Original text
editA half-cell is a structure that contains a conductive electrode and a surrounding conductive electrolyte separated by a naturally occurring Helmholtz double layer. Chemical reactions within this layer momentarily pump electric charges between the electrode and the electrolyte, resulting in a potential difference between the electrode and the electrolyte. The typical anode reaction involves a metal atom in the electrode dissolved and transported as a positive ion across the double layer, causing the electrolyte to acquire a net positive charge while the electrode acquires a net negative charge. The growing potential difference creates an intense electric field within the double layer, and the potential rises in value until the field halts the net charge-pumping reactions. This self-limiting action occurs almost instantly in an isolated half-cell; in applications two dissimilar half-cells are appropriately connected to constitute a Galvanic cell.
A standard half-cell, used in electrochemistry, consists of a metal electrode in a 1 molar (1 mol/L) aqueous solution of the metal's salt, at 298 kelvin (25 °C).[1] The electrochemical series, which consists of standard electrode potentials and is closely related to the reactivity series, was generated by measuring the difference in potential between the metal half-cell in a circuit with a standard hydrogen half-cell, connected by a salt bridge.
The standard hydrogen half-cell:
- 2H+(aq) + 2e- → H2(g)
The half-cells of a Daniell cell:
References
edit
Category:Electrochemistry Category:Electrolysis
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Revised text
editA half-cell is a structure that contains a conductive electrode and a surrounding conductive electrolyte separated by a naturally occurring Helmholtz double layer. In other words, a half-cell is either the anode or cathode portion of a Electrochemical cell. Each half-cell together makes up the entire electrochemical cell. Each half cell either has a reduction or an oxidation. An easy way to remember this is "An Ox" stands for, the anode oxidizes. And "Red Cat" stands for reduction occurs at the cathode. Chemical reactions within this layer momentarily pump electric charges between the electrode and the electrolyte, resulting in a potential difference between the electrode and the electrolyte. The typical anode reaction involves a metal atom in the electrode dissolved and transported as a positive ion across the double layer, causing the electrolyte to acquire a net positive charge while the electrode acquires a net negative charge. The growing potential difference creates an intense electric field within the double layer, and the potential rises in value until the field halts the net charge-pumping reactions. This self-limiting action occurs almost instantly in an isolated half-cell; in applications two dissimilar half-cells are appropriately connected to constitute a Galvanic cell.
Anodes In a Galvanic Cell the Anode is where the oxidation is occurring. In this half-cell, the solid portion of the electrode will be sending electrons from away from the half-cell and into the other. Sometimes this results in some of the electrode releasing positively charged ions into solution, such as the Zn(s), Zn(aq) half-cell. In a forward reaction, this half-cell would be considered negatively charged, and the black lead would be connected to the electrode. In an electrolytic cell however, the reaction is being driven by an outside source. This results in the Anode being positively charged and the red lead is attached to the electrode instead.
Cathodes In a Galvanic Cell the Cathode is where the reduction is occurring. In this half-cell, the solid portion of the electrode will be receiving elections from the lead and adding them to ions in solution, such as the Cu(s), Cu(aq) half-cell. In a forward reaction, this half-cell would be considered positively charged, and the red lead would be connected to the electrode. In an electrolytic cell however, the reaction is being driven by an outside source. This results in the cathode being negatively charged and the black lead is attached to the electrode instead.
Example of Galvanic Cell: |Zn(s)||Zn(aq)|Cu(aq)||Cu(s)|
Example of Half-Cells: |Zn(s)||Zn(aq)| (Anode) and |Cu(aq)||Cu(s)| (Cathode)
A standard half-cell, used in electrochemistry, consists of a metal electrode in a 1 molar (1 mol/L) aqueous solution of the metal's salt, at 298 kelvin (25 °C).[1] The electrochemical series, which consists of standard electrode potentials and is closely related to the reactivity series, was generated by measuring the difference in potential between the metal half-cell in a circuit with a standard hydrogen half-cell, connected by a salt bridge.
The standard hydrogen half-cell:
- 2H+(aq) + 2e- → H2(g)
The half-cells of a Daniell cell:
References
editCategory:Electrochemistry Category:Electrolysis
| | This article about analytical chemistry is a stub. You can help Wikipedia by expanding it. |