Redox Reactions and Electrochemistry Flashcards Preview

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Flashcards in Redox Reactions and Electrochemistry Deck (15):

What type of reactions ether produce or use electrical energy?

Electrochemical Reactions



Loss of electrons



Gain of electrons


Which chemical oxidizes another while being reduced?

Oxidizing agent (as opposed to Reducing Agent)


Galvanic Cell

Spontaneous (negative ΔG) reaction that produces an electric current, where the anode is considered negative. Energy harnessing involves separating the two half reactions into two half-cells connected by a salt bridge. The salt bridge replenishes the necessary ions to the solutions of each half-cell according to Le Châtelier's Principle for a limited time. In a Daniell Cell, solid Zn in a solution of ZnSO4 acts as the anode, while solid Cu in a solution of CuSO4 is the cathode.


What occurs in a Daniell Cell?

The Daniell Cell is a galvanic cell, where the oxidation if zinc leads to the production of Zn2+ and the reduction if Cu2+ builds on the copper solid. The reactions have to be separated because the aqueous copper would react with the zinc bar directly. A wire must be added to provide the electrons with a place to flow. The salt bridge fights the charge gradient for some time, with the negative ion flowing into the Zn half cell and the positive ion flowing into the Cu half-cell.


Electrolytic Cell

An electrochemical cell that requires electrical energy to drive a non-spontaneous (positive ΔG) reaction. The anode is positive, because it attracts anions from the solution while attached to the positive pole of the battery. Electrons go through the anode.



The electrode where reduction occurs. Cathode attracts the cations. Electrons flow towards the cathode.



Electrode where oxidation occurs. Electrons flow away from the anode.


Reduction Potential

Measurement of the tendency of a chemical to be reduced. The more positive the reduction potential, the greater the tendency toward reduction. The Standard Reduction Potential is measured at 1 M or 1 atm and 25°C. The Standard Hydrogen Electrode has a potential f 0.00 V.


Standard Electromotive Force

EMF = E°red + E°ox = E°cathode - E°anode


Faraday Constant

One Faraday is equal to the amount of charge in one mole of electrons.
1 F = 9.65 X 10^4 C/ mol e-.

It = nF


How are the EMF and Gibbs Free Energy related?

ΔG dependent on the amount of energy available.
ΔG = -nFEcell
ΔG must be expressed in J, not kJ if Faraday's Constant is in Coulombs (J/V)
ΔG° = -nFE°cell under standard conditions


How does concentration of a species affect EMF?

EMF varies with concentrations. The Nernst Equation determined the EMF at non standard conditions: Ecell = E°cell - (0.0592/n) logQ = E°cell - (RT/nF) lnQ


How does EMF relate either the equilibrium constant?

ΔG° = -RT lnKeq

nFE°cell = RT lnKeq