Equilibrium

Chemical Equilibrium

General Info on Equilibrium

• Concerned with how far a reaction goes. Why does it have a low or high % yield? Why do some of the reactants never become products?

• Not concerned w/ how fast (not looking at rates)• At equilibrium, a reaction:

– Macroscopically looks finished– Microscopically the forward reaction rate = the

reverse reaction rate– Has both reactants and products present but not

necessarily in equal concentrations or amounts.– Is in a closed system

What determines where equilibrium is reached? (Why do some reactions have low yields and others high

yields?)

• The equilibrium state is a compromise between a loss in enthalpy (↓ in ∆H) and a gain in entropy (↑ in ∆S).– ∆S = change in entropy. Entropy is a measurement of chaos.

• H20(l) H2O(g) the forward reaction has a gain in ∆S and a gain in ∆H. = equilibrium

• Reactions which do not have this compromise such as combustion reactions, have no measurable equilibrium. The reaction only goes forward.

• Ex. CH4(g) + 2O2(g) CO2(g) + 2H2O(g) this reaction loses enthalpy and no change in entropy when going in the forward direction. = no equilibrium

Quantitative Equilibrium

• The equilibrium math expression:– K = [Products]/[Reactants]

K= the equilibrium constant– Rules for writing the K expression:

i) do not include solids or liquid water

ii) Coefficients are written as exponents

iii) Substitute into the expression only equilibrium concentrations Kc or equilibrium pressures Kp.

K constant

• The K constant for a reaction stays the same unless temperature is changed. (Changing temp. will change the K constant because it changes the thermodynamics of the reaction system. A new compromise between ∆S and ∆H is reached)

• Larger K constant = greater concentration of the products at equilibrium. Therefore, there is a higher % yield, or the forward reaction is more favored.

K constant cont.

• Categories of K constants and expressions:

• Ka = Acids

• Kb =Bases

• Kp=gases using partial pressures

• Kc = using molar concentrations

Le Châteliers principle

• Any change forced upon a reaction system at equilibrium will cause the reaction to respond in such a way as to counter act that change and restore equilibrium.

• Concentration’s effects:– A change in concentration only changes the

position of equilibrium (shifts equilibrium) but does not change the K constant.

Le Châteliers principle cont.

• Concentrations effects (cont.)• An increase of the concentration of a reactant or a

decrease in the concentration of the product causes the reaction to shift in the forward direction to produce a product.

• A decrease in the concentration of a reactant or an increase in the concentration of the product causes the reaction to shift in the reverse direction.

Le châteliers principle cont.

• Pressure (similar to concentration) does not effect the K constant.

• It only affects gases• The side (reactants or products) which has the most moles of

gas will experience the effect of the pressure change to the greatest extent.

• An increase in pressure = an increase in concentration.• A decrease in pressure = a decrease in concentration.• An increase in pressure causes the reaction to shift in the

forward direction if there are more moles of gases as reactants

• A decrease in pressure causes the reaction to shift in the reverse direction with more moles of gases as reactants

Le Châteliers principle cont.

• Temperature: alters both the position and the constant– Endothermic: an increase in temperature

causes a shift in the forward direction and an increase in the K constant, because there will be more products than reactants present

– Exothermic: an increase in temperature causes a shift in the reverse direction and a decrease in the K constant. Less products present at equilibrium.

Le châteliers principle cont.

• Catalyst: has no effect on the position or the constant

Haber Process

• This reaction is used for the mass production of ammonia

• The reaction is: N2 + 3H2 2NH3 all (g) ΔH = -92kJ

• K= [NH3]2/ [N2] [H2]3

The Contact Process

The industrial process of making ammonia

Yields of NH3 at various pressures and temperatures

Summary of the Haber process

• Is a great example for the application of Le Chatelier’s principle

• Increases the pressure (300atm) to create a forward shift.

• Constantly supplies the reactants H2 & N2

• Removes the NH3 through cooling (NH3 has hydrogen bonding)

• Also applies principles of kinetics.– Uses a Fe catalyst which lowers the Ea

– Uses high temperatures to speed up the reaction rate. (The reaction will not reverse since there is no ammonia available)

The Contact Process

• This process consists of a series of reactions which produce sulfuric acid.S(s) + O2(g) SO2(g)

2SO2(g) + O2(g) 2SO3(g) ΔH -196 kJ/mol

SO3(g) + H2O (l) H2SO4(l)