Lets Review!!!: Chemical Equilibrium By: Emma and Gabe

Chemical Equilibrium: An Introduction Some chemical reactions DO NOT reach completion where all reactants are converted into products  reach a constant state instead Chemical Equilibrium= the state where concentrations of all reactants and products remain constant over time Terminology: Far to the right = at one time, more products exist than reactants Far to the left = a small amount of the reactants are converted into products

13.1: The Equilibrium Condition Remember: equilibrium is DYNAMIC on the atomic level… the net result is no change Why does equilibrium happen? As the reactions progresses, the concentrations of the products build up, initiating the reverse reaction Eventually, the concentrations reach levels where the forward and reversion reaction rates are EQUAL Factors affecting equilibrium: initial concentrations, energies of reactants/products, & “organization” *NOTE*= nature tries to achieve MAXIMUM disorder and LOWEST energy state

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13.2: The Equilibrium Constant Law of Mass Action: aA + bB  cC + dD Applies to systems in solution AND gas phase Remember that equilibrium constant is a RATIO and the concentrations can vary Equilibrium Expression: Use the equilibrium concentrations in this expression Do you guys remember how to use it?

Just for Grins… Let’s Practice Write the equilibrium expression for the following reaction: N2(g) + 3H2(g)  2NH3(g)

13.3: Equilibrium Expressions Involving Pressures Equilibrium expression in terms of pressure: Remember that these are the partial pressures of the gases The problem-solving process is still the same

13.4: Heterogeneous Equilibria Homogeneous Equilibria: both the reactants and products are in the same state Heterogeneous Equilibria: the reactants and products exists in different states For equilibrium calculations, the position does NOT include the amounts of pure solids or liquids  the concentrations of these substances can’t change

Let’s Practice… Again Write the equilibrium expression for this chemical reaction: 2H2O(l)  2H2(g) + O2(g)

13.5: Applications of the Equilibrium Constant Larger K value  reaction system will form mostly products Smaller K value  reaction system will form mostly reactants Remember: the size of K and the time required to reach equilibrium are NOT directly related Reaction Quotient (Q) = used to determine shift in a reaction Apply the law of mass action using initial concentrations rather than those at equilibrium Q equal to K: the system is already at equilibrium Q greater than K: the system shifts to the left Q less than K: the system shifts to the right Also remember how to do equilibrium pressure calculations…

Third Times a Charm… Let’s Practice For the synthesis of ammonia at 500°C, the equilibrium constant is 6.0 x 10-2. Predict the direction in which the system will shift to reach equilibrium at the following initial concentrations: [NH3]= 0.01 M; [N2]= 0.001 M; [H2]= .02 M At a defined temperature, a 1.00-L flask initially contained 0.301 moles of PCl3 gas and 0.00898 moles of PCl5 gas. After the system had reached equilibrium, 0.0021 moles of Cl2 gas was found in the flask. Calculate the equilibrium concentrations of all species and the K value. Hint: PCl5  PCl3 + Cl2

13.6: Solving Equilibrium Problems Problem-Solving Steps for Equilibrium Problems; Balance the equation. Using law of mass action, write the equilibrium expression. ICE chart it out. Calculate Q to find the direct of shift to reach equilibrium. Apply the change to the initial concentrations in the ICE chart. Solve for Keq with these new equilibrium concentrations & solve for the unknown.

Last Time… Everybody Practice Assume that hydrogen iodide gas is synthesized from hydrogen and iodine gas at a temperature with the equilibrium constant 100. In a reaction system, HI is at 0.500 atm, H2 at 0.0100 atm, and I2 at 0.00500 atm in a 5.00 L flask. Calculate the equilibrium pressures of all of the species.

13.7: Le Châtelier’s Principle Le Châtelier’s Principle: if a change is imposed on a system at equilibrium, the position of the equilibrium will shift in a direction that tends to reduce that change If the concentration of a component increases, the equilibrium position will shift in the opposite direction Addition of an inert gas does NOT effect the equilibrium position When the volume of the system is reduced, the system will decrease its total number of gas molecules Remember: The above changes do NOT affect the equilibrium constant… only temperature alters the K value