Acids buffers

Acids, pH, and Buffers:
Some Basic Chemistry for Biological
Science
Terry Platt
University of Rochester

1

Paul Mathews has just learned about acidity
in his chemistry class, and…
• surprises his parents with questions about
the acidity of common substances, such as:
– Coffee
– Vinegar

pH 5!
pH 3!

• Can you guess the pH of tomatoes? Cola or
beer? Lemon juice? Your stomach?

But what is an acid, and what is pH?
2

What is an Acid?
• An acid is a substance which, when dissolved in
water, releases protons.
• The extent of dissociation, that is, the amount of
protons released compared to the total amount of
compound, is a measure of the strength of the acid.

• For example, HCl (hydrochloric acid) is a strong acid,
because it dissociates completely in water,
generating free [H+] and [Cl-].
• Acidity can be measured on a scale called pH (more
scarily, “the negative logarithm of the hydrogen ion
concentration”).
3

CQ#1: An acid is a compound that in
aqueous solution will readily:
A. Shed a proton.
B. Shed an electron.
C. Gain a proton.
D. Gain an electron.
E. None of the above.

4

SOURCE: http://en.wikipedia.org/wiki/Image:PH_scale.png#file

[H+] M
100
10-1

A strong acid

10-2
10-3
10-4
10-6
10-7
10-8

• [H+] is controlled in all
biological organisms, and in
virtually all biochemical
experiments.

10-9

10-10
10-11
10-12
10-14

• Most living cells have a very
narrow range of tolerance for
pH, i.e. [H+].

• The [H+] concentration will be
important (either explicitly or
implicitly) for many other
topics in biology.

10-5

10-13

pH

A strong base

• Each pH unit represents a
factor of 10 difference in [H+].

The pH scale goes from 0 to 14—because [H+][OH-] = 10-14

5

CQ#2: In an aqueous solution where the
H+ concentration is 1 x 10-6 M, the OHconcentration must be:
A. 14 x 10-6 M
B. 1 x 10-6 M

C. 1 x 10-7 M
D. 1 x 10-8 M
E. 14 x 10-8 M

6

The Conceptual Problem with pH
• Because it’s a logarithmic scale, it doesn’t make
“sense” to our brains.
• But Paul explains it well—every factor of 10
difference in [H+] represents 1.0 pH units, and
• Every factor of 2 difference in [H+] represents 0.3
pH units.
• Therefore, even numerically small differences in
pH, can have profound biological effects…
7

How Can You Actually
Determine the pH of a Solution?
• Use a pH meter—read the number.
• Use pH paper (color patterns indicate pH).

• Titrate the solution with precise amounts of
base or acid in conjunction with a soluble
dye, like phenolphthalein, whose color
changes when a specific pH is reached.

8

In our story, the mother had to run home really
hard—let’s look at some of the effects:
• Is there a volunteer to help out?
• Instructions: blow gently through the straw
into the pink solution for 15-30 seconds.

• What happens?
• What does the “indicator” tell you?

• Why did it happen?
9

RESULTS
• The solution went from pink to clear,

• indicating a decrease in pH (the colorchanging dye is phenolphthalein),
• i.e., an increase of acidity.

What could this be due to?

10

The Hydration of Carbon Dioxide in Water
CO2 + H2O

H2CO3

⇌

HCO3- + H+

As carbon dioxide goes into solution, carbonic acid is
formed, which partially dissociates, liberating protons (H+)
and thus causing the solution to become more acidic, i.e.,
lowering the pH.
11

And you can apply what you just learned
about carbon dioxide…
• When Molly’s mom ran home so hard, what do you think
would have happened to the pH (acidity) of her blood?
• What would be the effect on pH once she stops running,
but continues to breathe hard for a few minutes? Why?

• For an athlete about to begin an event, what would the
consequence of hyperventilation be on her blood pH?
What advantage might this confer?

So, in water, CO2 forms a weak acid!
12

But What’s a Weak Acid?
• Some substances, like acetic acid (vinegar!)
dissociate poorly in water.
• Thus, they release protons, but only a small fraction
of their molecules dissociate (ionize).
• Such compounds are considered to be weak acids.
• Thus, while 1 M HCl is pH = 0 (why?), 1 M acetic acid
is only pH = 2.4…

13

Weak acids thus are in equilibrium
with their ionized species:
Governed by the Law of Mass Action, and
characterized by an equilibrium constant:

HA
Keq =

+
H

-

+A

+][A-]
[H

[HA]
14

Water: A Very Weak Acid

+

+
hydronium ion

2 H2O

H3 O+
(an acid)

hydroxide ion

+

OH(a base)

But this hardly happens at all: In fact, at equilibrium,
[H+] = [OH-] = 0.0000001 M = 10-7 M = pH 7
Indeed, only two of every 109 (1 billion) molecules in pure water are
ionized at any instant - Can you confirm this?
15

Comparative Equilibrium Constants
• Water:

Keq = 1.8 x 10-16

• Acetic acid

Keq = 1.7 x 10-5

A 100 billion-fold difference…
But still, of every 1000 acetic acid molecules in a 1
M solution of acetic acid, only 4 are ionized.

Can you figure out how to figure that out?
16

For biological systems:
• Ionization of a strong acid is TOO
• Ionization of water itself is way

BIG!

TOO LITTLE!

• Ionization of a weak acid is JUST RIGHT!

17

But Molly’s problem is with a weak acid…
• She has ingested far too much aspirin, i.e.,
acetylsalicylic acid, which has brought her to
the Emergency Room.
• The window of pH within which humans can
survive is between a blood value of about pH
6.8 to around pH 8.0.
• Outside that range brings coma and death.

• Thus rapid treatment for Molly is crucial.
18

CQ#3: By what factor does the [H+] of
Molly’s blood (pH 6.8) differ from normal
(pH 7.4)?
A.
B.
C.
D.
E.

0.25x
0.5x
0.6x
2.0x
4.0x

19

CO2 + H2O

H2CO3

What if you could
reverse this…
by removing carbon
dioxide?
HCO3- + H+

As carbon dioxide leaves the solution, carbonic acid is used
up, which by the Law of Mass Action shifts the equilibrium to
the left, using up protons (H+) and thus causing the solution
to become less acidic, i.e., raising the pH.
20

CQ#4: Why is Molly breathing so rapidly and
deeply when she arrives at the Emergency
Room, despite being nearly comatose?
A. The aspirin has inhibited her ability to
use oxygen effectively.
B. Her body is trying to rid itself of CO2.
C. She is out of breath from all she has
been through.
D. Her hemoglobin can’t deliver oxygen at
low pH.
21

Unfortunately, Molly’s breathing isn’t helping
very much—can you guess why that is?
• Think about the aspirin in her bloodstream and the
equilibrium between its acid (undissociated) and its
dissociated form.
• As protons are removed from solution by her heavy
breathing, is there a reservoir of others to take their
place…?
• Try to discuss and explain this to a person sitting near
you.

22

CQ#5: How else might you raise the pH
of Molly’s blood to get it back into the
normal range?
A.
B.
C.
D.
E.

Add HCl
Add acetylsalicylic acid
Add phosphate
Add bicarbonate
Add NaOH

23

Why should bicarbonate help?
• As just seen, bicarbonate is one of the ionization
products of carbonic acid—what is the other
one?
• By the Law of Mass Action, the addition of a lot of
bicarbonate should drive the reaction—in which
direction?
• Think about this: if you add baking soda to
vinegar, what will happen, and what does it
mean?

24

How Does This Work?
CO2 + H2O

H2CO3

⇌

- + H+
HCO3

Here, the addition of excess bicarbonate will soak up many
of the free protons, and drive the equilibrium to the left.
This will reduce the acidity, increasing the pH, and the
carbon dioxide produced will be blown off in the lungs. And
make Molly feel MUCH better!
25

How can this help Molly?
•

This will reduce the acidity, which means
increasing the pH, and …
the carbon dioxide produced will be blown off in the lungs.

•
•

This will make Molly feel much better.
And get her out of medical danger…

26

Weak acids, their conjugate
bases, and buffers…
• Weak acids have only a modest tendency to shed their
protons (definition of an acid).
• When they do, the corresponding negatively charged
anion becomes a willing proton acceptor, and is called
the conjugate base.
• The properties of a buffer rely on a balance between a
weak acid and its conjugate base.
• And a titration curve looks like this…
27

Titration of acetic acid with sodium hydroxide

pH 7

_

_
pKa = 4.76

Buffering range:
only small pH
changes result
from addition of
base or acid

_
_

_
_
_

0

50%
dissociation

equiv. of NaOH
added

1.0
28

How did this save Molly’s life?
• A buffer is a solution of a weak acid and its
conjugate base that resists changes in pH in
both directions—either up or down.
• A buffer works best in the middle of its
range, where the amount of undissociated acid
is about equal to the amount of the conjugate
base.
• One can soak up excess protons (acid), the
other can soak up excess hydroxide (base).
29

Ionization of acetic acid:
Resisting changes both ways
OH-

Acetic acid
(CH3COOH)

H2O

HAc

Ac-

Acetate
(CH3COO-)

H+
30

Insights for the Future
• pH control is important, as many enzymes have a
narrow range in which they function optimally.
• Buffering capability is essential for the well-being of
organisms, to protect them from unwelcome
changes in pH.
• For example, your stomach is about pH 1, yet the
adjacent portion of your intestine is near pH 7—
think about (or look up) how that might happen
[Hint: what is one function of the pancreas?].
• Many compounds and macromolecules in addition
to bicarbonate can serve a buffering function—
proteins comprise one of the major classes.
31

Acids buffers

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