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CHEM102 Lecture Notes - Lecture 10: Exothermic Process, Quadratic Equation, Equilibrium Constant

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23 Jun 2015

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CHEM102 Full Course Notes

CHEM102 Full Course Notes

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The flying mccoys by glenn and gary mccoy. Altering equilibrium conditions: concentration, adding/removing reactants/products, pressure, volume, temperature, catalyst. [h+] = [hcoo ] = k [hcooh] = 0. 014 m. Add (hcoo)na so that [hcoo ] = 0. 02 m. Result: more hcooh when equilibrium will be reestablished. If a change is imposed on a system at equilibrium, the position of equilibrium will shift in a direction that tends to reduce that change: effect of concentration. An example of concentration change, the dissociation of formic acid. Hcooh (aq) + h2o (l) h3o+ (aq) + hcoo (aq) [hcooh] since the source for both is the same hcooh. If a component is added to a system at equilibrium, the equilibrium position will shift in the direction that tries to lower back the concentration of that component. If the component is removed the reverse effect occurs. No change with addition/removal of solids provided some are left in the rxn flask.

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Related Questions

Any help/explanations with this Le Chatelier's Principle lab would be extremely helpful. I attemtped to do question A, but I'm not exactly sure if my answer is interpreted correctly. I have included the actual lab procedure with the 5 questions. Thank you again.

ENT ives To determine equilibrium concentrations for species in solution . Examine temperature effects on reaction equilibrium Investigate the shift in equilibrium from different stresses on a system kills . Use colorimeter to determine unknown solution concentrations ced .Preparing temperature baths to change reaction conditions tionEquilibrium is defined as a state in which the concentrations of both the reactants and products have no tendency to change with time. We can restate this by saying that the rate of the forward reaction is equal to the rate of the reverse reaction. Recall that with rate we are simply defining how fast a reaction is proceeding, a speed at which the reaction proceeds. Take note that when writing an equilibrium expression, pure liquids (denoted by l or 2) and pure solids (denoted by s) are omitted. Typically we express the equilibrium by K or Kp where c and p denote either concentration or partial pressure respectively. We refer to this value of K, or Kp as an equilibrium constant For any reaction, Kc can be written as the product of the concentration of products, raised to their stoichiometric coefficients, divided by the product of the concentration of the reactants raised to their stoichiometric coefficients. See Equation (5-1) below for an example: aA(aq) + bB(aq-cC(aq) + dD(aq) (5-1) A] [B]b As mentioned before, K, or K, are equilibrium constants meaning that they will have the same value unless temperature is changed. This brings in two important ideas 1) We can perturb the equilibrium concentrations of a reaction, say [A] for instance which will cause the other concentrations to shift so that K for the reaction remains constant. 2) If we change temperature, we will either increase or decrease our value of K depending on the specific reaction we are examining. Determining K, and Equilibrium Concentrations During lecture we often discuss the idea of using ICE tables as a method for organizing and calculating concentrations of reactions as they proceed towards equilibrium. In a typical ICE table you need to know initial concentrations and the magnitude of K to determine the equilibrium concentrations. Alternatively, if you know the equilibrium concentrations for a reaction you can calculate K. For this experiment we will be examining a complex ion formation, when iron(III) ion