Chemical Equilibrium Chapter 17 Chemical Equilibrium Equilibrium is

Chemical Equilibrium Chapter 17

Chemical Equilibrium § Equilibrium is a BALANCE between two opposing forces or processes § What does it take for two people to be in equilibrium on a teater totter? § When playin tug o’ war? § So, when does chemical equilibrium happen?

Chemical Equilibrium § When a reaction occurs until the concentrations of products and reactants no longer changes, the reaction is said to have reached equilibrium. § Does this mean that the reaction has stopped completely? § NO…it simply means that the RATE at which product is being produced and the reactants are being reformed is equal. § Does this mean that the NUMBER of products and reactants is equal? NO!

A Practical Example § Imagine 5 students inside the classroom and 25 students in the hallway. § If 2 students leave the classroom through one door when 2 students from the hallway enter the classroom, does the number of individuals in or out of the room change? § NO…The rate of entering is the same as the rate of leaving. § This is equilibrium…WITHOUT having equal numbers of students in and out

Equal movement in both directions means the rate is equal; therefore, equilibrium is met

A chemical example § Let’s look at a situation where chemical equilibrium is met… § When preparing a saturated solution of sodium chloride, equilibrium is met and can be written as follows: § Na. Cl (s) <--> Na+ (aq) + Cl- (aq) § Notice that the arrow is bidirectional § This MUST be a closed system…no more reactants or products are added from an outside source

§ We’ve discussed that equilibrium can exist only under conditions of constant temperature, pressure, volume, and concentration. § Henri Le. Chatelier examined what occurs when these factors do not remain constant § If one of these factors change, it is said to put “stress” on the reaction

Le. Chatelier’s Principle § When a stress is placed on a system in equilibrium, the system will adjust to remove the stress and to restore equilibrium in the system § 1. Changes in Concentration This principle allows us to predict the direction in which the equilibrium will shift when one or more of the concentrations of the products or reactants is altered.

Predicting Direction § You can use Le. Chatelier’s Principle to predict the direction of the “teeter” or “totter” § Which direction will the reaction move (toward the reactants or toward the products) in order to reestablish equilibrium? § Let’s Practice!

Changes in Concentration § CO(g) + 2 H 2 (g) <--> CH 3 OH (g) § § § What direction does the rxn shift if… More CO is added? Methanol is increased? Methanol is removed? Hydrogen gas source is reduced?

Le. Chatelier’s and Pressure § 2. The pressure of a system is directly proportional to the number of gas molecules present… § So the only way to reduce the pressure is to reduce the total number of molecules in the system § Increasing pressure on a gaseous system causes the equilibrium to shift to the side with the fewest number of molecules § So, if the opposite is true and pressure is decreased, then the eq shifts to the side with the greatest number of molecules

Try this § The following rxn has come to equilibrium in a container: § N 2 (g) + 3 H 2 (g) <--> 2 NH 3 (g) § In which direction will the rxn shift if the pressure on the system above decreases? § Left § Why? § Should the pressure on the above system be increased or decreased to produce more ammonia? Why?

Volume and Le. Chatelier’s § 3. When the volume of a rxn is reduced, the molecules are crowded together. § Decreasing the number of molecules can decrease the stress. § When the volume of a rxn involving gases decreases, the eq shifts to the side with the fewest number of molecules. (when the gas rxn volume increases, the eq shifts to the side with the greatest number of molecules. )

Try these: § Will the reaction shift toward the reactants (left) or the products (right) side if the VOLUME IS DECREASED: § PCl 5 (g) <--> PCl 3 (g) + Cl 2 (g) § Left § N 2 (g) + 3 H 2 (g) <--> 2 NH 3 (g) § Right § 2 CO (g) + O 2 (g) <--> 2 CO 2 (g) § Right

Temperature and Le. Chatelier’s § 4. Heating or cooling the reaction will result in shifting the eq to the left or right depending on whether the rxn is endothermic or exothermic. § Exothermic: § Increasing the temperature of an exothermic rxn § Reactants <--> products + heat energy is like increasing a product, so the shift will be to the left Decreasing will have the opposite effect

§ In an endothermic reaction however, increasing the temperature would be like increasing a reactant and would force the shift to the right § Reactants + heat energy <--> products

Relative Concentration § There is NO WAY you can tell what the concentrations of any of the products or reactants are by just looking at the equation; however, there is a ratio that scientists were able to come up with after much study: § The law of chemical equilibrium § Keq = [products] / [reactants] § Remember that the [brackets] mean “concentration of” whatever is inside

Keq § The calculation of Keq isn’t quite that simple, but conceptually it is. § For the reversible equation § a. A + b. B <--> x. X + y. Y § With the lower case letters being the number of moles § Keq is calculated: § [X]x[Y]y / [A]a[B]b

Keq § Try writing out the formula for Keq for the following reaction: § 4 NO(g) + 6 H 2 O(g) <--> 4 NH 3 (g) + 5 O 2 (g) § [NH 3]4[O 2]5 / [NO]4[H 2 O]6 = Keq § Practice: § Try a few problems from your teacher § Set up the equilibrium constants (Keq’s) for these equations

Practice Problems: § 1. 2 NO 2 (g) <--> N 2 O 4 (g) § 2. N 2 (g) + 3 H 2 (g) <--> 2 NH 3 (g) § 3. 2 NO (g) +2 H 2 (g) <--> N 2 (g) + 2 H 2 O (g)

In the Lab § Calculating the Keq in the lab involves writing the correct formula for Keq, substituting in the measured concentrations § H 2 + I 2 --> 2 HI § [H 2]=. 0056 M § [I 2]=. 00059 M § [HI]=. 0127 M

§ Keq = [HI]2 / [H 2][I 2] § Keq = (. 0127)2 / (. 0056)(. 00059) § Keq = 48. 8 § If you measured different concentrations in subsequent lab trials, calculate Keq and then average all of the Keq values.

Ksp § When an ionic solid is placed in water, an equilibrium is established between the ions in the saturated solution and the excess solid phase. § Ex. § Ag. Cl(s) <--> Ag+(aq) + Cl-(aq) § or Ag 3 PO 4 (s) <--> 3 Ag+(aq) + PO 4 -3(aq) § For each of these, we could write the Keq expressions. This is called the Ksp or solubility product § Do NOT include the solid (in Ksp or in Keq) § Write the Ksp for the above equations:

§ [Cl-][Ag+] = Ksp § And for the second reaction, [Ag+]3[PO 4 -3] = Ksp What does Ksp tell us? The larger the Ksp, the more soluble a salt is in water.

When doesn’t equilibrium occur? § All of the reactions we’ve discussed have been in EQUILIBRIUM, so do all reactions reach equilibrium? § What types of rxns don’t? § Strong acid ionization, precipitations, formation of a gas from an aqueous solution, and the formation of water as a product of the reaction § Why don’t these types of rxns reach eq?