Redox - Mel Flashcards
(29 cards)
Oxidation
- Loss of electrons
- Gain of oxygen / loss of hydrogen
- Increase in oxidation state numbers
Reduction
- Gain of electrons
- Loss of oxygen / gain of hydrogen
- Decrease in oxidation state numbers
Voltaic Cells
Voltaic cells generate electricity from chemical reactions that occur at the electrodes in the cell.
Voltaic Cells: Electrodes
Anode: Negative
Cathode: Positive.
Voltaic Cells: Salt Bridge
The salt bridge is there to allow ions to pass through from one half cell to the other so that the half cells do not develop an electrical charge.
Voltaic Cells: Disproportionation
A disproportionation reaction is a redox reaction in which a molecule, atom, or ion is simultaneously oxidised and reduced
Voltaic Cell: Flow
The direction from the cell with the most negative potential to the cell with the least negative (most positive) potential.
Voltaic Cell: Representing the cell
(solid/solution||solution/solid)
H2/H+||Cu2+/Cu
Cell potential (E)
Standard electrode potentials (E) are potential values determined by comparison with the Standard Hydrogen Electrode (SHE) which is the potential created by 1 mol dm-3 of hydrogen ions at 100kPa H2 at 198K, which has the E value of 0.00V.
(- It is equal to the difference between the electrode potentials of the two half cells.)
Cell potential (E) equation
- Turn the equation being oxidised around
- E= the more positive value - the more negative value
Standard Hydrogen electrodes - Diagram
- Inverted tube -> So that all the H2 reacts
- Inert electrode (platinum / graphite)
- 1 mol dm-3 (M) of H+ (aq) -> so for example, 1 M of HCl or 0.5 M of H2SO4 (because of the 2)
- 2H+ (aq) + 2e- <=> H2 (g)
- Cell potential (E) = 0.00V
- Strong acid as electrolyte
(- Salt bridge)
Data Booklet list
- Top right: Strongest reducing agents
- Bottom left: Strongest oxidising agents
Electrolytic Cells
This is, fundamentally, the reverse of a voltaic cell. In this case a greater electromotive force is applied from the external circuit using a battery or power source and this forces the species within the cell to perform the reverse reaction to what they would normally tend to do.
Electrolytic Cells: Electrodes
Anode: Positive
Cathode: Negative
Electrolytic Cells: Molten Solutions
Anode: Anion
Cathode: Cation
Electrolytic Cells: Aqueous solutions with inert electrodes
Electrode: Platinum or Carbon
Anode: a) If dilute: O2 is produced
b) If concentrated: Cl2 is produced
Cathode: H+ is produced, unless the metal in under H in the reactivity series.
Electrolytic Cells: Aqueous solutions with non-inert electrodes
Electrode: Copper
Anode: Electroplating takes place
Cathode: H+ is produced, unless the metal in under H in the reactivity series.
Electrolytic Cells: Electroplating
The ability of electricity force chemical change makes the deposition of metals possible in an electrolytic cell.
Electrode: Anything (Iron, Zinc, Silver)
Anode: Oxidation
Cathode: Reduction
Electrolytic Cells: Factors Affecting the amount of products
A) Charge on ion (Ag+ better than Cu2+)
B) Current flowing through the circuit
C)
Problems asking to find the mass produced
- Q=It
- Q / F => This is the moles of electrons
- Write the equation and find the mol ratio, and thus the mols.
- Do m= n * Mr to find the mass.
Connection of E with ΔG
A positive E cell corresponds to a negative ΔG making the reaction spontaneous, although the activation energy may still be too high.
Gibb’s Energy
- For redox reaction, Gibb’s energy is the electric energy, which, when properly set up in an electric cell, is the charge transferred (q) time the potential (E).
- Gibb’s energy is the maximum electric energy derived from a battery.
ΔG= -nFE
What is a battery
A battery is a portable electrochemical cell in which chemical energy is converted into electrical energy.
Primary Cells (battery)
- Carbon-Zinc dry cell
- Such as copper / zinc voltaic cell can only be used once as the redox reaction taking place is irreversible.
- Low cost / lightweight / convenient
- Environmental pollution / Can only be used once
