The effect of addition of acid or base to …
acid added
base added
an unbuffered solution
acid added
base added
or a buffered solution
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Characteristics of Buffers
• 1. Contains a weak acid HA and its CB A• 2. Resists changes in pH by reacting with any
added H+ or OH-, so these ions do not accumulate.
• 3. Any added H+ reacts,
• H+ + A- → HA
• 4. Any added OH- reacts,
• OH- + HA → H2O + A-
What is a buffer?
• Chemical system that can withstand pH changes when
(limited) amounts of acid or base are added.
• A buffer system consists of two components
– acid component to react with OH1- ions
– base component to react with H1+ ions
• The components should not consume each other
• A solution containing both:
a weak acid + its salt OR a weak base + its salt.
• Works on Le Châtelier’s principle and Common Ion
Effect
Common Ion Effect
• A solution of two dissolved solutes that contain
the same ion - common ion
e.g. In solution of CH3COOH + CH3COONa
CH3COOH(aq) <-----> CH3COO1- (aq) + H1+ (aq)
CH3COONa (s) + H2 O(l) ---> CH3COO1- ( aq) +
Na1+(aq)
Both solutes produce CH3COO1- ( aq) ions.
• According to Le – Chatelier’s principle,
if add acid, then reaction HA <---> H+ + A– goes
to left to absorb change; vice-versa if add base
Figure 19.2
How a buffer works
Buffer after addition of H3O+
Buffer with equal
concentrations of
conjugate base and acid
H3O+
H2O + CH3COOH
H3O+ + CH3COO-
Buffer after addition of OH-
OH-
CH3COOH + OH-
H2O + CH3COO-
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Buffering capacity: amount of H+ or OH- a
buffered solution can absorb without significantly
changing the pH. Capacity determined by
magnitude of [HA] and [A-].
Concentrations of the components should be high
enough to neutralize the added H1+ or OH1- ions
Buffers in Natural Systems
•Biological fluids, eg blood, contain buffers:
pH control essential because biochemical reactions
are very sensitive to pH
•Human blood is slightly basic, pH  7.39 – 7.45
•In a healthy person, blood pH is never more than 0.2
pH units from its average value
•pH < 7.2, “acidosis”; pH > 7.6, “alkalosis”
•Death if pH < 6.8 or > 7.8
Buffer System in Blood
“extracellular” buffer (outside cell)
H+ + HCO3– < -->
H2CO3 < -- >
H2CO3
H2O + CO2(g)
–Removal of CO2 shifts equilibria to right, reducing
[H+ ], ie, raising the pH
–The pH can be reduced by:
H2CO3 + OH–
<--- >
HCO3– + H2O
Another Blood Buffer
•phosphate buffer, present inside cells
(“intracellular” buffer)
•H2PO4– and HPO42– :
H2PO4– < -- >
H+ + HPO42–
from H3PO4 , a tribasic (triprotic) acid
Figure 19.3
The relation between buffer capacity and pH change
Problem: Which of the following solutions are
buffers?
a) KH2PO4/H3PO4
b) b) 0.30 mol CH3COOH/ 0.30 mol NaOH
• Solution:
• a) H3PO4 is a weak acid and KH2PO4 , a salt
of its conjugate base H2PO4 1-, is a weak
base. The system contains a conjugate acidbase pair and therefore is a buffer system
b) 0.30 mol CH3COOH/ 0.30 mol NaOH
Stoichiometric quantities of a weak monoprotic
acid and strong monohydroxy base will
produce a salt containing the conjugate base of
the weak acid, CH3COONa. This weak base
alone without its conjugate acid will not be a
buffer.
Auto-ionization of water
• 2 H2O ↔ H3O+ + OH• This solution is still neutral b/c there are
equal amounts of [H+] and [OH-], which is
1.0 x 10-7 M
Indicators and
Titration Curves
Acid – Base Titrations
• is a carefully controlled neutralization reaction
• acid + base  salt + water
• involves the progressive addition of one
reactant from a burette, to a known volume of
the other reactant in a flask.
• Standard – solution in which you know the
concentration
Acid-Base Indicators
• Indicators - Weak organic acids or
bases
– Distinctly different colors in two forms
– Two forms related to pH of solution
– Phenolphthalein – most often used
indicator for titrations
• Clear – acidic
• Pink – basic
Acid – Base Indicators
•
HIn (aq)   H+(aq) + In-(aq)
•
In acidic medium, equlibrium shifts to the
left and the predominant color is HIn color
In basic medium, equilibrium shifts to the
right and the predominant color is In1- color
•
Colors and approximate pH range of some
common acid-base indicators
The color change of the indicator bromthymol blue
basic
acidic
change occurs
over ~2pH units
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Titration
• Animation
Equivalence / End Pts
• Equivalence point – when neutralization occurs
# H+ ions donated = # H+ ions accepted
or
total moles H+ ions = total moles of OH- ions
• End Point – when the indicator changes color
Strong Acid-Strong base Titration
•pH
starts low and increases gradually as acid
is neutralized by the added base
•Close
to equivalence point pH rises steeply
•Beyond
this, pH increases slowly with
addition of more base
•Equivalence
point
• the mol OH1- = mol H1+
•The pH = 7
Strong base
titrated with a
strong acid
Equivalence pt
is 7
Strong acid
titrated with
a strong
base
Weak Acid-Strong Base Titration
•Initial pH of weak acid
•starts higher than pH
of a strong acid of the
same molarity
•increases
as the acid is neutralized by added
base
•Gradually
rising curve is called buffer region as
the weak acid reacts with the strong base to
produce its conjugate base
•At
the mid-point of buffer region half of original
acid has reacted with the base to produce conj.
base
Strong base
titrated with a
weak acid
Equivalence pt
above 7
Weak acid
titrated with
a strong
base
Weak Base-Strong Acid Titration
•Initial
pH of weak base
•starts higher than 7 but lower than pH of a
strong base of the same molarity
•Decreases
as the weak base is neutralized by
added acid
•Gradually
dropping curve is called buffer region as
the weak base reacts with the strong acid to
produce its conjugate acid
•At
the mid-point of buffer region half of original
base has reacted with strong acid to produce
conj.acid
Weak base
titrated with a
strong acid
Equivalence pt
below 7
Strong acid
titrated with
a weak base
Steps to a Titration
• Determine what indicator to use based upon strengths
of acid/base & known end points for indicators
• Fill a buret with standard solution & record initial
volume
• Start adding std solution slowly, with mixing, to the
solution of unknown concentration
• Continue adding std solution until equivalence point in
reached (equivalence & end point should be relatively
close)
• Utilize given and obtained data to determine
concentration of unknown.
Calculations
• Write the balanced equation for the neutralization
• Extract all the relevant information from the
question .
• MAVA = MBVB
– This can be used because your moles of hydronium and
hydroxide are equal at the equivalence point.
– CAUTION: If the acid is polyprotic or the base is
polyhydroxic, you must use the normality instead of
molarity in this formula.
• Normality is found by multiplying the number of hydronium or
hydroxide ions in the formula unit by the molarity.
Example
A 43.0mL volume of NaOH was titrated with 32.0 mL of
0.100M HCl. What is the molarity of the sodium
hydroxide solution?
What key information can be obtained from problem?
NaOH + HCl  NaCl + H2O
43.0mL NaOH
32.0mL 0.100M HCl
Use the formula and plug in the values.
MAVA = MBVB
(x)(43.0ml) = (0.10 M)(32.0 ml)
M(HCl) = 0.075 M
Example #2
A volume of 25 mL of 0.120 M H2SO4 neutralizes 41 mL
of a NaOH solution. What is the concentration of the
NaOH?
What key information can be obtained from problem?
2 NaOH + H2SO4  Na2SO4 + 2 H20
25 ml, 0.120M H2SO4
41ml, NaOH
Use the formula and plug in the values.
MAVA = MBVB
(0.120M x 2)(25 ml) = (x)(41 ml)
M(NaOH) = 0.14 M
Your turn….
1. If it takes 54 mL of 0.1M NaOH to
neutralize 125 mL of an HCl solution, what
is the concentration of HCl?
2. It takes 25 ml of a 0.05M HCl to neutralize
345mL of NaOH solution, what is the
concentration of the NaOH?