3/14 Redox Flashcards
Oxidation
loss of e-
reduction
gain of e-
O.N rules
1) uncombined elements have an O.N of 0
2) overall OS of compounds is 0
3) elemental ions ahve an O.S equal to charge
4) overall OS of moelcular ion equal to charge
5) group 1: always +1
6) group 2: always +2
7) aluminium: always +3
8) hydrogen: always +1 (except in metal hydrides, eg NaH where it is -1)
9) oxygen: always -2 (except in peroxides, eg H202 where it is -1)
10) group 7: always -1 (except when combined with a more electronegative element)
disproportionation
element simultaneously oxidised and reduced
oxidising agent
what has been reduced, electron acceptor
reducing agent
what has been oxidised, electron donor
redox in acidic conditions
any oxygen lost is H20 in products, add H+ ions to balance the H in H20
half cell
1) made up of a solid metal electrode and 1.0 mol dm-3 solution of X ions
2) an eq is set up between X(s) and X ions (aq)
3) eq can shift: forward reaction is reduction, reverse reaction is oxidation
combining half cells
1) any two half cells can be combined to create a potential difference (voltage)
2) put most negative E0 value at top
3) apply anticlockwise rule
4) high resistance voltmeter
5) salt bridge (AgNO3) completes circuit
LOAN
LHS
oxidation
anode
negative
RRCP
RHS
reduction
cathode
positive
convention
oxidised state on left and reduced state on right
double line to represent salt bridge
single line to represent different phase
half cells of ions in solution
Fe3+(aq) + e- –>
standard hydrogen electrode
1) the redox power of half cells is relative
2) redox potential of all half cells measured against the SHE
3) platinum:unreactive and porous: large SA fro e- transfer
4) SHE connected to another half cell and potential difference measured
5) 2H+(aq) + 2e- –> H2(g) 0.00v
6) 298K, 1.0mol dm-3 H+ ions eg HCL, H2 gas pumped in at 100kPa
down the electrochemical series
1) increase in oxidising power
2) eq lies to right
3) ions readily gain electrons
4) easily reduced
5) last one is strongest O.A Ag+ (aq)
up the electrochemcial series
1) increase in reducing power
2) eq lies to left
3) atoms readily lose e-
4) easily oxidised
5) first one is strongest R.A Li (s)
calculating emf
most positive - most negative
oxidised cell value - reduced cell value
emf always postivie
greatest EMF value
largest difference in E0 values
smallest EMF value
smallest difference in E0 values
Li(s)
stronger reducing agent
most easily oxidised
most negative E0 value
Ag+ (aq)
strongest oxidising agent
moest easily reduced
most postive E0 value
Predicting feasibility of reactions
predictin if reaction will occur (spontaneous)
only happens if reactants are top right and bottom left
limitations to predicting feasibility using E0 values
all E0 values are based on standard conditions: if any of these change then the Ecell values will also change
eg Cu will react with conc h2so4 (10moldm-3) but not with 1 moldm-3
kinetics (rate) of reaction may be so slow that no real reaction is seen due to high AE
effect of changing concentration
1) chateliers principle: when conc or temperature changed
2) increase conc of reactant and decrease conc of product: Ecell increases
3) decrease conc of reactant and increase conc of product: Ecell decreases
4) no E0 value because deviated from standard conditions
effect of changing temperature
1) vast majoity of cells with a postiive E0 cell value (spontaneous) will have a negative deltaH (exothermic)
2) increase in temp shifts eq to left favouring endothermic reaction (against the expected reaction) decreasing Ecell
3) decrease in temp shifts eq to right favouring the exothermic reaction (with the expected reaction) increasing Ecell
disproportionation and E0 values
Cu2+(aq) + Cu+(aq)–> Cu2+(aq) + Cu(s)
cu+ oxidises and reduces itself
of the reaction is feasible then the reaction can occur
storage cells (Zn, C batteries)
1) at the anode (LHS), zinc oxidises to Zn2+ and releases e-
2) e- flow through deice and through to the carbon cathode
3) from the carbon cathode flow through the graphite rod and leave into the electrolyte
4)electrolyte: NH4Cl
5) non rechargeable: zinc used up walls get thinner over time and may leak
6) battery runs out when NH4+ runs out
7)Zn2+ + 2e- –<>— Zn
2NH4+ + 2e- —> 2NH3 + H2
Relationship between Ecell and entropy
if E0 cell is positive, it is considered a spontaneous/feasible reaction - delta G must be negative
the more positive E0 cell the more negative delta G
the more positive delta S, the more negative delta G
the more negative delta G, the more positive E0 cell
therefore delta S and E cell are directionally proportional
relationship between E cell and lnK
lnK and E0 cella are directionally proprotional
the larger the lnK the more positive the E0 cell
the lithium ion battery: during use
1) copper anode on left (not involved in redox), next to it is graphite to allow flow of e- (store of delocalised e-)
2) aluminium cathode on right(not involved in redox), next to it is lithium cobalt oxide Li+(CoO2)-
3) lithium oxidises to lithium ions and loses e- which flow through graphite and cu anode through device
4) lithium ions move through electrolyte to the cobalt oxide near the right and gain e’s that flow to the aluminium cathode to form Li+(CoO2)-
recharge: opposite reactions
advantages of lithium ion battery
1) store lots of energy
2) small
3) high voltage - powerful
4) fast recharge
disadvantages of lithium ion battery
1) lithium very reactive
2) lithium very toxic
3) high production cost
4) becomes less efficient over time
fuel cells
reaction between hydrogen rich fuels such as ethanol and oxygen can be utlisied to generate a voltage
1)liquid ethanol and liquid water enter left into a redox reaction releasing hydrogen ions and e- and co2 emitted as a waste product
2)e- build up inside anode causing an electrode potential thus e- flow through device
3) hydrogen ions flow through proton membrane in the middle of the fuel cell
4) oxygen flows thorugh inlet on the rHS where it reacts with the hydrogen ions that have come through and the e- that have come through the cathode which creates water as a by product
5)oxidation at the anode
reduction at the cathode
6)C2H5OH + 3H2O –>12H+ + 12e- + 2CO2 Oxidation
3O2 + 12H+ + 12e- –>6H2O reduction
C2H5OH + 3O2 –> 3H20 + 2CO2
hyrdogen fuel cell alkaline electrolyte
1)H2(g) goes in at the anode
2)O2 gas goes in the cathode
3) at the anode, the hydrogen combines with hydoxide ions to form water and electrons (oxidation)
2H2+40H- –> 4h2O + 4e-
4) water is expelled as a by product whilst the e-‘s flow through anode through the device
the e-‘s come through the cathode and combine with oxygen and water to form hydoxide ions (reduction)
o2 + 2H2O + 4e- –> 4OH-
5) this is how OH- ions are produced for the reaction at the anode, these OH- ions act as an alkaline electrolyte which can move through the membrane
6) these OH- ions dont actually get used up as 2H2 + 02 –> 2H2O
hydrogen fuel cell for acidic electrolyte
1) anode: 4H+ +4e- —> 2H2O
Hydrogen ions move from left and right with acidic electrolyte compared to alkaline electrolyte where OH- ions move from right to left
Pros of hydrogen fuel cell
1) 0 emissions just water and no greenhouse gases
2)highly efficient
3)h2 + o2 readily available
4)maintain a constant voltage as voltage in traditional cells tend to decrease near end of life
5)number of options for hydrogen storage
liquefied under pressure
adsorbed on a solid medium
absorbed in a solid medium
Cons of a hydrogen fuel cell
1) very expensive to produce the fuel cells
2) toxic chemicals used in production
3) h2 comes from h2o via electrolysis which is expensive and energy used may come from fossil fuels or h2 may come from hydrocarbons
4) storage of h2 is dangerous
5) limited life span
