Redox Equilibria: Variable Oxidation States, Electrode Potentials, Electrochemical Series
Edited by Jamie (ScienceAid Editor), Taylor (ScienceAid Editor), vcdanht
Variable Oxidation States
On of the properties of transition metals is their variable oxidation states. This is crucial for looking at their reactions; the rules for working out the oxidation number is exactly the same in transition metals. The oxidation state of a metal can change in redox reactions. We represent this change of individual elements in half equations. Have a look at the example below.
When combining half equations; once you have them both (one for reduction and one for oxidation), you must make the number of electrons in each the same, then it is a simple matter of combining the two.
When a zinc wire is placed in a solution of zinc ions, an equilibrium of electrons leaving and joining zinc is set up. Zn (s) ==>> Zn2+ + 2e- By convention however, half equations are written as reductions, therefore the above equilibrium will be reversed. This also means that any enthalpy change will also be reversed.
If the equilibrium is more to the right and the zinc dissolves, a negative charge will build up. This arrangement is called a half cell, and if we could measure the charge, it would indicate how readily electrons are released and hence how good a reducing agent the metal is. This is the electrode potential.
It is possible however to form a circuit and measure the potential difference using a voltmeter. In order to do this, two half cells are made and connected via a salt bridge. A simple salt bridge would be filter paper soaked in salt solution. A wire cannot be used to complete the circuit because that would create another half cell.
Below the diagram is written the standard representation of cells. The single lines represent a change in phase and the double lines (sometimes dashed) represent the salt bridge. It can also be written if the oxidation state of a species changes, in this case a comma separates them. Or, if a platinum electrode is used because the particular element does not exist has a solid (as is the case with the standard hydrogen electrode) it will be written as below.
Pt | H2 (aq) | 2H+ (aq) || ...
In order to compare the electrode potential of different metals, it is important to use a standard cell to connect to other half cells to. For this, we use the hydrogen electrode.
Hydrogen is bubbled in to the solution and the black platinum wire is used to conduct electricity to the voltmeter. The potential of the hydrogen electrode is defined as zero, so when it is connected to another half cell the voltage, known as the electromotive force (e.m.f or E) can be measured - this is the electrode potential. However, the hydrogen electrode is difficult to use so it is much easier to use a secondary standard such as the calomel electrode or silver/silver chloride. These are calibrate against the hydrogen electrode. To calculate the standard E (Eθ) the second cell must be in standard conditions too: ions at 1 moldm-3, 298K.
A more positive Eθ means the metal is a strong oxidizing agent (a substance oxidizing another and itself reduced), and negative are strong reducing agents.
A list of electrode potentials is called an electrochemical series and these can be used to predict the direction of a redox reaction. It is possible to calculate the feasibility of a redox reaction using their E values the calculations involved are outlined below.
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Redox Equilibria: Variable Oxidation States, Electrode Potentials, Electrochemical Series. (2017). In ScienceAid. Retrieved May 25, 2019, from https://scienceaid.net/chemistry/physical/redeq.html
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Categories : Physical
Recent edits by: Taylor (ScienceAid Editor), Jamie (ScienceAid Editor)