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Electrochemistry One day sir, you may tax it - Michael Faraday, in response to a question posed about the practical uses of electricity.

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Chapter Overview and Introduction

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Chapter Learning Goals Balancing redox reactions (these are fun ). Finding voltages, gibbs free energy and equilibrium constants for electrolytic and galvanic cells. Relate standard and non-standard voltages Determining what will and will not undergo redox reactions And much much more!!!

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Real Life Explanation Goals Fuel cells Corrosion Electroplating If we have time: battery alternatives: bacteria, viruses and other cool things

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Electrochemistry: A Definition Electro chemistry is the intersection of electrical and chemical energy One species is reduced, another is oxidized- a transfer of electrons from one (or more) atoms to another.

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Zn(s)+Cu 2+ (aq) Zn 2+ (aq)+ Cu(s) Electrons go from the Zn to the Cu 2+

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Electrochemistry Definitions

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Learning Outcomes Define oxidation number, reduction, oxidation, reducing agent and oxidation agent Identify the oxidation number for species Identify the species being oxidized Identify the species being reduced Identify the reducing agent Identify the oxidation agent Before moving on to harder things, make sure being able to do these things is automatic!!!!

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Definitions Oxidation number: imaginary charge if the compound was broken down into ions Oxidation number of atom in its elemental state is zero Oxidation number of a monatomic ion is equal to its charge Oxidation numbers of individual atoms must add up to the charge on the whole molecule. Use this to find oxidation states of ions you dont know, or that may have multiple possibilities. +0 +1

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Oxidation state examples: Na 2 CrO 4 MnO 4 - Na= +1 (2)= +2 O= -2 (4)= -8 -6 Neutral compound so Cr=+6 O= -2 (4)= -8 -8 Charge is -1, so Mn=+7

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Definitions Oxidized : Losing e -, increased charge Reduced : Gaining e -, lowered charge Oxidizing Agent: Species which oxidizes other compounds (and is thereby reduced)** Reducing Agent : Species which reduces other compounds (and is thereby oxidized)** Anode : Where oxidation takes place Cathode : where reduction takes place +0 +1 Oxidized Reduced Reducing Agent Oxidizing Agent Anode Cathode Memory Tricks: LEO the lion goes GER RED CAT / AN OX **Note: The whole compound is the oxidizing/reducing agent.

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Learning outcomes Learn to balance redox reactions in acidic or basic solutions Using half reaction method Leaving the reactions together You should know how to do both methods: Leaving reaction together is always possible, but sometimes more difficult Half reaction method is sometimes not possible.

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Some methods are easier for different situations, be open to switching ways! Half reaction method : couple of different algorithms, pick one and stick with it. (I picked the one I think is easiest for students) Leaving the reaction together: great for when the species cant be separated. Sometimes its just easier. there will be one where you HAVE to do this on the exam. Redox Reactions- Balancing Two main methods

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Step 1: write oxidation numbers for every species. Step 2: identify which is being reduced and which is being oxidized Step 3: separate into two separate half reactions, do steps 4-7 for each. Step 4:First balance all elements EXCEPT H and O Step 5: Add H 2 O to balance oxygen Step 6: Add H + to balance H If its a basic solution add as many OH - to each side as needed to cancel H + into H 2 O (can also be done after steps 7, 8, or 9, instead. Do wherever is most convenient) Step 7: add electrons to balance charges (oxidation states) Step 8: Look at both half reactions and multiply each as needed to get the number of electons on both sides to be equal. Step 9: Combine reactions together. Half method rules: Use this as a checklist

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Examples: Balance the following, identify which species is oxidized, which is reduced, which is the oxidizing agent and which is the reducing agent. Basic Acid solution

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Balancing without separating Step 1: Write oxidation numbers. Step 2: Balance all species that ARE NOT H or O Step 3: Identify which species is being oxidized and which is reduced. Step 4: Decide how many electrons are transferred in each. Step 5: Multiply coefficients as needed to get the electrons transferred in the Red/Ox to be equal Step 6: Balance O by adding water Step 7: Balance H by adding H If its a basic solution add as many OH - to each side as needed to cancel H + into H 2 O

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Examples: Balance the following, identify which species is oxidized, which is reduced, which is the oxidizing agent and which is the reducing agent. Balance keeping it together in an acidic solution Rebalance using this method in an acidic solution

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Extra examples (podcasted) Cinnabar/vermillion Both Acidic and Basic

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Review Balancing redox reactions require that you account for the exchange of electrons as well. You must have the same number of electrons being transferred from one reaction to the other. Learn at minimum the method for leaving reactions together. The half reaction method is often very useful though too!

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Introduction to electrochemical cells

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Learning Outcomes Identify the components of a galvanic cell Identify the components of a cell. Determining which reaction happens in which part will be covered in the next section. Define the difference between a galvanic and electrolytic cell.

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Electrochemical cells Galvanic cell: an electrochemical cell where a spontaneous chemical reaction is used to generate an electric current Electrolytic cell: an electrochemical cell where a non-spontaneous chemical reaction is occurring due to a supplied electrical source. Anode Cathode Oxidation occurs at the Anode Reduction occurs at the Cathode Memory trick: Red Cat An Ox

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Galvanic cells: Questions Where is metal deposited? Copper Which direction do the electrons flow? Anode to cathode Why must we separate the reactions? In order to have the electrons do work Why is the salt bridge necessary? Allow flow of anions and cations Anode Cathode

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Review Galvanic cells separate reactions. This allows the electrons to flow from one area to the other doing work. You should be able to identify the anode and cathode given the reactions But how would we know which is which if not given the reactions occurring? How do we know how much voltage we can get?

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Cell Voltage.

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Learning Outcomes Define EMF (electromotive force) or voltage Identify the standard hydrogen electrode and how it is used as a standard. Use table of values and half reactions to calculate the EMF of a cell. Use the cell voltage to decide if a cell is spontaneous or not.

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Electromotive (EM) force Not really a force it is a voltage Synonymous with cell voltage Listed values are E o red aka reduction voltages E o red = - E o oxid These are standard values, 1atm or 1M

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Electromotive (EM) force Listed values are E o red aka reduction voltages E o red = - E o oxid EMF of a cell is the difference in potential energy between the anode and cathode.

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Standard Hydrogen Electrode We need a voltage to compare everything too. The H 2 half reaction cell was chosen. Why do we have a Pt electrode? Need an inert metal to conduct electrons Standard potentials are calculated against the standard hydrogen electrode!

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Standard Hydrogen Electrode Why do we have a Pt electrode? Need an inert metal to conduct electrons Standard potentials are calculated against the standard hydrogen electrode!

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Calculating Cell Voltage Two equivalent ways. Pick your favorite and stick with it. Be very careful not to interchange them!!! Why are these equivalent? Method 1: Method 2: E o red = - E o oxid positive E o cell is a spontaneous reaction (galvanic cell) negative E o cell is a non-spontaneous reaction (electrolytic cell)

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Calculating Cell Voltage Which is Anode/Cathode How do we decide? spontaneous: E o cell = positive non-spontaneous: E o cell =negative When reversing the sign of the anode (oxidation), they must add to be positive. Copper=cathode Zinc= anode E o cu2+/Cu = +0.34V E o Zn2+/Zn = -0.44V Calculate the cell voltage for a spontaneous Zn/Zn 2+, Cu/Cu 2+ cell.

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Calculating Cell Voltage: Calculate the cell voltage for a spontaneous Zn/Zn 2+, Cu/Cu 2+ cell Logic 1: Logic 2: E o cell =E o red_Cu2+/Cu - E o red_Zn E o cell =E o red_Cu2+/Cu + E o oxidation_Zn E o cell =(+0.34V) - (-0.44V)= +.078 V E o cell =(+0.34V)+(+0.44)= +0.78V

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A typical alkaline battery has the following half reactions, identify the cathode and anode and write the complete reaction and find E 0. E 0 red = -1.28V E 0 red = 0.15V One needs to be oxidized, when added together E red and E oxid need to add to be positive Reverse so ZnO (aka Zn 2+ ) needs to be oxidized and is therefore it is the anode E 0 ox = +1.28VE 0 red = 0.15V E 0 cell = 0.15V-(-1.28)= 1.43V Logic 1: Logic 2: E 0 red anode = -1.28V E 0 red cathode = 0.15V E 0 cell = 0.15V+(+1.28)= 1.43V Example: Calculating Cell Voltage

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Review Cell Voltages are calculated as the difference in potential energy between the anode and cathode. (Pick your favorite way of thinking about it, and stick with that) A spontaneous cell will have a positive E o cell, while a non- spontaneous cell will have a negative E o cell.

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Electrochemistry:What will react? By looking at the table, which species will react with each other.

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Learning Outcomes Use what we already know to determine how you can tell what species will readily react spontaneously by looking at the table of E o red values.

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Diagonal rule: Species on left reacts with any species on the right that is lower than it. Diagonal Rule Why? Easily reduced Easily oxidized

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Example: Lets take 5 reactions Draw arrows next to the table excerpt identifying the most likely substances to be reduced or oxidized, and the best oxidizing and reducing agents. Give one spontaneous combinations, and one non-spontaneous reactions, using the above reactions (lots of available options). Easily reduced Easily oxidized

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Review On a standard reduction table with the highest reduction values listed on top, species on the left will react with species on the right of a reaction that is lower than it. This is because species that are easily reduced make good oxidizing agents.

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Cell Notation

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Learning Outcomes Write the cell notation for a given cell. Given the cell notation, write the half reactions.

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Cell Diagram/Notation Single line denote phase change Double line denote salt bridge Start with Anode on far left, work forward in order youd encounter it

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Cell Diagram Examples Write the cell diagram for a cell consisting of Zn 2+ /Zn and Pt/H + /H 2 Anode Cathode

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Review Example Combining everything weve learned.

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Zinc/Tin Cell Example The standard reduction potential of a zinc electrode is -0.76 V. Given that the standard potential of the cell where Zn is oxidized and Sn 4+ is reduced to Sn 2+ is +0.91V find the standard potential of the Sn 4+ /Sn 2+ half reaction and write the cell diagram.

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Electrochemistry and Thermodynamics

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Learning Outcomes Using thermodynamic data (gibbs free energy or enthalpy and entropy) calculate the E o cell. Calculate the Gibbs free energy of a cell using E o cell. Calculate the equilibrium constant from E o cell. Calculate E o cell from the equilibrium constant.

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Thermodynamics: Gibbs Free Energy We can use this to find the Gibbs free energy of a reaction if we can measure/find E cell (or of course vice versa) n= number of molesF=faradays constant= 96,500 J/(V*mol)

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Example Calculate the standard free-energy change for the following reaction at 25 o C using standard reduction potentials.

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Thermodynamics: Gibbs Free Energy What else does G equal? How can we use this to relate K and G? Set them equal to each other. Cleaning the equation up a bit:

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Thermodynamics G and K Now put all the constants on one side Filling in constants and solving

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Example- Relating E and K; Calculate the pressure of H 2 in atm, required to maintain equilibrium with respect to the following reaction at 25 o C, Given that [Pb 2+ ]=0.035 and the solution is buffered at pH 1.6

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Nernst Equation Think way back to thermodynamics in Chem1B, how did we relate G and G 0 ? But using the relation between G and E.. Rearranging OR NOTE: The difference between E and E o or G and G o is that the o symbolizes standard state

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Example Calculate the reaction quotient Q for the cell reaction, given the measured values of the cell potential.

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Review We can relate between E, G and K by using one of the three equations we introduced in this section. For non standard conditions, use the Nernst equation.

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Applications of electrochemical cells.

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Learning Outcomes Introduction to two types of batteries that are commonly used. Specifics of these will not be tested. I.e., dont memorize the reactions. Identify anode and cathode of batteries given the reactions that are occurring. Introduction to fuel cells. Identify the anode and the cathode of a given fuel cell

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Dry Cell and Alkaline Batteries Dry Cell: Originally developed Problems with corrosion and unstable current and voltage lead to the development of the alkaline dry cell Used in a variety of applications, but are not reachargable, limiting their utility in many others Alkaline dry cells replace NH 4 Cl with NaOH or KOH Dry Cell Anode Cathode

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Lead Storage Batteries: Discharged Charged Anode Cathode Overall i.e car battery Grids provide large surface area, low specific energy- allows high current for short periods

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Batteries Primary cells: non-rechargable Secondary cells: rechargeable Specific energy: energy that can be generated divided by mass Dry cell and alkaline batteries Lead acid cell: car battery Grids provide large surface area, low specific energy- allows high current for short periods

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Fuel Cells Runs similarly to a battery However, reactants are supplied continuously and only products are H 2 O Leads to less waste and less weight Originally used space applications, now provides backup power to many industrial applications, as well as fuel cell cars.

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Review, Recap and Note There are many many types of batteries. Ive shown you two variations. Batteries can be tweaked to provide more voltage, more current or various other desired outcomes. Suggestion, look up some emerging research on batteries that interest you and see if you can pick out how the reactions work, and are similar or different to the ones we talked about.

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Electrolytic Cells Review and Applications

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Learning Outcomes Review the difference between an electrolytic and galvanic cell. Determine the products at the anode and cathode of an electrolytic cell.

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Electrolytic cells: Reminder A cell that requires an outside source of power. i.e. a non spontaneous cell. These can be separated, like in previous cells, but can also be located in one container. Power must be supplied!

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More about electrolytic cells Anode and Cathode are often in the same container and often only contain one electrolyte. Reduction still occurs at cathode and oxidation still occurs at the anode. Many uses, electrolysis, electroplating, ore purification, ect

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More about electrolytic cells Reminder: E, G, K are all calculated the same. Non-spontaneous so E is negative. Applications are the new: interesting part!

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Electroplating Anode Cathode

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Electroplating Example How many grams of aluminum can be deposited by the passage of 105 C through an electrolytic cell? How long does it take to deposit 0.63 g Ni on a decorative drawer handle when 8.7 A are passed through a Ni(NO 3 ) 2 solution Use dimensional analysis. Important things to know: 1 C= 1 amp*sec F=9.6485x10 4 C/mol

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Corrosion How can we stop this? Give the O 2 something else to react with!!!! Fe 3+ precipitates out as Fe 2 O 3 H 2 O

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Stopping Corrosion: Galvanizing Galvanizing: Put a Zinc coating on it. Zinc is more reactive, it becomes the anode, Aka gets oxidized instead of the Iron Zinc is sacrificial metal Boats, wire, roofing, anything where you need to stop corrosion. E o Zn2+/Zn = -0.76 E o Fe2+/Fe = -0.44 E o Fe3+/Fe2+ = +0.77

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Review Electrolytic cells are much like galvanic cells only because they are non-spontaneous, they must have a power source. The E o is negative. Electrolytic cells are used for applications such as galvanizing or electroplating metals.

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Many important biological reactions involve electron transfer. Because the pH of bodily fluids is close to 7,the biological standard potential of an electrode E, E*, is measured at pH=7. a) Calculate the biological standard potential for the reducion of hydrogen ions to hydrogen gas, and the reduction of nitrate ions to NO gas. Calculate the biological standard potential E* for the reduction of the biomolecule NAD+ to NADH in aqueous solution. The reduction half reaction under thermodynamic standard conditions is NAD + (aq)+H + (aq)+2e - NADH (aq), with E o = -0.099V. The pyruvate ion, CH 3 C(=O)CO 2 -, is formed during the metabolism of glucose in the body. The ion has a chain of three carbon atoms. The central carbon atom has a double bond to a terminal oxygen atom and one of the end carbon atoms is bonded to two oxygen atoms in a carboxylate group. Draw the Lewis structure of the pyruvate ion and assign a hybridization scheme to each carbon atom. The lactate ion has a similar structure to the pyruvate ion, except that the central carbon is no attached to an OH group: CH 3 CH(OH)CO 2 -. Draw the Lewis structure of the lactate ion and assign a hybrization scheme to the central carbon atom. During exercise the pyruvate ion is converted to lactate ion in the body by coupling to the half reaction for NADH given above. For the half reaction pyruvate+2H + +2e - lactate, E*=-0.190 V. Write the cell reaction for the spontaneous reaction that occurs between these two biological couples and calculate E* and E o for the overall reaction. Calculate the standard Gibbs free energy of reaction for the overall reactions in the above reaction. Calculate the equilibrium constant at 25 o C for the reaction. Long electrochem problem (done on podcast)