Transcript Chapter 16 : Acid-Base Equilibira

Chapter 16 :
Acid-Base Equilibria
Created by Lauren Querido
Table of Contents
16.1 Review
16.2 Brønsted-Lowry
Acids and Bases
16.3 Autoionization of
Water
16.4 pH Scale
16.5 Strong Acids and
Bases
16.6 Weak Acids
16.7 Weak Bases
16.8 Relationship
Between Ka and Kb
16.9 Acid-Base Properties
of Salt Solutions
16.10 Acid-Base Behavior
and Chemical
Structure
16.11 Lewis Acids and
Bases
16.1 Review
• Acids
– Sour in taste
– Litmus paper turns red
• Bases
– Bitter, slippery
– Litmus paper turns blue
• When acids and bases
mix, their properties
disappear!
Arrhenius Acids and Bases
• Svante Arrhenius
(1880)
– In aqueous solutions:
• Acids will increase the
concentration of H+ ions
when dissolved in
water.
• Bases will increase the
concentration of OHions when dissolved in
water.
16.2 Brønsted-Lowry Acids and Bases
• 1923 Brønsted and Lowry
made a more general definition
– Brønsted-Lowry Acid is a
substance that can transfer a
proton. It must have a
hydrogen atom that can be
lost as H+.
– Brønsted-Lowry Base is
asubstance that can accept a
proton. Must have a
nonbonding pair of
electrons to gain a H+ ion.
Conjugate Acid-Base Pairs
• Conjugate base- Removal of proton from the acid
• Conjugate acid- Addition of proton to the base
Relative Strengths of Acids and Bases
• The stronger the acid, the weaker its conjugate base.
• The stronger the base, the weaker its conjugate acid.
1. Strong acids completely transfer protons to water.
2. Weak acids partly dissociate in aqueous solutions and exist
as a mixture of acid molecules and component ions.
3. Negligible acidity contain Hydrogen but do not
demonstrate acidic behavior. Ex: CH4
• Position of equilibrium favors transfer of proton from
stronger acid to stronger base.
16.3 Autoionization of Water
• Ion product of water
–
=
– 1.0 x 10 –14 = [H+] [OH-]
– This is used to calculate concentrations of H+
and OH- .
• If [H+] = [OH-], than neutral equation
• If [H+] > [OH-], than acidic equation
• If [H+] < [OH-], than basic equation
16.4 The pH Scale
•
•
•
•
•
pH = -log [H+]
pH of 7 is neutral
Acidic solution 0 < pH < 7
Basic solution 14 > pH > 7
Other p scales are
– pOH = -log [OH-]
– pOH + pH = -log Kw = 14.0
Examples on the pH Scale
Measuring pH
• A pH meter consists
of a pair of electrodes
connected to a meter
which pH is generated
when placed in the
solution.
• An acid-base
indicator turns a
color if placed in acid
or base. Ex: litmus
paper
16.5 Strong Acids and Bases
• Strong Acids
– 7 most common strong acids are
• HCl, HBr, HI, HNO3, HClO3, HClO4, and H2SO4
– In acidic reactions, equilibrium lies entirely to
the right side.
– Completely dissociates
– Example:
• HNO3 => H+ + NO3-
Strong Bases
• Most common strong bases are ionic hydroxides
of alkali metals (1A) and heavier alkaline earth
metals (2A). Examples: LiOH, RbOH, CsOH,
NaOH, KOH, and Ca(OH)2, Sr(OH)2, and
Ba(OH)2.
• Other strong bases react with water to form OHsuch as Na2O, CaO.
• Also, anions O2-, H-, and N3- are stronger bases
than OH- and therefore remove a proton from
H2O.
– Example: N3- + H2O => NH3 + 3OH-
16.6 Weak Acids
• A weak acid only partially ionizes in aqueous
solutions.
• General weak acid equation
– HX  H+ + X- where H is Hydrogen
– Many weak acids contain some Hydrogen atoms
bonded to carbon atoms and oxygen atoms (organic
compounds).
– Ka is the acid dissociation constant.
– The larger the value of Ka , the stronger the acid.
Calculating Ka from pH
• Use and ICE box!
• Sample exercise
– A student prepared a .10 M solution of formic acid and
measures its pH which was 2.38.
• A) calculate Ka for formic acid
• B) what percentage of the acid is ionized in the .10M solution?
Answer
a) HCHO2  H+ + CHO2Ka = [H+][CHO2-]
[HCHO2]
pH= -log[H+]
=10 –2.38 = 4.2 X 10-3M
HCHO2

H+
Ka = [4.2 X 10–3 ][4.2 X 10–3]
[.10]
1.8 X 10-4 = [4.2 X 10–3 ][4.2 X 10–3]
[.10]
b) Percent Ionization =
Concentration of H+
Initial concentration of
component
CHO
= 4.2%
I
.10 M
0M
0M
C
-4.2 X 10–3
+4.2 X 10–3
+4.2 X 10–3
E
.10 - 4.2 X 10–3
+4.2 X 10–3
+4.2 X 10–3
Using Ka to Calculate pH
The best way to explain this is by an example.
Calculate the pH of a .30 M solution of
acetic acid at 25o C. (Ka = 1.8 X 10-5)
So… HC2H3O2  H+ + C2H3O2Ka = [H+][C2H3O2-] = 1.8 X 10-5
[HC2H3O2]
What now?
I
C
E
HC2H3O2

H+
C2H3O2-
.30 M
-x
.30-x
0M
+x
x
0M
+x
x
Ka = (x)(x) = 1.8 X 10-5
(.30 – x)
Either do the quadratic equation or in this case
you can take out x in the denominator.
[H+] = x = 2.3 X 10-3
pH = -log 2.3 X 10-3 = 2.64
Polyprotic Acids
• Polyprotic acids have more than one
ionizable Hydrogen atom.
• Example:
• H2SO3  H+ + HSO4• HSO4-  H+ + SO32-
• The second Ka (Ka2) is much smaller than
Ka1 because it is easier to remove the first
proton.
16.7 Weak Bases
•
•
•
Weak base + water => conjugate acid + hydroxide
ion
Kb is the base-dissociation constant (equilibrium in
which base reacts when H2O to form conjugate acid
and OH- ion).
Types of weak bases:
1. Neutral substances that have atoms with a non-bonding
pair of electrons that can serve as a proton acceptor.
–
Most of these contain amines, N-H which is sometimes replaced
with a bond between C or N Ex: NH2CH3
2. Anions of weak acids
–
Ex: ClO- + H2O  HClO + H+
–
ClO- is the weak base
16.8 Relationship Between Ka and Kb
• Reaction 1 + reaction 2 = reaction 3
Which leads to K1 x K2 = K3
Which leads to Ka x Kb = Kw
• Kw is the ion-product constant for water
– Kw = 1 x 10-14
• As the strength of the acid increases, the strength of
the base decreases and visa-versa.
• pKa + pKb = pKw = 14.00
16.9 Acid-Base Properties of Salt Solutions
• Hydrolysis is the process at which ions react with
water and produce H+ or OHX- + H2O HX + OH-
• Anions of strong acids do not influence pH
– Ex: NO3-
• Anions that still have ionizable protons are
amphoteric
– Ex: HSO3- from H2SO4
• Most cations (except 1A elements and Ca+2, Sr+2.
Ba+2) act as weak acids in solution.
Predicting the pH of a Solution
1. Salts derived from a strong acid and a strong base makes a
neutral pH (pH of 7).
• NaOH + HCl => NaCl + H2O
2. Salts derived from a strong base and a weak acid will yield
a pH of above 7 because the anion hydrolyzes to produce
OH- ions and the cation does not hydrolyze.
• NaOH + HClO => NaClO + H2
3. Salts derived from a weak base and a strong acid will result
in a pH that is below 7 because the cation hydrolyzes to
produce H+ ions and the anion does not hydrolyze.
• Al(OH)3 + 3HNO3 => Al(NO3)3 + 3H2O
4. Salts derived from a weak base and a weak acid will
yield a pH that is dependant on the constant value of
the constant dissociations (Ka and Kb).
•
if the base is more basic than the acid is acidic, then the
solution will have a pH that is greater than 7.
• If the acid is more acidic, than the pH will be less than 7.
•
•
•
•
NH4+ + CN-  NH4CN
NH4+
Ka = 5.6 X 10-10
CNKb= 2.0 X 10-5
Therefore, the pH of NH4CN is greater than 7
16.10 Acid-Base Behavior and
Chemical Structure
• Factors that effect acid strength
– If H-X bond is polarized (X is more electronegative)
the H acts as a proton acceptor.
– Non-polar bonds (CH4) produce neutral solutions.
– Weaker bonds dissociate more easily than very strong
bonds.
– HF is a weak acid because of this.
– The greater the stability of the conjugate base, the
weaker the acid.
– Ultimately, there are three factors effecting acid strength:
• Polarity of H-X bond
• Strength of H-X bond
• Stability of conjugate base, X-
Binary Acids
• Binary acids are composed of Hydrogen and a
non-metal.
– Ex: HCl, HF, H2S, etc.
• The more polar the bond,the stronger it is
• The weaker the bond, the stronger the acid.
• Strength of the bond decreases (acidity increases)
as the element increases in size or moves down a
group.
• Acid strength increases (acidity decreases)
moving from left to right
CH4
NH3
HF
No acid or base
properties
Weak base
SiH4
PH3
H2S
HCl
No acid or base
properties
Weak base
Weak acid
Strong acid
H2O
-------
Increasing acid strength
Increasing base strength
Weak acid
Increasing base strength
Period 3
7A
Increasing acid strength
Period 2
4A
Group
5A
6A
Oxyacids
• Oxyacids are acids with an OH group is
bound to a central atom.
– Example: H2SO4
–
OH- Bonding
• To determine if an OH group acts as an acid or base,
consider this:
• If Y is a metal than sources of OH- behave as bases.
• If Y is a non-metal than the compound will not readily lose
the OH- ion.
– The electronegativity will increase and so will the
acidity.
• The increasing number of Oxygen atoms stabilizes the
conjugate base and thus increases the strength of the acid.
Oxyacid Rules of Thumb
1. Oxyacids that have the same number of OH
groups and the same number of Oxygen atoms,
acid strength increases with increasing
elecronegativity of the central atom
•
Example: HClO > HBrO > HIO (> = more acidic)
For oxyacids with the same central atom, acid
strength increases with increasing number of
Oxygen atoms that are attached.
2.
•
Example: HClO < HClO2 < HClO3 < HClO4 ( < less acidic)
Carboxylic Acid
• Carboxylic acids are organic compounds.
•
•
-COOH is the functional group
-R is either a Hydrogen
or Carbon based group
– If an extra Oxygen is added than it stabilizes the
conjugate base and increases the acidity.
– If conjugate base has resonance structures, it spreads
the negative charge evenly over the compound.
– Acid strength of carboxylic acid increases as the
number of electronegative atoms increase.
6.11 Lewis Acids and Bases
• G.N. Lewis proposed this:
• Lewis Acids have an incomplete
octet of electrons. Function as
electron pair acceptors
• Lewis Bases act as electron pair
donators
Hydrolysis of Metal Ions
• Hydration is a process when when metals
attract unshared electron pairs of water
molecules.
–
–
–
–
The metal acts as Lewis acid
The water acts as Lewis base
Ex: Fe(H2O)6+3  Fe(H2O)5(OH)2+ + H+
So, general equation
• M(H2O)nc  M(H2O)n-1(OH)c-1 + H+