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(+))
Galvanic Cell
Also known as voltaic cell
Electrochemical cell in which chemical energy is converted into electrical energy
Electrical energy produced via spontaneous redox reactions within the cell
Parts of a galvanic cell
Two half cells: half reactions occur in separate half cells. Each half cell contains an electrode in contact with a solution (electrolyte).
Electrodes (anode and cathode): metal strips. Anode is where oxidation occurs, negative electrode. Cathode is where reduction occurs, positive electrode.
Solution (electrolyte): Contains ions. Conducts electricity as charged ions freely move. Soluble solutions that match electrodes ensure no competing reactions.
External circuit: Wire allows electron movement from anode to cathode.
Salt bridge: Contains ions that move to balance the charges formed in the half cells. Cations move towards cathode, and anions towards anode. Usually KNO(3) or NaNO(3).
Electrolytic Cell
Electrolysis occurs, producing a non-spontaneous redox reaction by the passage of electrical energy from a power supply through a conducting liquid.
Convert electrical energy into chemical energy (operate in reverse compared to galvanic cells)
Parts of an electrolytic cell
Porous barrier: half reactions occur in one container, porous barrier separates products
Solution (electrolyte): Contains a salt (made of positive and negative ions) that conducts electricity
Cathode: Reduction occurs, negatively charged
Anode: Oxidation occurs, positively charged
Voltage: Required voltage is higher than the voltage produced in an equivalent galvanic cell
! Understand that the ability of an atom to gain or lose electrons can be predicted from the atom’s position in the periodic table, and explained with reference to valence electrons, consideration of energy and the overall stability of the atom !
Redox conjugate pair
An electron donor and acceptor
Inert electrode
an electrode that serves only as a source or sink for electrons without playing a chemical role in the electrode reaction
E.g. carbon
