Redox Flashcards
Oxidation definitions
gain of oxygen
loss of electrons
loss of hydrogen
increase in oxidation state
Reduction definitions
loss of oxygen
gain of hydrogen
gain of electrons
decrease in oxidation state
Oxidation number
the charge that an atom would have if all the covalent bonds were broken so that the more electronegative element kept all the electrons in the bond
Rules for oxidation states (numbers)
- elements = 0
- simple ions = same as charge
- in compounds:
- Oxygen = -2 (except in peroxides where O = -1, OF2 = +2)
- Hydrogen = +1 (except in metal hydrides = -1)
- Chlorine = -1 (except when with O or F)
- in polyatomic ions the oxidation number of the atoms add up to the net charge on the ion
- oxidation state of a transition metal in a complex ion can be worked out using the charges on the ligands and the overall charge
Oxidising agent (oxidant)
will oxidise other molecules but itself is reduced
Reducing agent (reductant)
will reduce other molecules but itself is oxidised
Steps in balancing a redox equation
- separate the redox equation into 2 half equations (omitting spectator ions)
- balance all atoms (except O and H)
- balance O using H2O
- balance H using H+ (the acid used is sulfuric acid as the SO4 2- ion does not react unlike other acids)
- balance the charges by adding electrons to the more positive side
- multiply each equation by some factor that will cause the electrons to cancel
- cancel like terms
Titration - Iron with manganate (VII)
this redox titration uses KMnO4 in an acidic solution as the oxidising agent, which oxidises Fe2+ ions into Fe3+
during the reaction MnO4 - is reduced to Mn2+
the reaction is accompanied by a colour change from purple to colourless so reaction acts as its own indicator
Titration - iodine-thiosulfate
several different titrations use an oxidising agent to react with excess iodide ions to form iodine
the excess iodine is then titrated with sodium thiosulfate using starch as an indicator
the starch indicator added during the titration forms a deep blue colour by complexing I2
as the I2 is reduced to I- during the reaction, the blue colour disappears
Winkler method for calculating dissolved oxygen
dissolves oxygen content of water is an important indicator of its quality (as pollution increases dissolves O2 generally decreases)
the biological oxygen demand (BOD) is used as a means of measuring the degree of pollution (BOD = amount of oxygen used to decompose the organic matter in a sample of water)
a high BOD indicates a lower level of dissolved O2
calculated based on a series of redox titrations
1. the dissolved O2 in the water is ‘fixed’ by the addition of manganese (II) salt. reaction of this salt with O2 in basic solution causes oxidation of Mn(II) to higher oxidation states such as Mn(IV)
2. acidified iodide ions (I-) are added to the solution, and are oxidised by the Mn (IV)
3. the iodine is then titrated with sodium thiosulfate
so for every 1 mol of O2 in the water 4 mol of thiosulfate ions are used
Voltaic cells
convert energy from spontaneous, exothermic chemical processes to electrical energy
oxidation occurs at the anode (negative electrode) and reduction occurs at the cathode (positive electrode)
they have a salt bridge (which contains a strong electrolyte - one that conducts electricity) these maintain electrical neutrality and completes the circuit
electrons flow from anode to cathode
Electrolytic cells
convert electrical energy to chemical energy, by bringing about non-spontaneous processes
oxidation occurs at the anode (positive electrode) and reduction occurs at the cathode (negative electrode) therefore cations migrate to the cathode and anions migrate to the anode
both electrodes are submerged in the same electrolyte and connected to a DC power supply
as the electric current passes through the electrolyte, redox reactions occur at the electrodes removing the charges on the ions and forming electrically neutral products
Cell diagram convention - voltaic cells
has the following features:
- a single vertical line represents a phase boundary such as that between a solid electrode and an aqueous solution within a half-cell
- a double vertical line represents the salt bridge
- the aqueous solutions of each electrode are placed next to the salt bridge
- the anode is on the left and the cathode on the right so electrons flow from left to right
- spectator ions are emitted
Standard hydrogen electrode (SHE)
the electrode potential of any two half cells are measured against the SHE
all reactions are completes under standard conditions
and under these conditions the standard electrode potential of the hydrogen electrode = 0 V
the SHE consists of hydrogen gas bubbled over a platinum electrode in contact wth 1 moldm-3 acid
Standard electrode potential
of a half-cell is its electrode potential relative to the standard hydrogen electrode (SHE), measured under standard conditions (pressure 100kPa, temp. 298K, all solutions 1 moldm-3)
all are reduction potentials and are measure of how much the reduction half equation wants to occur
when a half-cell is above H in the activity series e flow from half-cell to SHE, and electrode is given a -ve value
if the half-cell contains a metal below H in the activity series e flow from the SHE to the half-cell and the electrode has a +ve value
