1. Equilibrium

2.  The state of a chemical reaction in which its forward and reverse reactions occur at equal rates so that the concentration of the reactants and products does not change with time.  A state of balance or equity  Dynamic state – lots of activity; collisions; interactions

3.  Represented with double arrows in opposite directions  Over time concentrations become constant, rates become equal. ◦ Rate forward = Rate reverse ◦ Takes time; not instantaneous  Depends on: ◦ Initial concentrations ◦ Relative energies of reactants and products ◦ Degree of organization of reactants and products  Want lowest possible energy state and highest number of interactions

5.  If reverse reaction, flip K statement (reciprocal)  If change coefficients in a balanced equation, will change K.  NO units on K – the all cancel out.  Primarily works for solution and gaseous equilibria. ◦ Pure liquids and solids left out of expression  Usually all experimental data is for a given temperature (constant)  May be many equilibrium positions (set of concentrations) but only 1 value of K ◦ (ex. p613)

6. If 0.1908 mol CO 2, 0.0908 mol CH 4, 0.0092 mol O 2 are present in a 2.0 L container, determine the equilibrium constant for the reaction. CH 4 + O 2 ⇆ CO 2 + H 2 O 1. Balance equation 2. Find K expression 3. Solve for K 4. Now find K for the reverse reaction = K’

7.  K = K c = equilibrium with concentrations (mol/L)  K p = equilibrium with partial pressures  K might be equal to K p but not always K p = K(RT) ∆n ∆n = ∑ coefficients of products - ∑ coefficients reactants

9.  Likelihood of reaction to occur depends on value of K (extent of reaction) ◦ If K is greater than 1 – reaction consists of mostly products (equilibrium lies to right) ◦ If K is less than 1 – reaction consists of mostly reactants (equilibrium lies to left)  If not sure if in equilibrium, need to do further calculations/comparison ◦ use reaction quotient to determine if reaction at equilibrium or if it needs to shift to get there. ◦ Reaction quotient = Q = determined using initial concentrations of reactants and products

10.  Determined just like K; same expression  If Q = K  system (reaction) in equilibrium; no shifting to occur  If Q > K  system shifts to left ◦ Too many products are produced so need to consume those and make more reactants  If Q < K  system shifts to right ◦ Too many reactants present, so need to consume reactants and form products

11.  [equilibrium] = [initial] +/- [change] ◦ +/- depends on whether reactant or product  Use ICE method: 1. balanced equation needed 2. Write K expression 3. Identify initial concentrations 4. Identify the change (+/-X) 5. Write equilibrium concentrations 6. Find value of X by plugging in to K expression. 7. Find K.

13.  Tricks: ◦ Use quadratic formula to solve for X as needed. ◦ If you know how to graph it to solve, go ahead. ◦ If K is really small (x10 -5 or more), you can ignore X and assume it is really small also. ◦ If dealing with pressures, remember Dalton’s Law of Partial Pressures

14.  Understanding factors that affect the position of equilibrium  If a change is imposed on a system at equilibrium, the system will shift to compensate for the change.  Need to predict the shift direction and effect of change  Changes may include: ◦ Adding or removing reactant or product ◦ Adding new inert gas (not involved in reaction) ◦ Changing volume of vessel/container ◦ Change in temperature

15.  Volume change: ◦ Need to look at number of molecules present on reactant and product sides ◦ If volume decreased, shifts to decrease volume of system ◦ If volume increased, shifts to increase volume of system  Heat change: ◦ Determine if reaction is endothermic (heat a reactant) or exothermic (heat a product)  Temperature change: ◦ Temp affects K.  Addition of reactant: ◦ Shifts to produce more product  Addition of product: ◦ Shifts to give more reactants  Addition of catalyst: ◦ Nothing – only speeds up reaction