CATALYST
IF YOU ADD 1.5 MOLS OF H 2 AND 0.38
MOLS OF N 2 TO A 2.25 L CONTAINER TO
SYNTHESIZE AMMONIA AND AT
EQUILIBRIUM HAVE 0.12 MOLS OF
AMMONIA…
• WHAT ARE THE CONCENTRATIONS
OF HYDROGEN AND
NITROGEN AT
EQUILIBRIUM?
• WHAT IS THE EQUILIBRIUM
CONSTANT?
• WHAT IS THE K P ?
PH SCALE
THE PH SCALE
Because [H+] values are typically VERY
small, the pH scale provides and easy way
to represent acidity.
pH = -log[H+]
Thus… for [H+] = 1.0 x 10-7
The pH = -log (1.0 x 10-7) = 7.00
THE POH SCALE
Similarly, the pOH scale provides and easy
way to represent basicity.
pH = -log[OH-]
Also…
pK = -log(K)
PH SIG FIGS!
[H+] = 1.0 x 10-7 : 2 sigfigs
pH = -log (1.0 x 10-7) = 7.00 : 2 decimal
places after the #
EXAMPLE
Calculate the pH and pOH for each of the
following solutions
•1.0 x 10-3 M OH•1.0 M H+
CONSIDER…
It is useful to consider the log form of
the expression:
Kw = [H+][OH-]
…
pKw = pH + pOH
…
pH + pOH = 14
EXAMPLE
The pH of a sample of human blood was
measured to be 7.41 at 25 °C. Calculate the
pOH, [H+], [OH-] for the sample.
BOOK
DISCLAIMER!
THINK CHEMISTRY!
Focus on the solution components and
their reactions. It will almost always be
possible to choose one reaction that is the
most important!
BE SYSTEMATIC
Acid-Base problems require a step-by-step
approach!
BE FLEXIBLE!
Although all acid-base problems are
similar in many ways…important
differences occur…
Teach each problem mas a separate entity.
Don’t try to force a given problem into
matching any you have solved before!
BE PATIENT
Pick apart the problem into workable steps
The complete solution to a complex
problem may not be immediately apparent
BE CONFIDENT
Let the problem guide you
Assume you can think it out!
DO NOT rely on memorizing solutions to
problems
UNDERSTAND AND THINK! Don’t
memorize!
ACIDS
CALCULATING THE PH OF
STRONG ACID SOLUTIONS
Step 1: FOCUS on the MAJOR SPECIES (the solution
components present in large amounts!)
• Ex. HCl contains MOSTLY: H+, Cl-, and H2O and only
a little (tiny tiny bit of) HCl, and OHCalculate the pH of a 0.1 M HCl solution.
• Two equations at work:
• HCl  H+ + Cl• H2O  H+ + OH• Do both contribute H+ ions?
• HCl – definitely yes!
• H2O – probably NOT (H+ ions from HCl will drive this
reaction left!)
EXAMPLE
Calculate the pH of a 0.1 M HNO3 solution
Calculate the pH of a 1.0 x 10-10 M HCl solution.
(Don’t be fooled here)
CALCULATING THE PH OF
WEAK ACID SOLUTIONS
WORK
SYSTEMATICALLY!!!!!
We will develop the necessary skills by
working through an ex problem
• Calculate the pH of a 1.00 M solution of HF
(Ka = 7.2 x 10-4)
CALCULATE THE PH OF A 1.00
M SOLUTION OF HF
What is Step 1?
Since this is a pH problem…what
compounds can generate H+ ions?
•H2O  H+ + OHKw = 1.0 x 10-14
•HF  H+ + FKa = 7.2 x 10-4
Which source of H+ is dominant? (usually
we can single one out!)
We can tell because the Ka (although small
itself) is still MUCH bigger than Kw
CALCULATE THE PH OF A
1.00 M SOLUTION OF HF
HF will be contributing the most H+ ions, thus
we will ignore the contribution from H2O.
• Write the equilibrium expression for HF
• Write an I-C-E table for HF
• Substitute equilibrium expressions into
the K equation
Simplify if possible! Solve for “x”
HOW VALID IS THE APPROXIMATION
THAT [HF] = 1.00 M?
Typically Ka values are known to an accuracy
of only about ±5%. It is reasonable to apply
this figure when determining if we can
approximate:
Step 1: Calculate x by assuming [HA]0 – x ≈
[HA]0
Step 2: Compare the sizes of x and [HA]0
through the expression:
x
´100%
[HA]0
Step 3: if it is ≤ 5%, x is small enough!
STEPS TO SOLVING WEAK ACID
EQUILIBRIUM PROBLEMS
1.
List major species in solution
2.
Choose the species that can product H+ and write their balanced
equations
3.
Use the equilibrium constants for the reactions you have written,
decide which equilibrium will dominate in producing H+
4.
Write the equilibrium expression for the dominant equilibrium
5.
List the initial concentrations of the species participating in the
dominant equilibrium.
6.
Define the change needed, x
7.
Write the equilibrium concentrations in terms of x
8.
Solve for x by assuming [HA]0 – x ≈ [HA]0
9.
Use the 5% rule
10. Calculate [H+] and pH
EXAMPLE
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
chlorine. In addition hypochlorite has a relatively
high affinity for protons (it is a much stronger
base than Cl-) and forms the weakly acidic
hypochlorous acid (HOCl, Ka = 3.5 x 10-8)
Calculate the pH of a 0.100M aqueous solution of
hypochlorous acid.
PH OF A MIXTURE OF WEAK
ACIDS
Calculate the pH of a solution that
contains 1.00 M HCN (Ka = 6.2 x 10-10)
and 5.00 M HNO2 (Ka = 4.0 x 10-4). Also
calculate the concentration of cyanide
ion (CN-) in this solution at equilibrium.
% DISSOCIATION
amountdissociated(mol / L)
PercentDissociation =
´100%
initialconcentration(mol / L)
For example:
We found that a 1.00 M HF solution has [H+] = 2.7 x
10-2 M. To reach equilibrium, 2.6 x 10-2 mol/L of the
original 1.00 M HF dissociates, so…
2.7 ´10-2 (mol / L)
%Dissociation =
´100% = 2.7%
1.00(mol / L)
For a given weak acid…the % Dissociation
INCREASES as it becomes more dilute!
CALCULATE THE %
DISSOCIATION
a. 1.00 M Acetic Acid (Ka = 1.8 x 10-5)
b. 0.100 M Acetic Acid (Ka = 1.8 x 10-5)
Notice: for Acetic acid…the % Dissociation
INCREASES as it becomes more dilute!
1.00 M vs. 0.100 M
CALCULATING KA FROM %
DISSOCIATION
Lactic Acid (C3H6O3) is a waste product that
accumulates in muscle tissue during exertion,
leading to pain and a feeling of fatigue. In a
0.100 M aqueous solution, lactic acid is 3.7%
dissociated. Calculate the value of Ka for this
acid.
BASES
PH OF STRONG BASES
Calculate the pH of a 5.0 x 10-2 M NaOH
PH OF WEAK BASES I
Calculate the pH for a 15.0 M solution of
NH3 (Kb= 1.8 x 10-5).
PH OF WEAK BASES II
Calculate the pH of a 1.0 M solution of
methylamine, CH3NH2 (Kb = 4.38 x 10-4)
POLYPROTIC
ACIDS
POLYPROTIC ACIDS
Have more than 1 available proton
(ex. H2SO4, H3PO4)
They dissociate in a stepwise manner
POLYPROTIC ACIDS
They dissociate in a stepwise manner…
Ex. H2CO3
H2CO3  H+ + HCO3-
Ka1 = 4.3 x 10-7
HCO3-  H+ + CO32-
Ka2 = 5.6 x 10-11
CALCULATE THE PH OF A 5.0 M
H3PO4 SOLUTION AND THE
EQUILIBRIUM CONCENTRATIONS
OF THE SPECIES H3PO4, H2PO4-,
HPO42-, PO43Table 14.4 pg 651
CALCULATE THE PH OF A 1.0 M
H2SO4 SOLUTION
CALCULATE THE PH OF A 1.00
X 10-2 M H2SO4 SOLUTION
POLYPROTIC ACID RULES
• Typically for Weak Polyprotic Acids the 1st Ka value
is much larger than the following, so t is the only one
that makes a significant contribution to the
equilibrium concentration of H+.
• For Sulfuric Acid (only strong polyprotic acid)
• High concentrations negate the 2nd Ka
• Low concentrations, you need to take the 2nd Ka into
account (cannot assume x is negligible)