Applications of Redox
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Chemistry I · original Chapter 7 · combined Chapter 7
section breakdown
- Galvanic Cells
- Batteries
- Corrosion
Section 4
Applications of Redox
So now that we touched the basics of redox reactions, lets get into more of the applied theory behind it. Usually when dealing with redox reactions, the most common application for it is in electrochemistry. Now electrochemistry means exactly like it sounds. It describes how a chemical change can result in electrical current, or oppositely, produce a chemical change.
Now the first application that we can see this in is galvanic cells. Just as a reminder to things, remember that oxidation involves the loss of electrons for a compound and that reduction is when a compound gains electrons. Now for example purposes, lets use the reaction between MnO4- and Fe2+ in an acidic solution:
Now, it should have gotten to the point that when looking at this reaction you can tell that Mn is the oxidizing agent, therefore it is being reduced and that Fe is the reducing agent since it is being oxidized. But if you forgot about that, maybe you should review Section 7.2 over again. In addition, from this, you should be able to write out the half reactions for both the oxidized and reduced compound. In case you forgot again, here is what it’ll look like:
Interactive 2.3 Galvanic Cell with Salt Bridge
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Salt Bridge
Usually contains some sort of ion solution that will help balance overall charges in the solutions (ex. Na+ and Cl -)
Interactive 2.2 : Simple Galvanic Cell
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Separation
The key is separating the two aqueous solutions so that the reaction can still occur, while harnessing the electron transfer individually.
Conductive Wire
This wire will be the medium in which electron transfer will occur and flow
Now, the only problem is that with this set up is that ultimately there will be such a large charge buildup due to the constant flow of electrons that the left beaker (receiving electrons) will become more negative while the right beaker (losing electrons) will become more positive. Creating such a drastic charge separation will not allow electrons to flow properly. So in order to counter this, usually a salt bridge is installed that allows for ions to constantly be supplied to the reagent solutions and allow for electrons to continue flowing.
Gallery 2.3 Examples of Batteries
Car batteries are usually made of lead batteries. It has six galvanic cells in series where electron transfer occurs between the solid lead anodes and cathodes and the sulfuric acid solution.
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Image 1
Car batteries are usually made of lead batteries. It has six galvanic cells in series where electron transfer occurs between the solid lead anodes and cathodes and the sulfuric acid solution.
Image 2
Calculator batteries and watch batteries are known as dry cell batteries. Zinc acts as an anode and depending on the type of battery (acidic or alkaline), the cathode constituents may vary (NH4Cl for acidic and KOH or NaOH for alkaline).
A battery contains one or more electrochemical cells that use redox reactions to deliver electrical energy. Its cell potential is a voltage, meaning energy per unit charge; it is not itself the total energy stored. Total delivered energy also depends on how much charge passes and on operating conditions. A galvanic cell separates oxidation and reduction so that electrons can flow through an external circuit.
The typical half reactions for a dry cell battery in an acidic solution are,
The middle graphite rod in the middle helps transfer the electrons to whatever device is in need of the overall electric potential. Now there is another application of redox reactions that are just as important to consider as it pertains to an everyday phenomenon. And you can see what this is on the next page.
Interactive 2.4 Mechanism of Iron Corrosion
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Formation of Rust
The Fe2+ ions that are trapped in the moisture travel to a cathodic surface (much like how a salt bridge acts) and react with O2 to form rust.
Oxidation of Fe
When solid iron is exposed to air and water, Fe will be oxidized and turn to Fe2+ and produce pits.
Now these are very important reactions to consider because we see it in some case examples such as aluminum corrosion but more specifically the corrosion of iron. THe corrosion of iron is usually known as rust. This is the orangy-brownish material that see on some metals. Iron in this state becomes weak structurally and can easily be broken.
You can usually see this in bridges and buildings, which are usually made up of steel (composed of iron and other metal alloys). It is important to consider because the formation of rust and corrosion of these objects could result in the collapse of a bridge or building.
Corrosion can be localized where electrochemical conditions make an area anodic, often influenced by coating damage, oxygen differences or the environment. Pits are not generally caused simply by a higher local percentage of iron. A zinc coating can provide a physical barrier and, when electrically connected under suitable conditions, sacrificial anodic protection: zinc oxidizes preferentially while helping protect the iron. Distinguish the observation of pits from the chemical mechanism used to explain them.
Glossary and deeper reading · 14 refreshers
These editorial refreshers supplement the reading. Definitions and examples are added guidance; original source wording and recorded answers remain in their original context.
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Electrochemical cell
An electrochemical cell separates oxidation and reduction into electrode processes. A galvanic cell delivers electrical energy from a favored reaction, while an electrolytic cell uses supplied electrical energy to drive an unfavored direction. For example, a zinc–copper galvanic model transfers electrons through the external wire while ions carry charge through the solutions.
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Corrosion
Corrosion is deterioration of a material through chemical or electrochemical reactions with its environment. Metal corrosion commonly involves oxidation coupled to a reduction process. For example, rust formation involves iron oxidation together with reactions involving oxygen and water.
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Redox reaction
A redox reaction combines oxidation of one species with reduction of another. Electron loss and gain must balance in the net equation along with atoms and total charge. For example, zinc reducing copper ions involves zinc oxidation and copper-ion reduction.
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Oxidation
Oxidation is electron loss, represented by an increase in oxidation number. It occurs together with reduction so that electrons are conserved in the overall reaction. For example, a zinc atom becoming a zinc ion with charge +2 undergoes oxidation.
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Reduction
Reduction is electron gain, represented by a decrease in oxidation number. It is paired with oxidation in an electron-transfer reaction. For example, a copper ion with charge +2 gaining two electrons to form copper metal undergoes reduction.
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Oxidizing agent
An oxidizing agent causes another species to lose electrons. The agent accepts those electrons and is itself reduced. For example, a copper ion acts as an oxidizing agent when it accepts electrons from zinc metal.
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Reducing agent
A reducing agent causes another species to gain electrons. The agent supplies those electrons and is itself oxidized. For example, zinc acts as a reducing agent when it transfers electrons to copper ions.
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Half-reaction
A half-reaction describes either the oxidation or reduction part of a redox process. It includes electrons explicitly so that charge and electron transfer can be balanced. For example, zinc forming a zinc ion and two electrons is an oxidation half-reaction.
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Salt bridge
A salt bridge provides an ionic conducting path between separated half-cells. Ion movement helps prevent bulk charge buildup while electrons move through the external circuit. For example, bridge anions can move toward a half-cell where metal cations are being produced.
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Ion
An ion is an atom or group of atoms with a net electric charge. Ordinary ion formation changes electron count without changing the element's proton count. For example, a sodium atom losing one electron becomes a sodium ion with charge +1.
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Anode
The anode is the electrode where oxidation occurs. This definition applies to both galvanic and electrolytic cells. For example, a zinc electrode releasing zinc ions and electrons is an anode in the zinc–copper galvanic model.
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Cathode
The cathode is the electrode where reduction occurs. It may deposit metal or host another reduction without necessarily gaining mass. For example, a copper electrode receiving electrons that reduce dissolved copper ions is a cathode.
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Cell potential
Cell potential is an electrical potential difference associated with a specified cell reaction and conditions. A reduction-potential table supplies reference values for half-reactions, and voltage is not multiplied when equation coefficients are scaled. For example, balancing twice as many transferred electrons changes reaction amount bookkeeping but does not double the cell voltage.
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Alloy
An alloy is a metallic material containing more than one element. Added atoms can replace host atoms or occupy spaces between them, changing properties such as hardness. For example, carbon atoms in steel affect how easily iron atoms can move relative to one another.