Chapter 20 - Electrochemistry Flashcards
Balancing Redox Reactions (Acidic)
- Divide equation into half reactions
- Balance Half reactions (Add H2O to balance Os and add H+ to balance Hs.
- Multiply each side by an integer to make sure electrons are equal
- Combine half reactions to form final redox reaction and cancel species
Balancing Redox Reactions (Basic)
1.Divide equation into half reactions
2. Balance Half reactions (Add H2O to balance Os and add H+ to balance Hs.
3. For every H+, add an OH- to each side of the equation. For every H+ and OH- pair, simplify it to H2O
4. Multiply each side by an integer to make sure electrons are equal
5. Combine half reactions to form final redox reaction and cancel species
Voltaic Cell
a device where the flow of electrons goes through an external circuit instead of directly between reactants in the same reaction vessel
Anode
(-), the electrode at which oxidation occurs
Cathode
(+), the electrode at which reduction occurs
Electron direction
Electrons always flow from Anode (where oxidation occurs) to Cathode (where reduction occurs) in spontaneous reactions
Salt Bridge
allows for ion migration between two substances to maintain neutrality of charge
Volt is equal to…
J/C
(Joules/coulombs)
Electron value
1.6 x 10^-16 C
1 mol = 96485 C
Cell potential
Written as Ecell, measures the potential difference between the two electrodes of a voltaic cell. It is measured in volts. This value should always be positive.
Standard Cell Potential
Written as E°cell, this measures the potential difference between the two electrodes of a voltaic cell in standard conditions (1atm, 0° C, 1 M).
The standard Cell Potential is the same as the Standard Reduction potentials, meaning that it can be used to find the values of the standard reduction potentials of the half reactions that make up the whole reaction, using the equation
E°cell = E°red.(cathode) - E°red.(anode)
Electrical Potential
Measures the potential energy per electrical charge. Changing the mole value of the equation does not change the potential energy value, as the ratio would remain the same.
Ex:
Zn + 2e- -> Zn E°red = -0.76 V
10Zn + 20e- -> 10Zn E°red = -0.76 V
Half Reaction Tables
Tables list half reaction with an assigned E°red value.
The higher the value, the half-reaction will have a higher tendency to become reduced but lower tendency to become oxidized.
Alternatively, the lower the value, the half-reaction will have a lower tendency to become reduced but a higher tendency to become oxidized.
Knowing this, we can also apply our knowledge of oxidizing and reducing agents to this. Those who get reduced are considered oxidizing agents while those who get oxidized are considered reducing agents. Looking at the table, since half-reactions with a high E°red value will be easiest to reduce but most difficult to oxidize, thus meaning that they are potent oxidizing agents. Alternatively, half-reactions with a low E°red value will be hardest to reduce but easiest to oxidize, thus meaning that they are potent reducing agents.
Keep in mind that the oxidizing agents will be on the reactants side of the equation (the one with the electrons) while the reducing agents will be on the products side of the equation.
Oxidizing Agents
Halogens (F, Cl, Br, I, At, Ts)
Oxygen (O2)
Oxyanions (anions containing oxygen that follows the format of AxOyZ-)
These species will have positive values of E°red and thus easily undergo reduction
Reducing Agents
H2
Alkali metals (Li, Na, K, Rb, Cs, Fr)
Alkaline-Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)
These species have negative values of E°red and thus can easily undergo oxidation
