Electrochemistry,+Redox

Hey there! You have probably just learned about electrochemistry, and you're probably like: "what the heck is this??" Well, that's what we're here for. But in order to understand this stuff, we have to start with the basics of reactions. The basics of redox reactions. Don't worry; you'll get it after you read about it. First we will start with some words that are important and pertain to redox reactions which will be followed by a breakdown of a problem and some answers to common questions. After we finish with redox reactions electrochemistry will be covered in the same format as the redox reactions.

Redox-Reactions, also known as oxidation-reduction reactions involve the transferring of electrons. Oxidation is when the atom loses electrons. Reduction is when the atom gains electrons. An easy way to remember the difference between oxidation and reduction is a simple acronym: //**O**//xidation //**I**//s //**L**//ose //**R**//eduction //**I**//s //**G**//ain or OIL RIG Befor you can begin to balance redox reactions the oxidation state of the elements needs to be determined. here are the steps to determine the oxidation number. Determining Oxidation States Balencing Rexod Reactions Electrolytical Cell __Nernst Equation__ Electrochemistry and The Real World //**Determining the Oxidation States **// 1. the oxidation number of an atom of a pure element is zero. For example O S8 Cu and Br2 all have an oxidation number of zero. 2. the oxidation number of a monatomic anion equals its charge. For example the oxidation number of Fe³+ us +3 while the oxidation number of Cl¯ is -1. 3. some common elements almost always have the same oxidation number like O is always -2 and H is always +1(these ones just need to be memorized there is no way around them) 4. the sum of the oxidation numbers in a neutral compound is zero. In a polyatomic ion the sum of the oxidation numer i equal to the charge of the ion. for example concider the ion Cr2O72- above. what is the oxidation number of Cr in this compound? We know that the sum of all oxidation numbers must be -2 so: 2(x) + 7 (-2) = -2. solving x=6, so the oxidation number of chromium in theis compound is 6. now that you know how to determine the oxidation state half reactions is the next step. Each reaction has two half reactions which break the whole reaction down into two parts which in the end make it easier to balence redox reactions. Here is a break down of a half reaction: [] Here is a practice problem with the answer listed below. Zn(s) + 2 H+(aq) + 2 Cl-(aq) à Zn2+(aq) + 2 Cl(aq) + H2(g)

answer to the net reaction is: Zn(s) + 2H+(aq) à Zn2+(aq) + H2(g) Here is a link to how to get the answer: []

now that we have gone over half reactions we can start balancing redox reactions using the oxidation numbers and half reactions that were discussed above. I am going to work through a problem with you and label it step by step. Here is the equation : Al(s) + MnO^4-(aq) à Al^3+(aq) + Mn^2+(aq) (acidic Conditions) Determine what is oxidized and what is reduced: Al and Mn are 0 and +7 respectively on the reactant side and +3 and +2 respectively on the product side. Since Al oxidation number is increasing it is oxidized and since Mn oxidation number is decreasing it is reduced. Break the atoms in each half reaction: Al(s) à Al^3+(aq) (oxidation half-reaction) MnO4-(aq) à Mn^2+(aq) (reduction half-reaction) Balance the half reactions: Al(s) à Al^3+(aq) MnO^4-(aq) + 8H^+(aq) à Mn^2+(aq) + 4H2O(l) Balance the reactions by charge: Al(s) à Al^3+(aq) + 3e- MnO^4-(aq) + 8H^+(aq) + 5e- à Mn^2+(aq) + 4H2O(l) Multiply the half reactions by the appropriate factors so that the oxidation agent accepts as many electrons as the reducing agent produces. 5Al(s) à 5Al^3+(aq) + 15e- 3MnO^4-(aq) + 24H^+(aq) + 15e- à 3Mn^2+(aq) + 12H2O(l) (For basic reactions you need to add OH- on both sides to neutralize the hydrogen ion the hydrogen and hydroxide combined will make an equal amount of water.) Add the half reactions together to get the overall reaction: 5Al(s) à 5Al^3+(aq) + 15e- 3MnO^4-(aq) + 24H^+(aq) + 15e- à 3Mn^2+(aq) + 12H2O(l) 5Al(s) 3MnO^4-(aq) + 24H^+(aq) à 5Al^3+(aq) + 3Mn^2+(aq) + 12H2O(l) Check the balanced equation to make sure the masses and charges are balanced.
 * //Balancing redox reactions //**

Click here to watch a youtube video for how to assign redox and oxidation. [] Contributed by Courtney Anderson

Click here for a video explaining how to balance redox reactions in both an [|acid] and a [|base]. Videos contributed by Perrin Duvall.


 * Electrochemical Cell **

The electrochemical cell is an arrangement of an oxidizing and reducing agents such that they can only react if electrons flow through an outside conductor. A salt bridge is required to allow ions to flow between the electrodes.

Photo Credit to [|http://educ.queensu.ca/~science/main/concept/chem/c12/Eric/Grade%2011%20Electrochemistry/index.html] Here is the equation of an electrochemical cell: Zn (s) + Cu^2+ (aq) à Zn^2+ (aq) + Cu (s) Two half-reactions occur in the two separate beakers. Here they are: Zn (s) à Zn^2+ (aq) + 2eˉ (Oxidation) Cu^2+ (aq) + 2eˉ à Cu (s) (Reduction)

Electrode is the general term for a conductor that conducts electrical current in or out of something. The electrodes in this situation are the Zinc and Copper strips. The site of reduction (Cu strip) is called the cathode. The site of oxidation (Zn strip) is called the athode. Electrons flow away from the anode and towards the cathode.

From Nate: A nonspontaneous redox reaction, where electricity is added to force the reaction to occur, is known as an electrolytic cell. This process is known as electrolysis, and a voltmeter will give a negative reading for this reaction. It is often used in the industrial production of chemicals. A galvanic cell, also known as the voltaic cell, is when the redox reaction is spontaneous. Galvanic cells produce electricyity spontaneously, which can be routed through appliances motors and more. Galvanic cells will give a positive voltmeter reading. [|Here] is an examples of a galvanic cell~from Katelyn.

From Drew: This is an excelent visualization of a galvanic cell that demonstrates how the electrodes and the salt bridge work. []

Galvanic corrosion is the process that degrades metals electrochemically. The damage is caused when two different materials are placed together in a corrosive electrolyte. One of the metals becomes the anode and the other one becomes the cathode. The anode corrodes more quickly then it would otherwise, and the cathode corrodes more slowly. Temperature, humidity, the size of the anode and cathode, and the type of electrolyte can all influence galvanic corrosion.~Catherine photo from []

From Stu Stu: Under standard conditions, the voltage of a cell is equal to the voltage of the redox reaction. Under nonstandard conditions, we have to use the Nernst Equation to find the cell voltage. Ecell=Eºcell – (RT/nF)*lnQ Ecell=cell potential under nonstandard conditions Eºcell=cell potential under standard conditions R=the gas constant, 8.31 (volt-coulomb)/(mol-K) T= temperature in Kelvin (K) n= number of moles of electrons exchanged in reacton (mol) F= Faraday's constant, 96,500 coulombs/ mole Q= the reaction quotient (same as K, but using inital concentrations not at equilibrium)
 * //Nernst Equation //**

At 25ºC, the Nernst equation can be written as: Ecell=Eºcell- (0.0592)/n*logQ [|This is an example of the Nernst Equation being used.] Video contributed by Stuart Taliaferro


 * Important* As the concentration of the products of a redox reaction increases, the voltage decreases. As the concentration of the reactants in a redox reaction increases, the voltage increases. This is Le Chatelier's Law. Hi its Rachael! To see an example of this click here!

Electrochemistry is useful in the real world for the fact that it can be used to electroplate objects and give them a better finish. Electroplating is a finshing process that uses electrical current to reduce cations of a desired material from a solution and coat a conductive object with a thin layer of the material, such Cr, Ni, Au, or Ag. Electroplating is primarily used for depositing a layer of material to give the object a desired property (abrasion and wear resistance, corrosion protection, lubricity, aesthetic qualities, etc.) to a surface that otherwise lacks that property. The process of electroplating is called electrodeposition. In one technique, the anode is made of the metal to be plated on the part. Both components are immersed in a solution called an electrolyte containing one or more dissolved metal salts as well as other ions that permit the flow of electricity. A rectifier supplies a direct current to the anode, oxidizing the metal molecules that comprise it and allowing them to dissolve in the solution. At the cathode, the dissolved metal ions in the electrolyte solution are reduced at the interface between the solution and the cathode, such that they "plate out" onto the cathode. The rate at which the anode is dissolved is equal to the rate at which the cathode is plated, vis-a-vis the current flowing through the circuit. In this manner, the ions in the electrolyte bath are continuously replenished by the anode, and the item is electroplated with the wanted metal. The outer panels on the car are electroplated gold.... very nice and very expensive (the paint job probably costs as much as the car) (Electrochemistry and the Real World was created by Skyler Kuhn)
 * Electrochemistry and the Real World**:

Rhodium Plating a Penny By Cynthia Nuñez [] __Bibliography:__ Zumdahl, Steven S., and Susan L. Zumdahl. __Chemistry__. 6th ed. Florence: Brooks Cole, 2002. Masterton, William L., and Cencile N. Hurley. __Chemistry: Principles and Reactions__. 5th ed. Florence: Brooks Cole, 2004. Wilson, David. __Kaplan AP Chemistry 2009__. 1st ed. New York: Kaplan, 2009. Jespersen Ph.D., Neil D. __Barron's AP Chemistry__. 4th ed. New York: Barron's Educatinal Series, 2007. J.W. Oldfield: "Electrochemical Theory of Galvanic Corrosion", in "Galvanic Corrosion" - ASTM STP 979 - H.P. Hack Ed., ASTM International, West Conshohocken (PA), 1988.