Chemical Equilibrium And Ph Biology Essay


In a chemical process, chemical is a state in which cocentration of reactants and the concentration of products does not change overtime i.e when the forward reaction proceeds at the same rate as backward reaction and it exists in dynamic equilibrium.

The laws of chemical equilibrium define the direction in which a chemical reaction will proceed, as well as the quantities of reactants and products that will remain after the reaction comes to an end. An understanding of chemical equilibrium and how it can manipulated is essential for anyone involved in Chemistry and its applications. The fundamental equation provides the basis for understanding chemical equilibrium.

pH is Quantitative measure of stength of acidity and alkalinity of solution. When the concentration of H+ and OH- ions in aqueous solution are frequently very small numbers and there fore inconvenient to work with, Soren Sorensen in 1909 proposed a more practical measure called pH. It is defined as the negative of the hydrogen ion concentration.

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pH= -log[H+].

Acidic and basic are two extremes that describe a chemical property chemicals. Mixing acids and bases can cancel out or neutralize their extreme effects. A substance that is neither acidic nor basic is neutral.

The pH scale measures how acidic or basic a substance is. The pH ranges from 0 to 14. A pH of 7 is neutral. A pH less than 7 is acidic. A pH greater than 7 is basic.


The concept of chemical equilibrium develop after Berthoket (1803) found that some chemical are reversible. For any reaction to be equilibrium, rate of forward reaction is equal rate of backward reaction, so at equilibrium nearly all the reactant are used up & for to left it hardly any product formed from reactant.

In 1864 Guldberg and waage showed experimentally that in chemical reactions an equillibrium is reached that can be approached from either direction. They were apparently the first to realize that there is a mathematical relation between the concentration of reactants and products at equillibrium. In 1877 van`t hoff suggested that in the equillibrium expressions the concentrationof each reactant should appear to the first power, corresponding with the stochiometric numbers in the balanced chemical equation.

The concept of p[H] was first introduced by Soren Peder Lauritz Sorensen at the carlsberg laboratory in 1909 and revised to the modern pH in 1924 after it became apparent that electromotive force in cells depended on activity rather than concentration of hydrogen ions.


Physical equillibrium:- Equillibrium between two phases of the same substance is called Physical equillibrium because the changes that occur are physical processes.

e.g; The vaporization of water in a closed container at a given temperature is an example of physical equilibrium. In this instance the number of H2O molecules leaving and the number of returning to the liquid phase are equal.

H2O(l) = H2O(g)

The study of physical equilibrium Yields useful information such as the equilibrium vapour pressure. If a reaction is:-

aA + bB = cC + dD

where a, b, c and d are the stoichiometric coefficients for the reacting species A, B, C , and D. For the reaction at a particular temperature

K = (C)C (D)d

(A)a (B)b

where K is the equilibrium constant.

This equation was formulated by cato Guldberg and Petr Wage in 1864. It is the mathematical expressions of their law of mass of action, which holds that for a reversible reaction at equilibrium and a constant temperature, a certain ratio of reactant and product concentration has a constant value K. and the equilibrium constant does depend on the volume, concentration, catalyst, pressure e.t.c. It only depends upon temperature.

The equilibrium constant ,then, is defined by a quotient, the numerator of which is obtained by multiplying together the equilibrium concentration of the products, each raised to a power equal to its stoichiometric coefficient in the balanced equation denoted by Q.

If Q < K then reaction takes forwad reaction.

If Q > K then reaction takes backward reaction.

If Q = K then reaction is at equilibrium state.

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The magnitude of equilibrium constant tells whether the reaction favors the products or reactant.

If K >> 1 , the equilibrium will lie to the right and favours the products.

If K << 1 , the equilibrium will lie to the left and favours the reactants.

Homogeneous equillibrium:- The term homogeneous equilibrium applies to reactions in which all reacting species are in the same phase.

E.g; An example of homogenous gas phase equilibrium is the dissociation of N2O4.

N2O4 (g) = N2O4 (g)

aA(g) = bB(g)

then, Kc = (B)b/(A)a

the concentrations of reactants and the products in gaseous reaction can also be expressed in terms of their partial pressures. At constant temperature the pressure P of gas is directly related to the concentrations in mol/ L of the gas. i.e; P =(n/V)RT.

And the expressions can be given by Kp = PbB/PaA where Pa and P b are the partial pressure of A and B. After substiuting these relations in to the expressions, we get


If b-a is equal to zero, then reaction is at equilibrium.

If b-a > 0 , then backward reaction is favourable.

If b-a < 0 , then forward reaction is favourable.

Heterogeneous equilibrium:- A heterogeneous equilibrium results from a reversible reaction and products that are in different phases.

e.g; when calcium carbonate is heated in a closed vessel, the following equilibrium is attained.

CaCO3(S = CaO(S) + CO2

If a reaction aA(S) + bB(s) = cC(s) + dD(g),

the concentration of a solid , like it`s density , is an intensive property and does not depend on how much of the substance is present. In thermodynamics , the activity of pure solid us 1. Thus yhe concentration terms for A,B and C are unity. So,

Kc = [D]d

Similarly , the activity of a pure liquid is also 1. Thus if a reactant or a product is a liquid, we can`t omit in the equilibrium constant expressions.

Also we can express the equilibrium constant as

Kp = P(D)d .

Factor Affecting Equilibrium:-

Chemical equilibrium represents a balance between forward and reverse reactions. In most cases, this balance is quit delicate. Change in experimental condition may disturb the balance and shift the equilibrium position so that variable can be controlled experimentally are concentration, pressure, volume and temperature.

Effect of change of temperature:- A change in concentration, pressure, or volume may alter the equilibrium position , that is, the relative amounts of reactants and products but it dose not change the value of equilibrium constant .Only a change in temperature can alter the equilibrium constant. At equilibrium at a certain temperature, the heat effect is zero because there is no net reaction .If we treat heat as though it were a chemical reagent, than a rise in temperature "adds" heat to the system and a drop in a temperature "removes" heat from the system .As with a change in any other parameter ,(concentration, pressure, or volume), the system shift to reduce the effect of the change. Therefore, a temperature increase favours the endothermic direction (from left to right of the equilibrium equation ) and a temperature decrease favours the exothermic direction.

In summary, a temperature increase favours an exothermic direction and a temperature decrease favours the exothermic reaction. The effect of changing temperature on an equilibrium constant is given by the van 't Hoff equation

d ln K/ d T = H/RT2

Thus, for exothermic reactions, (ΔH is negative) K decreases with an increase in temperature, but, for endothermic reactions, (ΔH is positive) K increases with an increase temperature. An alternative formulation is

d ln K/ d(1/T) = - H/R

At first sight this appears to offer a means of obtaining the standard molar enthalpy of the reaction by studying the variation of K with temperature.

Effect of change of pressure and volume:- Change in pressure ordinarily do not effect the concentration of the reacting species in condensed phases(say, in an aqueous solution) because liquids and solids are virtually in compressible. On the other hand, concentration of gases are greatly affected by change in pressure. The greater the pressure, the smaller the volume, and vice versa. Note, too, that the term (n/V) is the concentration of the gas in mol/L, and it varies directly with pressure.

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In general an increase in pressure (decrease in volume) favours the net reaction that decreases the total numbers of moles of gases (the reverse reaction, in this case), and a decrease in pressure (increase in volume) favours the net reaction that increases the total numbers of moles of gases(here the forward reaction).For reaction in which there is no change in the numbers of moles of gases, a pressure (or volume)change has no effect on the position of equilibrium.

Effect of catalyst at equilibrium :-

Catalyst does not affect the the equilibrium constant. And it can not shift the equilibrium positon. Catalyst only enhance the rate of reaction to reach the equilibrium position.

Effect of concentration of reactant:- Effect of concentration change the position of equilibrium. If we increase the concentration of the products shifts the equilibrium to left, and decreasing the concentration of the product shifts the equilibrium to right. These results are just predicted by Le Chatelier principle.

Le chatelier principle:- If an external stress is applied to a system at equilibrium, the system adjusts in such a way that the stress is partially off set as the system reaches a new equilibrium position.


Because the concentrations of H+ and OH- in aqueous solution are frequently very small numbers and therefore inconvenient to work with, Soren Sorensen in 1909 proposed a more practical measure called pH.

The pH of a solution is defined as the negative logarithm of hydrogen ion concentration in (mol/L):

pH = -log [H3O+] or pH = -log [H+]

The negative logarithm gives us a positive number for Ph, which otherwise would be negative due to small value of [H+]. Furthermore the term [H+] permits only to the numerical part of the expression for hydrogen ion concentration, for we can't take the logarithm of units. Thus, like the equilibrium constant, the pH of a solution is a dimensionless quantity.

Acidic solutions: pH is less than 7.00.

Basic solutions: pH is more than 7.00.

Neutral solutions: pH is equal to 7.00.

Notice that pH increases as [H+] decreases.

For real solutions, activity usually differs from concentrations, sometimes appreciably. Knowing the solute concentration, there are reliable ways based on thermodynamics for estimating its activity, but the details are beyond the scope of text.

Keep in mind, therefore, that the measured pH of a solution is usually mot same as that , because the concentration of the H+ ion in molarity is not numerically equal to its activity value. Although we will continue to use concentration in our discussion, it is important to know that this approach will give us only an approximation

Of the chemical process that actually take place in solution phase. In the laboratory, pH of a solution is measured with a pH meter.

List of some pH of a number of common fluid, the pH of a body fluid varies greatly, depending on location and function. The low pH (high acidity) of gastric juices facilitates digestion whereas a higher pH of blood is necessary for transport of oxygen.


A pH indicator is substances that change colour around a particular pH. It is weak acid or weak base and the colour changes occur around 1pH unit either side of its acid dissociatian constant.

For.ex. Naturaly occuring indicator litmus is red in acid solution and blue in alkaline solution.


P.W. Atkins, Physical Chemistry, third edition, Oxford University Press, 1985.

F.van Zeggeren and S.N Steery,the computation of chemical equilibrium,1920.

W.R.Smith and R.W.Mission,chemical equilibrium Analysis.

RAYMOND CHANG, Williams college.