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Redox Processes (Ira and Tanishi) - Coggle Diagram
Redox Processes (Ira and Tanishi)
9.2: Electrochemical Cells
Voltaic cells
Uses a redox reaction to produce an electric current
Comprises two half-cells, each with one electrode (cathode or anode depending on the solution in the half-cell). These electrodes are connected by an external circuit through which the voltage produced is measured. The circuit is completed by a salt bridge which allows ions to flow and complete the circuit.
half-cells: an electrode (usually a metal) in a solution of its own ions.
Salt bridge: can be made from a strip of filter paper soaked in saturated KNO3 or a glass tube with agar gel inside.
Voltage produced depends on reactivity of the two metals in the half-cells
More reactive metals have a greater tendency to lose electrons (become oxidised) as they are stronger reducing agents. In a voltaic cell, the more reactive metal 'drives' the redox process
Cell potential: also known as the EMF (electromotive force) is the difference in electrode potentials of the half-cells in a voltaic cell. Measured in volts.
The further the two metals are apart on the reactivity series, the greater the voltage produced by the cell
Oxidation at negative anode, reduction at positive cathode
Anions migrate from the salt bridge to the anode, cations migrate from the salt bride to the cathode
More reactive metal at the anode (greater reducing agent), and less reactive metal (weaker reducing agent) at the cathode
Exothermic reaction, resulting in the release of heat. Since the two reactions are separated through the use of half-cells and linked by a circuit, electrical energy can be produced.
Cell diagram convention: representations of voltaic cells.
Characteristics:
• Phase boundary between solid and aqueous solution
• salt bridge represented by two vertical lines
• anode on the left, cathode on the right
• two types of ions in the same half-cell are differentiated with a comma
• spectator ions present in the reaction are not included
Chemical energy to electrical energy (KEY DIFFERENCE BETWEEN VOLTAIC AND ELECTROLYTIC CELLS)
Electrolytic cells
Key definitions
Electrolysis: process by which a compound is broken down into its constituent elements using electricity
Inert materials do not take part in the chemical reaction
Electrical conductivity of mobile ions: mobile ions in aqueous solutions move in a specific direction in an electrical field
Theory supported by an experiment in a U-tube involving the movement of coloured positive and negative ions in an electric field
Composed of a molten or aqueous electrolyte, a battery, and two electrodes (usually made of graphite or platinum because they're inert)
Anode is positively charged + attract anions, cathode is negative charged + attracts cations (KEY DIFFERENCE BETWEEN VOLTAIC AND ELECTROLYTIC CELLS)
Cations/anions reach the surface of the electrode --> undergo oxidation or reduction--> discharged from solution
Oxidation at positive anode, reduction at negative cathode
Endothermic, non-spontaneous reaction (KEY DIFFERENCE BETWEEN VOLTAIC AND ELECTROLYTIC CELLS)
Electrical energy to chemical energy (KEY DIFFERENCE BETWEEN VOLTAIC AND ELECTROLYTIC CELLS)
Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte
9.1: Oxidation and Reduction
Oxidation: gain of oxygen, loss of hydrogen, loss of electrons, increase in the oxidation state
Reduction: loss of oxygen, gain of hydrogen, gain of electrons, decrease in the oxidation state
Redox reaction: oxidation + reduction are taking place
Half equations: represent the two processes involved (oxidation + reduction) - can be combined into an overall equation
Oxidising agent: substance that oxidises other substances - readily accepts electrons and is reduced
Reducing agent: substance that reduces other substances - readily donates electrons and is oxidised
Oxidation states: give a measure of the control an atom has over the electrons in the compound; written with a +/- sign (+2 not 2+)
Negative - gained electron control
Positive - lost electron control
Rules
Free elements - oxidation state 0
Sum of the oxidation states = overall charge on the compound
Alkali metals = +1
Fluorine = -1
Alkaline earth metals = +2
Hydrogen is always +1 unless it's with certain metals hydrides (eg NaH) = -1
Oxygen is always -2 except in peroxides = -1 or when with fluorine
Chlorine is -1 unless it's with oxygen or fluorine
The sum of the oxidation states in a polyatomic ion must add up to the charge on the ion
Systematic names - roman numerals are used to show the oxidation state when there is more than one possible oxidation state for an atom in a compound - e.g. iron(II) oxide
Disproportionation: same species is oxidised + reduced simultaneously in the reaction
Activity series: lists metals in order of their strength as reducing agents
Useful to predict a displacement reaction (element will be displaced if it is less reactive than the one displacing it)
Rust prevention - redox reaction
Requires the presence of oxygen + liquid water
Sacrificial protection: waterproof layers can protect iron from rusting
Galvanising: protection involving a more reactive metal
Redox in acidic solutions
Determine what is being reduced and oxidised
Balance the atoms (all apart from hydrogen + oxygen)
Balance for oxygen by adding H2O
Balance for hydrogen with H+
Balance each half equation with electrons
Balance the two half-equations for electrons
Cancel what appears on both sides
Winkler method: measure the concentration of dissolved oxygen in a water sample
Used to determine the biochemical oxygen demand (BOD) - amount of dissolved oxygen required to biologically decompose the organic matter in a water sample over a set time period (usually five days)
Polluted water with a high BOD without means of replenishing oxygen won't be able to sustain aquatic life - lead to eutrophication
