+ buffers. + buffer any substance or mixture of compounds that, added to a solution, is capable of...
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Buffers
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Buffer any substance or mixture of compounds that, added to a solution, is capable of neutralizing both acids and bases without appreciably changing the original acidity or alkalinity of the solution.
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+A buffer is created when both a weak acid and its conjugate base are present in solution (or a weak base and its conjugate acid)
i.e. CH3COOH and sodium acetate
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Equilibria in Acidic Buffer Solutions
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Acidic Buffer Weak acid and its conjugate base
Both in relatively high concentrations
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to write the equilibrium equation for an acidic buffer, write the IONIZATION of the WEAK ACID in water
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Example 1 A buffer solution is prepared that contains 1M CH3COOH and 1M NaCH3COO
Step 1: Determine the Ionization Equation
CH3COOH + H2O CH3COO- + H3O+
Step 2: Determine the effect on the equilibrium Equation increase the H3O+ / decrease the H3O+
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+Focus on the H3O+ in Acid Buffers
When acid or base is added to the buffer it changes the amount of H3O+
present in the solution
The equilibrium will shift either left or right to counteract the change in H3O+
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Example 2 A small amount of HCl is added to an unbuffered solution, and the pH changes from 7.0 to 5.0. If the same amount of HCl is added to a buffer solution with the same volume as the unbuffered solution, what would the resultant change in pH be?
Draw a graph to demonstrate
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the pH goes DOWN when acid is added
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Equilibria in Basic Buffer Solutions
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Basic Buffer Weak base and its conjugate acid (in a salt)
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Example 1 A buffer solution is prepared containing 1M NH3 and 1M NH4Cl.
Step 1: Determine the equilibrium equation
NH3 + H2O NH4+ + OH-
Step 2: Determine the effect on the system
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+Focus on OH- in Basic Buffers
When acid or base is added to the buffer it changes the amount of OH-
present in the solution
The equilibrium will shift either left or right to counteract the change in OH-
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Example 2 A small amount of NaOH is added to an unbuffered solution, and the pH changes from 7.0 to 9.0. If the same amount of sodium hydroxide is added to a buffer solution with the same volume as the unbuffered solution, how would the net pH of the solution respond?
Draw a graph
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Adding an Acid to an Acidic Buffer
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A small amount of HCl is added to a buffer solution containing 1M acetic acid CH3COOH and 1M sodium acetate NaCH3COO.
The equilibrium equation is:
CH3COOH + H2O H3O+ + CH3COO-
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What Happens?1. The addition of HCl immediately
increases the H3O+ and decreases the pH
2. To counteract the increase in H3O+, the equilibrium shifts to the left
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+3. As the equilibrium shifts to the left,
CH3COO- reacts with H3O+. Because there initial [CH3COO-] is high, there is enough of it to react with much of the excess H3O+
4. The H3O+ goes down close to, but not as low as, its original value
5. In the overall process, the final H3O+ is slightly higher and the pH is slightly lower than it was before the acid was added
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Adding a Base to an Basic Buffer
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Example 4 A small amount of NaOH solution is added to a buffer solution consisting of 1M NH3 and 1M NH4Cl.
The buffer equilibrium equation is:
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When NaOH is added to the buffer, the [OH-] immediately ______________, and the pH immediately _______________.
This would cause a shift to the ____________, and the [OH-] would gradually ____________ during the shift. The pH of the solution would gradually ____________ during this shift.
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+ Draw a graph showing what happens to the [OH-] from before the time NaOH is added until the shift is complete and a new equilibrium mixture exists
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Adding an Base to an Acidic Buffer
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+A small amount of NaOH is added to an acetic acid / acetate buffer
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Adding an Acid to a Basic Buffer
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+A small amount of HCl is added to an ammonia / ammonium buffer
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Buffer Limitations
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Moles Matter Once all the buffer is used up, buffering will no longer occur and the effect of the added acid/base will be noticeable
It all depends on the amount of moles present in solution!
i.e. if you add 6moles of HCl to a 2mole basic buffer, the buffer will only handle 2moles of HCl.
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+Buffer Uses
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+ Stabilizing pH in hot tubs and swimming
Maintaining pH for pharmaceuticals
Controlling pH in wines and foods
Calibrating pH meters
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Blood Buffers CO2 is a product of metabolism in human cells.
The CO2 dissolves in blood plasma (water) to form H2CO3:
CO2(g) + H2O(l) H2CO3(aq)
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+ H2CO3 ionizes to form H3O+ and HCO3-:
H2CO3(aq) + H2O(l) H3O+(aq) + HCO3-(aq)
This process is summarized in the body as:
CO2(g) + 2H2O(l) H3O+(aq) + HCO3-(aq)
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+ In this case, CO2(g) is the weak base,HCO3-(aq) is the conjugate acid.
This buffer helps to keep the blood pH level close to 7.35
During periods of high physical exertion, CO2 is controlled by breathing and [HCO3-] is controlled by the kidneys