Experiment 5

Experiment 5 - EXPERIMENT 5 Determination of an Equilibrium...

Info iconThis preview shows pages 1–3. Sign up to view the full content.

View Full Document Right Arrow Icon
EXPERIMENT 5 Determination of an Equilibrium Constant 41 E XPERIMENT 5 Determination of an Equilibrium Constant 5.1 Purpose In experiment 5 , the equilibrium constant for the formation of a transition metal complex ion will be determined. 5.2 Background Many transition metal ions form complex ions , which are composed of a central metal cation to which other groups called ligands are bonded. Ligands can either be ionic species, typically anions, or neutral molecules, such as water or ammonia. Substances containing complex ions are generally referred to as coordination compounds . Since many transition metal complexes have a distinctive color, the concentration of transition metal complexes can easily be determined using spectrophotometric measurements. The iron(III) cation forms a complex with the thiocyanate anion SCN - that is of a deep blood-red color. Thiocyanate is used in photographic processes, and the iron(III)thiocyanoto complex {Fe(SCN)} 2+ is often used to determine the concentration of thiocyanate solutions. The chemical equation for the complex forming reaction is shown in equation 5.1: The reaction of equation 5.1 represents a chemical equilibrium , and the corresponding equilibrium constant K C is defined in equation 5.2: However, when measuring the equilibrium constant for the complex forming reaction of equation 5.1, a few complications arise. When an iron(III) salt, such as Fe(NO 3 ) 3 , is dissolved in water, the iron(III) cation forms a hexaquo complex {Fe(H 2 O) 6 } 3+ . This complex is acidic in nature, and eventually leads to the formation of insoluble iron(III) hydroxide, equation 5.3. Fe 3+ + SCN - {Fe(SCN)} 2+ (5.1) [{Fe(SCN)} 2+ ] [Fe 3+ ][SCN - ] K C = (5.2)
Background image of page 1

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

View Full DocumentRight Arrow Icon
EXPERIMENT 5 Determination of an Equilibrium Constant 42 The chemical reactions outlined in equation 5.3 would strongly influence the concentration of free iron(III) cation. In order to prevent the precipitation of insoluble iron(III)hydroxide, the formation of the iron(III)thiocyanoto complex is carried out in acidic solution. The presence of an acid shifts the equilibria of equation 5.3 to the left hand side, thus preventing the formation of any hydroxide complexes. Preferably, the acid employed has the same conjugated-base anion as the iron salt used. Since in this experiment a solution of iron(III) nitrate is used, nitric acid HNO 3 is the acid of choice. A second problem arises in the construction of Beer’s law plot. Since the formation of the iron(iii) thiocyanato complex is an equilibrium process, it does not proceed to completion. Thus, the concentration of the species involved in complex formation is not known a priori . Once again, the corresponding equilibrium will be accordingly shifted employing LeChatelier’s principle. When preparing solutions for Beer’s
Background image of page 2
Image of page 3
This is the end of the preview. Sign up to access the rest of the document.

This note was uploaded on 04/29/2008 for the course CHEM 118 taught by Professor Jacobsen during the Spring '08 term at Tulane.

Page1 / 6

Experiment 5 - EXPERIMENT 5 Determination of an Equilibrium...

This preview shows document pages 1 - 3. Sign up to view the full document.

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