Redox wace Flashcards
Define oxidation-reduction reactions
What is the new and old theory associated with redox reactions?
Old theory of loss/gain of hydrogen and oxygen has been changed to the transfer of one or more electrons from one species to another.
Oxidation-losing electrons
Reduction-gaining electrons
How is a redox reaction kept constant/ion conc kept constant
ions are produced at anode at the same rate they are consumed by cathode
Define what is meant by half equation
redox reactions in 2 parts-oxidation and reduction
Define oxidation numbers and the rules for their identification
Oxidation numbers can be used to determine whether a reaction not involving formation of ions can be classified as redox reaction
1.Elements (Na, C, Cl2, P4)=0
2.Monatomics ions (Na+, Cl+)=charge on the ion
3.ON of oxygen in a compound is -2 (e.g. H2O) except in peroxides (e.g. H202 and BaO2=-1) and (OF2=+2)
4.ON of hydrogen is +1 except in metal hydrides (-1)
5.Polyatomic species=charge on ion
6.Neutral molecules – addition of ON’s = 0
Explain oxidants and reductants
Oxidant/oxidising agent: undergo reduction + gain an electron in a redox reaction= electron acceptor
Reductant/reducing agent: undergo oxidation + lose an electron in a redox reaction = electron donor
Describe features of metals on the table at back of yellow booklet
Metals higher on the list are less reactive therefore weaker reducing ability (ability to act as reducing agents).
Metals have small number of valence electrons, small energy required to move, more readily to act as agent.
Describe metal displacement
What occurs when metals are below Zn and below Cu
Act as reductants, donate electrons to other substances. More reactive metals will oxidise by donating electrons to less reactive metal ions
More reactive metals will displace less reactive.
Metals below Zn will react with water, and metals a below Cu will react with acid
Describe halogen displacement
Halogens in group 17 become oxidants as they remove e- from another substance
Decrease in oxidising strength as periodic table go down i.e. fluorine strongest, iodine, weakest
Halogen will oxidise a halide ion lower on list e.g. chlorine can oxidise bromide or iodide but not fluorine
Describe combustion
Not look like redox reaction. If we analyse the oxidation state, we can see if redox has occurred e.g.
CH4(g) + 2O2(g) CO2(g) + 2H2O (l)
-4+1 + 0 +4-2 + +1-2
CH4 has been oxidised and 2O2 has been reduced
Describe the electrochemical cells and explain the two types
Electrochemical cells: allow transformation of energy between chemical energy and electrical energy
There are two types:
-Galvanic cell: transforms chemical energy into electrical energy to produce an electric current to form a spontaneous (natural) redox reaction e.g. battery, fuels
-Electrolytic cell: transforms electrical into chemical energy to provide the energy to allow a non-spontaneous redox reaction to occur. e.g. electroplating, coating metals
Describe a galvanic cell in terms of the reactions occurring at the anode and cathode, the role of the electrolyte, the salt bridge, ion migration and electron flow.
Works by transferring electrons through a wire (external current/circuit) from reductant to oxidant. Doesn’t allow for direct contact between reactants, two half reactions occur in both compartments
-Galvanic cell therefore consists of two half cells
-Each contains an electrode in contact with solution
-Species present consists of oxidising agent and corresponding reduced form
The electrolyte provides the path for the flow of electrons or ions inside the cell.
At anode/negative: oxidation releases electrons at electrode flow to cathode through wire
At cathode/positive: electrons are accepted by cations when they collide with electrode, reduction occurs
Anions: migrate through salt bridge to anode due to the electrode’s loss of electrons, opposites attract, therefore negative ion migrates to more positive electrode.
Cations: migrate through salt bridge to cathode due to the electrodes gaining of electrons, opposites attract, therefore positive ions migrates to more negative electrode
How are galvanic cells different for involving gases
Inert electrodes are required (picture with entrance tube)
Define the rules for cell notation with an example
e.g. anode: Ni(s)Ni2+(aq)+ 2e- cathode: Fe3+ (aq)+ e- Fe2+(aq)
cell notation= Ni(s)|Ni2+||Fe3+, Fe2+|Pt (s)
-Anode reaction always to left vs cathode always right of salt bridge (||)
-Electrodes always far left+right of notation
-Different phases separated by single |
-Species of same phase separated by ,
Define cell potential
What is meant when there is a greater difference in cell potential
Cell potential/electromotive force (emf)=the flow of electricity in galvanic cell due to the difference in electrical potentials of it’s two half cells=voltage of cell
-Also referred to as potential difference because Volt (V)= unit of elec +cell potential
-Greater difference in value=larger eqm constant and reaction more likely to proceed in predicted direction
What are limitations to direct redox predictions?
- Only applies to (aq) solutions
- Value of E° depends on STP
- Half cell potentials can vary under other conditions
- E° has no indication of reaction speed (large E° does not mean large reaction rate)
