Redox II

• an electrochemical cell contains:

  • two half cells connected by a wire.
  • a salt bridge connecting the cells.
  • 1 moldm3 solutions (standard concentration)

  
• salt bridge - piece of filter paper soaked in salt solution (usually potassium nitrate):

  • This salt solution allows the flow of ions between the cells (balance out charges as electrons flow through the voltmeter and cause redox reactions)
  • the salt should be unreactive with the electrodes and electrode solutions.
  • a wire isn't used as this would set up its own electrode system
• electrochemical cells are often represented using a cell diagram. In a cell diagram:

  • a solid vertical line represents a change in phase (e.g. between the solid electrode and the electrolyte solution).
  • a double vertical line represents the salt bridge.
  • the voltage produced is indicated.
  • the more positive half cell (less reactive one since it's gets oxidised/loses electrons LESS easily) is on the right

  
• in an electrochemical cell, the more reactive metal gets oxidised more readily (releases electrons). These electrons then formulate on the electrode, making it more negative.
• in a spontaneous reaction, electrons flow from the negative to the positive electrode.
• electrode potential - the tendency for an electrode to gain or lose electrons in a redox reaction/how easily a half cell undergoes reduction or oxidation.
• The voltmeter has a very high resistance (to stop current flowing). If a current was allowed to flow, too much of a redox reaction would occur. An equilibrium would then be reached as the concentrations of reactants and products reach a point where both electrodes would have the same potential. Thus there's no potential difference which leads to a voltage of 0V.

  
• using a voltmeter in the electrochemical cell gives the potential difference between the two half cells. It is impossible to measure the potential of an individual half-cell, instead it must be done by comparing to a reference (standard hydrogen electrode).
• the potential of all electrodes is measured by comparing their potential to that of the standard hydrogen electrode (SHE). The SHE is assigned the potential of 0 Volts.
• To make the standard hydrogen electrode (SHE) conditions must be:

  • hydrogen gas at 100kPa pressure.
  • a solution containing $H^+$ ions at $1\ moldm^{-3}$
  • temp of 298K

  (*tip- the spontaneous reaction between this metal and the standard hydrogen electrode may not occur if its protected by an oxide layer — this can be solved by sanding the metal)

  
• standard hydrogen electrode cell diagram:

  
• when an electrode system is connected to the hydrogen electrode system (and standard conditions apply), the potential difference measured is called the standard electrode potential. The conditions are:

  • ion solutions at $1\ moldm^{-3}$
  • no current flowing
  • gases at 100kPa pressure.
  • temp of 298K

  electrode potentials are always written as reductions:

  
• the more positive electrode potential (reduction reaction/gain of electrons) takes place at the positive electrode (less reactive substance)— right hand side of the cell diagram.

  the more negative electrode potential (oxidation reaction/loss of electrons) takes place at the negative electrode (more reactive substance)— left hand side of the cell diagram.
• the more positive the electrode potential for a reduction, the more that that half equation will favour reduction over oxidation.

  The more negative half cell will always oxidise (go backwards)

  The more positive half cell will always reduce (go forwards).

  
• $E_{cell}\$can be calculated using the standard electrode potentials for the half cells used.

  
  

  E cell (RHS) is which ever substance will undergo reduction in the reaction (can hence yield an overall negative value and not be spontaneous)

  
• the more negative the $E_{cell\ }$value, the more likely it (the reduced form) is to undergo oxidation itself, and hence the stronger the reducing agent it is.
• the more positive the $E_{cell\ }$value, the more likely it (the oxidised form) is to undergo reduction itself, and hence the stronger the oxidising agent it is.
• predicting whether a reaction occurs ($Al(s)\ +$$\ Ag^+$$(aq)$)

  1. electrons move spontaneously from the more negative to the more positive half cell.
  2. the strongest reducing agent is the reduced form in the more negative half (where oxidation is occurring): $Al\left(s\right)$ -- this can give away its electron easily.
  3. the strongest oxidising agent is the oxidised form in the more positive half cell (where reduction is occuring): $Ag^+$

  hence the 2 will react.
• reactions with a positive $E_{cell}$ value will be spontaneous as they give a negative Gibbs free energy change value:

  
• fuel cell - uses the energy from the reaction of a fuel (hydrogen, hydrocarbons, and alcohols) with oxygen to create a voltage.

  They can be up to 70% more efficient at generating electrical energy.
• alkaline hydrogen fuel cell

  (oxygen gains electrons and loses hydroxide)

  (hydrogen loses electrons and gains hydroxide)

  
  

  negative electrode:

  $4\ H_2O\ (l)\ +$$4e^-\leftrightarrow\ 2\ H_2\ (g)\ +$$4\ OH^-(aq)$

  positive electrode:

  $2\ H_2O\ (l)\ +$$O_2\ (g)\ +$$\ 4e^-\ \rightarrow\ 4\ OH^-(aq)$

  
• acidic hydrogen fuel cell:

  • $H_2$ which enters at the negative electrode is a reducing agent, and hence generates electrons
  • (*IN THIS EQUATION, EVERYTHING SHOULD ACTUALLY BE DOUBLED)

  
  

  (oxygen gains electrons and gains H+)

  (hydrogen loses electrons and loses H+)

  
  

  $2\ H^+$$(aq)\ +$$\ 2e^-\ \leftrightarrow\ H_2(g)$

  • $O_2$ which enters at the positive electrode is an oxidising agent, and hence takes electrons:

  $4\ H^+$$(aq)\ +$$\ 2\ O_2(g)\ +$$4e^-\ \rightarrow\ 2H_2O$

  
• for both acidic and alkaline fuel cells, when the cell is discharged (current flows and reaction happens), the overall equation is:

  $O_2(g)\ +$$\ 2\ H_2(g)\ \rightarrow\ 2H_2O(l)$
• the voltage of a fuel cell stays constant over time because the concentrations are always being replenished, and there is a constant supply of fresh reactants.
• methanol fuel cell:

  • in these cells, the reaction taking place is a combustion.

  
• n ,F and T are all constants under standard conditions. Hence, the more positive the $E_{cell}$ value, the greater the increase in total entropy of the reaction (hence the more negative ΔG is)

  
• n (number of $e^-$transferred in redox reaction), F, R, and T are all constants under standard conditions. Hence, the more positive the$E_{cell}$ value, the higher the value of $\ln\left(k\right)$. Hence the equilibrium lies further to the right of the equation.

  
• reactions with a positive $E_{cell}$ value will be feasible/spontaneous as they will yield a negative value for ΔG:

  
• in the following cell, hydrogen peroxide disproportionates because one molecule of it is the strongest oxidising agent, and the other molecule is the strongest reducing agent:

  $2\ H_2O_2\ (aq)\ \rightarrow\ 2\ H_2O\ (l)\ +$$\ O_2\ (g)$

  
• in the following cell, it's very difficult to isolate copper (I) ions in aqueous solution because:

  $Cu^+$$(aq)$ disproportionates (its both the strongest oxidising and reducing agent)

  
• half cell comparison diagram:

  
• advantages of fuel cells for powering vehicles:

  • more efficient
  • lighter
  • no toxic products
• disadvantages of fuel cells for powering vehicles:

  • compressed gases (danger)
  • solid adsorbants are expensive.
  • fuel cell manufacture is costly
• platinum electrodes are typically made by coating porous material with platinum rather than by using platinum rods to increase the surface area of the electrode.
• in a fuel cell, the alkali/acid is an electrolyte which allows the movement of ions between the electrodes
• Redox Calculation Example :

  
• Redox Titration Calculation Example :

  
• $E_{cell\ }$doesn't effect the rate of a reaction:

  A reaction with a positive $E_{cell}$ value may still have a very low rate because the activation energy is very high.
• fuel cells (fundamentals + memory aid)

  
  

  * the fuel is always being oxidised