The Determination of an Equilibrium Constant Purpose: The purpose of this lab is to determine the value of the equilibrium constant, K eq eq, for " the reaction between iron (III) ions and thiocyanate ions, C! # $e%&(aq) & C!" """' $eC!&(aq) This will be done by measurin absorbencies of different solutions with varyin concentrations of thiocyanate ions and discoverin the relationship between absorption and concentration of thiocyanate ions* This relationship can then be used to find equilibrium concentrations of each component and thus the equilibrium constant* This lab will also ascertain the concentration of an un+nown thiocyanate ion solution by usin the relationship we discovered before* Hypothesis: I hypothesie that the equilibrium constant, K eq eq, for the reaction between iron (III) " ions and thiocyanate ions, C! will be a lare positive number reater than one because the product of this reaction ($eC!&) is more stable than the separated ions* This means that products are favored over reactants and it leads to an equilibrium constant reater than one*
eq uilibrium constant for a reaction, one must first be Introduction: In order to calculate the equilibrium familiar with the notion of o f dynamic equilibrium* - dynamic equilibrium occurs in a reaction when product and reactant concentrations cease to chane and both are bein produced.consumed at an equal rate* In other words, the forward reaction and the reverse reaction proceed at the same rate* In the equilibrium reaction between iron (III) ions and thiocyanate ions, iron (III) and thiocyanate combine to form thiocyanatoiron, and at the same time thiocyanatoiron disassociates into iron and thiocyanate ions* There are two types of dynamic equilibrium# homoeneous equilibrium and heteroeneous equilibrium* In a homoeneous equilibrium, all the components of the reaction are in the same phase, ph ase, whether that be a as or a solution* /n the other hand, components in a heteroeneous equilibrium are in different phases* The reaction between iron (III) ion solution and thiocyanate ion solution results in a homoeneous equilibrium, because both reactants and products are aqueous solutions* -t any iven point durin a reaction, the equilibrium can be characteried by the equilibrium constant, K* -n equilibrium constant is obtained by lettin a reaction proceed to equilibrium and then measurin the concentrations of each component involved in that reaction and creatin a ratio of products to reactants* The equilibrium constant will remain the same for each reaction irreardless of initial concentrations* $or the eneric reaction a-&b0"""'cC&dD, the equilibrium constant is written as# K c 1 2C34c2D34d.2-34a2034b 2C34c2D34d.2-34a2034b The rate constant of a reaction is important, because it can tell whether the reaction is product favored or reactant favored* If K'5, the reaction is product favored* 6owever, if K75, the reaction is reactant favored* The value of K only chanes with a chane chan e in temperature* It is important to note that only concentrations of aqueous solutions and ases are ta+en into account in the equilibrium constant e8pression, solids and liquids are not* - reat deal of this lab involves usin a colorimeter to find absorbencies of different solutions at equilibrium* - colorimeter colorimeter measures the absorbance ab sorbance of particular wavelenths of liht for different solutions* The absorption of these solutions is directly variable to the concentration of solutes in the solution* The red $eC!& solution created in this lab absorbs blue liht, which is why the colorimeter must be set to 9:;nm in order to emit blue liht and measure the absorbency of each solution*
Procedure: In the first part of this lab, five solutions will be prepared with varyin amounts of C!"* They will be mi8ed with a constant volume of $e(!/%)% solution in all five test tubes in order to have the concentration of $e(!/%)% be a control* There will be a colorimeter connected to a ith all equilibrium concentrations +nown, the value of the equilibrium constant, K, can then be calculated* Data:
0ea+er
-bsorbance
5 % 9 A Bn+nown, =art II
*;? *95 *%@; *9@: *A:A *%5@
0est"fit line equation for the =art I standard solutions# y1:998&;*;@:9 oom Temperature# % derees Celsius Test Tube !umber 5 % 9 A
-bsorbance ;*;; ;*59% ;*5% ;*A@ ;*?
!et absorbance """""""""""""""""""""""""""""""""""" ;*59% ;*5% ;*A@ ;*?
Observations: The mi8ture turned red as soon as the two solutions combined • It was difficult to et the cap on the test tube riht away • The test tube was open for about a minute at times while the cuvette was bein rinsed and • filled
•
=aper towels used to wipe cuvette for finerprints are unreliable in cleanin up the simplest of messes*
Data Analysis:
Discussion of Theory: - dynamic chemical equilibrium in a reaction occurs when both reactants and products cease to chane concentrations and both the forward and reverse rates of reaction are equal* The equilibrium point of a reaction can be characteried by equilibrium constant, K* $or the eneral reaction a-&b0"cC&dD, this constant can be e8pressed as# Kc 1 2C34c2D34d.2-34a2034b The oal of this lab was to ultimately calculate the equilibrium constant of the reaction between iron(III) ions and thiocyanate ions# $e%&(aq) & C!" """' $eC!&(aq) The equilibrium constant only chanes with a chane in temperature, which is why the temperature of the room was initially recorded before mi8in an y chemicals* The temperature of the room was % derees Celsius* In order to calculate the constant, the concentrations of the reactants and products at equilibrium needed to be ascertained first* This was done by measurin absorbency values of mi8tures at a state of equilibrium that varied in their initial concentration of C!" ions* The pre"lab e8ercise showed the concentration o f $eC!& ions that would be in each test tube if the reaction ran to completion* This could be directly correlated to the concentration of C!" because one mol of C!"1 one mole of $eC!&* The first five solutions were prepared by addin e8actly A*;;m< of ;*;; $e(!/%)% solution into each test tube* Enouh water was also added to these test tubes so that the total volume of $e(!/%)%, C!", and water would yield A;*;m<* The first test tube contained no amount of C!" solution and its absorbency would function as a base for the rest* 0efore beinnin any tests, the calorimeter was calibrated with diluted water placed in the cuvette* Then, for the first test, the solution containin no C!" was used to rinse the cuvette thorouhly two times* The third time, the cuvette was filled about F with the solution* Durin this process of rinsin the cuvette and fillin it with solution, the test tube of solution would often be left e8posed to air for the entire period of time* This error would have resulted in the forward reaction proceedin at a rate faster than the reverse reaction, because conditions around the reaction vessel werenGt held constant throuhout* This would cause the concentration of $eC!& to be reater than its equilibrium concentration, and the equation for the best"fit line to be at a reater slope than it should be because the more than predicted $eC!& will absorb more liht than it should at a predicted concentration* >hen the un+nown solution is tested at the end and it is capped immediately, the resultin absorption value would correspond to a concentration less than its actual value* -fter fillin the cuvette with solution, it was then wiped for finerprints in order to ensure that the ma8imum amount of liht would enter the solution* everal times, the pape r towels used to wipe the cuvette clean were unreliable and didnGt appear to completely remove traces of finerprints* This would result in less liht enterin the cuvette and an absorption value that is less than e8pected* This lower absorption value would correspond to a lower than predicted concentration* -fter wipin down the cuvette, the colorimeter was set to blue liht, 9:;nm, because $eC!& absorbs this wavelenth of liht* /nce the absorbency readin for the solution remained stable at ;*; ;, the value was recorded on the
eventually predict a lower concentration value for the un+nown C!" solution than it should b e* /nce the chemicals were mi8ed and as the reaction proceeded, $e(!/%)% disassociated into nitrate ions and $e%& ions that react with C!" ions* The concentration values of C!" in each test tube.bea+er obtained in the pre"lab were used to correlate the absorbency values with concentration of C!" in a linear relationship* This line of best fit could be described as# y1:998&;*;@:9 In this equation, y represents absorbency and 8 represents the concentration of $eC!& initially* /ne purpose of this lab was to determine the concentration of an un+nown C!" solution* This was done by reactin it with the same amount of $e(!/%)% solution used in part one and measurin the absorbency* This absorbency value measured with the calorimeter could be traced on the line of best fit to find the correspondin concentration of thiocyanate* The e8perimental concentration of C!" was ascertained to be 5*;A85;4"9 * The actual molarity of the C!" solution was *;J5;4"% * The percent error in this lab was calculated to be "9*?
=art III of the lab involved findin absorption values for solutions similar to those in part I and II, e8cept this time the total volume of the solution would only be 5;*; m< instead of A;*;m<* /nce the absorption values were acquired they could be plued into the equation determined in part I to calculate for the concentration of $eC!& at equilibrium* The initial concentrations of the reactants could be determined by usin dimensional analysis to multiply the molarity before mi8in by the volume"to"volume ratio of the solute to the total solution respectively* >ith the initial concentrations of reactants determined, alon with the equilibrium concentration of the product, the concentrations of the reactants at equilibrium could be determined by usin a simple I*C*E table (Initial Chane Equilibrium)* ince the product, $eC!& started out at an initial concentration of e ro and only increased by one multiple of concentration chane (8), the equilibrium molarity of $eC!& would represent the chane in molarity of all components in the reaction equation* -lso, since all the coefficients in the reaction are simply one, the absolute values of chane in concentrations are equal for all components* $or e8ample, the reaction occurrin in test two had an initial $e%& concentration of @*;J5;4"9 molar, initial C!" concentration of 9*;J5;4"9 molar, and an equilibrium $eC!& concentration of 9*9J5;4"A molar* The I*C*E table for this reaction would loo+ li+e#
ince, the equilibrium concentration of $eC!& has been determined usin the absorption value and equation for line of best fit, the value of 8 has been ascertained to be that as well* The equilibrium concentrations of $e%& and C!" can be determined by subtractin the value of 8 from the initial concentrations of these reactants* The equilibrium concentrations for $e%& and C!" are calculated to be#
This process was used to calculate the equilibrium concentrations of all reactants and products in all five test tubes of =art III* The final piece to solvin the equilibrium constant at % derees Celcius was pluin in all equilibrium concentrations into the equation for Kc# Kc12C34c2D34d.2-34a2034b The equilibrium constants for test tubes "A were calculated to be 5;, 5;, ;, and 5; respectively* The overall constant for the reaction was calculated by averain these values outH it was calculated to be 5;* Conclusion: The purpose of this lab was to determine the value of the equilibrium constant, Keq, for the reaction between iron (III) ions and thiocyanate ions, C!" by measurin absorbencies of different solutions with varyin concentrations of thiocyanate ions and discoverin the relationship between absorption and concentration of thiocyanate ions* This relationship can then be used to find equilibrium concentrations of each component and thus the equilibrium constant* The equilibrium constant for this reaction at % derees Celcius was ascertained to be 5;* This labGs oal was also to calculate the concentration of an un+nown thiocyanate ion solution by usin the relationship we discovered before* The e8perimental concentration was discovered to be 5*;J5;4"9, which has a "9*? error from the actual concentration of *;J5;4"% *
