Chemical Equilibrium
Chemical Equilibrium
occurs when a reaction and its reverse proceed at the same rate
The Concept of Equilibrium
as a system approaches equilibrium, both the forward and reverse reactions are occurring
at equilibrium, the forward and reverse reactions are proceeding at the same rate
A system at equilibrium
once equilibrium is achieved, the amount of each reactant and product remains constant
The Equilibrium Constant
N2O4—>←-2NO2
Forward reaction: N2O4—>kf[N2O4]
Reverse reaction: kr[NO2]²
At equilibrium: Ratef=Rater
kf=[N2O4]=kr[NO2]²
rewriting this becomes…
kf/kr=[NO2]²/[N2O4]=Constant (Kc)
Consider the generalized reaction
aA+bB—>←-cC+dD
the equilibrium expression for this reaction would be
Kc=[C}^c[D]^d/[A]^a[B]^b
[products]/[reactants]
Sample Exercise 15.1
Evaluating Kc
equilibrium constant (Kc) does NOT depend on initial concentrations of reactants
change in temperature, alters the Kc, as rate constant is sensitive to change in temperature
Equilibrium Constant for Gas
Consider generalized reaction involving gases
aA+bB—>←-cC+dD
Since pressure is proportional to concentration for gases in a closed system, the equilibrium expression can also be written
kp=(Pc^c)(PD^d)/(PA^a)(PB^b)
P-partial pressure of respective gases
Kc and Kp conversions
From the ideal-gas law we know that
PV=nRT
Rearranging it we can get
P=[M]RT
plugging this into the expression for Kp, for each substance, the relationship between Kc and Kp becomes
Kp=Kc(RT)^delta n
delta n = (moles of gaseous product)-(moles of gaseous reactant)
Sample Exercise 15.2
Equilibrium Constants and Units
the ratio of [NO2]² to [N2O4] remains constant at this temperature no matter what the initial concentrations of NO2 and N2O4 are.
Equilibrium can be reached from either direction
it doesn’t matter whether we start with N2 and H2 or whether we start with NH3, we will have the same proportions of all three substances at equilibrium
Evaluating Kc
equilibrium constant (Kc) does not depend on initial concentrations of reactants
What Does the Value of K Mean?
If K>1, the reaction is product favored; product predominates at equilibrium
If K<1, the reaction is reactant favored; reactant predominates at equilibrium
Manipulating Equilibrium Constants
The equilibrium constant of a reaction in the reverse reaction is the reciprocal of the equilibrium constant of the forward reaction
N2O4—>←-2NO2
Kc=[NO2]²/[N2O4] = 0.212 at 100 degrees celsius
2NO2—>←-N2O4
Kc=[N2O4]/[NO2]² = 4.72 at 100 degrees Celsius
the equilibrium constant for a net reaction made up of two or more steps is the product of the equilibrium constants for the individual steps
The Concentrations of Solids and Liquids are Essentially Constant
both the concentrations of solids and liquids can be obtained by multiplying the density of the substance by its molar mass-and both of these are constants at constant temperature
therefore, the concentrations of solids and liquids DO NOT APPEAR in the equilibrium expression
PbCl2(s)—>←- Pb²+(aq) + 2Cl^- (aq)
Kc = [Pb²+][Cl^-]²
ex:
CaCO3(s) —>←- CO2 (g) +Ca) (s)
as long as some CaCO3 or CaO (solids) remain in the system, the amount of CO2 (gas) above the solid will remain the same
An Equilibrium Problem
a closed system initially containing 1.000×10^-3 M H2 and 2.000 × 10^-3 M I2 at 448 degrees Celsius is allowed to reach equilibrium. Analysis of the equilibrium mixture shows that the concentration of HI is 1.87 × 10^-3 M. Calculate Kc at 448 degrees Celsius for the reaction taking place
H2(g) + I2(s)—>←- 2HI (g)
The Reaction Quotient (Q)
same as Kc
Q gives the same ratio the equilibrium expression gives, but for a system that is NOT at equilibrium
To calculate Q, one substitutes the initial concentrations on reactants and products into the equilibrium expression
If Q = K, the system is at equilibrium
Q<K (the reaction will form products (shift right) because there are too many reactants)
Q = K (the reaction is in equilibrium)
Q>K (the reaction will form reactants (shift left) because there is too much product)
Le Chatelier’s Principle
“If a system at equilibrium is disturbed by a change in temperature, pressure, or the concentration of one of the components, the system will shift its equilibrium position so as to counteract the effect of the disturbance”
how to counteract
Concentration
adding or removing a reactant or product
if reactant is added, the reaction will shift to the right in order to produce more product.
if more product produced, then the equation will shift left, in order to equal out more reactants.
Pressure
changing pressure by changing the volume
at a constant temperature, reducing the volume of a gaseous equilibrium mixture causes the system to shift in the direction that reduces the number of moles of gas (shift to direction with less moles)
initial volume + added pressure (creates a smaller volume) = system will shift to direction of fewer moles of gas
Temperature
if the temperature of a system at equilibrium is increased, the system reacts as if we added a reactant to an endothermic reaction or a product to an exothermic reaction. The equilibrium shifts in the direction that consumes the “excess reactant”, namely heat.
Endothermic
increasing T = reaction shifts right
decreasing T = reaction shifts left
Exothermic
increasing T = reaction shifts left
decreasing T = reaction shifts right
Haber Process
the transformation of nitrogen and hydrogen into ammonia (NH3) is tremendous significance in agriculture, where ammonia-based fertilizers are important.
if H2 is added to the system, N2 will be consumed and the two reagents will form more NH3
this apparatus helps push equilibrium to the right by removing the ammonia from the system as a liquid
decrease volume = increase pressure
increase volume = decrease pressure
Catalyst
increase the rate of both the forward and reverse reactions
when one uses a catalyst, equilibrium is achieved faster, but the equilibrium composition remains unaltered