Kaplan General Chemistry Chapter 12 : Electrochemistry

What are the charges of anodes and cathodes in galvanic and electrolytic cells?

*Galvanic cell* anode: (-), oxidation cathode: (+), reduction *Electrolytic cell* (reversed charge) anode: (+), oxidation cathode: (-), reduction But anions are still attracted to the anode in both cases

Nernst equation

Ecell= E°cell - (RT/nF) (lnQ) Diagram is the simplified version of the equation @ 298K

Anode

Electrode at which oxidation occurs during a cell's redox reaction. Electrons always flow from the anode in an electrochemical cell. Can be positive or negative depending on type of cell

Cathode

Electrode at which reduction occurs during a cell's redox reaction. Can be positive or negative depending on type of cell

Electromotive Force and Thermodynamics

Electromotive force and change in free energy always have opposite signs

Charging

Electrons flow nonspontaneously as in an electrolytic cell

What matters in electrolytic cells?

Really only the electrolyte because the that's the substance being broken down to generate charge. The electrodes can be anything so long as they can resist the high temp and corrosion of process

Electrolytic Cells

Similar to galvanic cells except emf < 0 and delta g > 0, aka nonspontaneous Not separated because no work needs to be done, requires external voltage source e.g. rechargeable batteries

Concentration Cells

Specialized form of a galvanic cell in which both electrodes are made of the same material, rather than a potential difference causing the movement of a charge it is the concentration gradient b/w the 2 solutions

Galvanic Cells - Voltaic Cells

Spontaneous (delta G < 0 and emf > 0) Redox reaction is used to generate a flow of current Separation of reaction allows work to be done ex: daniel cell, most non-rechargeable batteries

Electrodeposition Equation (How do you calculate how much metal is deposited on a electrode in an electrochemical cell?)

Summarizes the faraday constant process and *helps determine the number of moles of elements being deposited on a plate. Mol M = (I)(t)/(n)(F) Mol M = moles of metal produced I = Current t = Time n = number of electron equivalents F = Faradays Constant Mnemonic: Calculating moles of Metal, it is not fun.

Cell Potentials

The likelihood to be reduced or oxidized

Daniel Cell

consists of half-cells where redox reactions occur, joined by a circuit and a salt bridge separation of reaction allows work to be done

Calculating delta G from E

delta G standard = -n*F*Ecell

Common things shared by all electrochemical cells

e- flows from anode to cathode (I/current flows cathode to anode) anode = oxidation cathode = reduction

Reduction Potentials

is a measure of the tendency of a chemical species to acquire electrons and thereby be reduced More positive number = more wanting to be reduced

potentiometer

kind of voltmeter that draws no current and gives a more accurate reading of the difference in potential between two electrodes

Energy Density

measure of a battery's ability to produce power as a function of its weight higher is better

Surge Currents

periods of large current early in the discharge cycle good for things that require rapid responses (i.e. remote control)

Standard Hydrogen Electrode - SHE

reduction potential of H2, by convention = 0V

Nickel-Metal Hydride Battery - NiMH

replaced Ni-Cd batteries because they have higher energy density, are more cost effective, and less toxic

Standard Electromotive Force -EMF

same as E cell Don't multiply

Electromotive Force - EMF

the force that results in the motion of electrons due to a difference in potential

galvanization

the process of protecting metals by coating them with a thin layer of zinc naturally results in a galvanic cell

voltmeter

A device used to measure voltage, or electrical potential energy difference

Electrodes

A metal strip that conducts electricity

Electrolysis

A process by which an electric current breaks chemical bonds. Occurs in electrolytic cells through the use of redox reactions

Salt Bridge

A tube that allows the slow transfer of ions and maintains the neutrality of the electrolyte solutions. circuit causes the flow of e- and the buildup of charge on one side (which would eventually cause the rxn to cease); salt bridge neutralizes the charge

Faraday Constant

Amount of charge per mole of electron ~10^5 C/mole e-

What is the flow of current

By convention, flow of positive charge Thus, cathode to anode

Electrochemical Cells

Cells that convert chemical energy into electrical energy stored in charges.

Nickel-Cadmium Battery

Discharging: Anode = Cd Cathode = NiO(OH) Charging: Anode = Cd(OH)2 Cathode = Ni(OH)2 higher energy density than lead-acid batteries

Oxidation potential

Never given, is the opposite sign of the standard reduction potential

Rechargeable Cell/Battery

One that can function as both a galvanic and electrolytic cell

Discharging

The state of a rechargeable electrochemical cell that is providing an electromotive force by allowing electrons to flow spontaneously from anode to cathode. Galvanic cell

Standard Reduction Potential

The tendency of a species to be reduced, as measured at 25 C when reacting species are of 1M concentration or 1atm partial pressure, aka standard conditions More positive number standard reduction potential means more likely to be reduced

Nernst Equation

Used to calculated potential generated by concentration difference Ecell= E°cell - (RT/nF) (lnQ)

Lead-Acid Battery

When fully charged (i.e. able to discharge) = galvanic Discharging: Anode = Pb Cathode = PbO2 When fully discharged (i.e. able to charge) = electrolytic Charging: Anode = PbSO4 Cathode = PbSO4

Cell Diagram

anode | anode solution || cathode solution | cathode ex: Zn(s) | Zn2+ (1M) || Cu2+ (1M) | Cu(s)