Redox Flashcards
Single Replacement Reaction
When one element replaces another element in a compound
A metal will replace a less reactive metal ion
Element + Compound —> New Element + New Compound
Single Replacement Reaction Example:
Mg (s) + CuCl(2) (aq)
(Balanced Equation, Complete Ionic Equation, Net Ionic Equation)
Balanced Equation:
Mg (s) + CuCl(2) (aq) —> MgCl(2) (aq) + Cu (s)
Complete Ionic Equation:
Mg (s) + Cu(2+) (aq) + 2Cl(-) (aq) —> Mg(2+) (aq) + 2Cl(-) (aq) + Cu (s)
Net Ionic Equation:
Mg (s) + Cu(2+) (aq) —> Mg(2+) (aq) + Cu (s)
Redox Reaction
A redox reaction involves the transfer of electrons from one chemical species to another
Reduction
When a chemical species gains electrons
Oxidation
When a chemical species loses electrons
Half Equation
A half equation represents either an oxidation or reduction half of a redox equation including a loss or gain of electrons
Reduction Half Equation Example:
O(2) (g)
O(2) (g) + 4e(-) —> 2O(2-) (s)
Oxidation Half Equation: Mg(2+) (s)
Mg (s) —> Mg(2+) (s) + 2e(-)
To balance with reduction half equation:
2Mg (s) —> 2Mg(2+) (s) + 4e(-)
Complete Half Equation Example:
2Li (s) + Br(2) (l) —> 2LiBr (s)
Oxidation Half Equation:
Li(s) —> Li+ (s) + e(-)
Reduction Half Equation:
Br(2) (l) + 2e(-) —> 2Br(-) (s)
To Balance Half Equations:
2Li(s) —> 2Li+ (s) + 2e(-)
Full Equation:
Br(2) (l) + 2Li(s) —>2BrLi (s)
Oxidising Agent or Oxidant
A chemical species that causes another chemical species to be oxidised
Reducing Agent or Reagent
A chemical species that causes another chemical species to be reduced
SEP Table
Metals at the top are most reactive , most likely to oxidise, strongest reducing agents
Reactivity for a Redox Reaction to Occur
A more reactive metal will be oxidised by a less reactive metal cation (the metal donates its electrons and the cation receives the electrons). Therefore, for a spontaneous redox reaction to occur, the metal ions must be less reactive than the solid metal
Example of Reactivity for a Redox Reaction to Occur:
Copper wire is placed in silver nitrate solution
Cu is more reactive and therefore more likely to oxidise than Ag, so the redox reaction will take place.
Reduction Half Equation:
Ag(+) (aq) + e(-) —> Ag (s)
Oxidation Half Equation:
Cu (s) —> Cu(2+) (aq) + 2e(-)
Full Equation:
2Ag(+) (aq) + Cu (s) —> 2Ag (s) + Cu(2+) (aq)
Potential Difference
Potential difference exists between two half-cells connected by an external wire.
Potential difference has the symbol E and unit of volts. It is measured using a voltmeter under standard conditions
Standard Conditions
Pressure 100k Pa
1M concentration for solutions
Temperature 298K (25 degrees)
Oxidation States
Oxidation states represent the charge that an atom would have if it was an ion
Determining whether a redox reaction has occurred using oxidation states
If there is no change in oxidation numbers for all atoms in a reaction it is not a redox reaction
Increase in oxidation state
An increase in oxidation state means an atom has been oxidised
Decrease in oxidation state
A decrease in oxidation state means an atom has been reduced
Oxidation rule for a free element
Oxidation state of a free element is 0
Oxidation rule for a simple ion. E.g. Na(+)
The oxidation number of a simple ion is equal to the charge of the ion
Oxidation rule for compounds (main group metals, hydrogen, oxygen)
Main group metals have an oxidation number equal to the charge of their ions
Hydrogen has an oxidation number of +1 in compounds with non-metals
Oxygen has on oxygen number of -2
Balancing half equations using the half reaction method
- Write the skeleton oxidation and reduction half equation
- Balance all the elements except for hydrogen and oxygen in the half equations
- Balance the oxygen atoms using H(2)O
- Balance the hydrogen atoms using H(+)
- Balance the charges by using electrons. The total charge on the left-hand side should equal the total charge on the right-hand side
- Add the reactions together, cancelling the electrons and any other elements that are present on both sides (e.g. H(2)O and H(+))
