Dynamic equilibrium:
Many chemical reactions don’t go to completion
Once the products are formed, the reverse reaction may take place to reform the reactants
In a closed system with constant temp, the concentrations of all the reactants and products will eventually become constant
When the system reaches this state, it is in dynamic equilibrium/chemical equilibrium
Rate of forward reaction = rate of reverse reaction (or rate of vaporization = rate of condensation)
It follows that the macroscopic properties (like colour, density, pH) are constant, despite the continual interconversion of reactants and products
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Equilibrium constant is merely products/reactants with concentrations at equilibrium of each raised to their stoichiometric coefficients
If all reactants and products are in the same phase, then it is a homogeneous reaction
Solids and liquids are omitted because of fixed density
Though apparently you include liquids if the reaction is completely composed of liquids….
Units according to fact that each concen. is mol/L
Kc has reciprocal if reaction is flipped
Also, equilibrium constant is related to the position of equilibrium
Kc >> 1 – reaction goes nearly to completion
Kc << 1 – reaction hardly proceeds
Kc between 10-2 and 102 then reactants/products will be present in noticeable amounts
Law of Multiple Equilibrium:
Equations are added – multiply Kc
Equations are subtracted – divide Kc
Equation is doubled – square Kc
Equation is halved – square root of Kc
Equation is reversed – reciprocal of Kc value
Le Chatelier’s Principle (minimizes effect of change on equilibrium):
Increase concentration – equilibrium shifts to opposite side
Decrease concentration – equilibrium shifts to that side
Increase pressure – equilibrium shifts to side with least moles of gas
Decrease pressure – equilibrium shifts to side with most moles of gas
Increase temperature – equilibrium shifts to endothermic direction
Decrease temperature – equilibrium shifts to exothermic direction
Catalyst – no effect (increases both forward AND reverse reactions)
Kc ONLY CHANGES WHEN TEMPERATURE IS INCREASED/DECREASED IN THE REACTION (rate constant changes)
Endo: Increase T, Kc increases
Exo: Increase T, Kc decreases
Le Chatelier’s Principle is a memory aid; it is not an explanation for changes; rates of forward/reverse reactions should be considered (see p188, textbook)
Industrial Processes:
Goal: minimize waste and input of energy, therefore kinetics and equilibrium needed
Haber Process:
Combination of nitrogen (liquid air) and hydrogen (natural gas) to produce ammonia (fertilizers, nitric acid)
Iron catalyst (small pieces so surface area maximizes efficiency)
N2 + 3H2 -> 2NH3 (exothermic)
High pressure and low temperature favour products, but low temperature = low rate so compromise needed (ie. Optimum temperature)
Usually high temps like 700K and high pressures like 200-1000atm
Not enough time for reaction to proceed to equilibrium, so a lot of waste – only 20% of reactants converted to ammonia
Therefore reactants are cooled and recycled (ammonia can also H-bond)
Contact Process:
Oxidation of sulfur to produce sulfuric acid (fertilizers, paints, detergents, feedstock for other chemicals)
2SO2 + O2 -> 2SO3 (exothermic)
Sulfur dioxide from burning sulfur or sulfide ores, oxygen from liquid air
High pressure needed to favour products (less moles of gas on right than left)
Low temperature needed since exothermic; optiimum temperature needed
However, only 2atm of pressure is needed to achieve the desired flow rate in the reactor; high yield at 2 atm and 450 celsius, so no need for higher pressures
Waste gases are often passed through converters to take out as much sulfur dioxide as possible so exhaust can be directly released into the air
V2O3 catalyst (small pieces)
Sulfur trixoxide then reacted with water to produce sulfuric acid
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HIGHER LEVEL
rate of vaporization = rate of condensation
Molecules of the liquid will escape from the surface and become a vapour. Then some of the vapour molecules strike the surface of the liquid and condense back.
Pressure exerted by particles in the vapour phase is vapour pressure of the liquid
Altering the area of the liquid does not have an overall effect on the equilibrium because it affects both rates equally; however enlarged area may cause equilibrium to take longer
Vapour pressure depends on IMF strength and temperature
Vapourization is endothermic since it requires the overcoming of IMFs and do work against the atmosphere; amount of energy needed for phase change is enthalpy of vapourization
If IMFs really strong, enthalpy of vapourization = high, vapour pressure = lower, boiling point = higher
At higher temperature, more molecules will have required kinetic energy to escape into vapour; therefore rate of vapourization rises; also increases vapour pressure
A liquid boils when vapour pressure = pressure on the surface of liquid because it allows bubbles of vapour to form
It follows that reduced pressure = reduced boiling point, pressure cooker (high pressure) = higher boiling point (graph on p45, study guide)
Enthalpy of vapourization – amount of energy required to convert one mole of the substance from the liquid to the gaseous state
Water has good H-bonding, so it has low vapour pressure but high enthalpy of vapourization
Equilibrium Calculations:
Use ICE table
Homogeneous equilibrium – reactants and products are in same phase
Heterogeneous equilibrium – two or more phases
Make sure concentrations are at equilibrium when substituted into Kc formula. If the concentrations aren’t at equilibrium, then you have Qc (reaction quotient).
Qc > Kc – value of Qc must fall; products must be converted to reactants; system must shift to left (reverse direction) until Qc = Kc
Qc < Kc – opposite of above