BIOENERGITICS

BIOENERGETICS

Biochemical thermodynamics or biochemical energetic or bioenergetics’ concerned with the transformation and use of energy by living cells, The chemical reactions occurring in the living beings are associated with the liberation of energy as the reacting system moves from a higher to a lower energy level.
In non biologic system the heat energy may be transformed into mechanical or electrical energy. Since biologic systems are exothermic, the heat energy cannot be used to drive the vital processes (such as e.g. Active transport, nerve conduction, muscular contraction etc) obtain energy by chemical linkage to oxidation reactions. In biological systems the energy production and utilization is always done in coupling where a spontaneous reaction takes place and same time non-spontaneous reaction takes place by utilizing the energy produced.

The concept of free energy
Change in free energy is the portion of the total energy change in a system which is available for total energy change in a system which is available for doing work, i.e., useful energy. The energy coupling occurs by coupling of Exergonic and endergonic reactions and liberation of heat.
If any reactions go from right to left with an intermediate which has a similarity with the structure of the reactants and productents

The simplest type of coupling may be represented as

A+C -------I ----------- B+D

Some Exergonic and endergonic reactions in biologic systems are coupled in this way. It should be appreciated that this type of system has a built in mechanism for biologic control of the rate at which oxidative processes are allowed to occur since the existence of the common obligatory intermediate for both the Exergonic and endergonic reactions.

An alternative method of coupling an Exergonic an endergonic process is to synthesize a compound of high-energy potential in the exergonic reaction and to incorporate this new compound into the endergonic reaction. Here an intermediate high energy compound is formed which is utilized by endergonic reaction eg formation of ATP for the metabolic cycles and utilization by the endergonic reactions. Transduction of energy through a common high-energy compounds is adenosine triphosphate.






The laws of thermodynamics

There are three fundamental laws of thermodynamics, which for historical reasons are known as the zero. The first, second, and third laws. The ones of immediate relevance to biochemistry are the first and second laws. These can be articulated in a variety of ways, but for our purposes:

First law of thermodynamics

The first law of thermodynamics sates that the total amount of energy is neither created nor destroyed. I another words the amount of energy is constant in universe. It states that energy is used to do work or is converted from one form to another form. The mathematical expression of first law is

DE = EB -- EA = Q -- W

DE Change in internal energy
EB Energy of a system at the start of the process.
EA Energy of a system at the end of the process.
Q Heat absorbed by the system.
W Work done by the system.

Second law of thermodynamics

The second law of thermodynamics says that the entropy in a closed system increases. The entropy of a system must increase if

ΔS(system)+ ΔS(surroundings)>0



Third law of thermodynamics

This is the combination of two laws of thermodynamics and the representation of the equation
ΔG = ΔH - TΔS
Josiah Gibbs articulated the concept of free energy (sometimes called Gibbs free energy), which is related to entropy and enthalpy by ΔG = ΔH - TΔS
The change in free energy when a reaction occurs is
ΔG = ΔH – TΔS

ΔG the change in free energy of a reacting system,
ΔH the change in heat content or enthalpy of this system
T the absolute temperature at which the process is taking place
ΔS the change in entropy of this system.

The TΔS is a fraction of ΔH which cannot be put to useful work. The ΔG
Indicates the free energy change or the theoretically available use full work.

ΔH = ΔE + PΔV
As the volume change is very small for biochemical reactions , hence ΔH is nearly equal to the change in internal energy, ΔE therefore the equation is modified to ΔG= ΔE- TΔS

If ΔG is negative reaction is spontaneous with loss of energy, ( exergonic) and irreversible reaction.

Entropy
Entropy represents the extent of disorder or randomness of the system increases as a system approaches towards equilibrium under constant condition of temperature and pressure.
We've already said that entropy is a measure of the disorder in a system. The second law of thermodynamics says that in general the entropy of a closed system will increase.
(What happens when molecules go into solution? The solute molecules usually undergo an increase in entropy, because they become free to dissociate from one another, and in the case of ionic solutes the cations can separate from the anions. On the other hand, the solvent molecules frequently become more organized in the vicinity of the solute molecules than they had been before the introduction of the solute, so their contribution to total change in entropy is frequently negative. The net effect is often slightly negative, i.e. the solution has slightly lower entropy than the separated components.)The entropy of single molecule can be characterized by statistical-mechanical methods if the molecule is simple enough. Zubay's Principles of Biochemistry breaks the entropy of liquid propane into translational, rotational, vibrational, and electronic components:
Enthalpy
The enthalpy is the heat content of the reacting system. It reflects the number and kinds of chemical bonds in the reactants and products when a chemical reaction releases heat, it is said to be exothermic, the heat content of the product is less than that of the reactants and

The relation between enthalpy and entropy are given equation are

ΔG= ΔH- TΔS

Relationship between free energy and equilibrium

Any biochemical reaction a+b ---------------à c+d
The free energy of the above reaction is calculated by the equation
ΔG = ΔGo+ RTln([c]c[d]d)/([a]a[b]b)


At equilibrium the ΔG = 0 0 = ΔGo + ([c]c[d]d)/([a]a[b]b)where Keq is the equilibrium constant of the reaction. In a bimolecular reaction a + b -> c + dthis equilibrium constant is Keq = + ([c]c[d]d)/([a]a[b])so

ΔGo = -RTlnKeq

Thus if a reaction is just barely spontaneous, i.e. ΔGo = 0, then Keq = 1.
If ΔGo <> 1, there will be more products than reactants at equilibrium. If ΔGo > 0 then Keq < 1, there will be more reactants than products at equilibrium.
Reactions in which ΔG o < 0 are called exergonic;
Reactions in which ΔGo > 0 are called endergonic.


Types of reactions:

Exergonic/ Exothermic: These reactions will liberate free energy phosphate groups these are spontaneous reactions and provide some energy for performing some work. eg. ATP is hydrolyzed to form adenosine diphosphate & phosphoric acid. This reaction provides -7300 calories/mole of free energy at pH 7.

Endergonic /Exothermic: This reaction mechanism is not spontaneous energy has to be provided for these reactions. eg. Glucose is phosporylated glucose-6-phosphate where for this reaction 5500 calories of energy has to be supplied.


Two examples of such coupling are:
Active transport. Spontaneous ATP hydrolysis (negative DG) is coupled to (drives) ion flux against a gradient (positive DG). For an example, see the discussion of SERCA
ATP synthesis in mitochondria. Spontaneous H+ flux across a membrane (negative DG) is coupled to (drives) ATP synthesis (positive DG). See the discussion of the ATP Synthase



High Energy Compounds
ATP is used by cells to drive many energy consuming reactions. It is a high energy compound, because it has a large negative free energy of hydrolysis under physiological conditions. Compounds with DeltaG more negative than 7 kcal/mole may be regarded as high energy compounds.
ATP is present in cells at 1 to 10mM, it is anionic, carrying four negative charges at pH 7.0, and is neutralized by complexing with Mg2+. ATP is sometimes described as the universal energy currency of living cells - an exaggeration. In bacteria such as E. coli, energy is provided by:
ATP - most biosynthetic reactions, some transport systems
GTP - e.g. protein synthesis
Thioesters - e.g. fatty acid synthesis
PEP - e.g. phosphotransferase system
Proton Motive Force - e.g. flagella, some transport systems
These other energy sources do not necessarily depend on ATP for their production. In fact, ATP may be generated from these other forms of energy.
It can mean the phosphate-phosphatebonds formed when compounds such as adenosine diphosphate and adenosine triphosphate are created. It can mean the compounds which contain these bonds, which include the nucleoside diphosphates and nucleoside triphosphates, and the high energy storage compounds of the muscle, the phosphagens. When people speak of a high energy phosphate pool, they speak of the total concentration of these compounds with these high energy bonds.
High energy phosphate bonds are pyrophosphate bonds, acid anhydride linkages, formed by taking phosphoric acid derivatives and dehydrating them. As a consequence, the hydrolysis of these bonds is exothermic under physiological conditions, releasing energy.
Energetic of High Energy Phosphate Reactions
Reaction
Δ G in kilojoules per mole
ATP + H2O → ADP + Pi
-36.8
ADP + H2O → AMP + Pi
-36.0
ATP + H2O → AMP + PPi
-40.6
PPi → 2 Pi
-31.8
AMP + H2O → A + Pi
-12.6
Except for PPi → 2 Pi, these reactions are generally not allowed to go uncontrolled in the human cell, but generally are coupled to other processes needing energy to drive them to completion. So, high energy phosphate reactions can
provide energy to cellular processes, to allow them to run by coupling processes to a particular nucleoside, allow for regulatory control of the process drive the reaction to the right, by taking a reversible process and making it irreversible.
The one exception is of value because it allows a single hydrolysis, ATP + 2H2O → ADP + PPi, to effectively supply the energy of hydrolysis of two high-energy bonds, with the hydrolysis of PPi being allowed to go to completion in a separate reaction. The AMP is regenerated to ATP in two steps, with the equilibrium reaction ATP + AMP ↔ 2ADP, followed by regeneration of ATP by the usual means, oxidative phosphorylation or other energy-producing pathways such as glycolysis.
Often, high energy phosphate bonds are denoted by the character '~'. In this notation, ATP becomes A-P~P~P.
Adenosine triphosphate (ATP) is considered by biologists to be the energy currency of life. It is the high-energy molecule that stores the energy we need to do just about everything we do. It is present in the cytoplasm and nucleoplasm of every cell, and essentially all the physiological mechanisms that require energy for operation obtain it directly from the stored ATP. (Guyton) As food in the cells is gradually oxidized, the released energy is used to re-form the ATP so that the cell always maintains a supply of this essential molecule. ATP is remarkable for its ability to enter into many coupled reactions, both those to food to extract energy and with the reactions in other physiological processes to provide energy to them. In animal systems, the ATP is synthesized in the tiny energy factories called mitochondria.
The structure of ATP has an ordered carbon compound as a backbone, but the part that is really critical is the phosphorous part - the triphosphate. Three phosphorous groups are connected by oxygens to each other, and there are also side oxygens connected to the phosphorous atoms. Under the normal conditions in the body, each of these oxygens has a negative charge, and as you know, electrons want to be with protons - the negative charges repel each other. These bunched up negative charges want to escape - to get away from each other, so there is a lot of potential energy here.
If you remove just one of these phosphate groups from the end, so that there are just two phosphate groups, the molecule is much happier. This conversion from ATP to ADP is an extremely crucial reaction for the supplying of energy for life processes. Just the cutting of one bond with the accompanying rearrangement is sufficient to liberate about 7.3 kilocalories per mole = 30.6 kJ/mol. This is about the same as the energy in a single peanut.
Living things can use ATP like a battery. The ATP can power needed reactions by losing one of its phosphorous groups to form ADP, but you can use food energy in the mitochondria to convert the ADP back to ATP so that the energy is again available to do needed work. In plants, sunlight energy can be used to convert the less active compound back to the highly energetic form. For animals, you use the energy from your high energy storage molecules to do what you need to do to keep yourself alive, and then you "recharge" them to put them back in the high energy state. The oxidation of glucose operates in a cycle called the Krebs cycle in animal cells to provide energy for the conversion of ADP to ATP
Potentially two high energy bonds can be cleaved, as two phosphates are released by hydrolysis from ATP (adenosine triphosphate), yielding ADP (adenosine diphosphate), and ultimately AMP (adenosine monophosphate). This may be represented as follows (omitting waters of hydrolysis):
AMP~P~P ® AMP~P + Pi (ATP ® ADP + Pi)
AMP~P ® AMP + Pi (ADP ® AMP + Pi)
Alternatively, as discussed above:
AMP~P~P ® AMP + P~Pi (ATP ® AMP + PPi)
P~P ® 2 Pi
Artificial ATP analogs have been designed that are resistant to cleavage of the terminal phosphate by hydrolysis, e.g., AMPPNP, depicted at right.
Such analogs have been used to study the dependence of coupled reactions on ATP hydrolysis. In addition, they have made it possible to crystallize an enzyme that catalyzes ATP hydrolysis with an ATP analog at the active site.
Many organisms store energy as inorganic polyphosphate, a chain of many phosphate residues linked by phosphoanhydride bonds. It may be represented as: P~P~P~P~P... Hydrolysis of Pi residues from polyphosphate may be coupled to energy-dependent reactions. Depending on the organism or cell type, inorganic polyphosphate may have additional functions. For example, it may serve as a reservoir for Pi, a chelator of metal ions, a buffer, or a regulator.
Why do phosphoanhydride linkages have a high free energy of hydrolysis? Contributing factors for ATP and PPi are thought to include:
Resonance stabilization of the products of hydrolysis exceeds resonance stabilization of the compound itself. See Fig. 16-22 p. 568.
Electrostatic repulsion between negatively charged phosphate oxygens favors separation of the phosphates.
Phosphocreatine (also called creatine phosphate), another compound with a "high energy" phosphate linkage, is used in nerve and muscle cells for storage of ~P bonds.
Creatine Kinase catalyzes the reversible reaction:
phosphocreatine + ADP « ATP + creatine
Phosphocreatine is produced when ATP levels are high. During exercise in muscle, phosphate is transferred from phosphocreatine to ADP, to replenish ATP. Phosphocreatine may also be used to transport "high energy" phosphate from one compartment of a cell to another.
A reaction that is important for equilibrating ~P among adenine nucleotides within a cell is that catalyzed by Adenylate Kinase:
ATP + AMP « 2 ADP
The Adenylate Kinase reaction is also important because the substrate for ATP synthesis, e.g., by the mitochondrial ATP Synthase, is ADP, while some cellular reactions dephosphorylate ATP all the way to AMP.
The enzyme Nucleoside Diphosphate Kinase (NuDiKi) equilibrates ~P among the various nucleotides that are needed, e.g., for synthesis of DNA and RNA. NuDiKi catalyzes reversible reactions such as:
ATP + GDP « ADP + GTP , ATP + UDP « ADP + UTP , etc.
Phosphoenolpyruvate (PEP), involved in production of ATP in Glycolysis, has a larger negative DG of phosphate hydrolysis than ATP.
Removal of phosphate from the ester linkage in PEP is spontaneous because the enol product spontaneously converts to a ketone.
The ester linkage in PEP is an exception. Generally phosphate esters (formed by splitting out water between a phosphoric acid and a hydroxyl group) have a low but negative DG of hydrolysis. Examples, shown below, include:
the linkage between the first phosphate of ATP and the ribose hydroxyl
the linkage between Pi and a hydroxyl group in glucose-6-phosphate or glycerol-3-phosphate.
ATP has special roles in energy coupling and phosphate transfer. The free energy of hydrolysis of phosphate from ATP is intermediate among the examples listed in the table below (more complete table p. 566). ATP can thus act as a phosphate donor, and ATP can be synthesized by transfer of phosphate from other compounds, such as phosphoenolpyruvate (PEP).
Compound
DGo' of phosphate hydrolysis (kJ/mol)
Phosphoenolpyruvate (PEP)
- 61.9
Phosphocreatine
- 43.1
Pyrophosphate
- 33.5
ATP (to ADP)
- 30.5
Glucose-6-phosphate
- 13.8
Glycerol-3-phosphate
- 9.2

Some other "high energy" bonds:
A thioester forms between a carboxylic acid and a thiol (SH) group, e.g., the thiol of coenzyme A (abbreviated CoA-SH).
Thioesters are "high energy" linkages. In contrast to phosphate esters, thioesters have a large negative DG of hydrolysis.

The thiol of coenzyme A can react with a carboxyl group of acetic acid (yielding acetyl-CoA) or a fatty acid (yielding fatty acyl-CoA).
The spontaneity of thioester cleavage is essential to the role of coenzyme A as an acyl group carrier. Like ATP, acyl-coenzyme A has a high group transfer potential.

Coenzyme A includes b-mercaptoethylamine, in amide linkage to the carboxyl group of the B vitamin pantothenate.
The hydroxyl of pantothenate is in ester linkage to a phosphate of ADP-3'-phosphate.
The functional group is the thiol (SH) of b-mercaptoethylamine.

3',5'-Cyclic AMP (abbreviated cAMP), shown at right and below, is used by cells as a transient signal.
Adenylate Cyclase (Adenylyl Cyclase) catalyzes cAMP synthesis:ATP ® cAMP + PPi. The reaction is highly spontaneous due to the production of PPi, which spontaneously hydrolyzes.
Phosphodiesterase catalyzes catalyzes hydrolytic cleavage of one of the phosphate ester linkages (in red), converting cAMP ® 5'-AMP. This is a highly spontaneous reaction, because cAMP is sterically constrained by having a phosphate with ester linkages to two hydroxyls of the same ribose. The lability of cAMP to hydrolysis makes it an excellent transient signal. Signal roles of cAMP will be discussed separately.

Distinction between thermodynamics and kinetics: A high activation energy barrier usually causes hydrolysis of a "high energy bond" to be extremely slow in the absence of an enzyme catalyst. This "kinetic stability" is essential to the role of ATP and other compounds with ~ bonds. If ATP would rapidly hydrolyze in the absence of a catalyst, it could not serve its important roles in energy metabolism and phosphate transfer. Phosphate is removed from ATP only when the reaction is coupled to some other reaction useful to the cell, such as transport of ions or phosphorylation of glucose.
Oxidation & reduction will be covered in a later class. A brief introduction to selected topics will be presented here.
Oxidation of an iron atom involves loss of an electron (to some acceptor atom):
Fe++ (reduced) ® Fe+++ (oxidized) + e-
Oxidation of carbon is a spontaneous (energy yielding) reaction.
NAD+ (Nicotinamide Adenine Dinucleotide) functions as an electron acceptor in catabolic pathways.
The nicotinamide ring of NAD+, which is derived from the vitamin niacin, accepts 2 e- and one H+ (a hydride) in going to the reduced state, as NAD+ becomes NADH. See also p. 461 & 555.
NADP+/NADPH is similar, except for an additional phosphate esterified to a hydroxyl group on the adenosine ribose. NADPH functions as an electron donor in synthetic pathways.

The electron transfer reaction may be summarized as:
NAD+ + 2 e- + H+ « NADH
It may also be written as:
NAD+ + 2 e- + 2H+ « NADH + H+
FAD (Flavin Adenine Dinucleotide) also functions as an electron acceptor. The portion of FAD that undergoes reduction/oxidation is the dimethylisoalloxazine ring, derived from the vitamin riboflavin. See p. 556.
FAD normally accepts 2 e- and 2 H+ in going to its reduced state: FAD + 2 e- + 2 H+ « FADH2
NAD+ is a coenzyme, that reversibly binds to enzymes.
FAD is a prosthetic group, that usually remains tightly bound at the active site of an enzyme

pH, pka, Henderson Hasselbalch Equation

Definition of Keq, Kw and pH
Definition of pKa
The Henderson-Hasselbalch Equation
Definition of Buffering
Significance of Blood Buffering
Ampholyte, Polyampholyte, pI and Zwitterion
Definition of Solvation and Hydration
Kidneys and Acid-Base Balance
Sodium Bicarbonate Reabsorption
Excretion of Acid
Ammonia Secretion
Neurotoxicity of Ammonia
Acidosis and Alkalosis

Ionic Equilibrium

Keq, Kw and pH
As H2O is the medium of biological systems one must consider the role of this molecule in the dissociation of ions from biological molecules. Water is essentially a neutral molecule but will ionize to a small degree. This can be described by a simple equilibrium equation:
H2O <-------> H+ + OH- Eqn. 1
This equilibrium can be calculated as for any reaction:
Keq = [H+][OH-]/[H2O] Eqn. 2
Since the concentration of H2O is very high (55.5M) relative to that of the [H+] and [OH-], consideration of it is generally removed from the equation by multiplying both sides by 55.5 yielding a new term, Kw:
Kw = [H+][OH-] Eqn. 3
This term is referred to as the ion product. In pure water, to which no acids or bases have been added:
Kw = 1 x 10-14 M2 Eqn. 4
As Kw is constant, if one considers the case of pure water to which no acids or bases have been added:
[H+] = [OH-] = 1 x 10-7 M Eqn. 5
This term can be reduced to reflect the hydrogen ion concentration of any solution. This is termed the pH, where:
pH = -log[H+] Eqn. 6
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pKa
Acids and bases can be classified as proton donors and proton acceptors, respectively. This means that the conjugate base of a given acid will carry a net charge that is more negative than the corresponding acid. In biologically relavent compounds various weak acids and bases are encountered, e.g. the acidic and basic amino acids, nucleotides, phospholipids etc.
Weak acids and bases in solution do not fully dissociate and, therefore, there is an equilibrium between the acid and its conjugate base. This equilibrium can be calculated and is termed the equilibrium constant = Ka. This is also sometimes referred to as the dissociation constant as it pertains to the dissociation of protons from acids and bases.
In the reaction of a weak acid:
HA <-----> A- + H+ Eqn. 7
the equlibrium constant can be calculated from the following equation:
Ka = [H+][A-]/[HA] Eqn. 8
As in the case of the ion product:
pKa = -logKa Eqn. 9
Therefore, in obtaining the -log of both sides of the equation describing the dissociation of a weak acid we arrive at the following equation:
-logKa = -log[H+][A-]/[HA] Eqn. 10
Since as indicated above -logKa = pKa and taking into account the laws of logrithms:
pKa = -log[H+] -log[A-]/[HA] Eqn. 11
pKa = pH -log[A-]/[HA] Eqn. 12
From this equation it can be seen that the smaller the pKa value the stronger is the acid. This is due to the fact that the stronger an acid the more readily it will give up H+ and, therefore, the value of [HA] in the above equation will be relatively small. __________________________________________________________


Clinical Significance of Blood Buffering
The pH of blood is maintained in a narrow range around 7.4. Even relatively small changes in this value of blood pH can lead to severe metabolic consequences. Therefore, blood buffering is extremely important in order to maintain homeostasis. Although the blood contains numerous cations (e.g., Na+, K+, Ca2+ and Mg2+) and anions (e.g., Cl-, PO43- and SO42-) that can, as a whole, play a role in buffering, the primary buffers in blood are hemoglobin in erythrocytes and bicarbonate ion (HCO3-) in the plasma. Buffering by hemoglobin is accomplished by ionization of the imidazole ring of histidines in the protein.
The formation of bicarbonate ion in blood from CO2 and H2O allows the transfer of relatively insoluble CO2 from the tissues to the lungs, where it is expelled. The major source of CO2 in the tissues comes from the oxidation of ingested carbon compounds.
Carbonic acid is formed from the reaction of dissolved CO2 with H2O. The relationship between carbonic acid and bicarbonate ion formation is shown in equations 16 and 17.

CO2 + H2O <----> H2CO3 Eqn. 16
H2CO3 <----> H+ + HCO3- Eqn. 17
The reactions shown in equations 16 and 17 occur predominately in the erythrocytes, since nearly all of the CO2 leaving tissues via the capillary endothelium is taken up by these cells. This reaction is catalyzed by carbonic anhydrase. Ionization of carbonic acid then occurs spontaneously (as shown in equation 17), yielding bicarbonate ion.
Carbonic acid is a relatively strong acid with a pKa of 3.8. However, carbonic acid is in equilibrium with dissolved CO2. Therefore, the equilibrium equation for the sum of equations 16 and 17 requires a conversion factor, since CO2 is a dissolved gas. This factor has been shown to be approximately 0.03 times the partial pressure of CO2 (PCO2). When this is entered into the Henderson-Hasselbalch equation:

pH = 6.1 + log [HCO3-/(0.03)(PCO2)] Eqn. 18
where the apparent pKa for bicarbonate formation, 6.1, has been introduced into equation 18.
The PCO2 in the peripheral tissues is approximately 50mm Hg, whereas in the blood entering the peripheral tissues it is approximately 40mm Hg. This difference results in the diffusion of CO2 from the tissues into the blood in the capillaries of the periphery. When the CO2 is converted to H2CO3 within the erythrocytes and then ionizes, the hydrogen ions (H+) are buffered by hemoglobin. The production of H+ ions, within erythrocytes, and their subsequent buffering by hemoglobin results in a reduced affinity of hemoglobin for oxygen. This leads to a release of O2 to the peripheral tissues, a phenomenon is termed the Bohr effect.

Representation of the transport of CO2 from the tissues to the blood with delivery of O2 to the tissues. The opposite process occurs when O2 is taken up from the alveoli of the lungs and the CO2 is expelled. All of the processes of the transport of CO2 and O2 are not shown such as the formation and ionization of carbonic acid in the plasma. The latter is a major mechanism for the transport of CO2 to the lungs, i.e. in the plasma as HCO3-. The H+ produced in the plasma by the ionization of carbonic acid is buffered by phosphate (HPO42-) and by proteins. Additionally, some 15% of the CO2 is transported from the tissues to the lungs as hemoglobin carbamate as shown in Eqn. 19.

As CO2 passes from the tissues to the plasma a minor amount of carbonic acid takes form and ionizes. The H+ ions are then buffered predominantly by proteins and phosphate ions in the plasma. As the concentration of bicarbonate ions rises in erythrocytes, an osmotic imbalance occurs. The imbalance is relieved as bicarbonate ion leaves the erythrocytes in exchange for chloride ions from the plasma. This phenomenon is known as the chloride shift which is also shown in the diagram above. Therefore, the majority of the bicarbonate ion formed as CO2 leaves the peripheral tissues is transported by the plasma to the lungs.
Around 15% of CO2 transport from the tissues to the lungs occurs through a reversible combination with non-ionized amino groups (-NH2) of hemoglobin forming what is termed hemoglobin carbamate.
Hemoglobin-NH2 + CO2 <---> Hemoglobin-NH-COO- + H+ Eqn. 19
The formation of hemoglobin carbamate results in a reduced affinity of hemoglobin for O2 thus favoring dissociation of bound oxygen in the tissues where the concentration of CO2 is high. The process is reversed when the erythrocytes enter the lungs and the partial pressure of O2 is elevated.
The partial pressure of O2 (PO2) in the pulmonary alveoli is higher than the PO2 of the entering erythrocytes that contain predominantly deoxygenated hemoglobin. This increased PO2 leads to oxygenation of hemoglobin and release of H+ ions from the hemoglobin. The released H+ ions combine with the bicarbonate ions to form H2CO3. Cellular carbonic anhydrase then catalyzes the reverse of reaction 17, leading to release of CO2 from erythrocytes. Owing to the PCO2 gradient (described above), the CO2 diffuses from the blood to the alveoli where it is expelled.
The great utility of bicarbonate as a physiological buffer stems from the fact that if excess acid is added to the blood the concentration of bicarbonate ion declines and the level of CO2 increases. The CO2 then passes from capillaries in the pulmonary alveoli and is expelled. As a consequence, the H+ ion concentration drives reaction 17 to the left and bicarbonate ion acts as a buffer until all of the hydrogen ion is consumed. Conversely, when excess base is added to the blood, CO2 is consumed by carbonic acid and replaced by metabolic reactions within the body.
If blood is not adequately buffered, the result may be metabolic acidosis or metabolic alkalosis. These physiological states can be reached if a metabolic defect results in the inappropriate accumulation or loss of acidic or basic compounds. These compounds may be ingested, or they may accumulate as metabolic by-products such as acetoacetic acid and lactic acid. Both of these will ionize, thereby increasing the level of H+ ions that will in turn remove bicarbonate ions from the blood and alter blood pH. The predominant defect in acid or base elimination arises when the excretory system of the kidneys is impaired. Alternatively, if the lungs fail to expel accumulated CO2 adequately and CO2 accumulates in the body, the result will be respiratory acidosis. If a decrease in PCO2 within the lungs occurs, as during hyperventilation, the result will be respiratory alkalosis. Bicarbonate as a buffer
The major buffering systems in the body are proteins, particularly those with the amino acids histidine and cysteine exposed to the outside environment, phosphate, and bicarbonate. All three of these are weak acids with pKa values lower than physiological pH. As a consequence, buffering capacity increases as the pH is lowered from the physiological range. This meets the needs of most organisms because physiological pH excursions generally occur in the acid direction. Hence, the low pKa of these buffering systems is poised to respond to metabolic acidosis.
Of these three, only the bicarbonate system, which is critical for buffering extracellular fluids such as blood, is in steady-state between production and removal. Thus, pH changes via this dynamic bicarbonate system are taking place on a background provided by the more static protein and phosphate systems.
Production of Bicarbonate Cells generate and excrete large quantities of carbon dioxide (CO2) during aerobic metabolism of glucose and fats. CO2 is subsequently converted to carbonic acid (H2CO3), which serves as the basis for the bicarbonate buffering system. Hence, the body is not dependent on ingestion of exogenous compounds or complex syntheses to maintain this buffering system.
Removal of Bicarbonate The bicarbonate buffering system is in volatile equilibrium (via breathing) with the external environment (lungs and air). Thus it is able to respond rapidly to endogenous alterations. It can also be affected, either positively or negatively, by environmental manipulation.
Movement into and out of cells
The acid components of the bicarbonate system (i.e. H+ and CO2) cross biological membranes rapidly, thus do not depend on complex transport kinetics. The base component (HCO3-), on the other hand, is transported rapidly in all cells via anion exchange. Consequently, the bicarbonate buffering system helps to maintain both intracellular and extracellular pH.
Role of other blood components
The acid components of the bicarbonate system are transported from the tissues to the lungs by hemoglobin. Thus, this important protein participates in both the production and removal of metabolic acid.
Clinical Correlates: Acidosis & Alkalosis
CO2 produced by metabolism is normally balanced by CO2 expired from the lungs, resulting in no net production of H2CO3. However, as detailed in the table below, certain medically significant circumstances can throw the equation out of balance.
respiratory acidosis
apnea or impaired lung capacity, with a build-up of CO2 in the lungs
metabolic acidosis
ingestion of acid, production of ketoacids in uncontrolled diabetes, or kidney failure
Note: These conditions all result in build-up of H+ from sources other than excess CO2.


Condition
Possible causes
respiratory alkalosis
hyperventilation, with a net loss of CO2 from the blood.
metabolic alkalosis
ingestion of alkali, prolonged vomiting leading to a loss of HCl, or extreme dehydration leading to kidney retention of bicarbonate.
Note: The common thread here is the loss of H+ for reasons other than depletion of CO2.


Ampholytes, Polyampholytes, pI and Zwitterion
Many substances in nature contain both acidic and basic groups as well as many different types of these groups in the same molecule. (e.g. proteins). These are called ampholytes (one acidic and one basic group) or polyampholytes (many acidic and basic groups). Proteins contains many different amino acids some of which contain ionizable side groups, both acidic and basic. Therefore, a useful term for dealing with the titration of ampholytes and polyampholytes (e.g. proteins) is the isoelectric point, pI. This is described as the pH at which the effective net charge on a molecule is zero.
For the case of a simple ampholyte like the amino acid glycine the pI, when calculated from the Henderson-Hasselbalch equation, is shown to be the average of the pK for the a-COOH group and the pK for the a-NH2 group:

pI = [pKa-(COOH) + pKa-(NH3+)]/2 Eqn. 20

For more complex molecules such as polyampholytes the pI is the average of the pKa values that represent the boundaries of the zwitterionic form of the molecule. The pI value, like that of pK, is very informative as to the nature of different molecules. A molecule with a low pI would contain a predominance of acidic groups, whereas a high pI indicates predominance of basic groups. ­­­­­­­­­­­­­­­­­­ _______________________________________________________________

Solvation and Hydration shells
Depending on the pH of a solution, macromolecules such as proteins which contain many charged groups, will carry substantial net charge, either positive or negative. Cells of the body and blood contain many polyelectrolytes (molecules that contain multiple same charges, e.g. DNA and RNA) and polyampholytes that are in close proximity. The close association allows these molecules to interact through opposing charged groups. The presence, in cells and blood, of numerous small charged ions (e.g. Na+, Cl-, Mg2+, Mn2+, K+) leads to the interaction of many small ions with the larger macroions. This interaction can result in a shielding of the electrostatic charges of like-charged molecules. This electrostatic shielding allows macroions to become more closely associated than predicted based upon their expected charge repulsion from one another. The net effect of the presence of small ions is to maintain the solubility of macromolecules at pH ranges near their pI. This interaction between solute (e.g. proteins, DNA, RNA, etc.) and solvent (e.g. blood) is termed solvation or hydration. The opposite effect to solvation occurs when the salt (small ion) concentration increases to such a level as to interfere with the solvation of proteins by H2O. This results from the H2O forming hydration shells around the small ions. ___________________________________________________________

Role of the Kidneys in Acid-Base Balance
The kidneys function to filter the plasma that passes through the nephrons. Filtration of the plasmas occurs in the glomerular capillaries of the nephron. These capillaries allow the passage of water and low molecular weight solutes (less than 70 kDa) into the capsular space. The filtrate then passes through the proximal and distal convoluted tubules where reabsorption of water and many solutes takes place. In the course of glomerular filtration and tubule reabsorption the composition of the plasma changes generating the typical composition of urine. From a biochemical standpoint the kidneys serve important roles in the regulation of plasma acid-base balance and the elimination of nitrogenous wastes.
Sodium Bicarbonate Reabsorption
Regulation of plasma acid-base balance is primarily effected within the kidney through control over HCO3- reabsorption and secretion of H+. Secretion of H+, in excess of its capacity to react with HCO3- in the tubular lumenal fluid, requires the presence of other buffers (see below). The generation of HCO3- and H+ occurs by dissociation of carbonic acid (H2CO3), formed in the tubule cells from H2O and CO2, through the action of carbonic anhydrase. Secretion of H+ into the lumen of the tubule is accompanied by an exchange for Na+. This reabsorption of Na+ occurs by an antiport mechanism during the exchange for H+. Reduction in the intracellular concentration of Na+ occurs by an active transport process involving a Na+/K+-ATPase pump which pumps the excess Na+ into the interstitial fluid. The intracellular HCO3- then diffuses from the tubule cell into the interstitial fluid.
The capacity of the kidney to secrete H+ is regulated by the maximal H+ gradient that can form between the tubule and lumen and still allow transport mechanisms to operate. This gradient is determined by the pH of the urine which in humans is near 4.5. The capacity to secrete H+ would be rapidly reached if it were not for the presence of buffers within the interstitial fluid. The H+ secreted into the tubular lumen can undergo three different fates depending upon the concentration of the three primary buffers of the interstitial fluid. These buffers are HCO3-, HPO42- and NH3. Reaction of H+ with HCO3- forms H2O and CO2 which diffuse back into the tubule cell. The net result of this process is the regeneration of HCO3- within the tubule cell. This process is termed reabsorption of sodium bicarbonate. The reabsorption of sodium bicarbonate takes place primarily within the proximal convoluted tubules.

Excretion of Acid
As the concentration of HCO3- in the tubular lumen drops, the pH of the fluid drops due to an increasing concentration of H+. The pH of the tubular fluid gradually approaches the pKa for the dibasic/monobasic phosphate buffering system (pKa = 6.8). The excess H+ reacts with dibasic phosphate (HPO42-) forming monobasic phosphate (H2PO4-). The H2PO4- so formed is not reabsorbed and its excretion results in the net excretion of H+. The greatest extent of H2PO4- formation occurs within the distal convoluted tubules and the collecting ducts.
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Acidosis and Alkalosis
The kidneys play an important role in the control of acidosis by responding with an increase in the excretion of H+. When H+ is excreted as a titratable acid such as H2PO4- or when the anions of strong acids such as acetoacetate are excreted there is a requirement for simultaneous excretion of cations to maintain electrical neutrality. The principal cation excreted is Na+. As the level of excretable Na+ is depleted excretion of K+ increases. In conditions of acidosis the kidney will increase the production of NH3 from tubular amino acids or amino acids absorbed from the plasma. As indicated the NH3 can diffuse across the tubule cell membrane where it will react with H+ to form the excretable ammonium ion without a concomitant requirement for cation excretion. This demonstrates that an inability of the kidney to generate NH3 would rapidly lead to fatal acidosis.
When the kidneys fail to modulate HCO3- excretion, metabolic alkalosis will develop. Alkalosis is normally countered quite effectively by the kidney allowing HCO3- to freely escape. Alkalosis generally only becomes problematic if the kidneys are restricted in their ability to secrete HCO3-. This situation can occur in patients taking diuretics since several of this class of drug cause a reduction in the ability of the kidney to reabsorb an anion (e.g. Cl-) concomitant with the reabsorption of Na+.