Unit 3, Topic 2: Redox Reactions Flashcards
Redox reactions
- involve simultaneous reduction and oxidation reactions
- They involve the transfer of electrons from one chemical species to another
Oxidation
Losing electrons
Reduction
Gaining electrons
Oil Rig
Oxidation Is Loss Reduction Is Gain
Metals
Tend to oxidise - This is due to their differing electronegativities
Non-metals
Tend to reduce - This is due to their differing electronegativities
The number of electrons lost
determined by the electron configuration – tries to reach most stable electron configuration
Oxidising agent
An oxidising agent or oxidant enables another chemical to be oxidised – they themselves are reduced
Reducing agent
A reducing agent or reductant enables another chemical to be reduced – they themselves are oxidised
Oxidation number
- Oxidation numbers and oxidation states are interchangeable terms
- The +- sign is before the number i.e. +2 and not 2+
- Based off the assumption that all bond in the compound are ionic
Oxidation numbers in use
- Oxidation number can be used to determine whether the species will oxidise or reduce – an element in its highest oxidation state can only be reduced, vice versa is true – element in lowest oxidation state can only be oxidised
Oxidisation (oxidation state)
increase in oxidation state
Reduction (oxidation state)
decrease in oxidation state
Oxidation number of free elements (e.g. O2)
zero
Oxidation number of a simple ion
equal to charge e.g. oxidation number of Na+ is +1
Oxidation number of main group metals
Main group metals have oxidation number equal to charge
Hydrogen oxidation state
Hydrogen is normally +1 except in metal hydrides
Oxygen oxidation state
Oxygen is normally -2 except in fluorine where it has a positive number and in peroxides where oxidation number is -1
Fluorine oxidation state
Fluorine is always -1
Sum of all oxidation numbers in a neutral compound
zero
The sum of all oxidation numbers in a polyatomic ion
equal to the charge
The most electronegative element
assigned the lowest oxidation number e.g. OF2 where O is +2 and F is -1+2 − (2 × −1) = 0
Balancing redox reactions
- Identify what has been oxidised and what has been reduced – half equations
- A) Balance equations for everything but oxygen and hydrogen
B) Balance oxygen using H2O(l)
C) Balance hydrogen using H+(aq) - If necessary, multiply both equations by an integer so that the number of electrons is equal on both sides
- Add two half equations cancelling out spectator ions etc.
- Check for balancing
- (If in basic solution) cancel out H+ ions by adding OH- ions to both sides of the equation
Galvanic cells
- Galvanic cells (voltaic) are a type of electrochemical cell where chemical energy is converted to electrical energy
- Designed so that the half cells (oxidation and reduction half equations) are kept separately and are connected by an external circuit – facilitates transfer of electrons
anode - Galvanic
negative
cathode - galvanic
positive
Salt bridge
Helps balance charges
Half-Cells
- Each half cell is an electrode in contact with a solution
- The chemical species in the half cells form a conjugate redox pair
- If one member of the pair is metal, it is likely to be the electrode
- If no metal is present, an inert electrode is used like platinum or graphite
Salt Bridge
- The salt bridge houses ions that are free to move
- The cations and anions in the bridge are able to move to either half-cell to balance electrical charge
- Known as an internal circuit
From an oxidised species to reduction species:
- The more positive the E° value is, the stronger it is as an oxidising agent
- The more negative the E° value is, the stronger it is as a reduction agent
E˚
determined off a standard half-cell with a hydrogen cell
E˚ of hydrogen cell
0 V
E𝐶𝑒𝑙𝑙° formula
E𝐶𝑒𝑙𝑙° = E𝑅𝑒𝑑𝑢𝑐𝑡𝑖𝑜𝑛° − E𝑂𝑥𝑖𝑑𝑎𝑡𝑖𝑜𝑛°
Cell voltage greater than 0
spontaneous reaction
cell voltage less than 0
nonspontaneous reaction
Electrolytic Cell
Type of electrochemical cell where electrical energy is converted to chemical energy
(opposite of Galvanic Cells)
- Electrolytic cells have two electrodes in contact with the electrolyte connected to an external power supply
Electrolysis
involves passing electrical energy through a conducting liquid o Causes non-spontaneous redox reactions to occur
Cathode - electrolytic
reduction
anode - electrolytic
oxidation
power supply - electrolytic cells
The power supply draws electrons from the positive electrode to the negative electrode
anions - electrolytic cells
Anions move towards anode
cations - electrolytic cells
cations move toward the cathode
Electrolysis of a Molten Salt
- Inert electrodes (platinum or graphite) are used – doesn’t interfere with the reaction
- Electrons are supplied to the negative electrode and cycles through to the positive electrode
Electrolysis of an aqueous solution
- In aqueous solutions, the water itself may be oxidised/reduced
- Therefore, all four electrode potentials must be considered
- The combination of oxidation and reduction with the more positive potential will occur
Concentrated vs dilute
- In a concentrated NaCl solution, the secondary oxidation reaction (2𝐶𝑙(−𝑎𝑞) → 𝐶𝑙2(𝑔) + 2𝑒−) will occur
- This is because the original electrochemical series is determined off standard conditions with 1 mol solutions
- Also, in a concentrated solution, there are more chloride ions vs hydroxide ions – therefore their production will dominate
Electroplating
- industrial use of electrolysis
- Involve depositing a layer of metal on the surface of another metal
- object to be plated is connected to the negative terminal - it becomes the cathode
- the positive electrode (anode) is made of the metal to be plated onto the cathode and the solution contains ions of said metal
- these ions will be reduced into the cathode
Electrorefining
- Involves extracting pure metals from ores
- When copper is extracted from ores, it is 98% pure and is further purified by electrorefining
- Impure copper is at the anode and a pure copper sheet is the cathode
- Oxidation of the copper occurs forming Cu2+ ions o Attracted to the cathode – essentially plating copper plates
- Impurities left (zinc, nickel) are more reactive and become oxidised – stays as ions
- Less reactive (gold, silver) aren’t oxidised
– collects at the bottom of the solution
