(aq) + H2O(l) ⇌ H3O+ (aq) + H2PO4-(aq)
H2PO4-(aq) + H2O(l) ⇌ H3O+ (aq) + HPO42-(aq)
acid 1
base 2
acid 2
base 1
Strongest
acids
back
Weakest
acids
bases
Acid
HClO4
H2SO4
HI
HBr
HCl
HNO3
H3O+
HSO4H2SO3
H3PO4
HNO2
HF
CH3CO2H
H2CO3
H2 S
NH4+
HCN
HCO3HSH2 O
NH3
OH-
Base
ClO4HSO4IBrClNO3H2O
SO42HSO3H2PO4NO2FCH3CO2HCO3HSNH3
CNCO32S2 OHNH2O2-
Weakest
bases
Strongest
Leveling Effect
All acids above H3O+ in the table are
strong acids, which dissociate completely
in aqueous solution.
 All bases below OH- in the table are
strong bases, which dissociate completely
in aqueous solution.
 The table can be used to make predictions,
based on the principle that the stronger
acid reacts with the stronger base to form
a weaker acid and a weaker base.

Predicting Acid-Base Reactions


HCl + HSO3- ⇌ H2SO3 + Clstronger stronger
acid

base
weaker weaker
acid
base
We must also consider H2O as a
possible acid or base. Thus, HNO3 will
transfer its proton to H2O, not to Clbecause H2O is a stronger base than Cl-.
Group Work

Write an equation showing the position
of equilibrium for the following
mixtures. Remember that H2O can also
be either an acid or a base.
HSO4- and FHS- and HCO3-


HSO4- + F- ⇌ SO42- + HF
HS- + HCO3- ⇌ H2S + CO32-
14.2 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


HA(aq)
H+(aq) + A-(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.

Back to Pairs

Strong acids

Weak acids

Ka is large

Ka is small
[H+] is equal to
[HA]
 A- is a weaker
base than water

[H+] <<< [HA]
 A- is a stronger
base than water

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
 2H2O(l)
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-]

14.3 pH Scale
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

Measuring pH

litmus or pH paper

color changes of indicators
voltage generated by
electrodes (pH meter)

pH Indicators
Relationships
KW = [H+][OH-]
 -log KW = -log([H+][OH-])
 -log KW = -log[H+]+ -log[OH-]
 pKW = pH + pOH
So…

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 through equilibrium relationships (Kw)
[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-]
14.5 pH and Ka of Acid 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. Apply good chemistry!

Strong Acid: HNO3

HNO3 is completely dissociated into the
ions, H3O+ and NO3- (work sample 14.7)
Strong Acids

HBr, HI, HCl, HNO3, H2SO4, HClO4
ALWAYS WRITE THE MAJOR
SPECIES
 Completely dissociated
 [H+] = [HA]i
 [OH-] is going to be small because of
equilibrium

Kw = 10-14 = [H+][OH-]
Changes in pH with Dilution

Group Work: pH for factors of 10 dilution?

What is the pH of 1.0 M HCl?
pH = 0.00
What is the pH of 0.10 M HCl (a 1:10 dilution)?
pH = 1.00
What is the pH of 0.010 M HCl?
pH = 2.00





Changes in pH with Dilution
What is the pH of 1.0 x 10-3 M HCl?
 pH = 3.00
 What is the pH of 1.0 x 10- 4 M HCl?
 pH = 4.00
 What is the pH of 1.0 x 10-5 M HCl?
 pH = 5.00
 What is the pH of 1.0 x 10-6 M HCl?
 pH = 5.996

Changes in pH with Dilution






What is the pH of 1.0 x 10-7 M HCl?
pH = 6.791
What is the pH of 1.0 x 10-8 M HCl?
pH = 6.996
Why does the pH stop changing at a value of
about 7?
Water has a pH of 7 due to autodissociation,
so it is never possible to get a pH higher than
7 by addition of water.
– If [HA] < 10-7 water contributes the H+
14.5 pH of Weak Acids


Except for the strong acids, most acids
do not ionize completely. These acids
are called weak acids.
HF(aq) + H2O(l) ⇌ H3O+(aq) + F-(aq)
14.5 pH of 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
 Rest is just like chapter 13.

Example


Calculate the pH of 2.0 M acetic acid
HC2H3O2 with a Ka 1.8 x10-5
Calculate pOH, [OH-], [H+]
A mixture of Weak Acids
The process is the same.
 Determine the major species.
 The stronger acid will predominate.
 Bigger Ka if concentrations are
comparable


Calculate the pH of a mixture 1.20 M
HF (Ka = 7.2 x 10-4) and 3.4 M HOC6H5
(Ka = 1.6 x 10-10)
Some Weak Acids
Percent dissociation
amount dissociated
x 100
initial concentration
 For a weak acid percent dissociation
increases as acid becomes more dilute.
 Calculate the % dissociation of 1.00 M
and 0.100 M Acetic acid (Ka = 1.8 x 10-5


=
As [HA]0 decreases [H+] decreases but
% dissociation increases.
– Le Chatelier principle with dilution (pg 642)
The other way
What is the Ka of a weak acid that is
8.1 % dissociated as 0.100 M solution?
 Compare to Sample Exercise 14.11

pH and Ka
One way to find the value of Ka for a weak
acid is using concentration & pH data.
 The pH of 0.500 M HNO2 is 1.827. What is
Ka of HNO2?
 [H3O+] = 10-1.827 = 0.0149 M

HNO2 + H2O  H3O+ + NO2I 0.500
-0
0
C -0.0149
+0.0149 +0.0149
E 0.485
0.0149 0.0149
 Ka = [H3O+][NO2-]/[HNO2]
 Ka = (0.0149)2/0.485 = 4.58 x 10-4

14.6 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.
 NH3 + H2O
NH4+ + OH It 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- ]
[B]
Strength of Bases
Hydroxides are strong.
 Others are weak.
 Smaller Kb = weaker base.


Consider a solution of 4.0 M pyridine
(Kb = 1.7 x 10-9)
N:
– What are the major species present in solution?
– Determine the [OH-] and pH of the solution.
14.7 Polyprotic acids
Always dissociate stepwise.
 Denoted Ka1, Ka2, Ka3, etc… (Table
14.4)
 H2CO3 + H2O
H+ + HCO3

HCO3- + H2O
Ka1= 4.3 x 10-7
H+ + CO3-2
Ka2= 4.3 x 10-10

Conjugate base in 1st step is acid in 2nd.
Polyprotic Acid Observations
The first H+ comes of much easier than
the second.
 Ka for the first step is much bigger than
Ka for the second.

Calculate the Concentration…
…of all the ions in a solution of 5.00 M
sulfurous acid.
Ka1 = 1.5 x 10-2
Ka2 = 1.0 x 10-7
 Find the pH of the solution


See Sample Exercise 14.15 for another
example.
Sulfuric acid, an interesting case

Calculate the pH and concentration of all
species in a 1.5 M solution of H2SO4
– Pg 651, Table 14.4 contains Ka’s for polyprotics
Calculate the concentrations in a
1.5 x 10-2 M solution of H2SO4
 See Sample Exc. 14.16 & 17
 Summary pg. 655

– In first step H2SO4 is a strong acid.
– 2nd step, it is a weak acid: Ka2 = 1.2 x 10-2
14.8 Acid/Base Properties of Salts

Salts are ionic compounds.
– Recall Model: Ionic compounds dissociate into
ions in solution. The ions move around
independently.
These cations and anions can (but do
not have to!) act as acids or bases
depending on how they react with
water (referred to as a hydrolysis
reaction)
 Ch 14, #99 discuss on WB’s

Salts that make Neutral Solutions
Salts containing the cation from a
strong base and the anion from a strong
acid are neutral.
 Hydrolysis is not observed with ions
derived from strong acids or bases:

Cations of group I and II (except Be2+)
Anions: Cl-, Br-, I-, NO3-, ClO4-
for example NaCl, KNO3
 There is no equilibrium established
from these ions.

Basic Salts
If the anion of a salt is the conjugate base
of a weak acid the solution will be basic.
 Consider an aqueous solution of NaF
 The aqueous species are Na+, F-, and H2O
 F- + H2O
HF + OH The equilibrium constant for hydrolysis
is just a Ka or Kb, depending on the type
of hydrolysis.

Basic Salts

Kb =[HF][OH-]
[F- ]
but Ka = [H+][F-]
[HF]
 Ka x Kb = [HF][OH-]
[F- ]

x [H+][F-]
[HF]
Basic Salts

Ka x Kb = [HF][OH-]
[F- ]
x [H+][F-]
[HF]
Basic Salts

Ka x Kb = [HF][OH-]
[F- ]
x [H+][F-]
[HF]
Basic Salts
Ka x Kb = [HF][OH-]
[F- ]
 Ka x Kb =[OH-] [H+]

x [H+][F-]
[HF]
Basic Salts
Ka x Kb = [HF][OH-]
[F- ]
 Ka x Kb =[OH-] [H+]
 Ka x Kb = Kw

x [H+][F-]
[HF]
pH of Basic Salts
The anion of a weak acid is a weak base.
 The CN- ion competes with OH- for the
H+
 Relative base strength:
OH- > CN- > H2O
 Calculate the pH of a solution of 1.00 M
NaCN.
 Ka of HCN is 6.2 x 10-10

Acidic salts
If the cation of a salt is the conjugate acid
of a weak base the solution will be acidic.
 The same development as bases leads to


Ka x Kb = KW
Calculate the pH of a solution of 0.40 M
NH4Cl (the Kb of NH3 1.8 x 10-5).
 Other acidic salts are those of highly
charged metal ions.

Acidic Salts – Metal ions




Ions are often modified when dissolved in
solution. Examine the photos of Fe(III)
salts and solutions. Why do they have
different colors?
Fe(NO3)3.6H2O contains pink Fe(H2O)63+
Solutions may hydrolyze to give yellow
Fe(H2O)5OH2+ or even reddish brown
Fe(H2O)3(OH)3
FeCl3.6H2O contains ions such as yellow
Fe(H2O)5Cl2+
Hydrolysis of Metals

Hydrolysis is more
important for more
highly charged ions
and smaller ions.

Hydrogen from
water easier to
remove.
Hydrolysis of Metals

Table below gives values of Ka for metal ions:
Ion
Radius
Ka
Na+
Li+
Be2+
Mg2+
Ba2+
Cr3+
Zr4+

95 pm
60 pm
31 pm
65 pm
135 pm
69 pm
78 pm
3.3 x 10-15
1.5 x 10-14
3.2 x 10-7
3.8 x 10-12
1.5 x 10-14
9.8 x 10-5
6.0 x 10-1
Greater values of Ka for ions with larger charge
and smaller size.
Hydrolysis of Metals

Calculate the pH of a 0.20 M CrCl3
solution. (From the table, the Ka value
for Cr3+ is 9.8 x 10-5)
Anion of weak acid, cation of weak base
If both of the ions have acid/base
properties, compare the Ka of the cation
to the Kb of the anion
 Ka > Kb
acidic
 Ka < Kb
basic
 Ka = Kb
neutral
 Predict whether an aqueous solution of
ammonium cyanide will be acidic,
basic, or neutral.
 Summary Table 14.6, pg 660

14.9 Structure and Acid base
Properties


Draw lewis structures (that obey octet rule) for
the following oxyacids.
HClO4
Ka = Large (~107)
HClO3
Ka = ~1
HClO2
Ka = 1.2 x 10 –2
HClO
Ka = 3.5 x 10 –8
Can you explain the relative strengths of these
acids using your knowledge of atomic structure
and bonding?
Strength of oxyacids
Electron Density
O
O
O
Cl
O
H
Strength of oxyacids
Electron Density
O
Cl
O
O
H
Strength of oxyacids
Electron Density
O
Cl
O
H
Strength of oxyacids
Electron Density
Cl
O
H
Strength of oxyacids
The more oxygen bonded to the central
atom, the more acidic the hydrogen.
 The oxygens are electronegative
 Pulls electron density away from
hydrogen
 HClO4 > HClO3 > HClO2 > HClO

– Remember that the H is attached to an oxygen
atom.
14.9 Structure and Acid base
Properties


Draw lewis structures (that obey octet rule) for
the following oxyacids.
HOCl
Ka = 4 x 10 –8
HOBr
Ka = 2 x 10 –9
HOI
Ka = 2 x 10 –11
HOCH3
Ka = ~ 10 –15
Can you explain the relative strengths of these
acids using your knowledge of atomic structure
and bonding?
14.9 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.

14.10 Acid-Base Properties of
Oxides
Non-metal oxides dissolved in water
can make acids.
 SO3 (g) + H2O(l)
H2SO4(aq)
 Ionic oxides dissolve in water to
produce bases.
 CaO(s) + H2O(l)
Ca(OH)2(aq)

14.11 Lewis Acids and Bases
Most general definition.
 Acids are electron pair acceptors.
 Bases are electron pair donors.

F
H
B
F
F
:N
H
H
Lewis Acids and Bases

Boron triflouride wants more electrons.
F
H
B
F
F
:N
H
H
Lewis Acids and Bases
Boron triflouride wants more electrons.
 BF3 is Lewis acid NH3 is a Lewis base.

H
F
F
F
B
N
H
H
Lewis Acids and Bases
Al+3+ 6
Al
( )
( )
H
O
H
H
O
H
6
+3