Electrochemistry

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Electrochemistry is the branch of chemistry that deals with the chemical action of electricity and the production of electricity by chemical reactions. Of particular importance are the interface of an electronic conductor (typically a metal electrode) and an ionic conductor (the electrolyte).

Electrolytic Cell

An electrolytic cell for leak examination.

Electrolytic cells are used in the batteries. They are changing chemical energy into electronic energy. The electrolytic cells have two kinds of chemical reaction, reduction and oxidation. These two reactions are spontaneous reactions and are driven by electrolysis.[1]

Voltaic Cell

Voltaic cells or Galvanic Cells are the cells that could help change chemical energy into electrical energy through the redox reactions.

Usually we can see zinc and copper in a normal battery, and these two metals are the electrodes, which can transfer electrons to or from a substance. The oxidation reactions occur on the anode (the positive electrode). It is the electrode to which anions flow (negatively charged ions) to be "oxidized" by electrons on the anode. The cathode is where reduction occurs (negative electrode). It is the electrode where cations flow (positively charged ions) to be "reduced" by reacting with (negatively-charged) electrons on the cathode. Therefore, the electrons flow from the anode to the cathode. Electrons are moving all the time to keep the reaction balanced. (Wilbraham, p665)

When reduction or oxidation occurs, the metal ion in each part of the reaction is a half-cell, and these half-cells are connected by a salt bridge. The salt bridge can make the ions move from one half-cell to the other but not completely mixed together. We usually use the metal like platinum as the conducting wire.

The activity series of metals determined which half-cell oxidized and usually the more active metal is the one oxidized during the reaction.

Dry Cell

People use voltaic cells as energy sources, and a dry cell is a voltaic cell in the pasted electrolyte. The flashlight battery is one common dry cell. In the dry cell, the moist paste contains MnO₂, ZnCl₂, NH₄Cl, and H₂O. The normal flashlight battery uses zinc as the container and zinc is also the anode; the graphite rod exists in the moist paste and it is the cathode.

This is the half-reaction for the cells:

Oxidation: Zn--->Zn²⁺+2e^- (anode reaction)

Reduction: 2MnO₂+2NH₄⁺+2e^---->Mn₂O₂+2NH₃+H₂O (cathode reaction)

The graphite rod only serves as conductor and does not experience reduction in the ordinary dry cells.

Lead Storage Batteries

We use lead storage batteries in the cars. There are six voltaic cells in one 12V car battery. Each voltaic cell produced 2V electric power. In the lead storage battery lead(IV) oxide is the cathode, it filled in the lead grid. The lead grid which filled with spongy lead is the anode side.

The half reactions:

Oxidation：Pb+SO₄--->PbSO₄+2e^-

Reduction:PbO₂+4H⁺+SO₄+2e^---->PbSO₄+2H₂O (Wilbraham, p669)

Electrical Potential

People use electrical potential to measure the ability of a cell can produce an electric current. We cannot measure two separate half-cells, but we can measure the difference between them, which called the cell potential, by using this formula:

Cell Potential= reduction potential in reduction - reduction potential in oxidation or E⁰_cell= E⁰_red - E⁰_oxid

Standard Cell Potential

In standard cell potential, the E⁰_red represent the standard reduction potential for reduction, and E⁰_oxid represent the standard reduction potential for oxidation.

The standard reduction potential of hydrogen electrode at 25 degree is 0.00V. Usually shows as E⁰_H⁺=0.00V [2]

Standard Reduction Potential

Usually when we need to determine the standard reduction potential of some half-cells, can use the standard reduction potential of reduction subtract the standard reduction potential of oxidation, which shows in this formula:

E⁰_cell= E⁰_red- E⁰_oxid [3]

Electrolytic vs. Voltaic Cells

Electrolytic cells can use electrical energy to change the chemical properties. The electrons inside the electrolytic cells need an outside power to push them completing the flowing in the reaction, but the voltaic cells do not need a outside power to do this. (Wilbraham, p679)

Electrolysis of Water

When we add a low concentration acid into pure water, the electrolysis might happen and the products of the electrolysis of water are the hydrogen gas (H₂) and oxygen gas (O₂).

Reduction: 2H₂O+2e^---->H₂+2OH^- (cathode)

Oxidation: 2H₂O--->O₂+4H⁺+4e^- (anode)

The region that around the cathode and anode turns to basic and acidic because the OH^- and H⁺ we got from the half reactions. (Wilbraham, p680)

Uses of Electrolysis

People usually use electrolysis in metal processing. (Wilbraham, p682)

Electroplating

There are a lot of plated stuff around our life. Especially some jewelries, usually they were plated with gold or platinum to make them looks have more value and have better looking. And basically electroplating is just put a thin metal layer outside the products and this layer might help the object from corrosion. (Wilbraham, p682)

Electrowinning

Electrowinning is a way to purify metals using electrolytic cells. (Wilbraham, p682)

Electorefining

By using electorefining we could produce pure metals and determine the impure ones. (Wilbraham, p682)

References

• Prentice Hall Chemistry. Wilbraham, Staley, Matta, Wateram. 2008 by Pearson Education Inc., publishing as Pearson Prentice Hall, Boston, Massachusetts 02116.
• Electrolytic Cells by Bodner Research Web
• Voltaic Cells by R.H. Logan, Instructor of Chemistry, DCCCD
• Demonstration of a Voltaic Cell by Blackgold
• Dry cell by Wikipedia
• The Dry-Cell Battery by Library.kcc.hawaii.edu

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See Also

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