Supplemental Study Guide
pH and buffering
pH
-
-
-
a measurement of the activity of hydrogen ions (H+)
is expressed as a number between 0 and 14
pH can be derived from the number of H+ ions in a given solution.
Values from 0 through 6 are considered to be acidic.
Values above 7 and through 14 are considered basic or alkaline.
A solution with a pH of 7 is considered to be neutral, neither acid nor base. For
example, a solution with a pH of 2 is acidic; a solution with a pH of 13 is
basic.
The pH of a solution can be derived from [H+].
If the pH of a solution is increased from a value of 0 to 3, for example, the total
concentration decreases by a factor or 1000. If the pH of a solution is
decreased from a value of 13 to 11, the total concentration increases by a factor
of 100.
A buffering solution retains a constant pH value when small amounts of acid or
bases are added; buffers are composed of weak acids or bases
An increase in pH from 1 to 2 means a decrease in H+ concentration by 10. An
increase pH from 1 to 3 would result in a decrease in H+ by 100 (each place
represents a power of 10).
pOH
- a measurement of the activity of hydroxide ions (OH-)
- pOH can be measured on the pH scale – just opposite; 7 is neutral, above 7 is
an acidic pOH; below 7 is an alkaline pOH
Titration
- a procedure that allows the concentration of an unknown acid or base solution
to be determined
- uses a buffering solution and a complicated apparatus that includes flasks,
beakers, mass balances, pipettes, burettes, ring stands, pH meters and a Bunsen
burner to retain a constant pH when small amounts of acid or
base are added to a solution
- A titrant (or chemical reagent) drips into the reactant until a permanent color
change is achieved and the reactant is neutralized.
- Color changes to the analyte (another name for the reactant which contains the
unknown compound) will be permanent
- The volume at the time of the permanent color change in the reactant is called
the endpoint. when the endpoint in a titration is complete.
Acid-Base Theory
- based on the Arrhenius, Bronsted-Lowry, and Lewis definitions
- conjugate acid – the acid member of a pair of two compounds (product) that
form as a result of the a gain or loss of a proton in a reaction
- conjugate base – the base that results from the reaction
- protonation – a chemical reaction in which an atom, ion, or molecule accepts a
proton (H+)
Summary:
Acid-base theory
Arrhenius
Definition
Acids yield an H+ ion and a negatively
(-) charged ion; bases yield hydroxide
ions (OH-) and a positively charged (+)
ion.
Limitations
Only aqueous solutions
An example of an Arrhenius acid-base is
HNO3 ↔ H+ + NO3-.
Bronsted-Lowry
Acids are proton donors; bases are
proton acceptors.
None
A Bronsted-Lowry acid-base can be
defined as HA + B ↔ A- + BH+.
Bronsted-Lowry acids and bases form
conjugate pairs.
The conjugate base for
HCl + NH3 ↔ Cl- + NH4+ is Cl-.
The conjugate base for
HA + B ↔ A- + BH+ is A-..
Lewis
Acids are electron pair acceptors and
bases are electron pair donors.
Lewis acids and bases can be shown
with a Lewis electron dot diagram.
Lewis acids or bases are not limited to
aqueous solutions.
None
An example of a Lewis acid-base is
NH3 + H+ ↔ NH4.
The protonation of ammonia is an
example of a Lewis acid-base reaction.
Water is an amphoteric substance because it can act as an acid or a base
pH and pOH
- pH is a measure of the activity of hydrogen ions in a solution.
- Buffering agents can be either weak acids or weak bases. They can be added to water
to form buffering solutions.
- Buffering agents are added to substances that are intended to be put into acidic or
alkaline environments.
- Buffering solution are used to keep a solution at a constant pH.
- Example of a buffering solution - HA(aq) + H2O(l) ↔ H3O+(aq) + A- (aq); here we
see that HA is a weak acid.
- Buffering solutions are formed from weak acids and weak bases as long as they can
form conjugate acids and bases in sodium hydroxide.
- If a small amount of acid is added to a solution, the weak base will consume the
hydrogen ions and the conjugate acid of the weak base will be formed
- buffering agents are added to aspirin to keep them at a constant pH and not upset the
stomach.
- A solution with a pOH of 11.0 is acidic. A solution with a pOH of 3 is alkaline.
- Water becomes self-ionized when two water molecules react together to produce a
hydronium (hydrogen) ion and a hydroxide. This occurs because of the amphoteric
nature of water.
Some model problems to study:
1. A liquid solution exists at 1.00 x 10-8 mol/L at room temperature. Calculate
the pH value of this water.
pH = -Log10[H+]
Substitute the given concentration: pH = -Log10[1.00 x 10-8]
Solve with your calculator:
Type 1
Push the EE button (10 to the …)
Type 8
Push the reverse sign button
Push the = button
Get 0.00000001 (1 x 10-8)
Push the Log button
Push the reverse sign button
Push the = sign
Get 8.00
The pH is 8.00.
This solution is slightly alkaline.
2. A solution of nitric acid (NO3-) has an [H+] of 1.20 x 10-2 mol/L. Determine the
pH of this solution.
pH = -Log10[H+]
Substitute the given concentration: pH = -Log10[1.20 x 10-2]
Solve with your calculator:
Type 1.20
Push the EE button
Type 2 (this is the exponent)
Push the reverse sign button
Push the = button
Get 0.012
Push the Log button
Push the reverse sign button
Push the = sign
Get 1.92
The pH is 1.92.
This solution is a strong acid.
3. A solution of sodium hydroxide (strong base) has an [OH-] value of 0.80 x 10-3
mol/L. Calculate the pOH of this solution.
pOH = -Log10[OH-]
Substitute the given concentration: pOH = -Log10[0.80 x 10-3]
Solve with your calculator:
Type 0.80
Push EE
Type 3 (this is the exponent)
Push the reverse sign button
Push the = button
Get 0.0008 (0.8 x 10-3)
Push the Log button
Push the reverse sign
Push the = sign
Get 3.10
The pOH is 3.10
This is a strong base.
4. Determine the [H+] of a substance that has a pH of 5.5.
pH = -Log10[H+]
Substitute the pH: 5.5 = -Log10[H+]
Solve: Multiply both sides by a negative sign to get -5.5 = Log10[H+]
Isolate the H+ by making each side an exponent of 10
10-5.5 = 10Log[H+] (the property of Log makes 10 raised to the Log of any
number equal to that number)
10-5.5 = H+
Use your calculator:
The [H+] = 3.16 x 10-6 mol/L
Type 10
Push the yx button
Type 5.5
Push the reverse sign button
Push the = sign
Get 0.000003162 (3.16 x 10-6)
5. If we are given an [OH-] of 2.00 x 10-3 mol/L, what is the pOH.
pOH = -Log10[OH-]
Substitute the given concentration: pOH = -Log10[2.00 x 10-3]
Solve with your calculator:
Type 2.00
Push EE
Type 3 (this is the exponent)
Push the reverse sign button
Push the = button
Get 0.002 (2 x 10-3)
Push the Log button
Push the reverse sign
Push the = sign
Get 2.70
The pOH is 2.70
This is a strong base.
6. Given that a substance has a pH of 2.5, what is the pOH?
The pH + the pOH is equal to 14.
Subtract 2.5 from 14 to get the pOH.
The pOH is 11.5
7. Given that the pH of a solution is 0.5, what is the [OH-]?
pH = 0.5
14 – 0.5 = 13.5
pOH = 13.5
pOH = -Log10[OH-]
Substitute in the value of the pOH to get
13.5 = -Log10[OH-]
Rewrite as
10-13.5 = [OH-]
Use your calculator to solve:
Type 10
Push the yx button
Type 13.5
Push the reverse sign button
Push the = sign
Get 3.16227766-14
The [OH-] = 3.16 x 10-14 mol/L