Group+B


 * ﻿ ﻿ ﻿ ﻿ This will be the space for the second group's answers.**


 * This should include:**
 * The original posed question
 * The solution
 * A step-by-step walkthrough of how to solve the problem
 * Background history explaining the chemistry of the problem (i.e. when in nature or in industry does said reaction occur, who first discovered the reaction, what are some of the common concerns with the reaction, with references)
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** Authors: ** Andi C. and Eric P.  ** Due Date: **  October 18th ** Question: ** An electrochemical cell consists of a copper electrode in an acidic solution of 1.00 molar Cu + connected by a salt bridge to a second compartment with a iron electrode in an acidic solution of 1.00 molar Fe 2+.

(a) Write the equation for the half-cell reaction occurring at each electrode. Indicate which half-reaction occurs at the anode. Cu+ //(aq)// + 1e- →  Cu//(s)// = 0.521 V  Fe2+//(aq)// + 2e- →  Fe//(s)// = -0.440 V (reverse)

Since iron has a negative reduction potential (V), it is identified as the anode.

(b) Write the balanced chemical equation for the overall spontaneous cell reaction that occurs when the circuit is complete. Calculate the standard voltage, //E//°, for this cell reaction.

2Cu+ (aq) + 2e- →  2Cu(s) Fe(s) →  Fe2+(aq) + 2e- Cancel out the e- and add the half cells 2Cu+ //(aq)// + Fe//(s)// <span style="color: #1f497d; font-family: 宋体; font-size: 12pt; line-height: 115%;">→ <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> Fe2+//(aq)// + 2Cu//(s)// Form the balanced chemical equation //<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt;">Ecell = Ered - Eox // //<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt;">Ecell = //<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt;"> 0.521 V - (-0.440) = 0.961 V

(c) Calculate the equilibrium constant for this cell reaction at 298 K. <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">Use the Gibbs free energy function to solve for <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">equilibrium constant <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">ΔG = change in Gibbs free energy <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">n = number of electrons per mole product = 2 <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">F = Faraday constant <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">= 96485 J/mol*V <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">//<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">E° // <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;"> = //Ecell =// .961 V <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">ΔG° = -nFE° <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">= -185444.17 J/mol
 * -2 |||| 96485 J || 0.961 V ||
 * |||| mol*V || ﻿ ||

<span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%;">ΔG° = - RT lnK <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%;">K = equilibrium constant <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%;">R = gas constant = 8.314 J/mol*K <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%;">T = temperature = 298 K <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%;">lnK = ΔG°/(-RT) <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%;">= 74.8
 * -185444.17 J |||| mol*K ||  ||
 * mol |||| -8.314 J || 298K ||

<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">lnK = 74.8 <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">K = e^74.8 <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">K = 3.06 * 10^32

(d) A cell similar to the one described above is constructed with solutions that have initial concentrations of 0.10 molar Cu + and 0.0200 molar Fe 2+. Calculate the initial voltage, //E//°, of this cell. <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%; margin: 0in 0in 0pt;">Use Nernst equation to solve for initial voltage E = E° - (RT/nF)(lnQ) <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%; margin: 0in 0in 0pt;">R = gas constant = 8.314 J/mol*K, T = temperature = 298 K <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%; margin: 0in 0in 0pt;">n = number of electrons per mole product = 2, F = Faraday constant = 96485 J/mol*V, E° = Ecell = .961 V <span style="color: #123b6d; font-family: 'Times New Roman',Times,serif; font-size: 120%; margin: 0in 0in 0pt;">Q = reaction quotient = product^(coeffient of the product)/ reactant^(coeffient of the reactant) = [Fe2+]^1/[Cu+]^2

<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">RT/(nF) = <span style="color: #123b6d; font-family: 'Times New Roman','serif'; font-size: 120%; line-height: 115%; margin: 0in 0in 0pt;">= 1.28 * 10^-2 V
 * 8.314 J || 298 K || mol*V ||~  ||
 * mol*K || 2 || 96485 J ||~  ||
 * ~ Q = || [Fe2+]^1 ||~ = || [. 0200 molar] ||~ =2 ||
 * ~  || [Cu+]^2 ||~   || [.10 molar]^2 ||~   ||

<span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">E = 0.961V - 0.0128V * lnQ <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">E = 0.961V - 0.0128V * ln2 <span style="color: #1f497d; font-family: 'Times New Roman','serif'; font-size: 12pt; line-height: 115%; margin: 0in 0in 10pt;">E = 0.952V

From Splung.Com // __History__- In it's most simple form a battery can be regarded as a pump that provides the energy to move charge around a circuit. In order to provide a [|potential difference], or [|electro-motive force (EMF)] a store of energy is required. One such method is a battery or cell. The common usages of the term battery is any device that converts chemical energy into electrical energy. However, strictly speaking, the term battery is used when several electrical cells are connected together to provide a source of a potential difference in a circuit. If it is just a single chemical source then it is called a cell. //

// In 1791, Galvani noticed that a circuit created with two different metals, when touched on the ends of the leg of a dead frog, would cause it twitch. The two metals were creating an electric current within the frog's leg, causing the muscles to contract. Early batteries were an improvement of this method transfering chemical energy into electrical energy. // // The first battery was invented in 1793 by Alessandro Volta. Just as the two different metals touching the wet skin of a frog's leg, caused an electrical current to flow, early batteries increased the voltage that could be produced by stacking a pile of discs made from silver and zinc sandwiched between paper soaked in a salt water solution as shown in Figure 2. In honour of Volta, we use the Volt as the unit of potential difference and EMF (1) //

__Copper electrode__- The Copper electrode is used for measuring [|electrochemical potential] and is the most commonly used reference electrode for testing [|cathodic protection] [|corrosion] control systems. **Cathodic protection** (**CP**) is a technique used to control the [|corrosion] of a metal surface by making it the [|cathode] of an [|electrochemical cell][|[2]]. The simplest method to apply CP is by connecting the metal to be protected with another more easily corroded metal to act as the anode of the [|electrochemical cell]. Cathodic protection systems are most commonly used to protect [|steel], water or fuel [|pipelines] and [|storage tanks] , steel pier [|piles] , water-based vessels including yachts and powerboats[|[3]], offshore [|oil platforms] and onshore [|oil well] casings. Cathodic protection can, in some cases, prevent [|stress corrosion cracking]

From Sciencedirct.com // __Iron electrode__- The Iron electrode is used for potentiometric detection of ascorbic acid. A chemically modified electrode constructed by incorporating iron (II) phthalocyanine [Fe(II)Pc] into carbon-paste matrix was used as a sensitive potentiometric sensor for detection of ascorbic acid. The resulting electrode exhibits catalytic properties for the electro oxidation of ascorbic acid, and lowers the over potential for the oxidation of this compound. The faster rate of electron transfer results in a near-Nernstian behavior of the modified electrode, and makes it a suitable potentiometric sensor for detection of ascorbic acid. A linear response in concentration range from 10−6 to 10−2 M (0.18–1800 μg ml−1) was obtained with a detection limit of 5 × 10−7 M for the potentiometric detection of ascorbic acid. The modified electrode was used for the determination of ascorbic acid in vitamin preparations. The recovery was 97.2–102.4% for the vitamin added to the preparations with a relative standard deviation of less than 5%. The modified electrode exhibited a fast response time (<10 s),had good stability, and had an extended lifetime. (4) //

__Citations__ 1. [] 2. A.W. Peabody, Peabody's Control of Pipeline Corrosion, 2nd Ed., 2001, NACE International. p.6, [|ISBN 1575900920] 3. [] 4 .http://www.sciencedirect.com/science?_ob=ArticleURL&_udi=B6W9V-45BC70N-B2&_user=10&_coverDate=03%2F15%2F2001&_rdoc=1&_fmt=high&_orig=search&_origin=search&_sort=d&_docanchor=&view=c&_acct=C000050221&_version=1&_urlVersion=0&_userid=10&md5=175d0fab3c6390eac36bba6bfff698f8&searchtype=a