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CHEM 1A Lecture Notes - Lecture 19: Supersaturation, Sodium Acetate, Equilibrium Constant

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6 Apr 2016

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Solubility of non-ionic solids in liquid solvents is an equilibrium process described by a solubility constant, ks. As always, the state of equilibrium involves matching rates: in this case the rates of dissoluion and precipitaion (crystallizaion) of a solute at the soluion/solute interface. Since pure compounds (solids and liquids) do not enter into equilibrium constant expressions, the magnitude of ks corresponds to the molar solubility of a solute in a given solvent. Ks also describes solubility of one liquid in another, or miscibility. What maters in such cases is the relaive magnitude of the pair-wise interacions of the two liquids, a-a, b-b and a-b. The solubility of gases in liquids is described by the henry"s law equilibrium constant, kh, which is equal to the raio of the solubility of a gas (mol/l) to the parial pressure of the gas above the liquid.

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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