No Observable Changes As Time Goes By Biology Essay


Equilibrium is a state in which there are no observable changes as times goes by. When a chemical reaction has reached the equilibrium state, the concentration of reactant and products remain constant over time, and there are no visible changes in the system. However, there is much activity at the molecular level because reactant molecule to form product molecule while the product molecule react to yield reactant molecule. This dynamic situation is the subject of term equilibrium .As few chemical reactions proceed in only one direction. Most are reversible, at least to some extent. At soon as some product molecule are formed, the reverse process begins to take place and the react molecule reactant molecule are formed from product molecule. Chemical equilibrium is achieved when the rates of forward and reverse reactions are equal and the concentrations of the reactants and the product remain constant.

CHEMICAL EQUILIBRIUM is a dynamic process. As such, it can be linked to the movement of skiers at a busy sky resort, where the number of skiers carried up the mountain on the chair lift is equal to the number coming down the slopes. Although there is constant transfer of skiers, the numbers of people at the top and the number at the bottom of the slope do not change.

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Note that the chemical equilibrium involves different substances as reactant and product. Equilibrium between two phase of the same substances is called physical equilibrium because the changes that occur are physical process.


The concept of chemical equilibrium develops 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 reactants are used up & for to left it hardly any product formed from reactant.

In 1864 Guldberg and wage showed experimentally that in chemical reaction an equilibrium 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 equilibrium. In 1877 van't hoff suggested that in the equilibrium expressions the concentration of 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.



The equilibrium involving only physical changes called physical equilibrium. Examples are evaporation of water, melting ice, dissolution of ice dissolution of sugar in water etc.


It can be established only in case of closed system i.e. the system should neither gain matter from surrounding nor lose matter to surrounding.

Process is dynamic in nature, not stop but changes take place in forward or backward direction with same speed.

The measurable property of system become constant at equilibrium.

They may in form of :-

Solid- liquid Equilibrium.

Liquid- Gas Equilibrium.

Solid- Vapour Equilibrium.


Equilibrium setup in chemical reaction called chemical reaction. Example reaction between nitrogen and hydrogen, Decomposition of calcium carbonate, reaction between sulphur and oxygen.


It is dynamic in nature.

The observable properties of the system become constant equilibrium and remain unchanged thereafter.

The equilibrium can be approached from either direction.

The equilibrium attain only in the system is closed one.

The free energy at constant pressure and temperature is zero.

They are of two types-

Irreversible reaction.

Reversible reaction.


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

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


It is applicable only when concentration of reactant and product have attained their equilibrium state, the value of equivalent constant is independent of the original concentration of reactant. The value of equilibrium constant is the same irrespective of direction from which equilibrium has been attained, The presences of catalyst cannot alter the position of equilibrium so the value of equilibrium constant, and it also not alter in presence of inert material.


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Predicting the extent of a reaction:- The magnitude of a equilibrium constant K indicates the extent of reaction can go. The larger the value of K, greater will be equilibrium concentration of components on right hand side of reaction to those on left hand side i.e. the proceed to greater extent, The value of K is very small which means that the forward reaction has proceed to small extent only.

Predicting the direction of reaction:- The equilibrium constant helps in predicting the direction in which reaction can proceed at any stage. If Q is reactive Quotient , K is equilibrium constant.

Q is more than K= Reaction in forward direction.

Q is less than K= Reaction in back ward direction.

Q is equal to K= Reaction is at 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 . We will examine the effect of a catalyst


When concentration of any constituent increases, the system will try to reduce the increases in concentration. Thus, increases in concentration of reactant favours forward reaction and decreases in concentration of product favours backward reaction.


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.

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.


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.


We know that a catalyst enhance the rate of a reaction by lowering the energy of the forward reaction and reverse reactions to the same extant .We can therefore conclude that the presence of a catalyst does not alter the equilibrium constant, nor does it shift the position of an equilibrium system .Adding a catalyst to a reaction mixture that is not at equilibrium will simply cause the mixture to reach Equilibrium sooner.


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):

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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.


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