Simple Redox Reactions

Simple Redox Reactions


Learning Experiences


Pre-Lab Questions

  1. - What is the crystalline structure of tungsten trioxide?
  2. - Why can small ions be readily inserted into the crystalline lattice of tungsten trioxide?


Simple Redox Reactions

Laboratory Exercise

INTRODUCTION:

In this lab., you will carry out a simple redox reaction. In a redox reaction, one component is oxidized, and the other is reduced to maintain balance. The simplest redox reaction is the oxidation of metals, e. g. the oxidation of copper giving copper oxide:

Cu + 1/2 O2 = CuO

In this reaction the copper is oxidized, and the oxygen is reduced. In an oxidation reaction, the valence (or oxidation state) of the oxidized specie increases, whereas that of the reduced specie decreases.

Looking at the above reaction, we can consider it as two half-reactions:

Cu = Cu2+ + 2e- oxidation

1/2 O2 + 2e- = O2- reduction

Thus, in an oxidation half-reaction electrons are generated, whereas in reduction reactions they are consumed.

A slightly more complex reaction is the highly exothermic reaction between sodium metal and water in which hydrogen gas is evolved:

Na + H2O = NaOH + 1/2 H2

In this case the sodium is oxidized and the hydrogen reduced. The driving force in all these reactions is a reduction in energy, and you will see in the electrochemistry section of your text how you can measure this energy in an electrochemical cell. Conversely all electrochemical cells involve redox reactions. Thus, all batteries, fuel cells and electrochromic windows are based on redox processes just as is the burning of fuel in your car.

A simple solid state redox reaction is that between lithium and titanium disulfide:

xLi + TiS2 = LixTiS2

for which the half-reactions are:

Li = Li+ + e- oxidation

and

TiS2 + e- = TiS2- reduction

In this case the negative charge is not localized on the titanium or the sulfur but is delocalized in the conduction band of the TiS2. This reaction occurs rapidly at room temperature, and is used in some lithium batteries. TiS2 has a layer structure, as shown in figure 1, in which every other metal layer is vacant. The lithium ions can be inserted into these sites through a simple expansion of the crystalline structure perpendicular to the TiS2 sheets. This type of reaction is also known as an intercalation reaction, and LixTiS2 as an intercalation compound. (Also commonly known as a sandwich compound).

Fig. 1 Structure of TiS2 and its reaction with lithium.

Because there is no strong chemical bonding, just weak van der Waals bonding between the TiS2 sheets, the sheets can be readily separated. Your teaching assistant will demonstrate this with large single crystals of molybdenite, the mineral form of MoS2, using just a piece of scotch tape. Because these sheets can be separated and thereby slide across each other easily, molybdenite is used as a solid lubricant in a number of applications. Graphite is similarly used.


Your instructor will assign you one of the following two experiments.

EXPERIMENTS


1. Hydrogen Insertion into Tungsten Trioxide.


Tungsten trioxide has a very simple structure consisting of WO6 octahedra joined at their corners. It may also be considered as having the perovskite structure of CaTiO3 with all the calcium sites vacant. In this experiment you are going to place hydrogen on these sites.

Fig. 2 Structure of HWO3. Red is oxygen, green is hydrogen.

Place 500 mg of WO3 into a 150 ml beaker. What is its color? Carefully pour about 50 ml dilute hydrochloric acid (3 molar) onto the WO3. Does anything happen? Now very carefully drop some zinc filings into the acid and observe what takes place. After all reaction has ceased, note the colors of all the products, and then filter off the solid product. After washing, place it on a filter paper and observe what happens when you leave it exposed to the air. Write balanced equations for both the reduction of WO3 and its subsequent re-oxidation.

Change in conductivity upon intercalation. Make pellets of WO3 and HxWO3 in a pellet press as you did in the superconductivity experiment. Place the WO3 pellet on a microscope slide, and measure the electrical resistance across its diameter using an ohmmeter (adjust the meter to read in the 10-100kW range). Repeat the measurement on the HxWO3 sample. How much difference is there? How do these conductivities compare with that of the superconductor?

The color and conductivity changes you have observed are due to the intercalation of protons into the cavities in the WO3 structure, and the donation of their electrons to the conduction band of the WO3 matrix. These electrons make the material behave like a metal, with both its conductivity and color being derived from free electron behavior. This coloration reaction is now being used in electrochromic displays for sun glasses, rear view mirrors in cars, and in the future for the glass panes in large buildings. In these applications the coloring reaction, i. e. your redox reaction, is carried out in an electrochemical cell.


2. Lithium Insertion into Vanadium Pentoxide.

(An experiment contributed by Professor T. Mallouk of Pennsylvania State University)


Vanadium pentoxide is a layered semiconductor which can be reductively intercalated according to the reaction shown below:

V2O5 + xLi+ + xe- = LixV2O5 (x less than or equal to 1)

In pure V2O5, the valence band is full and the conduction band is empty, so that it is a semiconductor. The band gap determines the color of V2O5. What color is your V2O5? Thus, what colors are absorbed by this semiconductor (remember, the color of the powder indicates the light which is reflected from it). Estimate the band gap of V2O5 from its color. When lithium is intercalated, the electrons initially enter levels between the valence and conduction band, and eventually the conduction band itself. Thus, the conductivity of V2O5 is enhanced on intercalation and its color changes. In this laboratory we cannot handle lithium metal, so you will work with lithium iodide. How thermodynamically stable is LiI? Based on the reactions you are about to perform what can you say about the energy of formation of LixV2O5?

Weigh out 0.5 g V2O5, and put it in a 50 ml beaker with 15 ml deionized water. Allow the suspension to settle for a minute or two. Weigh out 2.0 g lithium iodide (LiI), and carefully add the LiI to the suspension without stirring. Note any color changes which occur in the bottom of the beaker. After about 10 minutes, stir the solution briefly, and allow it to stand for another 10 minutes. Suction filter the solution through a Buchner funnel, and wash two times with 50 ml portions of water. Allow the solid to air-dry completely before proceeding further.

Note the color of the solid, and of the solution which passed through the filter. What is the source of these colors? Write out a balanced equation for the reaction of LiI and V2O5.

Change in conductivity upon intercalation. Make pellets of V2O5 and LixV2O5 in a pellet press as you did in the superconductivity experiment. Place the V2O5 pellet on a microscope slide, and measure the electrical resistance across its diameter using an ohmmeter (adjust the meter to read in the 10-100kW range). Repeat the measurement on the LixV2O5 sample. How much difference is there? How do these conductivities compare with that of the superconductor?

You will now re-oxidize the LixV2O5. Make a solution consisting of 20 ml bleach (grocery store bleach is approximately 5% sodium hypochlorite, NaOCl) and 80 ml water in a 150 ml beaker. Add the pellet of LixV2O5. Stir the suspension, stopping occasionally to let the solid settle so that you can observe any color changes. Suction filter and wash with water as before. What has happened to the LixV2O5? Write a balanced equation for its reaction with hypochlorite.

V2O5 has a layered structure with V=O---V bonds between the layers. Upon intercalation the weak O---V is stretched and the sheets buckle as the layers are separated to allow the lithium to be incorporated into the structure. At low lithium contents such as used in this experiment, the reaction is reversible. At higher lithium contents, the bonds are broken and the reaction becomes irreversible.


For more information please contact Stan Whittingham: stanwhit@binghamton.edu


Copyright © 1989-1995 M. Stanley Whittingham and Thomas E. Mallouk