chemwiki book 30 dec 2015 3

Chemwiki• Voltaic Cells

• Case Study: Battery Types

• Batteries: Electricity though chemical reactions

• Rechargeable Batteries

• Discharging Batteries

• Electrochemical Cells under Nonstandard Conditions

• Concentration Cell

• Electrochemical Cell Conventions

• The Cell Potential

• Writing Equations for Redox Reactions

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Voltaic Cells

Voltaic Cells

In redox reactions, electrons are transferred from one species to another. If the reaction is spontaneous, energy is released,which can then be used to do useful work. To harness this energy, the reaction must be split into two separate half reactions: theoxidation and reduction reactions. The reactions are put into two different containers and a wire is used to drive the electronsfrom one side to the other. In doing so, a Voltaic/ Galvanic Cell is created.

Introduction

When a redox reaction takes place, electrons are transferred from one species to the other. If the reaction is spontaneous, energyis released, which can be used to do work. Consider the reaction of a solid copper (Cu(s)) in a silver nitrate solution (AgNO3(s)).

\[2Ag^+_{(aq)} + Cu_{(s)} \leftrightharpoons Cu^{2+}_{(aq)} + 2Ag_{(s)}\]

The \(AgNO_{3\;(s)}\) dissociates in water to produce \(Ag^+_{(aq)}\) ions and \(NO^-_{3\;(aq)}\) ions. The NO3-(aq) ions

can be ignored since they are spectator ions and do not participate in the reaction. In this reaction, a copper electrode is placed

into a solution containing silver ions. The Ag+(aq) will readily oxidize Cu(s) resulting in Cu2+

(aq), while reducing itself to Ag(s).

This reaction releases energy. When the copper electrode solid is placed directly into a silver nitrate solution, however, theenergy is lost as heat and cannot be used to do work. In order to harness this energy and use it do useful work, we must split thereaction into two separate half reactions; The oxidation and reduction reactions. A wire connects the two reactions and allowselectrons to flow from one side to the other. In doing so, we have created a Voltaic/ Galvanic Cell.

Voltaic Cells

A Voltaic Cell (also known as a Galvanic Cell) is an electrochemical cell that uses spontaneous redox reactions to generateelectricity. It consists of two separate half-cells. A half-cell is composed of an electrode (a strip of metal, M) within a solution

containing Mn+ ions in which M is any arbitrary metal. The two half cells are linked together by a wire running from oneelectrode to the other. A salt bridge also connects to the half cells. The functions of these parts are discussed below.






http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Redox_Chemistry

Figure 1 Voltaic Cell

Half Cells

Half of the redox reaction occurs at each half cell. Therefore, we can say that in each half-cell a half-reaction is taking place.When the two halves are linked together with a wire and a salt bridge, an electrochemical cell is created.

Electrodes

An electrode is strip of metal on which the reaction takes place. In a voltaic cell, the oxidation and reduction of metals occursat the electrodes. There are two electrodes in a voltaic cell, one in each half-cell. The cathode is where reduction takes placeand oxidation takes place at the anode.

Through electrochemistry, these reactions are reacting upon metal surfaces, or electrodes. An oxidation-reduction equilibriumis established between the metal and the substances in solution. When electrodes are immersed in a solution containing ionsof the same metal, it is called a half-cell. Electrolytes are ions in solution, usually fluid, that conducts electricity through ionicconduction. Two possible interactions can occur between the metal atoms on the electrode and the ion solutions.

1. Metal ion Mn+ from the solution may collide with the electrode, gaining "n" electrons from it, and convert to metalatoms. This means that the ions are reduced.

2. Metal atom on the surface may lose "n" electrons to the electrode and enter the solution as the ion Mn+ meaningthat the metal atoms are oxidized.

When an electrode is oxidized in a solution, it is called an anode and when an electrode is reduced in solution. it is calleda cathode.

• Anode: The anode is where the oxidation reaction takes place. In other words, this is where the metal loses electrons. Inthe reaction above, the anode is the Cu(s) since it increases in oxidation state from 0 to +2.


http://chemwiki.ucdavis.edu/@api/deki/files/159/voltaic_cell_diagram.png

• Cathode: The cathode is where the reduction reaction takes place. This is where the metal electrode gains electrons.Referring back to the equation above, the cathode is the Ag(s) as it decreases in oxidation state from +1 to 0.

Remembering Oxidation and Reduction

When it comes to redox reactions, it is important to understand what it means for a metal to be “oxidized” or “reduced”. Aneasy way to do this is to remember the phrase “OIL RIG”.

OIL = Oxidization is Loss (of e-)

RIG = Reduction is Gain (of e-)

In the case of the example above \(Ag^+_{(aq)}\) gains an electron meaning it is reduced. \(Cu_{(s)}\) loses two electrons thusit is oxidized.

Salt Bridge

The salt bridge is a vital component of any voltaic cell. It is a tube filled with an electrolyte solution such as KNO3(s) or KCl(s).The purpose of the salt bridge is to keep the solutions electrically neutral and allow the free flow of ions from one cell toanother. Without the salt bridge, positive and negative charges will build up around the electrodes causing the reaction to stop.

Flow of Electrons

Electrons always flow from the anode to the cathode or from the oxidation half cell to the reduction half cell. In terms of Eocellof the half reactions, the electrons will flow from the more negative half reaction to the more positive half reaction.

Cell Diagram

A cell diagram is a representation of an electrochemical cell. The figure below illustrates a cell diagram for the voltaic shownin Figure 1 above.


Figure 2 Cell Diagram. The figure below illustrates a cell diagram for the voltaic shown in Figure 1.

When drawing a cell diagram, we follow the following conventions. The anode is always placed on the left side, and the cathodeis placed on the right side. The salt bridge is represented by double vertical lines (||). The difference in the phase of an elementis represented by a single vertical line (|), while changes in oxidation states are represented by commas (,).

Constructing a Cell Diagram

When asked to construct a cell diagram follow these simple instructions. Consider the following reaction:

\[2Ag^+_{(aq)} + Cu_{(s)} \rightleftharpoons Cu^{2+}_{(aq)} + 2Ag_{(s)}\]

Step 1: Write the two half-reactions.

\[Ag^+_{(aq)} + e^- \rightleftharpoons Ag_{(s)}\]

\[Cu_{(s)} \rightleftharpoons Cu^{2+}_{(aq)} + 2e^-\]

Step 2: Identify the cathode and anode.

\(Cu_{(s)}\) is losing electrons thus being oxidized; oxidation occurs at the anode.

• Anode (where oxidation occurs): \(Cu_{(s)} \rightleftharpoons Cu^{2+}_{(aq)} + 2e^-\)

\(Ag^+\) is gaining electrons thus is being reduced; reduction happens at the cathode.

• Cathode (where reduction occurs): \(Ag^+_{(aq)} + e^- \rightleftharpoons Ag_{(s)}\)

Step 3: Construct the Cell Diagram.

\[Cu_{(s)} | Cu^{2+}_{(aq)} || Ag^+_{(aq)} | Ag_{(s)}\]


The anode always goes on the left and cathode on the right. Separate changes in phase by | and indicate the the saltbridge with ||. The lack of concentrations indicates solutions are under standard conditions (i.e., 1 M)

Cell Voltage/Cell Potential

The readings from the voltmeter give the reaction's cell voltage or potential difference between it's two two half-cells. Cellvoltage is also known as cell potential or electromotive force (emf) and it is shown as the symbol Ecell.

Standard Cell Potential: Eocell = Eo

right(cathode) - Eoleft(anode)

The Eo values are tabulated with all solutes at 1 M and all gases at 1 atm. These values are called standard reductionpotentials. Each half-reaction has a different reduction potential, the difference of two reduction potentials gives the voltage

of the electrochemical cell. If Eocell is positive the reaction is spontaneous and it is a voltaic cell. If the Eocell is negative, thereaction is non-spontaneous and it is referred to as an electrolytic cell.

Practice Problems

Consider the following two reactions:

\[Cu^{2+}_{(aq)} + Ba_{(s)} \rightarrow Cu_{(s)} + Ba^{2+}_{(aq)}\]

\[Al_{(s)} + Sn^{2+}_{(aq)} \rightarrow Al^{3+}_{(aq)} + Sn_{(s)}\]

1. Split the reaction into half reactions and determine their standard reduction potentials. Indicate which would be theanode and cathode.

2. Construct a cell diagram for the following each reactions.3. Determine the \(E^o_{cell}\) for the voltaic cell formed by each reaction.

*Solution are given below.

Solutions

1.a) Ba2+(aq) + 2e- → Ba(s) Eo = -2.92 V Anode

Cu2+(aq) + 2e- → Cu(s) Eo = +0.340 V Cathode

1.b) Al3+(aq) 3e-→ Al(s) Eo = -1.66 V Anode

Sn2+(aq) → Sn(s) +2e- Eo = -0.137 V Cathode

2.a) Ba2+(aq) | Ba(s) || Cu(s) | Cu2+(aq)


2.b) Al(s) | Al3+(aq) || Sn2+

(aq) | Sn(s)

3.a) Eocell = 0.34 - (-2.92) = 3.26 V

3.b) Eocell = -0.137 - (-1.66) = 1.523 V

Links• http://www.youtube.com/watch?v=nNG5PMlHSoA&feature=related

• http://www.youtube.com/watch?v=L1NU7Hzp_ZQ&feature=related

• http://www2.ohlone.edu/people/jklent/labs/101B_labs/Voltaic%20Cells.pdf

• http://butane.chem.uiuc.edu/cyerkes/Chem104ACSP08/Genchemref/standpot.html This is a link that shows the standardreduction potentials of all common half-reactions:

References1. Brady, James E., Holum, John R. “Chemistry: The Study of Matter and Its Changes”, John Wiley & Sons Inc 19932. Brown, Theodore L., LeMay, H. Eugene Jr. “Chemistry: The Central Science” Third Edition, Prentice-Hall, Inc.

Englewood Cliffs, N.J. 07632 19853. Brown, Theodore L., LeMay, H. Eugene Jr., Bursten, Bruce E. “Chemistry: The Central Science” Fifth Edition,

Prentice-Hall, Inc. Englewood Cliffs, N.J. 07632 19914. Gesser, Hyman D. “ Descriptive Principles of Chemistry”, C.V. Mosby Company 19745. Harwood, William, Herring, Geoffrey, Madura, Jeffry, and Petrucci, Ralph, General Chemistry: Principles and

Modern Applications, Ninth Edition, Upper Saddle River,New Jersey, Pearson Prentice Hall, 2007.6. Petrucci, Ralph H. Genereal Chemistry: Principles and Modern Applications 9th Ed. New Jersey: Pearson

Education Inc. 2007.7. Vassos Basil H. Electroanaytical Chemistry. New York: Wiley-Interscience Publication. 1983.8. Zumdahl, Steven S. Chemistry 7th Ed. New York: Houghton Mifflin Company. 2007.

Contributors• Shamsher Singh, Deborah Gho


http://www.youtube.com/watch?v=nNG5PMlHSoA&feature=related

http://www.youtube.com/watch?v=L1NU7Hzp_ZQ&feature=related

http://www2.ohlone.edu/people/jklent/labs/101B_labs/Voltaic%20Cells.pdf

http://butane.chem.uiuc.edu/cyerkes/Chem104ACSP08/Genchemref/standpot.html

Case Study: Battery Types

Case Study: Battery Types

Ranging from the very crude to the highly sophisticated, batteries come in a plethora of variety. Batteries in short areelectrochemical cells that produce a current of electricity via chemical reactions. More specifically, batteries produce electricalenergy from oxidation-reduction reactions. A collection of electrochemical cells wired in series is properly called a battery. Aflashlight battery is really a single electrochemical cell, while a car battery is really a battery since it is three electrochemicalcells in series.

Introduction

Electrochemical cells have been in use longer than what was once thought. Discovered in Khujut Rabu of modern day Iraq anddating from the Parthian (250 B.C.-A.D. 224) and Sassanid (A.D. 224-600)

periods, the Baghdad Battery is the first known battery in the world. Consisting of acopper cylinder, an iron rod, an asphalt stopper, and a small earthenware jar, theBaghdad Battery was filled with an unknown electrolytic solution and may have beenused for electroplating.

About 2000 years later the Voltaic Pile, a stack of individual cells of zinc and copperdisks immersed in sulfuric acid, was created by the Italian Count Volta and effectivelyreplaced the use of the Leyden Jar, an instrument that stored static electricity for futureuse. Volta's battery is considered the first electrochemical cell and the reaction forwhich is as follows:

oxidation half-reaction: \(Zn \rightarrow Zn^{2+} + 2e^-\)

reduction half-reaction: \(2H^+ + 2e^- \rightarrow H_2\)

Because zinc is higher in the electrochemical series, the zinc anode reacts with sulfate anions and is oxidized whilst protons arereduced to hydrogen gas. The copper cathode remains unchanged and acts only as electrode for the chemical reaction. Becausethe Voltaic pile was unafe to use and the cell power diminished over time, it was abandoned.






http://chemwiki.ucdavis.edu/@api/deki/files/160/baghdad_battery.png

http://chemwiki.ucdavis.edu/@api/deki/files/160/baghdad_battery.png

Electrochemical cells typically consist of an anode (the negative electrode where oxidation occurs), a cathode (the positiveelectrode where reduction occurs), and an electrolyte (the medium conducting anions and cations within a reaction) allcontained within a cell. Electrons flow in a closed circuit from the anode to the cathode. Depending on the configuration of thecell and the electrolyte used, a salt bridge may be necessary to conduct ions from one half cell to another as an electric charge iscreated when electrons move from electrode to another. The difference created would keep electrons from flowing any further.Because a salt bridge permits the flux of ions, a balance in charge is kept between the half cells whilst keeping them separate.

Types of Batteries

The two main categories of batteries are primary and secondary. Essentially, primary cells are batteries which cannot berecharged while secondary cells are rechargeable. The distinction begs the question as to why primary cells are still in use today,and the reason being is that primary cells have lower self-discharge rates meaning that they can be stored for longer periodsof time than rechargeable batteries and maintain nearly the same capacity as before. Reserve and backup batteries present aunique example of this advantage of primary cells. In reserve, or stand-by, batteries components of the battery containing activechemicals are separated until the battery is needed, thus greatly decreasing self-discharge. An excellent example is the Water-Activated Battery. As opposed to inert reserve batteries, backup batteries are already activated and functional but not producingany current until the main power supply fails.

Biobatteries

Devices that generate electric energy via the digestion of carbohydrates, fats, and protiens by enzymes. The most commonbiobatteries are the lemon or potato battery and the frog or ox-head battery better described as a "muscular pile". In a lemoncell, the energy for the battery is not produced by the lemon but by the metal electrodes. Usually zinc and copper electrodes areinserted into a lemon (the electrolyte being citric acid) and connected by a circuit. The zinc is oxidized in the lemon in orderto reach a preferred lower energy level and the electrons discharged provide the energy. Using zinc and copper electrodes, alemon can produce about 0.9 Volts.

While not technically a biobattery, an Earth battery is comprised of two different electrodes which are either buriedunderground or immersed in natural bodies of water which tap into Telluric currents to produce electric energy.

Dry-Cell Batteries

During the 1860s, a French man named George Lelanche developed the Lelanche cell also known today as the dry-cell battery.A dry-cell battery is a battery with a paste electrolyte (as opposed to a wet-cell battery with a liquid electrolyte) in the themiddle of its cylinder and attached are metal electrodes. A dry-cell battery is a primary cell that cannot be reused. In order tofunction, each dry-cell battery has a cathode and an anode. Some examples of dry-cell batteries used in everyday objects todayare remote controls, clocks, and calculators.


Figure: Modern dry cell batteries. Image used with permission from Wikipedia

Types of dry-cell batteries are zinc-carbon batteries, alkaline-cell batteries, and mercury batteries. Before zinc-carbon batterieswere used, mercury batteries were the main resource. It was not until mercury was known to become harmful that zinc-carbonbatteries replaced it. Batteries may produce the following potential problems or hazards:

• Pollute the lakes and streams as the metals vaporize into the air when burned.• Contribute to heavy metals that potentially may leach from solid waste landfills.• Expose the environment and water to lead and acid.• Contain strong corrosive acids.• May cause burns or danger to eyes and skin.

Dry-cell batteries are the most common battery type used today. Essentially, the battery is comprised of a metal electrode (orgraphite rod) surrounded by a moist electrolyte paste that is enclosed in a metal cylinder. 1.5 volts is the most commonly usedvoltage for dry-cell batteries. The sizes of dry-cell batteries vary, however, it does not change the voltage of the battery.

Zinc-carbon cells

Zinc-carbon cells were the first really portable energy source. These cells have a short lifetime and the zinc casings becomeporous as the zinc is converted to zinc chloride. The substances in the cell that leak out are corrosive to metal and can terminallydestroy electronic equipment or flashlights. Zinc-carbon cells produce 1.5 volts.


For a dry-cell battery to operate, oxidation will occur from the zinc anode and reduction will take place in the cathode. Themost common type of cathode is a carbon graphite. Once reactants have been turned into products, the dry-cell battery willwork to produce electricity. For example, in a dry-cell battery, once \(Zn^{2+}\) has been oxidized to react with \(NH_3\), itwill produce chloride salt to insure that too much \(NH_3\) will not block the current of the cathode.

\[Zn^{2+}_{(aq)} + 2NH_{3\;(g)} + 2Cl^-_{(aq)} \rightarrow [Zn(NH_3)_2]Cl_{2\; (s)}\]

How does the reaction work? While the zinc anode is being oxidized, it is producing electrons that will be captured byreducing Maganese from an oxidation state of +4 to a +3.

• Reduction of Maganese: \(2MnO_{2\;(s)} + H_2O_{(l)} + 2e^- \rightarrow Mn_2O_{2\; (s)} + 2OH^-_{ (aq)}\)

The electrons produced by Zinc will then connect to the cathode to produce it's product.

• Oxidization of Zinc: \(Zn_{(s)} \rightarrow Zn^{2+}_{(aq)} +2e^-\)

Alkaline cells

Recently, the most popular dry-cell battery to be used has been the alkaline-cell battery. In the zinc-carbon battery shown above,the zinc is not easily dissolved in basic solutions. Though it is fairly cheap to construct a zinc-carbon battery, the alkaline-cell battery is favored because it can last much longer. Instead of using \(NH_4Cl\) as an electrolyte, the alkaline-cell batterywill use \(NaOH\) or \(KOH\) instead. The reaction will occur the same where zinc is oxidized and it will react with \(OH^-\)instead.

\[Zn^{2+}_{(aq)} + 2OH^-_{(aq)} \rightarrow Zn(OH)_{2\; (s)}\]

Once the chemicals in the dry-cell battery can no longer react together, the dry-cell battery is dead and cannot be recharged.Alkaline electrochemical cells have a much longer lifetime but the zinc case still becomes porous as the cell is discharged andthe substances inside the cell are still corrosive. Alkaline cells produce 1.54 volts.

Mercury cells

Mercury batteries are small, circular metal batteries that were used in watches. Mercury cells offer a long lifetime in a smallsize but the mercury produced as the cell discharges is very toxic. This mercury is released into the atmosphere if the cells are


incinerated in the trash. About 90% of the 1.4 million pounds of mercury in our garbage comes from mercury cells. Mercurycells only produce 1.3 Volts.

\[HgO + Zn + H_2O \rightarrow Hg + Zn(OH)_2\]

Mercury batteries utilize either pure mercuric oxide or a mix of mercuric oxide with manganese dioxide as the cathode. Theanode is made with zinc and is separated from the cathode with a piece of paper or other porous substance that has been soakedin the electrolyte (which is generally either sodium or potassium oxide).

In the past, these batteries were widely used because of their long shelf life of about 10 years, and also because of theirstable, steady voltage output. Also, they had the highest capacity per size. They were popular for use in button-type batteryapplications, such as watches or hearing aids. However, the environmental impact for the amount of mercury present in thebatteries became an issue, and the mercury batteries were discontinued from public sale.

lead-acid batteries

The lead-acid battery used in cars and trucks consists of six electrochemical cells joined in series. Each cell in a lead-acidbattery produces 2 volts. The electrodes are composed of lead and are immersed in sulfuric acid. The negative electrodes arespongy lead metal and the positive electrodes are lead impregnated with lead oxide. As the battery is discharged, metallic leadis oxidized to lead sulfate at the negative electrodes and lead oxide is reduced to lead sulfate at the positive electrodes. When alead-acid battery is recharged by an alternator, electrons are forced to flow in the opposite direction which reverses the reaction.

\[Pb + PbO_2 + 2 H_2SO_4 \rightarrow 2 PbSO_4 + 2 H_2O\]

Nickel-cadmium cells

Nickel-cadmium cells can also be regenerated by reversing the flow of the electrons in a battery charger. The cadmium isoxidized in these cells to cadmium hydroxide and the nickel is reduced. Nickel-cadmium cells generate 1.46 Volts.

\[Cd + NiO_2 + 2 H_2O \rightarrow Cd(OH)_2 + Ni(OH)_2\]

A Nickel-metal Hyride battery is a secondary cell very similar to the nickel-cadmium cell except that it uses a hydrogen-absorbing alloy in place of cadmium. The Nickel-metal Hyride battery has 2-3 times the capacity of a nickel-cadmium cell.


Questions1. Why are lead-acid batteries used in cars?2. What is the specific, more scientific name for a battery?3. how do batteries generally work?4. What kind of reaction occurs at the Cathode?5. What kind of reaction occurs at the anode?6. What type of battery can be recharged? Which cannot?7. If mercury batteries are so long lasting and efficient, then why are they not used anymore?8. What is the purpose of a salt bridge?

Answers1. Lead acid batteries produce two volts each and last a relatively long time. This type of battery is good for cars

because of its ability to be recharged by the alternator. The batteries are connected in a series to produce thenecessary voltage. They are larger than most other batteries, but work well for large energy-consuming machinerysuch as vehicles.

2. Batteries are actually called electrochemical cells.3. Batteries work through a series of oxidation-reduction reactions that produces a waste product and a known

amount of energy.4. Reduction reaction.5. Oxidation reaction.6. Secondary batteries can be recharged. Primary batteries cannot be recharged because the reaction is not

reversible.7. Mercury batteries are often disposed incorrectly, which causes large amounts mercury to seep into the

environment. Most or the mercury present in our environment today is the result of improper mercury batterydisposal.

8. Salt bridges conduct ions from one half cell to another to balance changing charges that could cause a halt to theflow of electrons in an electrochemical cell.

Outside Links• http://en.wikipedia.org/wiki/Rechargeable_battery

• Dry-cell battery picture: http://www.science-projects-resource...es/Drycell.gif

• http://library.kcc.hawaii.edu/extern...y_battery.html

• http://www.wisegeek.com/what-is-a-drycellbattery.htm

• http://en.wikipedia.org/wiki/Mercury_battery

• http://www.ehow.com/how_4865930_build-homemade-battery.html


http://en.wikipedia.org/wiki/Rechargeable_battery

http://www.science-projects-resources.com/images/Drycell.gif

http://library.kcc.hawaii.edu/external/chemistry/everyday_battery.html

http://www.wisegeek.com/what-is-a-drycellbattery.htm

http://en.wikipedia.org/wiki/Mercury_battery

http://www.ehow.com/how_4865930_build-homemade-battery.html

References1. Petrucci, Ralph. General Chemistry: Principles and Modern Applications. 10th ed. Upper Saddle River, NJ:

Pearson Canada, 2011. Print.2. Root, Michael. The TAB Battery Book: An In-depth Guide to Construction, Design, and Use. New York, NY:

McGraw-Hill, 2011. Print.

Contributors• Richard Banks (Boise State University)

• Erica Chen (UCD)


http://chemistry.boisestate.edu/people/richardbanks/

Batteries: Electricity though chemical reactions

Batteries: Electricity though chemical reactions

Batteries consist of one or more electrochemical cells that store chemical energy for later conversion to electrical energy.Batteries are used in many day-to-day devices such as cellular phones, laptop computers, clocks, and cars. Batteries arecomposed of at least one electrochemical cell which is used for the storage and generation of electricity. Though a variety ofelectrochemical cells exist, batteries generally consist of at least one voltaic cell. Voltaic cells are also sometimes referred to asgalvanic cells. Chemical reactions and the generation of electrical energy is spontaneous within a voltaic cell, as opposed to thereactions electrolytic cells and fuel cells.

Introduction

It was while conducting experiments on electricity in 1749 that Benjamin Franklin first coined the term "battery" to describelinked capacitors. However his battery was not the first battery, just the first ever referred to as such. Rather it is believed thatthe Baghdad Batteries, discovered in 1936 and over 2,000 years old, were some of the first ever batteries, though their exactpurpose is still debated.

Luigi Galvani (for whom the galvanic cell is named) first described "animal electricity" in 1780 when he created an electricalcurrent through a frog. Though he was not aware of it at the time, this was a form of a battery. His contemporary AlessandroVolta (for whom the voltaic cell and voltaic pile are named) was convinced that the "animal electricity" was not coming fromthe frog, but something else entirely. In 1800, his produced the first real battery: the voltaic pile.

In 1836, John Frederic Daniell created the Daniell cell when researching ways to overcome some of the problems associatedwith Volta's voltaic pile. This discovery was followed by developments of the Grove cell by William Robert Grove in 1844;the first rechargeable battery, made of a lead-acid cell in 1859 by Gaston Plante; the gravity cell by Callaud in the 1860s; andthe Leclanche cell by Georges Leclanche in 1866.

Until this point, all batteries were wet cells. Then in 1887 Carl Gassner created the first dry cell battery, made of a zinc-carboncell. The nickel-cadmium battery was introduced in 1899 by Waldmar Jungner along with the nickel-iron battery. HoweverJungner failed to patent the nickel-iron battery and in 1903, Thomas Edison patented a slightly modified design for himself.






http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrolytic_Cells

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Case_Studies/Case_Study%3A_Fuel_Cells

A major breakthrough came in 1955 when Lewis Urry, an employee of what is now know as Energizer, introduced the commonalkaline battery. The 1970s led to the nickel hydrogen battery and the 1980s to the nickel metal-hydride battery.

Lithium batteries were first created as early as 1912, however the most successful type, the lithium ion polymer battery used inmost portable electronics today, was not released until 1996.

Voltaic Cells

Voltaic cells are composed of two half-cell reactions (oxidation-reduction) linked together via a semipermeable membrane(generally a salt bath) and a wire (Figure 1). Each side of the cell contains a metal that acts as an electrode. One of the electrodesis termed the cathode, and the other is termed the anode. The side of the cell containing the cathode is reduced, meaning itgains electrons and acts as the oxidizing agent for the anode. The side of the cell containing the anode is where oxidationoccurs, meaning it loses electrons and acts as the reducing agent for the cathode. The two electrodes are each submerged in anelectrolyte, a compound that consists of ions. This electrolyte acts as a concentration gradient for both sides of the half reaction,facilitating the process of the electron transfer through the wire. This movement of electrons is what produces energy and isused to power the battery.

The cell is separated into two compartments because the chemical reaction is spontaneous. If the reaction was to occur withoutthis separation, energy in the form of heat would be released and the battery would not be effective.

Figure 1: A Zinc-Copper Voltaic cell

The voltaic cell is providing the electricity needed to power the light-bulb.


http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Redox_Chemistry/Oxidation-Reduction_Reactions

http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Electrodes/Standard_Hydrogen_Electrode

http://chemwiki.ucdavis.edu/Physical_Chemistry/Physical_Properties_of_Matter/Solutions/Solution_Basics/Electrolyte_Solutions

Types of Batteries

Figure 2: Primary versus Secondary Batteries

Primary batteries (left) are non-rechargeable and disposable. Secondary batteries (right) are rechargeable, like this cellularphone battery.

Primary Batteries

Primary batteries are non-rechargeable and disposable. The electrochemical reactions in these batteries are non-reversible. Thematerials in the electrodes are completely utilized and therefore cannot regenerate electricity. Primary batteries are often usedwhen long periods of storage are required, as they have a much lower discharge rate than secondary batteries.

Use of primary batteries is exemplified by smoke detectors, flashlights, and most remote controls.

Secondary Batteries

Secondary batteries are rechargeable. These batteries undergo electrochemical reactions that can be readily reversed. Thechemical reactions that occur in secondary batteries are reversible because the components that react are not completely usedup. Rechargeable batteries need an external electrical source to recharge them after they have expended their energy.

Use of secondary batteries is exemplified by car batteries and portable electronic devices.

Battery Cell Types

Wet Cells

Wet cell batteries contain a liquid electrolyte. They can be either primary or secondary batteries. Due to the liquid nature of wetcells, insulator sheets are used to separate the anode and the cathode. Types of wet cells include Daniell cells, Leclanche cells(originally used in dry cells), Bunsen cells, Weston cells, Chromic acid cells, and Grove cells. The lead-acid cells in automobilebatteries are wet cells.


Figure 3: A lead-acid battery in an automobile.

Dry Cells

In dry cell batteries, no free liquid is present. Instead the electrolyte is a paste, just moist enough to allow current flow. Thisallows the dry cell battery to be operated in any position without worrying about spilling its contents. This is why dry cellbatteries are commonly used in products which are frequently moved around and inverted, such as portable electronic devices.Dry cell batteries can be either primary or secondary batteries. The most common dry cell battery is the Leclanche cell.

Battery Performance

The capacity of a battery depends directly on the quantity of electrode and electrolyte material inside the cell.

Primary batteries can lose around 8% to 20% of their charge over the course of a year without any use. This is caused by sidechemical reactions that do not produce current. The rate of side reactions can be slowed by lowering temperature. Warmertemperatures can also lower the performance of the battery, by speeding up the side chemical reactions. Primary batteriesbecome polarized with use. This is when hydrogen accumulates at the cathode, reducing the battery's effectiveness.Depolarizers can be used to remove this build up of hydrogen.

Secondary batteries self-discharge even more rapidly. They usually lose about 10% of their charge each month. Rechargeablebatteries gradually lose capacity after every recharge cycle due to deterioration. This is caused by active materials falling offthe electrodes or electrolytes moving away from the electrodes.

Peukert's law can be used to approximate relationships between current, capacity, and discharge time. This is represented bythe equation


, where I is the current, k is a constant of about 1.3, t is the time the battery can sustain the current, and Qp is the capacity whendischarged at a rate of 1 amp.

Current, Voltage, and Standard Reduction Potential

There is a significant correlation between a cell's current and voltage. Current, as the name implies, is the flow of electricalcharge. Voltage is how much current can potentially flow through the system. Figure 4 illustrates the difference between currentand voltage.

Figure 4: The difference between voltage and current.

Water is flowing out of a hose and onto a waterwheel, turning it. Current can be thought of as the amount of water flowingthrough the hose. Voltage can be thought of as the pressure or strength of water flowing through the hose. The first hosedoes not have much water flowing through it and also lacks pressure, and is consequently unable to turn the waterwheel veryeffectively. The second hose has a significant amount of water flowing through it, so it has a large amount of current. The thirdhose does not have as much water flowing through it, but does have something blocking much of the hose. This increases thepressure of the water flowing out of the hose, giving it a large voltage and allowing the water to hit the waterwheel with moreforce than the first hose.

Standard reduction potential, Eo, is a measurement of voltage. Standard reduction potential can be calculated with the

knowledge that it is the difference in energy potentials between the cathode and the anode: Eocell = Eo

cathode − Eoanode. For

standard conditions, the electrode potentials for the half cells can be determined by using a table of standard electrode potentials.

For nonstandard conditions, determining the electrode potential for the cathode and the anode is not as simple as looking at a

table. Instead, the Nernst equation must be used in to determine Eo for each half cell. The Nernst equation is represented by

, where R is the universal gas constant (8.314 J K-1 mol-1), T is the temperature in Kelvin, n is the number of moles of electrons

transferred in the half reaction, F is the Faraday constant (9.648 x 104 C mol-1), and Q is the reaction quotient.


http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Nernst_Equation

Different Sizes of Batteries and Some Additional Facts

Batteries vary both in size and voltage due to the chemical properties and contents within the cell. However, batteries of

different sizes may have the same voltage. The reason for this phenomenon is that the standard cell potential does not depend

on the size of a battery but rather on its internal content. Therefore, batteries of different sizes can have the same voltage (Figure

5). Additionally, there are ways in which batteries can amplify their voltages and current. When batteries are lined up in a series

of rows it increases their voltage, and when batteries are lined up in a series of columns it can increases their current.

Figure 5: Four batteries of different sizes all of 1.5 voltage

Hazards

Batteries can explode through misuse or malfunction. By attempting to overcharge a rechargeable battery or charging it at anexcessive rate, gases can build up in the battery and potentially cause a rupture. A short circuit can also lead to an explosion.A battery placed in a fire can also lead to an explosion as steam builds up inside the battery. Leakage is also a concern,because chemicals inside batteries can be dangerous and damaging. Leakage emitted from the batteries can ruin the device theyare housed in, and is dangerous to handle. There are numerous environmental concerns with the widespread use of batteries.The production of batteries consumes many resources and involves the handling of many dangerous chemicals. Used batteriesare often improperly disposed of and contribute to electronic waste. The materials inside batteries can potentially be toxicpollutants, making improper disposal especially dangerous. Through electronic recycling programs, toxic metals such as leadand mercury are kept from entering and harming the environment. Consumption of batteries is harmful and can lead to death.

Homemade Batteries

Any liquid or moist object that has enough ions to be electrically conductive can be used to make a battery. It is even possible togenerate small amounts of electricity by inserting electrodes of different metals into potatoes, lemons, bananas, or carbonated


cola. A voltaic pile can be created using two coins and a paper dipped in salt water. Stacking multiple coins in a series canresults in an increase in current.

Practice Problems

Problems

1. Yes/No

1. Will adding batteries that are lined up in a row amplify the overall voltage of the batteries?2. Do electrolytic cells undergo non-spontaneous chemical reactions?3. Are rechargeable batteries also known as disposable batteries?4. Can batteries of different sizes have the same voltage?

2. T/F

1. In primary cells all of the components in the electrodes are almost always completely used.2. Primary and secondary cells differ in their cathode and anode properties.3. Redox reactions play a critical role in the cells within batteries.4. The cathode in a voltaic cell gains electrons.

3. Determine the standard electrode potential of a voltaic cell within a Leclanche (Dry) cell with half cell voltages of .875V atthe graphite cathode and .253V at the zinc anode.

4. Determine the standard electrode potential with given half cell voltages of .987V at the cathode and .632V at the anode.

5. Explain why rechargeable batteries might be advantageous over disposable batteries.

Solutions

1. Yes/No

1. Yes2. Yes3. No4. Yes

2. T/F

1. True2. False3. True4. True


3. E0cell=E0(cathode)-E0(anode)

E0celll= 0.875V - 0.253V = 0.622v

4. E0cell=E0(cathode)-E0(anode)

E0cell= 0.987V -0.632V = 0.355V

5. Even though disposable batteries are cheaper initially and easier to make, the longer lifespan of rechargeable batteries isoften more efficient and useful. Rechargeable batteries mean less waste, as less batteries need to be made and less are disposedof in land-fills or through recycling programs. Rechargeable batteries are also more convenient as changing batteries is nolonger required. This is especially beneficial in portable electronic devices. Also, because the components in a secondary cellare reusable, rechargeable batteries will generally cost less than disposable batteries in the long run.

References1. Harwood, William, Herring, Geoffrey, Madura, Jeffry, and Petrucci, Ralph. General Chemistry: Principles and

Modern Applications. Ninth Edition. Upper Saddle River, New Jersey: Pearson Prentice Hall, 2007.2. Kiehne, H.A. Battery Technology Handbook. Second Edition. Renningen-Malsheim, Germany: Expert Verlag,

2003.

Outside Links• Batteries: http://en.wikipedia.org/wiki/Battery_%28electricity%29

• History of the battery: http://en.wikipedia.org/wiki/History_of_the_battery

• Video on battery recycling: http://www.youtube.com/watch?v=Pd-RhoTogHA

• How to make a homemade battery: http://www.youtube.com/watch?v=ax3iMxqu3ks&feature=channel

Contributors• Abheetinder Brar (UCD)


http://en.wikipedia.org/wiki/Battery_(electricity)

http://en.wikipedia.org/wiki/History_of_the_battery

http://www.youtube.com/watch?v=Pd-RhoTogHA

http://www.youtube.com/watch?v=ax3iMxqu3ks&feature=channel

Rechargeable Batteries

Rechargeable Batteries

Rechargeable batteries (also known as secondary cells) are batteries that potentially consist of reversible cell reactions thatallow them to recharge, or regain their cell potential, through the work done by passing currents of electricity. As opposed toprimary cells (not reversible), rechargeable batteries can charge and discharge numerous times.

Introduction

Secondary cells encompass the same mechanism as the primary cells with the only difference being that the Redox reaction ofthe secondary cell could be reversed with sufficient amount of energy placed into the equation. The figure below illustrates themechanism of a charging secondary cell. The Charger shown on the top of the diagram is pulling the negative charges towardthe right side of the separator. This makes it seem like the positive charges are compiling on the other side of the cell which isnot allowed to pass the separator. This disequilibrium is the representation of the cell potential that, when allowed, could onceagain approach equilibrium through the transferring of the electrons.






Different secondary batteries provide various functions. For long-term use (followed by discharging and charging), longstorage time when not in use, remote activation, and use under harsh weather conditions are just a few obstacles of creatingsuch secondary cells. Unfortunately, there are no batteries that are capable of encompassing all functions mentioned above.Therefore, the user must decide which application is the most important for a specific task in order to determine the mostcompatible version of rechargeable batteries.

Lead-Acid Batteries

Lead-acid batteries are one of the most common secondary batteries, used primarily for storing large cell potential. Theseare commonly found in automobile engines. Its advantages include low cost, high voltage and large storage of cell potential;and disadvantages include heavy mass, incompetence under low-temperatures, and inability to maintain its potential for longperiods of time through disuse. The reactions of a lead-acid battery are shown below:

Reduction: PbO2(s) + 3 H+(aq) + HSO4-(aq) + 2e- → PbSO4(s) + 2H2O(l)

Oxidation: Pb(s) + HSO4-(aq) → PbSO4(s) + H+(aq) + 2e-

Overall: PbO2(s) + Pb(s) + 2H+(aq) + 2HSO4-(aq) → 2PbSO4(s) + 2H2O (l)

Discharging occurs when the engine is started and where the cell potential equals 2.02V. Charging occurs when the car isin motion and where the electrode potential equals -2.02V, a non- spontaneous reaction which requires an external electricalsource. The reverse reaction takes place during charging.

Nickel-Cadmium Battery

The nickel-cadmium (NiCd) battery is another common secondary battery that is suited for low-temperature conditions witha long shelf life. However, the nickel-cadmium batteries are more expensive and their capacity in terms of watt-hours perkilogram is less than that of the nickel-zinc rechargeable batteries.

Reduction: \(NiO_2\,(s) + H_2O\,(l)+ 2e^- \rightarrow Ni(OH)_2\,(s) + OH^-\,(l)\)

Oxidation: \(Cd\,(s) + 2OH^-\,(aq) \rightarrow Cd(OH)_2\,(s)+ 2e^-\)

Overall: \(Cd\,(s) + NiO_2\,(s) + 2H_2O\,(l) \rightarrow Ni(OH)_2\,(s) + Cd(OH)_2\,(s)\)

Advantages of the nickel-zinc battery are its long life span, high voltage, and the sufficient energy to mass to volume ratio.These characteristics make the nickel-zinc battery more attractive than some. However, it is not yet made in sealed form.


Silver-zinc batteries

Silver-zinc batteries, a less commonly used rechargeable battery, is capable of providing high currents, high voltage, and isequivalent in watt-hour capacity to six lead-acid batteries. These are commonly seen as the little silver buttons in hearingaids, tiny flash lights and so on. Because of its high energy density, silver-zinc batteries are used in missiles and torpedoes,electronics, satellites, and compact portable devices. Although it can provide high energy with a rather small mass, can be usedin a low-temperature condition, and encompasses a sufficient shelf life, it is expensive and has a shorter use time compared toother secondary batteries. In most cases, the silver-zinc battery is used when space and weight are the most important.

Overall, the silver-cadmium battery is high energy, smaller in size and weight, and resistance to overcharge. But its bigdisadvantage is its high cost. Silver-cadmium batteries are often found in satellites where space and weight are importantfactors.

Outside links• Scientific American - How Rechargeable Batteries Work

References1. Petrucci, Ralph H. General Chemistry: Principles & Modern Applcations 9th Ed. New Jersey: Pearson Education,

Inc.. 2007.2. Crompton, T.R. The Battery Reference Book. Boston. Butterworth-Heinemann, 1995.

Contributors• Ling Xie (UCD)


http://www.sciam.com/article.cfm?id=how-do-rechargeable-that

Discharging Batteries

Discharging Batteries






Electrochemical Cells under Nonstandard Conditions

Electrochemical Cells under Nonstandard ConditionsThis page was not added to the PDF due to the following tag(s): article:topic-guide






Concentration Cell

Concentration CellThis page was not added to the PDF due to the following tag(s): article:topic-guide






Electrochemical Cell Conventions

Electrochemical Cell Conventions

Using chemical reactions to produce electricity is now a priority for many researchers. Being able to adequately use chemicalreactions as a source of power would greatly help our environmental pollution problems. In this section of electrochemistry, wewill be learning how to use chemical reactions to produce this clean electricity and even use electricity to generate chemicalreactions. In order to induce a flow of electric charges, we place a strip of metal (the electrode) in a solution containing thesame metal, which is in aqueous state. The combination of an electrode and its solution is called a half cell. Within the halfcell, metals ions from the solution could gain electrons from the electrode and become metal atoms;or the metal atoms from theelectrode could lose electrons and become metals ions in the solution.

Introduction

We use two different half cells to measure how readily electrons can flow from one electrode to another, and the device used formeasurement is called a voltmeter. The voltmeter measures the cell potential, denoted by Ecell, (in units of Volts, 1V=1J/C),which is the potential difference between two half cells. The salt bridge allows the ions to flow from one half cell to anotherbut prevents the flow of solutions.

As indicated in the diagram, the anode is the electrode whre oxidation occurs; Cu loses two electrons to form Cu+2. The

cathode is the electrode where reduction occurs; Ag+ (aq) gains electron to become Ag(s). As a convenient substitution for thedrawing, we use a cell diagram to show the parts of an electrochemical cell. For example above, the cell diagram is :






Zn(s) | Zn2+ (aq) || Cu2+ (aq)| Cu(s)

oxidation- (half-cell) (salt bridge) (half-cell)-reduction

Where we place the anode on the left and cathode on the right, "|" represents the boundary between the two phases, and "||"represents the salt bridge. There are two types of electrochemical cells:

A Galvanic Cell (aka Voltaic Cell) induces a spontaneous redox reaction to create a flow of electrical charges, orelectricity. Non-rechargeable batteries are examples of Galvanic cells.

• A Reaction is spontaneous when the change in Gibb’s energy, ∆G is < 0.• Electrons flow from the anode(negative since electrons are built up here) to the cathode (positive since it is

gaining electrons).

An Electrolytic cell is one kind of battery that requires an outside electrical source to drive the non-spontaneous redoxreaction. Rechargeable batteries act as Electrolytic cells when they are being recharged.

• A reaction is non-spontaneous when ∆G is > 0.• Must supply electrons to the cathode to drive the reduction, so cathode is negative.• Must remove electrons from the anode to drive the oxidation, so anode is positive.

Both Galvanic and Electrolytic cells contain:

• Two electrodes: the Anode and the Cathode (NOTE: Cathode does not mean +, and Anode does not mean -)• Volt meter: measures the electric current. In Galvanic cells, this shows how much current is produced; in Electrolytic

cells, this shows how much current is charging the system.• Electrolyte

◦ conducting medium◦ has contact with electrodes◦ usually in aqueous solution of ionic compounds

• Salt Bridge◦ joins the two halves of the electrochemical cell◦ filled with a salt solution or gel◦ keeps the solution separate◦ Completes the circuit

Basic Terminology• Electrochemical cells use a vast amount of terminology. Here is a brief definition of some of the more common

terms:• Voltage- The potential difference between two half cells, also the amount of energy that drives a reaction. Voltage is

an intensive property (amount of voltage does matter).• Current-The flow of electric charges (in units of electrons per second). It is an extensive property (amount of current

does matter). NOTE: High voltage does not mean high current.


• Primary Battery- non-rechargeable batteries. AA, AAA, etc.• Secondary Battery- Rechargeable batteries. Lithium, cell phone batteries, etc.• Tertiary Battery- Fuel cells. Although not always considered as batteries, these often require a constant flow of

reactants.

Galvanic Cell (aka Voltaic Cells)

Zn(s) | Zn2+ (aq) || Cu2+ (aq)| Cu(s)

A galvanic cell produces an electrical charge from the flow of electrons. The electrons move due to the Redox reaction. As we

can see, Zn oxidizes to Zn2+ , while Cu2+ reduces to Cu. In order to understand the redox reaction, Solve the Redox equation.

First, split the reaction into two half reactions, with the same elements paired with one another.

Zn(S) → Zn+2(aq) Oxidation Reaction: takes place at the Anode

Cu+2(aq) → Cu(s) Reduction Reaction: Takes place at the Cathode

Next, we balance the two equations.

Oxidation: Zn(S) →Zn2+(aq) + 2e- (Anode)

Reduction: 2e- + Cu2+(aq) → Cu(s) (Cathode)

?Spontaneous redox reaction releases energy; The system does work on the surroundings.)

Finally, we recombine the two equations. As you can see, this equation was already balanced. However, not all cells arenecessarily balanced. It is important to check each time. Galvanic cells are quite common. A, AA, AAA, D, C, etc. batteries areall galvanic cells. Any non-rechargeable battery that does not depend on an outside electrical source is a Galvanic cell.


Electrolytic Cell

Cu(s) | Cu2+(aq) || Ag+

(aq) | Ag(s)

An electrolytic cell is a cell which requires an outside electrical source to initiate the redox reaction. The process of howelectric energy drives the nanspontaneous reaction is called electrolysis. Whereas the galvanic cell used a redox reaction tomake electrons flow, the electrolytic cell uses electron movement (in the source of electricity) to cause the redox reaction. Inan electrolytic cell, electrons are forced to flow in the opposite direction. Since the direction is reversed of the voltaic cell, the

E0cell for electrolytic cell is negative. Also, in order to force the electrons to flow in the opposite direction, the electromotive

force that connects the two electrode-the battery must be larger than the magnitude of E0cell. This additional requirement of

voltage is called overpotential.

Electrolytic cell for the example above:

Oxidation: Cu(s) → Cu2+ (aq)+2e- (anode)

Reduction: Zn2+ (aq)+2e- → Zn(s) (cathode)

(Nonspontaneous redox reaction absorbs energy to drive it; The surroundings do work on the system. )

Galvanic: turns chemical energy into electrical energy

Electrolytic Cell: turns electrical energy into chemical energy

The most common form of Electrolytic cell is the rechargeable battery (cell phones, mp3's, etc) or electroplating. While thebattery is being used in the device it is a galvanic cell function (using the redox energy to produce electricity). While the batteryis charging it is an electrolytic cell function (using outside electricity to reverse the completed redox reaction).

Inert & Active Electrode

An inert electrode is a metal submerged in an aqueous solution of ion compounds that transfers electrons rather than exchangingions with the aqueous solution. It does not participate or interfere in the chemical reaction but serves as a source of electrons.


Platinum is usually the metal used as an inert electrode. An active electrode is an electrode that can be oxidized or reduced in

half reaction. For example, Cu; Cu can be oxidized to Cu2+ at the annode and one Cu2+ ion can also reduces to a Cu atom at

the cathode. Cu is transfered from anode to cathode through the solution as Cu2+ from the example above.

Anode: Fe2+(aq) → Fe3+

(s) + e-

Cathode: MnO4-(aq) + 8H+

(aq) + 5e- → Mn2+(aq) + 4H2O(l)

Outside links

• http://en.wikipedia.org/wiki/Electrochemical_cell

• http://encarta.msn.com/encyclopedia_761569809/Electrochemistry.html

References1. Petrucci, Harwood, Herring, and Madura. General Chemistry: Principles and Modern Applications: Ninth Edition.

New Jersey: Pearson, 2007.2. Professor Delmar Larsen. Lecture 2, 3, and 6. Spring 20103. Rieger, Philip. Electrochemistry. 2, Illustrated. Springer Us, 1994. 112-113. Print.4. Hamann, Carl, Andrew Hamnett, and Wolf Vielstich. Electrochemistry. 2, Illustrated. Vch Verlagsgesellschaft Mbh,

2007. 82. Print.

Contributors


http://en.wikipedia.org/wiki/Electrochemical_cell

http://encarta.msn.com/encyclopedia_761569809/Electrochemistry.html

The Cell Potential

The Cell Potential

The batteries in your remote and the engine in your car are only a couple of examples of how chemical reactions create powerthrough the flow of electrons. The cell potential is the way in which we can measure how much voltage exists between the twohalf cells of a battery. We will explain how this is done and what components allow us to find the voltage that exists in anelectrochemical cell.

Introduction

The cell potential, \(E_{cell}\), is the measure of the potential difference between two half cells in an electrochemical cell. Thepotential difference is caused by the ability of electrons to flow from one half cell to the other. Electrons are able to movebetween electrodes because the chemical reaction is a redox reaction. A redox reaction occurs when a certain substance isoxidized, while another is reduced. During oxidation, the substance loses one or more electrons, and thus becomes positivelycharged. Conversely, during reduction, the substance gains electrons and becomes negatively charged. This relates to themeasurement of the cell potential because the difference between the potential for the reducing agent to become oxidized andthe oxidizing agent to become reduced will determine the cell potential. The cell potential (Ecell) is measured in voltage (V),which allows us to give a certain value to the cell potential.

Electrochemical Cell

An electrochemical cell is comprised of two half cells. In one half cell, the oxidation of a metal electrode occurs, and in theother half cell, the reduction of metal ions in solution occurs. The half cell essentially consists of a metal electrode of a certainmetal submerged in an aqueous solution of the same metal ions. The electrode is connected to the other half cell, which containsan electrode of some metal submerged in an aqueous solution of subsequent metal ions. The first half cell, in this case, will bemarked as the anode. In this half cell, the metal in atoms in the electrode become oxidized and join the other metal ions in theaqueous solution. An example of this would be a copper electrode, in which the Cu atoms in the electrode loses two electrons

and becomes Cu2+ .






The Cu2+ ions would then join the aqueous solution that already has a certain molarity of Cu2+ ions. The electrons lost by theCu atoms in the electrode are then transferred to the second half cell, which will be the cathode. In this example, we will assumethat the second half cell consists of a silver electrode in an aqueous solution of silver ions. As the electrons are passed to the Ag

electrode, the Ag+ ions in solution will become reduced and become an Ag atom on the Ag electrode. In order to balance thecharge on both sides of the cell, the half cells are connected by a salt bridge. As the anode half cell becomes overwhelmed with

Cu2+ ions, the negative anion of the salt will enter the solution and stabilized the charge. Similarly, in the cathode half cell, asthe solution becomes more negatively charged, cations from the salt bridge will stabilize the charge.

How does this relate to the cell potential?

For electrons to be transferred from the anode to the cathode, there must be some sort of energy potential that makes thisphenomenon favorable. The potential energy that drives the redox reactions involved in electrochemical cells is the potentialfor the anode to become oxidized and the potential for the cathode to become reduced. The electrons involved in these cellswill fall from the anode, which has a higher potential to become oxidized to the cathode, which has a lower potential to becomeoxidized. This is analogous to a rock falling from a cliff in which the rock will fall from a higher potential energy to a lowerpotential energy.

Note

The difference between the anode's potential to become reduced and the cathode's potential to become reduced is the cellpotential.

\[Eô_{Cell}= Eô_{Red,Cathode} - Eô_{Red,Anode}\]

Note:


• Both potentials used in this equation are standard reduction potentials, which are typically what you find in tables(e.g., Table P1 and Table P2). However, the reaction at the anode is actually an oxidation reaction -- the reverse of areduction reaction. This explains the minus sign. We would have used a plus sign had we been given an oxidationpotential \(Eô_{Ox,Anode}\) instead, since \(Eô_{Red}=Eô_{Ox}\).

• The superscript "o" in Eô indicates that these potentials are correct only when concentrations are 1 M andpressures are 1 bar. A correction called the "Nernst Equation" must be applied if conditions are different.

Electrochemical cell

Here is the list of the all the components:

1. Two half cells2. Two metal electrodes3. One voltmeter4. One salt bridge5. Two aqueous solutions for each half cell

All of these components create the Electrochemical Cell.

How to measure the cell potential?


http://chemwiki.ucdavis.edu/Reference/Reference_Tables/Electrochemistry_Tables/P1%3A_Standard_Reduction_Potentials_by_Element

http://chemwiki.ucdavis.edu/Reference/Reference_Tables/Electrochemistry_Tables/P2%3A_Standard_Reduction_Potentials_by_Value

http://chemwiki.ucdavis.edu/@api/deki/files/186/electrode_potential_2.png

The image above is an electrochemical cell. The voltmeter at the very top in the gold color is what measures the cell voltage,or the amount of energy being produced by the electrodes. This reading from the voltmeter is called the voltage of theelectrochemical cell. This can also be called the potential difference between the half cells, Ecell. Volts are the amount of energyfor each electrical charge; 1V=1J/C: V= voltage, J=joules, C=coulomb. The voltage is basically what propels the electrons tomove. If there is a high voltage, that means there is high movement of electrons. The voltmeter reads the transfer of electronsfrom the anode to the cathode in Joules per Coulomb.

Cell Diagram

The image above is called the cell diagram. The cell diagram is a representation of the overall reaction in the electrochemicalcell. The chemicals involved are what are actually reacting during the reduction and oxidation reactions. (The spectator ions areleft out). In the cell diagram, the anode half cell is always written on the left side of the diagram, and in the cathode half cellis always written on the right side of the diagram. Both the anode and cathode are seperated by two vertical lines (ll) as seen inthe blue cloud above. The electrodes (yellow circles) of both the anode and cathode solutions are seperated by a single verticalline (l). When there are more chemicals involved in the aqueous solution, they are added to the diagram by adding a commaand then the chemical. For example, in the image above, if copper wasn't being oxidized alone, and another chemical like Kwas involved, you would denote it as (Cu, K) in the diagram. The cell diagram makes it easier to see what is being oxidizedand what is being reduced. These are the reactions that create the cell potential.

Standard Cell Potential

The standard cell potential (\(Eô_{cell}\)) is the difference of the two electrodes, which forms the voltage of that cell. To findthe difference of the two half cells, the following equation is used:

\[Eô_{Cell}= Eô_{Red,Cathode} - Eô_{Red,Anode} \tag{1a}\]

with

• \(Eô_{Cell}\) is the standard cell potential (under 1M, 1 Barr and 298 K).

• \(Eô_{Red,Cathode}\) is the standard reduction potential for the reduction half reaction occurring at the cathode

• \(Eô_{Red,Anode}\) is the standard reduction potential for the oxidation half reaction occurring at the anode

The units of the potentials are typically measured in volts (V). Note that this equation can also be written as a sum rather thana difference

\[Eô_{Cell}= Eô_{Red,Cathode} + Eô_{Ox,Anode} \tag{1b}\]


http://chemwiki.ucdavis.edu/Analytical_Chemistry/Electrochemistry/Redox_Chemistry/Standard_Reduction_Potential



where we have switched our strategy from taking the difference between two reduction potentials (which are traditionally whatone finds in reference tables) to taking the sum of the oxidation potential and the reduction potential (which are the reactionsthat actually occur). Since Eô_{Red}=-Eô_{Ox}, the two approaches are equivalent.

Standard Cell Potential Example

The example will be using the picture of the Copper and Silver cell diagram. The oxidation half cell of the redox equation is:

Cu(s) → Cu2+(aq) + 2e- EoOx= -0.340 V

where we have negated the reduction potential EoRed= 0.340 V, which is the quantity we found from a list of standard

reduction potentials, to find the oxidation potential EoOx. The reduction half cell is:

( Ag+ + e- → Ag(s) ) x2 EoRed= 0.800 V

where we have multiplied the reduction chemical equation by two in order to balance the electron count but we have not

doubled EoRed since Eo values are given in units of voltage. Voltage is energy per charge, not energy per reaction, so it

does not need to account for the number of reactions required to produce or consume the quantity of charge you are using tobalance the equation. The chemical equations can be summed to find:

Cu(s) + 2Ag+ + 2e- → Cu2+(aq) + 2Ag(s) + 2e-

and simplified to find the overall reaction:

Cu(s) + 2Ag+ → Cu2+(aq) + 2Ag(s)

where the potentials of the half-cell reactions can be summed

EoCell= Eo

Red,Cathode+EoOx,Anode

EoCell = 0.800 V + (-0.340 V)

EoCell = 0.460V

to find that the standard cell potential of this cell is 0.460 V. We are done.

Note that since Eô_{Red}=-Eô_{Ox} we could have accomplished the same thing by taking the difference of the reductionpotentials, where the absent or doubled negation accounts for the fact that the reverse of the reduction reaction is whatactually occurs.

EoCell= Eo

Red,Cathode-EoRed,Anode


EoCell = 0.800V - 0.340V

EoCell = 0.460V

Important Standard Electrode (Reduction) Potentials

The table below is a list of important standard electrode potentials in the reduction state. To determine oxidation electrodes, thereduction equation can simply be flipped and its potential changed from positive to negative (and vice versa). When using thehalf cells below, instead of changing the potential the equation below can be used without changing any of the potentials frompositive to negative (and vice versa):

EoCell= Eo

Red,Cathode - EoRed,Anode

Table:

Reduction Half-Reaction Eo, V

Acidic Solution

F2(g) + 2e- → 2 F-(aq) +2.866

O3(g) + 2H+(aq) + 2e- → O2(g) + H2O(l) +2.075

S2O82-(aq) + 2e- → 2SO42-(aq) +2.01

H2O2(aq) + 2H+(aq) +2e- → 2H2O(l) +1.763

MnO4-(aq) + 8H+(aq) + 5e- → Mn2+(aq) +4H2O(l)

+1.51

PbO2(s) + 4H+(aq) + 2e- → Pb2+(aq) +4H2O(l)

+1.455

Cl2(g) + 2e- → 2Cl-(aq) +1.358

Cr2O72-(aq) + 14H+(aq) + 6e- → 2Cr3+(aq)+ 7H2O(l)

+1.33

MnO2(s) + 4H+(aq) +2e- -> Mn2+(aq) +2H2O(l)

+1.23

O2(g) + 4H+(aq) + 4e- → 2H2O(l) +1.229


2IO3-(aq) + 12H+(aq) + 10e- → I2(s) +6H2O(l)

+1.20

Br2(l) + 2e- → 2Br-(aq) +1.065

NO3-(aq) + 4H+(aq) + 3e- → NO(g) + 2H2O(l)

+0.956

Ag+(aq) + e- → Ag(s) +0.800

Fe3+(aq) + e- → Fe2+(aq) +0.771

O2(g) + 2H+(ag) + 2e- → H2O2(aq) +0.695

I2(s) + 2e- → 2I-(aq) +0.535

Cu2+(aq) + 2e- → Cu(s) +0.340

SO42-(aq) + 4H+(aq) + 2e- → 2H2O(l) +SO2(g)

+0.17

Sn4+(aq) + 2e- → Sn2+(aq) +0.154

S(s) + 2H+(aq) + 2e- → H2S(g) +0.14

2H+(aq) + 2e- → H2(g) 0

Pb2+(aq) + 2e- → Pb -0.125

Sn2+(aq) + 2e- → Sn(s) -0.137

Fe2+(aq) + 2e- → Fe(s) -0.440

Zn2+ + 2e- → Zn(s) -0.763

Al3+(aq) + 3e- → Al(s) -1.676

Mg2+(aq) + 2e- → Mg(s) -2.356

Na+(aq) + e- → Na(s) -2.713

Ca2+(aq) + 2e- → Ca(s) -2.84


K+(aq) + + e- → K(s) -2.924

Li+(aq) + e- → Li(s) -3.040

Basic Solution

O3(aq) + H2O(l) + 2e- → O2(g) + 2OH-(aq) +1.246

OCl-(aq) + H2O(l) + 2e- → Cl-(aq) +2OH-(aq)

+0.890

O2(g) + 2H2O(l) +4e- → 4OH-(aq) +0.401

2H2O(l) + + 2e- → H2(aq) + 2OH-(aq) -0.0828

Outside links• http://en.wikipedia.org/wiki/Standar...rode_potential

• http://www.chem.purdue.edu/gchelp/ho...Potentials.htm

• http://hyperphysics.phy-astr.gsu.edu...electrode.html

• Cheng, K. L. "A New Concept for pH-Potential Calculations." J. Chem. Educ. 1999 76 1029.

• Wuchang, Z.; Jiaxing, L.; Hebal, S. "On the standard hydrogen electrode (LTE)." J. Chem. Educ. 1991, 68, 356.

• https://www.chem.wisc.edu/deptfiles/...ial/18_51.html

• http://www.corrosion-doctors.org/Cor...-potential.htm

Problems1. For this redox reaction \[Sn(s) + Pb^{2+}(aq) \rightarrow Sn^{2+}(aq) + Pb (s)\] write out the oxidation and reduction

half reactions. Create a cell diagram to match your equations.2. From the image above, of the cell diagram, write the overall equation for the reaction.3. If \(Cu^{2+}\) ions in solution around a \(Cu\) metal electrode is the cathode of a cell, and \(K^+\) ions in solution

around a K metal electrode is the anode of a cell, which half cell has a higher potential to be reduced?4. What type of reaction provides the basis for a cell potential?5. How is the cell potential measured and with what device is it measured?6. The \(E^o_{cell}\) for the equation \[ 4Al(s) + 3O_2(g) + 6H_2O(l) + 4OH^-(aq) \rightarrow 4[Al(OH)_4]^-(aq) \] is

+2.71 V. If the reduction of \(O_2\) in \(OH^-\) is +0.401 V. What is the reduction half-reaction for this reduction halfreaction? \[[Al(OH)_4]^-(aq) + 3e^- \rightarrow Al(s) + 4OH^-\]


http://en.wikipedia.org/wiki/Standard_electrode_potential

http://www.chem.purdue.edu/gchelp/howtosolveit/Electrochem/Electrochemical_Cell_Potentials.htm

http://hyperphysics.phy-astr.gsu.edu/Hbase/Chemical/electrode.html

http://www.jce.divched.org/Journal/Issues/1999/Jul/abs1029.html

http://www.jce.divched.org/Journal/Issues/1991/Apr/jceSubscriber/JCE1991p0356.pdf

https://www.chem.wisc.edu/deptfiles/genchem/netorial/ROttosen/tutorial/modules/electrochemistry/05potential/18_51.html

http://www.corrosion-doctors.org/Corrosion-Thermodynamics/Standard-potential.htm

Answers

1. oxidation: Sn(s) → Sn2+(aq) + 2e-(aq)

reduction: Pb2+(aq) + 2e-(aq) → Pb(s)

cell diagram: Sn(s) | Sn2+(aq) || Pb2+(aq) | Pb(s)

2. Cu(s) + 2Ag+(aq) → 2Ag(s) + Cu2+(aq)

3. Because the half cell containing the \(Cu\) electrode in \(Cu^{2+}\) solution is the cathode, this is the half cell wherereduction is taking place. Therefore, this half cell has a higher potential to be reduced.

4. The redox reaction.5. Cell potential is measured in Volts (=J/C). This can be measured with the use of a voltmeter.6. We can divide the net cell equation into two half-equations.

• Oxidation: {Al(s) + 4OH-(aq) → [Al(OH4)]-(aq) + 3e-} x4; -Eo= ? This is what we are solving for.

• Reduction: {O2(g) + 2H2O(l) + 4e- → 4OH-(aq)} x3 Eo= +0.401V

• Net: 4Al(s) + 3O2(g) + 6H2O(l) + 4OH-(aq) → 4[Al(OH)4]-(aq) Eocell = 2.71V

Eocell= 2.71V= +0.401V - Eo{Al(OH)4]-(aq)/Al(s)}

Eo{[Al(OH)4]-(aq)/Al(s)} = 0.401V - 2.71V = -2.31V

Confirm this on the table of standard reduction potentials

References1. Petrucci, Harwood, Herring, and Madura. General Chemistry: Principles and Modern Applications. 9th ed. Upper

Saddle River, New Jersey: Pearson Education, 2007.

Contributors• Katherine Barrett, Gianna Navarro, Joseph Koressel, Justin Kohn



Writing Equations for Redox Reactions

Writing Equations for Redox Reactions

This page explains how to work out electron-half-reactions for oxidation and reduction processes, and then how to combine them to

give the overall ionic equation for a redox reaction. This is an important skill in inorganic chemistry.

Electron-half-equations

The ionic equation for the magnesium-aided reduction of hot copper(II) oxide to elemental copper is given below :

\[Cu^{2+} + Mg \rightarrow Cu + Mg^{2+}\]

The equation can be split into two parts and considered from the separate perspectives of the elemental magnesium and of the

copper(II) ions. This arrangement clearly indicates that the magnesium has lost two electrons, and the copper(II) ion has gained them.

\[ Mg \rightarrow Mg^{2+} + 2e^-\]

\[Cu^{2+} + 2e^- \rightarrow Cu\]

These two equations are described as "electron-half-equations," "half-equations," or "ionic-half-equations," or "half-reactions."

Every redox reaction is made up of two half-reactions: in one, electrons are lost (an oxidation process); in the other, those electrons

are gained (a reduction process).

Working out electron-half-equations and using them to build ionic equations

In the example above, the electron-half-equations were obtained by extracting them from the overall ionic equation. In practice,the reverse process is often more useful: starting with the electron-half-equations and using them to build theoverall ionic equation.

Example 1: The reaction between Chlorine and Iron (III) Ions






Chlorine gas oxidizes iron(II) ions to iron(III) ions. In the process, the chlorine is reduced to chloride ions. Fromthis information, the overall reaction can be obtained. The chlorine reaction, in which chlorine gas is reduced tochloride ions, is considered first:

\[ Cl_2 \rightarrow Cl^-\]

The atoms in the equation must be balanced:

\[ Cl_2 \rightarrow 2Cl^-\]

This step is crucial. If any atoms are unbalanced, problems will arise later.

To completely balance a half-equation, all charges and extra atoms must be equal on the reactant and product sides. In order to

accomplish this, the following can be added to the equation:

• electrons

• water

• hydrogen ions (unless the reaction is being done under alkaline conditions, in which case, hydroxide ions must be added andbalanced with water)

In the chlorine case, the only problem is a charge imbalance. The left-hand side of the equation has no charge, but the right-

hand side carries 2 negative charges. This is easily resolved by adding two electrons to the left-hand side. The fullybalanced half-reaction is:

\[ Cl_2 +2 e^- \rightarrow 2Cl^-\]

Next the iron half-reaction is considered. Iron(II) ions are oxidized to iron(III) ions as shown:

\[ Fe^{2+} \rightarrow Fe^{3+}\]

The atoms balance, but the charges do not. There are 3 positive charges on the right-hand side, but only 2 on the left. To reduce

the number of positive charges on the right-hand side, an electron is added to that side:

\[ Fe^{2+} \rightarrow Fe^{3+} + e-\]

The next step is combining the two balanced half-equations to form the overall equation. The two half-equationsare shown below:


It is obvious that the iron reaction will have to happen twice for every chlorine reaction. This is accounted for in the following

way: each equation is multiplied by the value that will give equal numbers of electrons, and the two resulting equations are added

together such that the electrons cancel out:

At this point, it is important to check once more for atom and charge balance. In this case, no further work is required.

Example 2: The reaction between Hydrogen Peroxide and Magnanate Ions

The first example concerned a very simple and familiar chemical equation, but the technique works just as well for more

complicated (and perhaps unfamiliar) chemistry.

Manganate(VII) ions, MnO4-, oxidize hydrogen peroxide, H2O2, to oxygen gas. The reaction is carried out with potassium

manganate(VII) solution and hydrogen peroxide solution acidified with dilute sulfuric acid. As the oxidizing agent, Manganate(VII)is reduced to manganese(II).

The hydrogen peroxide reaction is written first according to the information given:

\[ H_2O_2 \rightarrow O_2\]

The oxygen is already balanced, but the right-hand side has no hydrogen.

All you are allowed to add to this equation are water, hydrogen ions and electrons. Adding water is obviously unhelpful: if water is

added to the right-hand side to supply extra hydrogen atoms, an additional oxygen atom is needed on the left. Hydrogen ions are

a better choice.


Adding two hydrogen ions to the right-hand side gives:

\[ H_2O_2 \rightarrow O_2 + 2H^+\]

Next the charges are balanced by adding two electrons to the right, making the overall charge on both sides zero:

\[ H_2O_2 \rightarrow O_2 + 2H^+ + 2e^-\]

Next the manganate(VII) half-equation is considered:

\[MnO_4^- \rightarrow Mn^{2+}\]

The manganese atoms are balanced, but the right needs four extra oxygen atoms. These can only come from water, so four water

molecules are added to the right:

\[ MnO_4^- \rightarrow Mn^{2+} + 4H_2O\]

The water introduces eight hydrogen atoms on the right. To balance these, eight hydrogen ions are added to the left:

\[ MnO_4^- + 8H^+ \rightarrow Mn^{2+} + 4H_2O\]

Now that all the atoms are balanced, only the charges are left. There is a net +7 charge on the left-hand side (1- and 8+), but only

a charge of +2 on the right. 5 electrons are added to the left-hand side to reduce the +7 to +2:

\[ MnO_4^- + 8H^+ + 5e^- \rightarrow Mn^{2+} _ 4H_2O\]

This illustrates the strategy for balancing half-equations, summarized as followed:

• Balance the atoms apart from oxygen and hydrogen.

• Balance the oxygens by adding water molecules.

• Balance the hydrogens by adding hydrogen ions.

• Balance the charges by adding electrons.

Now the half-equations are combined to make the ionic equation for the reaction.


As before, the equations are multiplied such that both have the same number of electrons. In this case, the least common multiple

of electrons is ten:

The equation is not fully balanced at this point. There are hydrogen ions on both sides which need to besimplified:

This often occurs with hydrogen ions and water molecules in more complicated redox reactions. Subtracting 10 hydrogen ions from

both sides leaves the simplified ionic equation.

\[ 2MnO_4^- + 6H^+ + 5H_2O_2 \rightarrow 2Mn^{2+} + 8H_2O + 5O_2\]

Example 3: Oxidation of Ethanol of Acidic Potassium Dichromate (IV)

This technique can be used just as well in examples involving organic chemicals. Potassium dichromate(VI) solution acidified with

dilute sulfuric acid is used to oxidize ethanol, CH3CH2OH, to ethanoic acid, CH3COOH.


The oxidizing agent is the dichromate(VI) ion, Cr2O72-, which is reduced to chromium(III) ions, Cr3+. The ethanol to ethanoic acid

half-equation is considered first:

\[ CH_3CH_2OH \rightarrow CH_3COOH\]

The oxygen atoms are balanced by adding a water molecule to the left-hand side:

\[ CH_3CH_2OH + H_2O \rightarrow CH_3COOH\]

Four hydrogen ions to the right-hand side to balance the hydrogen atoms:

\[ CH_3CH_2OH + H_2O \rightarrow CH_3COOH + 4H^+\]

The charges are balanced by adding 4 electrons to the right-hand side to give an overall zero charge on each side:

\[ CH_3CH_2OH + H_2O \rightarrow CH_3COOH + 4H^+ + 4e^-\]

The unbalanced dichromate (VI) half reaction is written as given:

\[ Cr_2O_7^{2-} \rightarrow Cr^{3+}\]

At this stage, students often forget to balance the chromium atoms, making it impossible to obtain the overall equation. To avoid

this, the chromium ion on the right is multiplied by two:

\[ Cr_2O_7^{2-} \rightarrow 2Cr^{3+}\]

The oxygen atoms are balanced by adding seven water molecules to the right:

\[ Cr_2O_7^{2-} \rightarrow 2Cr^{3+} + 7H_2O\]

The resulting hydrogen atoms are balanced by adding fourteen hydrogen ions to the left:

\[ Cr_2O_7^{2-} + 14H^+ \rightarrow 2Cr^{3+} + 7H_2O\]

Six electrons are added to the left to give a net +6 charge on each side.

\[ Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O\]

The two balanced half reactions are summarized:

\[ CH_3CH_2OH + H_2O \rightarrow CH_3COOH + 4H^+ + 4e^-\]

\[ Cr_2O_7^{2-} + 14H^+ + 6e^- \rightarrow 2Cr^{3+} + 7H_2O\]


The least common multiple of 4 and 6 is 12. Therefore, the first equation is multiplied by 3 and the second by 2, giving 12 electrons

in each equation:

Simplifying the water molecules and hydrogen ions gives final equation:

Balancing reactions under alkaline conditions

Working out half-equations for reactions in alkaline solution is decidedly more tricky than the examples above. As some curricula do

not include this type of problem, the process for balancing alkaline redox reactions is covered on a separate page.

Contributors

Jim Clark (Chemguide.co.uk)


http://www.chemguide.co.uk

chemwiki book 30 dec 2015 3

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