Definition for Chemical Equilibrium

a.) A reversible reaction is a chemical reaction that results in an equilibrium mixture of reactants and products. For a reaction involving two reactants and two products this can be expressed symbolically as
aA + bB cC + dD
A and B can react to form C and D or, in the reverse reaction, C and D can react to form A and B. This is distinct from reversible process in thermodynamics.Dynamic Equilibrium refers to a reversible reaction in which the rates of forward and reverse reactions become equal and there is no change in the concentration of the reactants and products.


b.) Le Chatelier's Principle can be used to predict the effect of a change in conditions on a chemical equilibrium.


c.) If a chemical system at equilibrium experiences a change in concentration, temperature, volume, or total pressure, then the equilibrium shifts to partially counter-act the imposed change.

Concentration
Changing the concentration of an ingredient will shift the equilibrium to the side that would reduce that change in concentration. The chemical system will attempt to partially oppose the change affected to the original state of equilibrium. In turn, the rate of reaction, extent and yield of products will be altered corresponding to the impact on the system.
This can be illustrated by the equilibrium of carbon monoxide and hydrogen gas, reacting to form methanol.
CO + 2 H2 ⇌ CH3OH
Suppose we were to increase the concentration of CO in the system. Using Le Châtelier's principle we can predict that the amount of methanol will increase, decreasing the total change in CO. If we are to add a species to the overall reaction, the reaction will favor the side opposing the addition of the species. Likewise, the subtraction of a species would cause the reaction to fill the “gap” and favor the side where the species was reduced. This observation is supported by the "collision theory". As the concentration of CO is increased, the frequency of collisions of that reactant would increase also, allowing for an increase in forward reaction, and generation of the product. Even if a desired product is not thermodynamically favored, the end product can be obtained if it is continuously removed from the solution.

Temperature
Let us take for example the reaction of nitrogen gas with hydrogen gas. This is a reversible reaction, in which the two gases react to form ammonia:
N2 + 3 H2 ⇌ 2 NH3 ΔH = −92kJ mol-1
This is an exothermic reaction when producing ammonia. If we were to lower the temperature, the equilibrium would shift in such a way as to produce heat. Since this reaction is exothermic to the right, it would favour the production of more ammonia. In practice, in the Haber process the temperature is instead increased to speed the reaction rate at the expense of producing less ammonia.

Changes in pressure owing to changes in volume
The equilibrium concentrations of the products and reactants do not directly depend on the pressure subjected to the system. However, a change in pressure due to a change in volume of the system will shift the equilibrium.
Once again, let us refer to the reaction of nitrogen gas with hydrogen gas to form ammonia:
N2 + 3 H2 ⇌ 2 NH3 ΔH = −92kJ mol-1
Note the number of moles of gas on the left hand side, and the number of moles of gas on the right hand side. When the volume of the system is changed, the partial pressures of the gases change. Because there are more moles of gas on the reactant side, this change is more significant in the denominator of the equilibrium constant expression, causing a shift in equilibrium.
Thus, an increase in pressure due to decreasing volume causes the reaction to shift to the side with the fewer moles of gas.[2] A decrease in pressure due to increasing volume causes the reaction to shift to the side with more moles of gas. There is no effect on a reaction where the number of moles of gas is the same on each side of the chemical system (or equation)
g.) The Haber synthesis was developed into an industrial process.Conditions:medium temperature (~500oC) very high pressure (~250 atmospheres, ~351kPa) a catalyst (a porous iron catalyst prepared by reducing magnetite, Fe3O4).Osmium is a much better catalyst for the reaction but is very expensive.
This process produces an ammonia, NH3(g), yield of approximately 10-20%.
The reaction between nitrogen gas and hydrogen gas to produce ammonia gas is exothermic, releasing 92.4kJ/mol of energy at 298K (25oC).
N2(g)nitrogen+3H2(g)hydrogen heat, pressure, catalyst 2NH3(g) ammonia
H = -92.4 kJ mol-1
By Le Chetalier's Principle:
Increasing the pressure causes the equilibrium position to move to the right resulting in a higher yeild of ammonia since there are more gas molecules on the left hand side of the equation (4 in total) than there are on the right hand side of the equation. Increasing the pressure means the system adjusts to reduce the effect of the change, that is, to reduce the pressure by having fewer gas molecules.
Decreasing the temperature causes the equilibrium position to move to the right resulting in a higher yield of ammonia since the reaction is exothermic (releases heat). Reducing the temperature means the system will adjust to minimise the effect of the change, that is, it will produce more heat since energy is a product of the reaction, and will therefore produce more ammonia gas as wellHowever, the rate of the reaction at lower temperatures is extremely slow, so a higher temperature must be used to speed up the reaction which results in a lower yield of ammonia.
The equilibrium expression for this reaction is:
Keq =[NH3]2[N2][H2]3


h.) Chain Disproportionation occurs when two radicals meet, instead of coupling, they exchange a proton. That gives two terminated chains, one saturated and the other with a terminal double bond.
2R-(CH2CH2)n-CH2CH2• → R-(CH2CH2)n-CH=CH2 + R-(CH2CH2)n-CH2CH3
http://www.scribd.com/doc/6618373/InDepth-Overview-and-Comparison-of-2006-2007-AL-Chem-Syllabus