Chemical Equilibrium
What is Equilibrium? A + B C + D C + D A + B We have assumed that when reactants combine, they react entirely to completely create the products, but this isn’t always the case. A + B C + D In some reactions, once enough of the products are made, they can combine and form some of the original reactants. C + D A + B
What is Equilibrium? There is a constant, or dynamic, back and forth nature (creating products that create reactants) of the reaction until the reaction reaches EQUILIBRIUM. This is the point when the amount of products being created equals the amount of reactants being created. The rates of the forward and reverse reaction are occurring at the same speed.
Chemical equilibrium is achieved when: the rates of the forward and reverse reactions are equal. the concentrations of the reactants and products remain constant as time passes.
Concentration vs Time Graph equilibrium equilibrium equilibrium Start with NO2 Start with N2O4 Start with NO2 & N2O4 N2O4 (g) 2NO2 (g) 14.1
Rate of Reaction vs Time Graph Constant 14.1
We represent a chemical equation at equilibrium with a double arrow because both the forward and reverse reactions are happening simultaneously. 14.1
Equilibrium & K So, how will you know when a reaction reaches this mysterious “Equilibrium”??? Equilibrium depends on a lot of things: concentrations, temperature, pressure, ect. But, if the conditions are known, we can calculate the equilibrium constant (K) for a reaction. *Note this “K” is different from “k” in the kinetics lesson, but they are similar in concept.
Equilibrium & K The equilibrium constant (K) is a number needed for calculation purposes. It’s like R in the Ideal Gas Law: PV = nRt R = 0.082 atmL molK Knowing this value will let us solve for all other parts of the equation. We’re doing a similar thing here.
aW + bX ⟺ cY + dZ The Law of Mass Action How we solve for K (equilibrium constant) is similar to when we solved for k, the rate constant. The lowercase letters = coefficients The capital letters = compounds aW + bX ⟺ cY + dZ
aW + bX ⟺ cY + dZ [W]a [X]b The Law of Mass Action Equilibrium constant K = [Y]c [Z]d [W]a [X]b The brackets [ ] represent the molar concentrations. This time the coefficients do matter. The coefficients become the exponents! aW + bX ⟺ cY + dZ
Solving for Equilibrium = K Let’s look at a chemical reaction and set up an equation for the equilibrium expression: What should be on top? What should be on the bottom? Where did the “2” come from? K = [C]c[D]d [A]a[B]b aA + bB cC + dD N2O4 (g) 2 NO2 (g) K = [NO2]2 [N2O4]
Kb = [NOCl]2 [NO]2[Cl2] Kc = [Ag(NH3)2+] [Ag+][NH3]2 Ka = [O2]3 [O3]2 Write the equilibrium expression for K for the following reactions: Kb = [NOCl]2 [NO]2[Cl2] Kc = [Ag(NH3)2+] [Ag+][NH3]2 Ka = [O2]3 [O3]2 Note: Substances in solid or pure liquid form do not affect the equilibrium concentrations and are left out of the expression!
What Does the Value of K Mean? If K >> 1, the reaction is product-favored; More of the products are present when equilibrium is reached. If K << 1, the reaction is reactant-favored; More of the reactants are favored when equilibrium is reached.
Practice Problem: [N2][H2]3 At equilibrium, 1.00 M of Nitrogen gas reacts with 1.00 M Hydrogen gas to produce 2.00 M of Ammonia. Write the equilibrium expression, then solve for the equilibrium constant. Write the Balanced Equation: N2 + 3 H2 2 NH3 2) Write the Equilibrium Expression: K = [NH3]2 [N2][H2]3
Practice Problem: [1.00][1.00]3 K = 4.00 At equilibrium, 1.00 M of Nitrogen gas reacts with 1.00 M Hydrogen gas to produce 2.00 M Ammonia. Write the equilibrium expression, then solve for the equilibrium constant. 3) Solve for Equilibrium Constant (K): K = [2.00]2 [1.00][1.00]3 K = 4.00 *Notice: This problem wasn’t so bad because the values at equilibrium were given…..