Expt1 - CH 142 Experiment I: Chemical Equilibria and Le...

Info iconThis preview shows pages 1–2. Sign up to view the full content.

View Full Document Right Arrow Icon
CH 142 Experiment I: Chemical Equilibria and Le Châtelier’s Principle A local theatre company is interested in preparing solutions that look like blood for their upcoming production of Lizzie Borden . They have hired New England Chemical Solutions, Inc. (NECSI), to help them investigate the aqueous reaction of potassium thiocyanate with iron(III) nitrate that they have heard other companies are using as fake blood. You will investigate this equilibrium for NECSI both qualitatively and quantitatively . The following useful information for these experiments is excerpted from reliable Web sites, and is reproduced with permission of the authors. You should also prepare for this experiment by reading about chemical equilibria and Le Châtelier’s Principle (chapter 15 in your course textbook). The Iron-Thiocyanate Equilibrium When potassium thiocyanate [KNCS] is mixed with iron(III) nitrate [Fe(NO 3 ) 3 ] in solution, an equilibrium mixture of Fe +3 , NCS , and the complex ion FeNCS +2 is formed (equation 1). The solution also contains the spectator ions K + and NO 3 . The relative amounts of the ions participating in the reaction can be judged from the solution color, since in neutral to slightly acidic solutions, Fe +3 is light yellow, NCS is colorless, and FeNCS +2 is red. If the solution is initially reddish, and the equilibrium shifts to the right (more FeNCS +2 ), the solution becomes darker red, while if the equilibrium shifts to the left (less FeNCS +2 ), the solution becomes lighter red or straw yellow. You will add various reagents to this reaction at equilibrium to see if/how those reagents shift the equilibrium position of the reaction using the color of the resulting solution. Fe +3 + NCS FeNCS +2 (1) yellow colorless blood red Quantitatively, the relative amounts of the two reactants and the product are related by the equilibrium constant for the reaction; in this case, the formation constant K f , which is shown below. To precisely control the red color of the solution, it is necessary to know the value for K f . K f can be calculated through an experimental determination of the [FeNCS +2 ] eq using a standard curve (week 2) and deduction of the [Fe +3 ] eq and [NCS ] eq by subtracting the amount of FeNCS +2 produced from the known added initial amounts of Fe +3 and NCS (as that is how much was consumed during the reaction). [FeNCS +2 ] eq [Fe +3 ] eq [NCS ] eq Use of a Standard Curve In this technique a series of solutions with known concentrations is prepared and then a parameter such as absorbance is measured. This parameter is then plotted versus concentration to yield a standard curve, which is often a straight line (in this case how absorbance varies with concentration) with some degree of experimental error. Regression analysis of the data (e.g., with Excel) using the method of least squares allows determination of the best-fit line. A subsequent measurement of the absorbance for an unknown sample allows determination of its concentration using
Background image of page 1

Info iconThis preview has intentionally blurred sections. Sign up to view the full version.

View Full DocumentRight Arrow Icon
Image of page 2
This is the end of the preview. Sign up to access the rest of the document.

This note was uploaded on 02/09/2010 for the course CH CH242 taught by Professor Katz during the Spring '10 term at Colby.

Page1 / 6

Expt1 - CH 142 Experiment I: Chemical Equilibria and Le...

This preview shows document pages 1 - 2. Sign up to view the full document.

View Full Document Right Arrow Icon
Ask a homework question - tutors are online