Redox equilibria -5
- Created by: Shannon
- Created on: 02-03-15 11:17
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- Redox equilibria
- Redox potentials/Electrode potentials
- Tell us the power or potential of an aqueous species to oxidise or reduce
- A +ve E*/V will produce a spontaneous reaction
- The more +ve the value, the more likely it is to happen
- A +ve E*/V will produce a spontaneous reaction
- Occur under standard conditions
- 298 K
- 1 moldm-3
- 1 atm
- Hydrogen half cell is used as a reference point = 0.00 E*/V
- Hydrogen electrode attached to the electrode system being investigated
- Connected via a salt bridge
- Always the left hand component of a cell
- Hydrogen electrode attached to the electrode system being investigated
- The plate making contact with the solution must be platinum if the reduced form of the system we are studying i not a conducting metal
- E* values produce a way of comparing positions of equilibrium when elements lose electrons to form ions in solution
- The more -ve the E*, the further to the left the equilibrium lies
- More likely to lose electrons to form ions
- The more +ve the E* the further to the right the equilibrium lies
- Less likely to lose electrons to form ions
- The more -ve the E*, the further to the left the equilibrium lies
- Why is a reaction not feasible is E*/V +ve?
- Reaction takes a more favoured route than the 1 we are investigating
- Activation energy is too high
- Conditions are not standard
- Tell us the power or potential of an aqueous species to oxidise or reduce
- Cell short hand notations
- Oxidation takes place at the ANODE
- Reduction takes place at the CATHODE
- Phase boundary represented by '/'
- Salt bridge represented by '//'
- If two species are in the same phase, they are separated using a comma, e.g Cu+,Cu2+
- Conventionally, the hydrogen half cell is always written on the left
- Reduction & Oxidation
- Reduction is a gain of electrons
- A reducing agent donates electrons to the oxidising agent it reduces
- Oxidation is a loss of electrons
- An oxidising agent accepts electrons from the reducing agent it oxidises
- Reduction & oxidation can be represented by half equations
- These equations can be combined to give a full equation
- Reduction is a gain of electrons
- Hydrogen fuel cell
- Cell set up - polymer electrolyte membrane between anode and cathode, only allows H+ to pass through, e- have to travel through circuit to get to cathode - creating an electric current
- Advantages
- Efficiency almost 100% - more efficient over normal car batteries
- Only waste product is water - less pollution than normal car batteries
- Disadvantages
- Production of H2 requires more energy than the fuel cell provides
- Expensive and limited life-span
- Alcohol fuel cell
- Either uses methanol or ethanol - as H2 rich fuels
- Made from renewable biomass
- Electrons transported through an external circuit from anode -> cathode
- Advantages
- Higher hydrogen density the H2 fuel cell
- Don't need special refrigeration like H2 - as liquid at room temp
- Disadvantages
- Low efficiency as alcohols can pass through membranes
- Methanol is toxic
- Either uses methanol or ethanol - as H2 rich fuels
- Breathalysers
- Modern breathalysers VS Police breathalysers VS Old breathlysers
- Modern breathalysers are easily portable and accurate
- Ethanol fuel cell - amount of alcohol in breath in proportional to the current produced when breath is fed to the anode of the cell
- Police breathalysers use IR to detect the presence & quantity of ethanol
- Accurate but not easily portable
- Old breathlysers used the reduction of potassium dichromate
- Modern breathalysers are easily portable and accurate
- Modern breathalysers VS Police breathalysers VS Old breathlysers
- Redox potentials/Electrode potentials
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