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Unit 20 - Electrochemistry - Coggle Diagram
Unit 20 - Electrochemistry
20.1 - 20.2 Oxidation States and Balancing Redox Reactions
In electrochemical reactions, electrons are transferred from one species to another using oxidation and reduction.
OIL RIG
“Oxidation is Loss; Reduction is Gain”
An oxidizing agent (or oxidant) is one that promotes oxidation
A reducing agent (or reductant) is one that promotes reduction
Oxidation number (ON)
The ON of an element when uncombined is always zero.
The oxidation number is the same as the charge for a monatomic ion.
ON of metals of main group compounds have the charge associated with their group
Steps using the Half-Reaction Method (in acidic solutions)
Assign oxidation numbers to determine what is oxidized and what is reduced.
Write the oxidation and reduction half-reactions.
Balance each half-reaction.
Balance elements other than H and O.
Balance O by adding H2O.
Balance H by adding H+.
Balance charge by adding electrons.
Multiply the half-reactions by integers so that the electrons gained and lost are the same.
Add the half-reactions, subtracting things that appear on both sides.
Check that the equation is balanced according to mass.
Check that the equation is balanced according to charge.
Balancing Equations in a Basic Solution using the Half-Reaction Method
•If a reaction occurs in basic solution, we balance it as if it occurred in acid.
•Once the equation is balanced, add OH− to each side to “neutralize” the H+ in the equation and create water in its place.
•If this produces water on both sides, you might have to subtract water from each side.
20.3 Voltaic Cells
In spontaneous oxidation-reduction (redox) reactions, electrons are transferred and energy is released.
This energy is used to perform work through a voltaic (or galvanic) cell as the electrons are transferred through an external pathway (wire)
In a voltaic (or galvanic) cell, oxidation always occurs at the anode and reduction happens at the cathode.
The driving force that pushes electrons through a voltaic cell is called electromotive force (emf) which is a type of potential difference or voltage.
E cell > 0 for all spontaneous reactions.
The strongest oxidizers have the most positive reduction potentials.
The strongest reducers have the most negative reduction potentials.
The greater the difference between the two, the greater the voltage of the cell.
20.4 Cell Emf Under Standard Conditions
For voltaic cells, the reactions must be spontaneous, having a positive emf, for the cell to work and generate electricity.
Eo = Eredo(reduction process) + (- Eredo(oxidation process))
When E is positive, the reaction is a spontaneous process.
When E is negative, the reaction is nonspontaneous.
Therefore, this pattern helps explain the activity series of metals based on ease of reduction and thermodynamic favorability.
ΔGo = -nFEo
n = number of moles of electrons transferred
F = Faraday’s constant = 96,485 J/V mol
E = emf (V)
20.5 Free Energy and Redox Reactions
At standard conditions, Q = 1 and a voltaic cell produces emf. However, as the concentrations change and the reaction proceeds towards equilibrium, the emf decreases until E = 0.
Nernst Equation (for nonstandard conditions):
20.6 Cell Emf Under Nonstandard Conditions
E = Eo - RT/nF * Ln Q
It is possible to build a voltaic cell that has the same substance at both the anode and cathode, just at different concentrations. These types of voltaic cells are called concentration cells.
20.7 - 20.9 Applications of Redox Reactions
Electrolytic cells rely on outside electricity sources to work and allow nonspontaneous reactions to occur.
An electrolytic cell consists of two electrodes in a molten salt or a solution (to allow for the flow of ions). An outside source of emf (like a battery) works as an electron pump.
Electroplating
Electric Current and Electrochemistry
Electric Current = Electric Charge / time
I = Q/t
Q = charge in Coulombs(C)
T = time in seconds (s)