The following points highlight the three important types of chemical aspects. The types are: 1. Buffer 2. Oxidation-Reduction 3. Measurement of Electrode Potentials and the Hydrogen Scale.
Chemical Aspect: Type # 1.
Buffer:
A solution that contains a weak acid with its conjugate base, or a weak base with its conjugate acid will minimize the pH change that results from the addition of a standard quantity of strong acid or strong base.
These solutions are called buffer solutions and the buffer capacity of a solution is its effectiveness in minimizing the pH change. For example, acetic acid and the acetate ion form an acid-base conjugate pair. Acetic acid is a weak acid and, therefore, dissociates slightly in solution. If we consider the addition of OH– to a solution of acetic acid (HA).
HA + OH– = A– + H2O
The hydroxyl ions are neutralized by the free hydrogen ions produced by acetic acid and this neutralization causes more of the acetic acid to ionize so as to restore the original hydrogen ion concentration (pH). As more OH– is added, more acetic acid ionizes until all of the acetic acid is ionized.
At this point, any further addition of OH– will cause an abrupt rise in pH. If hydrogen ions (H+) are added, acetate ions rapidly unite with them to form un-dissociated acetic acid and, therefore, there will be no change in hydrogen ion concentration until all of the acetate ions have been converted. At this point, further addition of H+ ions causes an abrupt drop in pH.
In biological systems, amino acids and proteins act as the primary buffers and play a vital role. Particularly, enzyme activity and reaction rates are kept at their maximum by the cell’s buffer action. A small change in pH can dramatically affect the structure and functions of biological molecules. So maintaining a relatively constant pH is of paramount importance for living system.
Chemical Aspect: Type # 2.
Oxidation-Reduction:
Oxidation-reduction reactions are those in which there is transfer of electrons from an electron donor to an electron acceptor. The electron donor is the reducing agent or reductant and the electron acceptor is the oxidizing agent or oxidant.
Oxidation and reduction are simultaneous, i.e., there will be no oxidation without concomitant reduction. The reduced and oxidized forms of the compound constitute a conjugate pair which is called the oxidation-reduction couple or redox couple. Thus,
where n is the number of electrons (e) liberated from one molecule of the reductant. The reduction of an oxidant produces its conjugate reductant,
Any oxidation-reduction reaction must involve two redox couples that differ in their affinity for electrons. The couple that shows the greater electron affinity assumes the oxidizing role, and its oxidant will be reduced. The couple with lesser affinity for electrons is the oxidizing couple.
The overall reaction can be considered to be the resultant of two half-reactions as follows:
If, by the isothermal transfer of electrons from reda to oxb, the Gibbs free energy of the system is decreased, the overall reaction is thermodynamically spontaneous from left to right. When the Gibbs free energy content of the system reaches its maximum value, net reaction ceases and the system is at equilibrium.
Since the oxidation and reduction involves electron transfer from reda to oxb, the reaction could perform electrical work. This would be equal to its – ΔG (change in free energy) if the reaction could be carried out at constant temperature and pressure in a thermodynamically reversible manner. Thus the value of – ΔG for an oxidation-reduction reaction can be measured experimentally.
Different reducing agents differ in their tendency to lose electrons. The intensity of the tendency of a reducing agent to lose electrons or an oxidizing agent to accept electrons is measured as the standard oxidation-reduction potential.
It is defined as the electromotive force (e.m.f.) in volts given by a half-cell having both oxidant and reductant at 1.0 M concentration, at 25 °C and pH 7.0, at equilibrium. The half-cell possesses an electrode which can reversibly accept electrons from the reductant. The standard oxidation-reduction potential of the hydrogen electrode reaction (2H+ + 2e– ↔ H2) is used as the reference potential.
The standard hydrogen electrode consists of H2 gas in equilibrium with H+ ions (1.0 M) in solution, at 1.0 atm pressure and 25°C. The electrons are transmitted to the redox couple via an inert platinum electrode. The reference potential is set at 0.0V. The standard oxidation-reduction potential of the hydrogen electrode becomes – 0.42V.
In redox pairs with a more negative potential than the 2H+ – H2 couple, the reductant has a greater tendency to lose electrons than molecular hydrogen, and in pairs with a more positive potential the reductant has a lesser tendency to lose electrons than hydrogen (Fig. 1.2).
The potential difference of an electrical cell expresses the difference that exists between its half-cells in their affinity for electrons.
This electron affinity demonstrated by a half-cell (redox couple) is termed as electrode potential. The greater is a half-cell’s electron affinity the larger (more positive) is the value of its electrode potential. The electrode potential is, therefore, a measure of the tendency of a redox couple to undergo reduction.
Oxidized form + ne– → reduced form.
This is often called the reduction electrode potential of a half-cell. If two half-cells of different electrode potentials are joined to form an electrochemical cell, the spontaneous cell reaction (i.e., the discharge) will consist of a reduction reaction in the oxidizing half-cell (more positive electrode potential), coupled with an oxidation reaction in the reducing half-cell (less positive, more negative electrode potential).
The magnitude of the electrode potential of any half-cell is determined by its chemical composition (concentrations of the components in solution) and the prevailing temperature.
The composition of an electrical cell is expressed in the form of a cell diagram in which the half-cell constituents are displayed on either side of a pair of vertical, parallel lines representing the bridge between them. Conventionally the components are ordered from left to right indicating the cell reaction under consideration. Thus in the cell,
The potential of the cell (ΔE) measures the difference in electrode potentials of its component half-cells,
i.e., ΔE = (Electrode potential of the half-cell on the right of the bridge) – (Electrode potential of the half-cell on the left of the bridge).
Thus, in the above example of the electric cell P, Q || R, S
ΔE = E (R/S couple) – E (Q/P couple)
If the electrode potential of the R/S couple is more positive than the electrode potential of the Q/P couple, then,
(i) ΔE will have a positive value
(ii) The cell reaction P + R → Q + S will have a negative value of AC; (since ΔG = – nFΔE), and will, therefore, be spontaneous
(iii) Spontaneous discharge of the cell will be accomplished by left to right flow of electrons in the external circuit.
However, if the electrode potentials of the R/S couple are less positive (i.e., more negative) than that of Q/P couple, then,
(i) ΔE will have a negative value
(ii) The cell reaction will have a positive value of ΔG. The spontaneous reaction with a negative ΔG will, therefore, be the reverse reaction, viz., S + R → Q + P;
(iii) There will be right to left flow of electrons in the external circuit.
The biologists are mainly concerned with assessing the feasibility of reactions between redox couples of known electrode potentials. Thus, the redox couple of smaller electrode potential is purposely placed on the left hand of the bridge to ensure a flow of electrons from left to right in the external circuit driven by a positive cell potential difference.
The cell created by placing the couple of lesser electrode potential on the left of the bridge shows the positive potential difference by the following equation:
ΔE = E (right-hand couple) – E (left-hand couple).
For example, the reduction of pyruvate by NADH:
The redox potential of the NAD+: NADH couple is – 0.32 V, whereas that of pyruvate: lactate couple is – 0.19V. By convention, redox potentials of half-reactions are written as:
Oxidant + e– → reductant. Hence,
Putting values in the above equation:
E = – 0.1 9V – (- 0.32V) = + 0.13V
The Gibbs free energy (ΔG) for the reduction of pyruvate by NADH can be calculated. The standard free energy change ΔG is related to the change in redox potential ΔE by
ΔG = – nF ΔE
where n is the number of electrons transferred per molecule from reductant to oxidant, F is the caloric equivalent of the Faraday (23.062 kcal V-1), ΔE is e.m.f., i.e., the difference between the standard electrode potentials of the contributory redox couples in volts, and ΔG is in kilocalorie per mole. For the reduction of pyruvate n = 2 and so
ΔG = – 2 x 23.062 x 0.1 3 = – 6 kcal/mol
So, the positive value of ΔE signifies an exergonic reaction under standard conditions.
The redox potentials of many biologically important redox couples are known and given in the following table:
The span of the respiratory chain is 1.14 volts which corresponds to approximately 53 kcal. For oxidative phosphorylation the driving force is the electron transfer potential of NADH or FADH2 to O2.
The half-reactions are:
Electron-transfer reactions are of great biological importance.
In respiratory chain electrons are transferred from NADH to O2 via a series of electron transporting components of increasing reduction potentials creating a cascade of free energy that is utilized in the formation of ATP from ADP and Pi. NADH functions as an energy-rich electron transporting coenzyme.
The oxidation of one NADH to NAD+ supplies sufficient free energy to generate 3(or 2.5) molecules of ATP. NAD+ serves as an electron acceptor in many exergonic oxidative reactions.
Chemical Aspect: Type # 3.
Measurement of Electrode Potentials and the Hydrogen Scale:
A half-cell shows its electron affinity only when it is suitably coupled with another half-cell which donates or accepts electrons. There is no direct way of measuring the actual magnitude of the electrode potential of an isolated half-cell. It is, however, measured by comparing the electrode potential of one half-cell with that of any other.
Thus electrode potential magnitude of various half-cells can be assessed relative to one another. The observed values can be assigned on a scale whose zero is defined as being equal to the actual electrode potential of a reference half-cell. This device is very similar to the temperature scales. With the decreed zero the scale’s unit of measurement is arbitrarily defined.
The scale on which electrode potentials are measured is called the hydrogen scale as the electrode potential of standard hydrogen half-cell is taken as zero of the scale. Its scale division is the volt.
On this hydrogen scale of electrode potentials:
(a) The standard hydrogen electrode possesses at all temperatures an electrode potential of 0.00V;
(b) A half-cell which is more reducing than the standard hydrogen half-cell is credited with a negative electrode potential (Eh is – ve) whose magnitude measures the extent by which its electron affinity is less than that of the standard hydrogen electrode;
(c) a positive electrode potential (Eh is + ve) is shown by a half-cell which is more highly oxidizing than the standard hydrogen half-cell. Its potential magnitude measures the extent by which its electron affinity exceeds that of the standard hydrogen electrode.
This means that the (reduction) electrode potential of any redox couple ([ox] / [red]) is the e.m.f. in volts of the cell that formed when that couple acting as a half-cell is joined to a standard hydrogen electrode in the following manner:
Pt, H2 (1 atm.), H (a = 1) || [ox], [red], Pt.
It is the value of its electrode potential on the hydrogen scale (Eh) that is termed the redox potential of a redox couple.