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CHE 132 Lecture 7: CHE 132 Lecture 7

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viridianhare340

12 Feb 2019

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Stony Brook University

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Course

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Fernando Raineri

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CHE 132 verified notes

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CHE 132 Lecture 7: CHE 132 Lecture 7

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Document Summary

The concentration will not change by a large amount so we can neglect its value: reversing a reaction is when you go from reverse to forward or vice versa. The k value of the reverse can be discovered by 1 / k: remember to use stoichiometric coefficients when dealing with k. you put these numbers in the exponents in the reaction. If you combine two reactions, it is the same as multiplying the two k values: the reaction quotient (q) is similar to k except that you can use concentrations at any state for q. For k, you can only use concentrations at equilibrium. If k=q then the reaction is at equilibrium. If q < k then the products will be lower and you have more reactants. This will cause the reaction to shift to the right so that you can make more products. If q > k then the products will be greater than reactants.

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Related Questions

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