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
Look at your products.
*Strong acid + weak base, strong wins
*Weak acid + strong base, strong wins
*Strong acid with strong base, no winners,
*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
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

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

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”
• Fill in the salt concentration in the product
  “I” line under the conjugate the salt
• 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

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?
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
 Which combinations could make a
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
(pKa = -logKa)

Buffer with a pH>7. Select a base with a
  pKb as close as possible to the pOH you
(pKb = -logKb)
 Pick a weak acid or base for each
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
  0.1M HCl are added?
C. What is the pH of this buffer if 10mL of
  0.2M NaOH are added instead of the HCl
  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
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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