In chemistry, dynamic equilibrium is an equilibrium situation in which microscopic changes occur, but macroscopic changes do not. Note that when dynamic equilibrium exists all the observable properties are constant but at the molecular level there is a constant back and forth reaction between reactants and products which is in perfect balance.
The Characteristics of Equilibrium
a) rate of consumption of reactants= rate of production of reactants
b) [Reactants] differens from [products] in general c) [Reactants] is now constant in time and [Products] is constant in time
d) the forward and rverse rates do not change as time passes
e) A system which is not at equilibrium will tend to move toward a position of equilibriumhttp://wiki.ubc.ca/index.php?title=Chemical_Equilibrium&action=edit
Le Chatelier's Principle If a closed system at equilibrium is subjected to change, processes will occur that tend to counteract that change ie. whatever you do, the nature tries to undo
The effects of substrate concentration on Equilibrium
aA + bB ↔ cD + dD
Based on the reaction above, A and B act together to produce C and D; meanwhile, C and D act together to produce A and B
If we add more A to the reaction system, the reaction will shift to the right, to produce more C and D at the equilirium.
If we add more B to the reaction system, the reaction will shift to the right, to produce more C and D at the equilibrium.
If we add more A and B, the reaction will shift to the right to produce more C and D
Similarly, adding more C or more D will cause the reaction to shift to left to produce more A and B
...But What is the rational behind these observations?
When you add more A to the system, you're increasing the amount of reactant available to react and form products. Similarly, if you add more C or D to the system, you're increasing the amount of product available to react and form reactants. the system will adjust by moving to the left to reestablish equilibrium.
This is because there will be more collisions between A particles and B particles. Since we did no add any B to the system, the increased collisions among A and B particles , along with the increased production of C and D will tend to reduce the concentration of B at the equilibrium. After the equilibrium has shifted, there will be more A particles than there were before we added any more A. There will be fewer B particles than there were before we began, and therefore, more C and D particles.
Equilibrium constants of reactions involving gas mixtures
It is possible to work out the equilibrium constant for a chemical reaction involving a mixture of gases given the partial pressure of each gas and the overall reaction formula. For a reversible reaction involving gas reactants and gas products, such as:
aA + bB ↔ cD + dD
the equilibrium constant of the reaction would be:
KP = (PCc.PDd)/(PAa.PBb)
KP = the equilibrium constant of the reaction
a = coefficient of reactant A
b = coefficient of reactant B
c = coefficient of product C
d = coefficient of product D
PCc = the partial pressure of C raised to the power of c
PDd = the partial pressure of D raised to the power of d
PAa = the partial pressure of A raised to the power of a
PBb = the partial pressure of B raised to the power of b
For reversible reactions, changes in the total pressure, temperature or reactant concentrations will shift the equilibrium so as to favor either the right or left side of the reaction in accordance with Le Chatelier's Principle. However, the reaction kinetics may either oppose or enhance the equilibrium shift. In some cases, the reaction kinetics may be the over-riding factor to consider.
Ideal Gas Assumptions
Molecules of an ideal gas do not attract or repel each other.
Molecules of an ideal gas occupy zero volume
However, no gas ever acts completely as an ideal gas. In real gases, molecules attract each other slightly which causes them to strike the walls of their container with slightly less force than ideal gases. HOwever, most gases (especially lighter ones such as hydrogen and helium) uinder typical temperatures and pressures act enough as an ideal gas to make the concept useful.