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# Acid_Base Properties of Salts

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```									Acid/Base Properties of Salts
Conjugate acids and bases in salts
Salts are a metal paired with a conjugate
base (anion or negative ion) or a
conjugate acid paired with a negative ion
(normally a halogen, group 7 ion)

The Kb chart shows the conjugate acids for
all of the weak bases in the chart. This will
make them easier to spot.
Identify the conjugate in the salt
1. NH4Cl

2. NaC2H3O2

3. KF

4. CH3NH3F
Writing Net Ionic Salt reactions
Salt reactions are double replacement with water
as HOH. Look at the products. Identify them as
acid/base, strong/weak.

Strong acid products, cross off the back end, add a
positive charge to the front.

Strong base products, cross of the front part and
add a negative charge to the back part.
Deciding if it is acidic or basic
*Strong acid + weak base, strong wins
acidic
*Weak acid + strong base, strong wins
basic
*Strong acid with strong base, no winners,
neutral.
*Weak acid + weak base, no winners,
neutral (in AP you have to check which
weak is stronger and decide if it is acid or
base, we don’t do that)
Salt Reaction Examples
Write the reaction, decide if it is acidic or
basic.
1. NH4Cl + HOH ßà

2. NaC2H3O2 + HOH ßà

3. KCl + HOH ßà

4. CH3NH3F + HOH ßà
K a, K b & K w
Since salts can be acid or basic, we need to find
the Ka and Kb values for the conjugates. If H+ is
a product, look up Kb and find Ka. If OH- is
formed, look up a Ka and find Kb. We will use
the ionization constant for water (Kw).

HOH(l) ßà H+(aq) + OH-(aq)

Kw = [H+][OH-]
Kw = 1 x 10-14
Finding Ka & Kb examples
Determine the Ka or Kb for each reaction.
1. NH4Cl + HOH ßà

2. NaC2H3O2 + HOH ßà

3. KCl + HOH ßà

4. CH3NH3F + HOH ßà
Finding pH with salts
1. What is the pH of a 1.5M sodium acetate
solution?

2. What is the pH of a 2.2M NH4Cl solution?
Common Ion solutions
If a salt is dissolved in an acid or a base in
which the salt contains the conjugate of
the acid or the base it is a common ion
solution.

For example:
NaF dissolved in HF
CH3NH3Cl dissolved in CH3NH2
ICE with common ion solutions
• Write the reaction for the acid or the base.
• Fill in the acid or base in the reactant “I”
line.
• Fill in the salt concentration in the product
“I” line under the conjugate the salt
contains.
• Place a 0 on the “I” line under the H+ or
the OH-.
Example common ion ICE
1. What is the pH of a 0.5M potassium
acetate solution in a 0.75M acetic acid
solution?

2. What is the pH of 400mL of a 2.5M NH4Cl
solution when it is added to 500mL of a
1.75M NH3 solution?

3. What is the pH of 500mL of 1.3M HNO2
after 65g of KNO2 are added?
Websites
http://cwx.prenhall.com/petrucci/medialib/me
dia_portfolio/17.html
Buffers
Buffers are weak acids with a salt containing its
conjugate or weak bases with a salt containing
its conjugate that have the ability to maintain a
steady pH in a solution. Common ion solutions
make buffers

http://www.mhhe.com/physsci/chemistry/essentialc
hemistry/flash/flash.mhtml

http://www.nclark.net/Chemistry
Which combinations could make a
buffer?
1. NaClO3 / HClO3

2. NH4Cl / NH3

3. KBr/HBr

4. HNO2/NaNO3

5. NaHCO3 / Na2CO3
Making a Buffer
Buffer with a pH<7. Select an acid with a
pKa as close as possible to the pH you
need.
(pKa = -logKa)

Buffer with a pH>7. Select a base with a
pKb as close as possible to the pOH you
need.
(pKb = -logKb)
Pick a weak acid or base for each
buffer.
1. pH 4.5

2. pH 9.2

3. pH 7.3
Preparing a Buffer
You need to find the ratio of conjugate to
acid or base needed in your buffer. Use
the volume needed to find the moles of
each. Perform gram mole conversions
(for the conjugate you must add Na or K to
a conjugate base and a group 7 element
to a conjugate acid first).
How would you prepare a buffer?
1. 500mL of a buffer with a pH of 4.1?

2. 750mL of a buffer with a pH of 11.6?

3. How would you make 100mL of a buffer
with a pH of 4.7 from solid sodium acetate
and 6M acetic acid?
pH shifts in buffers
Write the reaction for the weak acid or base
(NEVER the salt)
Write the initial amount of moles for the weak and
its conjugate.
Find the moles of strong added. Decide which way
the reaction will shift and fill in the “C” line.
Fill in the “E” line, put an “x” under H+ or OH-.
Go back to molarity before plugging into Ka or Kb.
pH shift in a base buffer
A buffered solution contains 100mL of 0.5M
NH3 and 100mL of 0.8M NH4Cl.
A. What is the pH of this buffer?
B. What is the pH of this buffer if 10mL of
C. What is the pH of this buffer if 10mL of
in (b)?
pH shift in an acid buffer
500mL of a buffered solution contains 0.75M
HC2H3O2 and 0.9M NaC2H3O2.
A. What is the pH of the buffer?
B. What is the pH of 300mL of the buffer
after 15mL of 0.1M HCl are added?
C. What is the pH of the other 200mL of the
buffer after 20mL of 0.1M KOH are
Buffer Capacity
You could add more acid or base than a
buffer can handle and the pH will shift a
lot. The ability of a buffer to handle the
addition of acid or base is the buffer’s
capacity. The higher the concentration of
acid (or base) and its conjugate present in
the buffer the higher its capacity.
For each set, which has the bigger
capacity?
1. 0.5M HF / 0.7M NaF
0.05M HF / 0.07M NaF

2. 0.9M NH3 / 0.8M NH4Cl
1.6M NH3 / 1.3M NH4Cl

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