Electrochemistry

Electrochemistry (Chapter 19, pages: 817-847; Review chapter 4, section 4)

1. Be able to define the following concepts: a. Oxidation: always occurs in the presence of reduction. Oxidation is a loss of electrons. increase in oxidation state. Oxidation occurs at the Anode. A good way to remember this is "An Ox". Also, the chemical being oxidized is the reducing agent. b. Reduction Reduction is a gain of electrons.consequentely decreases in oxidation state. Reduction happens on the Cathode. A good way to remember this is "Red Cat". Also, the chemical being reduced is the oxidizing agent. c. Electrochemical cells A device that allows the exchange between chemical and electrical energy. The device is composed of two chambers (one holding the chemical to be oxidized and the other holding the chemical to be reduced), a salt bridge, an electrode in each chamber, and a wire or wires connecting the two electrodes. d. Voltaic cell/Galvanic cell: It is a cell in which a spontaneous chemical reaction generates a flow of electrons. A voltaic or galvanic cell has a potential greater than zero. Electrical cell that produces electricity e. Electrolytic cell An electrochemical cell used for electrolysis. An electrolytic cell is an electrochemical cell that releases electricity but won't happen without intervention, because the process is non-spontaneous. The potential for an electrolytic cell is negative. f. Fuel cell Electrical chemical cell that utilizes hydrogen with oxygen to produce electricity, are not really batteries but also supply electric energy via a spontraneous redox reaction. g. Half reaction:is a chemical equation representing only the oxidation or only the reduction of an oxidation-reduction reaction. Zn p Zn2+ + 2 e- Cu 2+ + 2 e-   ® Cu  " A half reaction is either the oxidation or reduction reaction component of a redox reaction. A half reaction is obtained by considering the change in oxidation states of individual substances involved in the redox reaction ." h. Anode: positively charged electrode by which electrons leave an electrical device. " The electrode where electrons are lost (oxidized) in redox reactions. " i. Cathode: negatively charged electrode by which electrons enter an electrical device. " The negatively charged component of an electrowinning cell, where valuable metals are deposited " h. Oxidation state A number assigned to an element in a compound according to some rules. These rules are: Fluorine always has an oxidation number of -1. Group 1A and 2A are +1 or +2, respectively. Hydrogen usually is +1. Oxygen is usually -2. Group 7A (other than F) is usually -1. (The difference here is that Fluorine ALWAYS is -1).

k. Oxidizing agent The oxidizing agent is the substance that causes another substance to lose electrons/become oxidized. So, the oxidizing agent itself will gain electrons/become reduced. l. Reducing agent The reducing agent is the substance that causes another substance to gain electrons/become reduced. The reducing agent itself will donate electrons/become oxidized. m. Disproportionation- A reaction in which the same element is both oxidized and reduced is called a disproportionation reaction. I n. Cell potential - Experimentally the difference in electrical potential between the anode and the cathode is measured by a voltmeter and the reading is called the cell potential o. Standard reduction potential (SRP) - the potential associated with a reduction half-reaction at an electrode when the ion concentration is 1 M and the gas pressure is 1 atm. p. Standard oxidation potential (SOP) - the potential associated with a oxidation (losing electrons). q. Voltage- A measure of how strongly a species pulls electrons towards itself. Also known as electromotive force. Measures the strength of electricity in contrast with amperage(current) which is the measure of speed within the circuit. r. Cell notation a shorthand way of expressing a certain reaction in an electrochemical cell. "ELECTRODE l (anion)REACTANT l (anion)PRODUCT ll (cation)REACTANT l (cation)PRODUCT l ELECTRODE" Example: Cd(s), Hg(s) | CdCl2(aq)(0.010M) || AgCl(s), Ag(s) s. Salt bridge- A device containing a strong electrolyte that allows ions to pass from beaker to beaker. Able to balthe Anode and Cathode by using any salt. For example, say KNO3 is the salt "on" the salt bridge. NO3- will go to the anode to make the charge 0 and the K+ will go to the Cathode side to also make the charge 0. The salt bridge provides a way to neutralize the ions in the anode/cathode. t. Electrodes: they are conductors through which an electric current enters or leaves a substance whose electrical characteristics are being measured, used, or manipulated. The electrode is always a solid metal. If a solid metal is present in the half reaction, that metal will be used as the electrode. If not, we must use Pt or Pd as the electrode. Gets bigger on the cathode size and smaller on the anode size if a solid is being formed. Also this is there to increase the surface area for the reactions to happen.

u. Voltmeter - A voltmeter is an instrument used for measuring the electrical potential difference between two points in an electric circuit. Measure the velocity of how elctrons are running away from the anode to the cathode. " instrument that measures cell potential by draw ing electric current through a known resistance " v. Power supply- ﻿A device that supplies electrical energy to a system of electric load.

2. Be able to determine the oxidation state of each element in a complex The rules for determining oxidation numbers are: Flourine always has an oxidation number of -1. Group 1A and 2A are +1 or +2, respectively. Hydrogen usually is +1. Oxygen is usually -2, except a few exceptions such as in peroxides [|H2O2] Group 7A (other than F) is usually -1. (The difference here is that Fluorine ALWAYS is -1).

3. Be able to identify a reaction as an oxidation reaction, reduction reaction or neither. An example of an oxidation reaction is: Fe2+ yields Fe3+ (an electron is lost) An example of a reduction reaction is: 14H+ + Cr2O72- yields 2Cr3+ + 7H20 (electrons are gained making it a reduction) 4. Be able to break down a balance redox reaction into half reactions MnO4-C(aq)+Cl( aq) yields Mn2+(ag) +Cl2 MnO4- error Mn2+ 2Cl-error Cl2

a. Oxidation reaction- In a redox reaction, the half reaction that supplies electrons. This reaction occurs in the Anode. b. Reduction reaction- In a redox reaction, the half-reaction that acquires electrons. This reaction occurs in the Cathode.

5. Be able to express a balanced redox reaction using cell notation AnodeIICathode Zn (s) yields ZNO( s) = anode, Ag2O yields Ag (s)=Cathode cell notation is : Zn(s)I ZnO II Ag2O(s)IAg(s) if the anode and cathode are ions, you need to use pt,or pd as the electrode. This is added to the cell notation as ptI Zn,ZnOIIAg2O IAgIpt 6. Be able to draw a electrochemical and galvanic cell based on a balanced redox reaction species that are being oxidized goes on the anode side and the species that are being reduced goes on the cathode side. add the electrode, wires , salt bridge ,and flow of electrons. you are just working backwards.

7. Using a electrochemical cell diagram, be able to predict: a. Anode The anode is where oxidation takes place; where the substance loses electrons. This will usually be on the left hand side of the diagram it seems, but it doesn't have to be I believe. b. Cathode The cathode is where reduction takes place; where the substance gains electrons. b. Half reduction reaction: when trying to balance an equation we must identify the compound or element being reduce c. Half oxidation reaction: in a chemical reaction be able to determine the compound or element being oxidize in order to balance the equation d. Move of electrons through the wire: as we did in the lab electron move from left to right from the anode to the cathode. e. Salt bridge- A dispenser of salt that puts the anions in the anode, and the cations in the cathode f. Movement of ions through the salt bridge: The ANion in a salt moves to the ANode and the CATion moves to the CAThode g. Reducing agent- a species that donates electrons. A substance that causes the reduction of another substance. l. Oxidizing agent- a species that accepts protons. A substance that causes the oxidation of another substance. m. Ions in solution are ionic compound that dissolved completely into ion. they are completely ionize in water n. Electrodes is an electrical conductor used to make contact with a nometalic part of a circuit are always solid metals. If a solid metal is not available in the half reaction, use Pt or Pd. o. Electrodes increasing in size (if any) The electrode in the cathode side will increase in size if it is not Pt or Pd, but keep in mind that Pt or Pd will also increase in size IF a solid is being formed in the cathode because that solid will attach to it as it grows. p. Electrodes decreasing in size (if any) The electrode in the anode side will decrease in size if it is NOT Pt or Pd.

8. Be able to describe an electrochemical cell using cell notation Cell notation is written as following using the anode followed by the cathode electrode|reactant, product||reactant, product|electrode For example: Pt|Mn2+, MnO4-||BiO3-, Bi3+|Pt

9. Be able to design a galvanic or electrolytic cell using given half reactions with their corresponded potentials Use the SRPs or SOPs and the formula Ecell = Ecathode - Eanode to get a positive Ecell for galvanic cells and a negative Ecell for electrolytic cells.

10. Be able to balance a redox reactions under acidic and basic conditions ACIDIC 1) Balance all the atoms with coefficients 2) Balance the Oxygen atoms utilizing H2O 3) Balance the remaining Hydrogen with H+ 4) Balance the present electrons. 5) In order to cancel out the electrons you must multiply both half reactions by the lowest common factor. 6) Cancel out any remaining common factors

BASIC (added on to the previous process) 7) Once everything is balanced for an acidic solution, you must match the present H+ with OH- in order to produce H2O. 8) Cancel out any further common factors that remain.

11. Be able to mathematically relate potential values to change in gibbs free energy under standard and non-standard conditions. -∆G = spontaneous reaction, and ∆G=-nFE cell, so E cell must be positive because both “n” and “F” are positive quantities.

12. Be able to describe the fundamentals behind battery design Part of the idea behind battery design is a complete circuit, it is simply an electrochemical cell. When a circuit is completed, one end of the battery (such as a AA) transfers the charge through the circuit back to the other end of the battery. The charge going through is what supplies the power to the bulb of a flashlight or the power for a radio to play.

13. Be able to describe the following applications of electrochemistry: a. Electroplating- s a plating process in which metal ions in a solution are moved by an electric field to coat an electrode. The process uses electrical current to reduce cations of a desired material from a solution and coat a conductive object with a thin layer of the material, such as a metal. Examples are, abrasion and wear resistance, corrosion protection, [|l] ubricity, aesthetic qualities, etc.

b. Electron transport ﻿- the bio-chemical reaction between an electron donor, and an electron acceptor to transfer electrons (e-) across a membrane, that result in the liberation of energy

c. Corrosion, etc: is the disintegration of an engineered material into its constituent atoms due to chemical reactions with its surroundings. In the most common use of the word, this means electrochemical oxidation of metals in reaction with an oxidant such as oxygen. Formation of an oxide of iron due to oxidation of the iron atoms in solid solution is a well-known example of electrochemical corrosion, commonly known as rusting.