Notes to the lecture : ICF, ABB and electrolytes
adapted from Voets : Biochemistry
Series of biochemical reactions, which are termed metabolic pathways, as well as the structures
of the enzymes that catalyze them are, for many basic processes, nearly identical from organism
to organism. This strongly suggests that all known life-forms are descended from a single
primordial ancestor in which these biochemical features first developed.
What we ask from biochemistry :
a) What are the chemical and three-dimensional structures of biological molecules and
assemblies, how do they form these structures, and how do their properties vary with them.
b) How do protein work; that is, what are the molecular mechanisms of enzymatic catalysis, how
do receptors recognize and bind specific molecules, and what are the intramolecular and
intermolecular mechanisms by which receptors transmit information concerning their binding
states.
c) How is genetic informationexpressed and how is it transmitted to future cell generation.
d) How are biological molecules and assemblies synthesized.
e) What are the control mechanisms that coordinate the large amount of biochemical reactions
that také place in cells and in organisms.
f) How do cells and organisms grow, differentiate, and reproduce.
Metabolism has usually divided to two major categories :
1. Catabolism or degradation, in which nutrients and cell constituents are broken down so as to
salvage their components and/or to generate energy.
2. Anabolism or biosynthesis, in which biomolecules are synthesized from simpler components.
The energy required by anabolic processes is provided by catabolic processes largely in the
form of adenosine triphosphate (ATP).
Biologically important elements.
Living matter consists of a relatively small number of elements C, H, O and N. Many
biomolecules also contain S and P. We speak about the macroelements (almost 99% of body
mass). A second group of biologically important elements (0.5% of body mass) are present in the
form of inorganic ions Na+, K+, Mg2+, Ca2+, Cl-. Other elements are present in so small
concentration that they are referred to as trace elements . Fe, Zn, Cu, Co and Mn, from nonmetals I and Se. So, there is no known biological requirements for 65 from 90 naturraly occuring
elements. Conversely with the exceptions of oxygen and calcium, the biological most abundant
elements are but minor constituents of Earth´s crust (earth´s crust = 47% of O, 28% of Si, 7.9%
of Al, 4.5% of Fe, and 3.5% of Ca).
The predominance of carbon in living matter is a result of its tremendous chemical versatility.
Carbon has the unique ability to form a virtually infinite number of compounds as a result of its
capacity to make as many as four highly stable covalent bonds (including single, double and
triple bonds) combined with its ability to form covalently linked C-C chains of unlimited extent.
Presentation
ad 2. exposure) It is known that life is based on morphological unit - cell. There are two major
classifications of cells eukaryotes and prokaryotes. Eukaryotes have a membrane-enclosed
nucleus and prokaryotes which lack this organelle. Eukaryotes can be multicellular as well as
unicellular. (Viruses are not classified as living because they lack the metabolic apparatus to
reproduce outside their host cells). The nucleus of eukaryotic cells is the depot of its genetic
information. This information is encoded in the base sequence of DNA molecules. The
Endoplasmic reticulum and the Golgi apparatus function to modify membrane bound and
secretory proteins. Mitochondria are the site of oxidative metabolism.
ad 3. and 4. exposures) The solar system was probably formed by the gravitionally induced
collapse of a large interstellar cloud of dust and gas. The major portion of this cloud was
composed predominantly from hydrogen and helium. This part condensed to form the sun. The
rising temperature and pressure at the center of the protosun eventually ignited the selfsustaining thermonuclear reaction (source of energy of the sun). The planets, which formed from
smaller clumps of dust, were not massive enough to support such a process. The smaller planets,
including Earth, consist of mostly heavier lements because their masses are too small to
gravitionally keep much hydrogen and helium. The primordial Earth´s atmosphere was quite
different from what it is today. It could not contain significant quantities of oxygen, a highly
reactive substance. Rather, in addition to the H2O , N2, and CO2 , the atmosphere probably
contained small amount of CO, CH4 , NH3, SO2 and possibly H2, all molecules that heve been
spectroscopically detected in interstellar space. It was the reducing atmosphere in contrast to
Earth´s present atmosphere which is an oxidizing atmosphere.
Oparin (1920) : UV radiation from the sun (now is absorbed by an ozone layer) or lightning
discharges caused the molecules of the primordial reducing atmosphere to react to form simple
organic compounds such as amino acids, nucleic acid bases, and sugars. We discriminate three
stages in the evolution of life. 1. Chemical evolution, in which biopolymers were formed from
small molecules. 2. Self organization, in which biopolymeres developed the capacity for selfreplication and 3. Biological evolution, in which primitive living cells generated sophisticated
metabolic systems and eventually the ability to form multicelullar organisms.
Oparin´s theory was first experimentally demonstrated in 1953 by S. Miller and H. Urey. On the
4. exposure (picture) is seen the effect of lightning storm, simulated by electric discharge, to the
refluxing mixture of H2O, CH4 , NH3, and H2 , for about a week. The resulting solution
contained significant amounts of water-soluble organic compounds (viz table). Condensation of
HCN → adenin, polymerization of formaldehyde → sugars.
The rates of synthesis of these complex polymers would have had to be greater than their
hydrolysis. Temperature of the "pond of compounds" may have been cold, possibly even below
0°C.
ad 5. exposure) The life has evolved in water and is still dependent on it. The physical and
chemical properties of water are, therefore, of fundamental importance to all living things.
The comparison with methane (similar mass and size). But the difference in the boiling point
250°C. As a result of this, water is liquid at earth´s temperature, and methane is gas. Uneven
distribution of electrons within the molecule of water, H-O-H bond is bent. In addition, the O-H
bonds are polarized due to the high electronegativity of oxygen. The spatial separation of
positive and negative charge gives the molecule of water the properties of electrical dipole.
When liquid water vaporizes , a large amount of energy has to be expended to disrupt these
interactions.
ad 6. exposure) Solubility depends on the ability of a solvent to interact with a solute more
strongly than solute particles interact with each other. The polar character of water makes it an
excellent solvent for polar and ionic compounds, which are said to be hydrophilic. On the other
hand, nonpolar substances are virtually insoluble in water are described like hydrophobic. Non
polar substances are soluble on the other hand in nonpolar solvents like for instance CCl4 and
hexane. Like dissolves like.
Task for you :
Let us take NaCl. These salts are held together by ionic forces (Coulomb´s law). Repeat what is
dielectric constant and try to explain what the force will be between NA+ and Cl- against the
molecules of hexane.
Explain hydrogen bond. What is Product of solubility. Why barium can be used in medical
practise even though barium is toxic element ?
ad 7. exposure) The effects of polar and apolar groups on the water-solubility of organic
compounds can be illustrated by the solubility of various fatty acids. The carboxylic groups of
fatty acids are ionized under physiological conditions and therefore well hydrated. But when
hydrophobic (apolar) hydrocarbon chain is prolonged, the solubility decrease. Acids with more
than 10 carbons are more or less insoluble in water.
The spontaneus separation of a mixture of oil and water into two distinct phases reduces
energetically unfavorable formation of clathrate structure. The large drop has a smaller surface
area than several smaller drops of the same volume.
Molecules that contain both polar and apolar groups are called amphipatic or amphiphilic.
Amphipatic substances in water tend to orient themselves so as to minimaze the area of contact
between the apolar regions of molecules and water. Monolayers have the "polar" head facing
toward the water. Soap bubbles consist of lipid bilayers with apolar parts facing outward and a
thin layer of water in the interior. Back-to-back arrangement of two monolayers, extended
bilayers arise (membranes). Vesicles are hollow membrane sacks. This structure plays a role in
the transport of substances within the cells and in body fluids.
Explanation to the lecture :
Electrolytes. Theory of acids and bases. Ion product of water. Ionization constants. pH.
Buffers.
9. Chemical equilibria
An equilibrium law and the equilibrium constant exist for each equilibrium.
aA + bB  cC + dD
Guldberg-Waage law :
Keq = [C]c [D]d / [A]a [B]b
Keq is called equilibrium constant. The size of Keq indicates the position of equilibrium .
Keq < 1, reactants are favored at equilibrium
Keq > 1, products are favored at equilibrium
Keq remains constant even when the equilibrium shifts.
9.1. Electrolytes
Solutions of ionic compounds in water readily conduct electrocity. All aqueous fluids of living
systems - plants or animals - contain dissolved ions and molecules. Blood for example, contains
sodium and chloride ions as well as molecules glucose and other molecular compounds.
When ionic compounds dissolve in water their ions, dissociate - we say that dissociation (or
ionization) occurs because oppositely charged ions are separated NaCl(s) → Na+ (aq)+ Cl(aq)).
Ionization is the formation of ions by a chemical reaction of a molecular compound with the
solvent. Hydrogen chloride e.g. undergoes ionization as it dissolves in water (H2O + HCl →
H3O+ + Cl-). Pure hydrogen chloride, either as a gas or a liquid, contains no ions. But if HCl
dissolves in water, almost 100% of its molecule react with water giving hydronium and chloride
ions.
Strong electrolytes give high concentration of ions in water and usually dissociate in 100% in
water (HCl, NaOH, KOH ).
Weak electrolytes are substances that generate ions in water only to a small percentage of its
molar concentration. Examples : aqueous ammonia, acetic acid and most of carboxylic acids and
so on.
Nonelectrolytes are that do not ionize and dissociate in water at all - essentially zero percent
ionization.
9.2. Ionic product of water
Important chemical equilibria, the self-ionization of water.
H2O  H+ (aq) + OH- (aq)
Its equilibrium law is
Keq = [H+] [OH-] / [H2O]
In pure water [H+] = [OH-]. Experimentally, each has a value 1.0 x 10-7 mol/L at 250C. Their
very low concentrations have no effect on the total concentration of water [H2O]. So we can
connect [H2O] with Keq and we get new constant KW, so called Ionic product of water.
KW = [H+] [OH-] = 10-14
In the logarithmic scale :
pH + pOH = pKW
9.3. Theory of acids and bases
Acids supply hydrogen ions, and bases neutralize hydrogen ions.
The three principal ion-producers in water are : acids, bases and salts.
Arrhenius theory of acids and basis ( Svante Arrhenius 1859-1927 - Swedish scientist) : All
acids produce hydrogen ions H+ in water. Bases produce OH- in water.
Bronsted theory of acids and bases (Johannes Bronsted 1879-1947 - Danish chemist) : Acids
are proton donors that means production of H+ , respectively H3O+ and bases are proton
acceptors. We can therefore say that in acidic solutions applies [H+] ˃[OH-], in neutral solution
[H+]=[OH-] and basic solutions [H+] <[OH-].
I will use often the symbol of H+ instead of H3O+ as a simpler way.
9.4. pH concept
At the molecular level of life we usually intersted in acid-base balance with weak acids and
bases and with very low concentrations of H+ and OH-. Because of these low concentrations, we
usually express the concentrations in logarithmic scale.
The pH of a solution is the negative logatithm of its [H+].
So we can define pH as
[H+] = 1 x 10-pH or
pH = -log [H+]
There are analogous equation for expressing low concentrations of OH[OH-] = 1 x 10-pOH or
pOH = -log [OH-]
The relationship between pH and pOH we can express
pH + pOH = 14
Acidic solutions have the pH values less than 7, and value more than 7 for a solution to be basic.
In pH terms, then, we have the following definition of acidic, basic, and neutral solutions when
temperatures are 250C :
acidic solution pH < 7
neutral solution pH = 7
basic solution
pH > 7
Small pH changes can mean large [H+] changes.
The problems of pH and calculations associated with it will be filled seminars.
9.5. Ionization (dissociation) constants
The strengths of weak acids are described quantitatively by their ionization (dissociation)
constant, KA. It is obtained by a simplification of the equilibrium law for the ionization
(dissociation) of the weak acid in water.
Dissociation constants (I will use in this article the word dissociation instead of ionization) are
very large for strong acids and we can ignore them. We assume that, for all practical purposes,
strong acids are 100% ionized (dissociated) in dilute aqueous solution. When we represent any
weak acid by the symbol HA, where A denotes the species that separates from H+ when HA
ionizes. The weak acids might electrically neutral like acetic acid, positively charged like NH4+,
or negatively charged like the bicarbonate ion. All these are proton donors, but they differ in acid
strengths.
HA (weak Bronsted acid) + H2O  H3O+ + A- (conjugate base)
The general form for the equilibrium law for this equilibria is :
Keq = [H3O+] [A-] / [HA] [H2O]
The value of [H2O] is essentialy constant, so we can now simplify this equation by coupling with
Keq and we get the equation for dissociation (ionization) constant :
KA = [H+] [A-] / [HA]
Weak, moderate, and strong Acids can be defined by KA more like by pKA values.
acid
formula
pKa1
acetic
HC2H3O2
4.74
acetylsalicylic
C8H7O2COOH
3.48
acrylic
HC3H3O2
4.26
aluminum ion
Al3+
5.01
arsenic
H3AsO4
2.30
arsenous
HAsO2
9.18
ascorbic
H2C6H6O6
4.10
boric
H3BO3
9.23
bromoacetic
HC2H2BrO2
2.90
butyric
HC4H7O2
4.83
carbonic
H2CO3
6.37
chloroacetic
CH2ClCOOH
2.85
chlorous
HclO2
1.92
chromium ion
Cr3+
4.0
cinnamic
HC9H7O2
4.44
citric
H3C6H5O7
3.13
cyanic
HCNO
3.46
cyanuric
HC3H2N3O3
6.78
dichloroacetic
HC2HCl2O2
1.26
fluoroacetic
HC2H2FO2
2.59
formic
HCOOH
3.75
germanic
H2GeO3
9.0
hydrazoic
HN3
4.72
hydrocyanic
HCN
9.23
pKa2
pKa3
7.03
11.53
11.80
10.25
4.76
12.4
6.40
hydrofluoric
HF
3.16
hydrogen peroxide
H2O2
11.65
hydroselenic
H2Se
3.89
11.0
hydrosulfuric
H2S
6.88
14.15
hydrotelluric
H2Te
2.64
10.80
hypobromous
HBrO
9.24
hypochlorous
HClO
7.55
hypoiodous
HIO
10.64
hyponitrous
H2N2O2
7.05
hypophosphorous
H3PO2
1.23
iodic
HIO3
0.80
iodoacetic
HC2H2IO2
3.18
iron(II)ion
Fe2+
6.74
3+
11.4
iron(III)ion
Fe
2.83
lactic
HC3H5O3
3.08
maleic
HOOCCH:CHCOOH
1.84
6.07
malonic
H2C3H2O4
2.82
5.70
nitrous
HNO2
3.14
oxalic
H2C2O4
1.23
phenol
HOC6H5
10.00
phosphoric
H3PO4
2.12
7.21
phosphorous
H3PO3
1.80
6.15
phthalic
H2C8H4O4
2.92
5.41
propionic
HC3H5O2
4.87
salicylic
HC7H5O3
1.96
-3
4.19
selenic
H2SeO4
1.66
selenous
H2SeO3
2.64
8.27
succinic
H2C4H4O4
4.21
5.64
sulfuric
H2SO4
none
1.92
sulfurous
H2SO3
1.89
7.21
thiophenol
HSC6H5
6.49
trichloroacetic
HC2Cl3O2
0.52
zinc ion
Zn2+
8.96
12.67
base
formula
pKb1
ammonia
NH3
4.76
aniline
C6H5NH2
9.37
codeine
C18H21O3N
6.05
diethylamine
(C2H5)2NH
4.51
dimethylamine
(CH3)2NH
3.23
ethylamine
C2H5NH2
3.36
hydrazine
N2H4
5.77
hydroxylamine
HONH2
9.04
methylamine
CH3NH2
3.38
morphine
C17H19O3N
6.13
piperidine
C5H11N
2.88
pyridine
C5H5N
8.70
quinoline
C9H7N
9.20
triethanolamine
C6H15O3N
6.24
triethylamine
(C2H5)3N
3.28
trimethylamine
(CH3)3N
4.20
pKb2
15.05
We can say that acid dissociation constants are used to classify acids as weak, moderate, or
strong according the following criteria:
KA < 10-3
weak acid
KA = 1 - 10-3 moderate acid
KA > 1
strong acid
Hydrolysis of the cation (reaction with water in the solution). It should be noted that most of the
transition metal cations and ammonium ion during of their hydrolysis in water solution, generate
H3O+ ions and lower the pH of the solution. For instance aqueous solution of ammonium
chloride, NH4Cl, its NH4+ generates in this equilibrium H3O+ (aq)
NH4+
+ H2O
 NH3 (aq) + H3O+ (aq)
Metal ions from group IA or IIA (except Be2+) do not hydrolyze.
Other metal ions as well as NH4+ to hydrolyze and generate H+
Base dissociation (ionization) constant let us compare the strengths of weak bases. Strong base
like sodium chloride, dissociate completely in water to release OH- ions.
NaOH (s)
→ Na+ (aq) + OH- (aq)
Other, even stronger bases, like the oxide ion in sodium oxide, react completely with water and
generate OH- ions: Na2O (s) + H2O → 2 Na+ (aq) + 2OH- (aq)
Weak bases, like ammonia or the bicarbonate ion, react incompletely with water, usually to a
small percentage, to make some OH- . An equilibrium is established in which the unchanged
base is favored.
NH3(aq) + H2O  NH4+ (aq) + OH- (aq)
HCO3- (aq) + H2O  H2CO3 (aq) + OH- (aq)
Weak bases vary considerably in their abilities to accept protons from water molecules and
generate hydroxide ions. To compare these abilities, we use base dissociation (ionization)
constant, KB. Base is symbolized by B, regardless of its electrical charge:
B (aq) (weak base) + H2O  BH+ (aq)(conjugate acid) + OH- (aq)
The base dissociation constant for this equilibrium, KB is defined by the following equation :
KB = [BH+] [OH-] / [B]
The smaller the KB, the weaker base and in turn, the larger the KB , the stronger the base.
Hydrolysis of anions. It should be noted again, that the reaction of an anion with water to
produce OH- and we call it the hydrolysis of the anion and generate an excess of OH- over H+.
HCO3- (aq) + H2O  H2CO3 (aq) + OH- (aq)
Anions whose conjugate acids are weak acids hydrolyze in water and tend to make the
solution basic.
Anions of strong acids do not hydrolyze.
With the previous rules we can generally predict correctly whether a given salt will affect the pH
of an aqueous solution.
pKA and pKB concept. For the same reason that the pH concept was devised, analogously pKA
and pKB have been defined.
pKA = - log KA
pKB = - log KB
For a conjugate acid-base pair, the product of KA and KB is KW A very simple relationship
exists between KA and KB when we work with conjugate acid-base pair.
KA KB = KW = 10-14
and thus we can write :
pKA + pKB = 14
9.6. Buffers
The pH of a solution can be held relatively constant if it contains a buffer - a weak and its
conjugate acid. So critical is the maintenance of the pH of the blood, which can not be allowed to
change by more than 0.2 to 0.3 pH units from its normal value 7.35.
Acidosis - if the pH of blood becomes lower - the acidity is increasing (untreated diabetes,
emphysema)
Alkalosis - means that the blood is tending to become more basic (overdose of bicarbonate,
exposure to to the low partial pressure of oxygen at high altitudes, prolonged hysteria)
Buffers prevent serious changes in pH to a minimum when strong acids or bases are added to an
aqueous solution. One part of the buffer system can neutralize H+, and the other part can
neutralize OH-. The blood and other body fluids include buffers, and much of the body´s work in
maintaining its acid-base status depends on buffers.
Buffers hold a pH steady, but not necessarily at pH 7. We´ll first explore the relationship
between the pH of a buffered solution and the concentrations and relative acid-base strengths of
its components. To make the discussion general, we assume that the buffer is made by dissolving
some weak acid, HA, together with some of its sodium salt, NaA, in water. We use a group IA
salt, like the sodium or potassium salt, because we generaly want a fully soluble salt, 100%
ionized - makes available the maximum amount of the conjugate base, A-, of the weak acid.
The HA/A- type of buffer system could involve any one of a number of weak acids of widely
varying KA values. So we cannot expect just one weak acid to work for the buffering acids of all
ranges of pH. We therefore need an equation to tell us at what pH a specific HA/A- buffer
system will work. KA for weak acid :
KA = [H+][A-]/ [HA]
Because we want to know [H+] and then pH, we aarange this equation to
[H+] = KA x [HA]/ [A-]
All molar concentrations are the concentrations at equilibrium after the solution was prepared
(not initial concentrations but we can use initial concentrations for simplicity.)
The pH of a buffer solution can also be found by Henderson/Hasslbalch equation (logaritmic
scale).
Henderson-Hasselbalch equation (H-H equation)
pH = pKA + log [anion]/[acid]
Buffers hold a pH steady, but not necessarily at pH 7.
The ratio [Anion]/[Acid] dominates the pH once a buffer pair is selected.
H-H equation very clearly shows that two factors govern the pH of a buffered solution. The first
is pKA of the weak acid in the buffer pair and the second is the ratio [anion]/[acid]. To decide
which weak acid and salt to use in a buffer, we first decide the pH that we want to protect. Then
we look for a weak acid with the pKA as close to it as possible.
The phosphate buffer is important within cells. It consists of the pair of ions, HPO42- and
H2PO4-. Any added OH- is neutralized by H2PO4- :
H2PO4- (aq) + OH- (aq) →
HPO42- (aq) + H2O
The proton acceptor or base in the phosphate buffer is the conjugate base of H2PO4-, the
HPO42- ion. It can neutralize H+ and so keep the pH from decreasing.
HPO42- (aq) + H+ (aq) →
H2PO4- (aq)
The carbonate buffer is the principal buffer in the blood. Traditionally, it has been described as
the conjugate pair, H2CO3 and HCO3- , carbonic acid and bicarbonate ion. Actually, the carbonic
acid in blood is almost entirely in the form of CO2 (aq). For every molecule of H2CO3, there are
400 molecules of CO2 (aq), however, is able to neutralize hydroxide ion directly by the
following reaction
CO2 (aq) + OH-(aq) →
HCO3- (aq)
Actually, this is forward step in an equilibrium
CO2 (aq) + OH-(aq)

HCO3- (aq)
The equilibrium of CO2 (aq) and OH- (aq) with HCO3- (aq) is catalyzed by what is one of the
most rapidly acting enzymes, carbonic anhydrase.Because of this enzyme’ s work, we can use
CO2(aq) as a stand/in for H2CO3(aq) in discussing the carbonate buffer. The CO2 (aq) of this
buffer can neutralize OH-, as we just said, and thus prevent an increase in pH. If there should be
some metabolic or respiratory problem that increases the level of conc. OH-, this OH- is
neutralized and alkalosis is prevented. The bicarbonate ion is the base of the blood’s carbonate
buffer. When the blood’s level of H+ increases, the H+ is neutralized by HCO3- and acidosis is
prevented.
HCO3- (aq) + H+ (aq) →
CO2 (aq) + H2O
When acid is neutralized by the carbonate buffer, the bloo’s level of CO2(aq) increases. During
moving of the blood through the capillaries in the lungs, gaseous CO2 is released from dissolved
CO2 and breathed out. Because the gas leaves, we cannot write this change as an equilibrium.
CO2 (aq) →
CO2 (g)
The loss of one molecule of CO2 (g) by this change means that the H+ ion neutralized by the
buffer is now permanently neutralized and ends in the molecule of water.
The blood thus neutralizes H+ by the work of carbonate buffer, and it uses physical process –
ventilation- to make this neutralization permanent. One control over ventilation is a site in the
brain called, the respiratory center, which monitors the CO2(aq) in the blood.
Buffers minimize but do not completely prevent pH changes. Dissolved buffer must be factored
into a Henderson-Hasselbalch equation and therefore H-H equation must be slightly modified.
The modified equation :
pH = pK’ + log [HCO3- (aq)] / [CO2(aq)]
We can’t use the usual symbol pKA, for two reasons. First, the bloodstream is not at 250C but at
the body temperature, 370C. Second, the buffer’s acid component is not carbonic acid but carbon
dioxide. So an apparent acid ionization (dissociation) constant is K’, and its pK’ has value 6. 1.
Then the pH of carbonate buffer in the body becomes:
pH = 6.1 + log[HCO3- (aq)] / [CO2(aq)]
Normally, in human arterial blood, [HCO3-] equals 24 mmol/L, and CO2 equals 1.2 mmol/L.
These data lead to pH value = 7.4
Let’s suppose that the blood is challenged with a sudden influx of acid in the equivalent of 10
mmol/L of HCl per liter of blood. This would neutralize 10 mmol/L and so reduce its
concentration from 24 mmol/L to 14 mmol/L. The HCO3- changes almost entirely to CO2(aq), so
10 mmol/L of new CO2(aq) appears in the blood. If this new CO2 could not be removed by
breathing, its level would increase by 10 mmol/L, thus to 11.2 mmol/L CO2. pH then would be
6.2.far, far too low to permit life to continue. But when all the CO2 will be breathed out, as CO2
(g). Although the level of HCO3- remains !$ mmol/L, the level of CO2(aq) quickly drops back to
1.2 mmol/L, so the pH will go to 7.2, which is stilltoo low but it doesn’ cause death. The body, if
healthy, responds quickly to a lowering of blood pH by increasing the rate of breathing, by
causing hyperventilation. This forces even more CO2 out of the blood into the exhaled air. Thus
two mechanisms have protected the system against the otherwise lethal problem. The buffer
system has neutralized the acid, and the respiratory system has readjusted the ratio [HCO3- (aq)] /
[CO2(aq)].
The situation just described is not actually back to normal in all respects. By normal metabolism,
the system will continue to produce CO2. This will increase the denominator in this ratio [HCO3(aq)] / [CO2(aq)], and so the pHwill gradually declaine again. The system must, therefore, also
raise the level of HCO3-, and this is done principally by the chemical work of the kidneys.
Calculations
A buffer solution was prepared from 0.085M formic acid, and sodium formate, dissolved in the
same solution at a concentration of 0.12 M. Calculate the pH of this solution. For formic acid,
KA = 1.8 x 10-4
A buffer has composition 0.1 M CH3COOH a 0.5 M CH3COOH. For acetic acid, KA = 1.8 x 10-5.
Calculate the pH of this solution.
Buffer was prepared by mixing of 6 volumetric parts of NaH2PO4 and of 4 volumetric parts of
Na2HPO4 . Define the pH of this solution if you know the dissociation constants of phosphoric
acids (KA)I = 7.52 x 10-3, (KA)II = 6.23 x 10-8, (KA)III = 1.78 x 10-12.
Adapted from
Holumn : Fundamentals of General. Organic and Biological Chemistry
Voet : Biochemistry
Koolman, Roehm: Color Atlas of Biochemistry