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Redox and Electrode Potentials Flashcards
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- 1What is a redox reaction?A reaction in which one substance is REDUCED, and another is OXIDISED
- 2What are reducing and oxidising agents?Oxidising agents = oxidise the other species, so are themselves reduced
Reducing agents = reduce the other species, so are themselves oxidised - 3What is oxidation and reduction?Oxidation = the GAIN of oxygen, but LOSS OF ELECTRONS (increase in oxidation number as more +vely charged when losing e-)
Reduction = the LOSS of oxygen, but GAIN OF ELECTRONS (decrease in oxidation number as more -vely charged when gaining e-)
OILRIG - 4Name some common oxidising and reducing agentsCommon oxidising agents = H2SO4, KMnO4, K2Cr2O7, MnO2, Cl2
Common reducing agents = H2, Zn, C, CO, LiAlH4, NaBH4 - 5What is a redox titration?The procedures are similar to acid-base titrations. They are used to find out the exact volume of oxidising/reducing agent is needed to react with and oxidise/reduce a substance
- 6What are the three common redox titrations to know about?OXIDISING AGENTS:
- Potassium Manganate (VII) KMnO4 (under acidic conditions)
- Potassium Dichromate (+6 charge) K2Cr2O7 under acidic conditions
REDUCING AGENTS:
- Sodium Thiosulfate Na2S2O3, for determining the concentration of I2 - 7Potassium Manganate (KMnO4) redox titrations with Fe2+ to Fe3+MnO4- + 8H+ + 5e- ---> Mn2+ + 4H2O (reduced)
Fe2+ ---> Fe3+ + e- (oxidised)
Its an oxidising agent as it is REDUCED from a charge of +7 to +2. So, it can be used to find the concentration of REDUCING AGENTS, such as Fe2+, and ethanedioc acid (COOH)2
The two equations can be combined, cancelling out the electrons, to produce an overall equation. Then, THIS equation can be used alongside titration data, in normal titration calculations. This can be done with ANY reducing agents, not just MnO4-
Redox titrations can also be used to calculate the value of x in calculations involving hydrated salts
The principles of redox titrations can also be applied to OTHER reducing/oxidising agents - 8What can be calculated with information from redox titrations?- Value of 'x' in hydrated salts
- The concentration of an unknown substance
- Percentage purity - 9KMnO4 redox titration procedure1. Add KMnO4 to the burette (record start point volume)
2. Add a measured volume of the solution being analysed to the conical flask
3. Add excess of dilute H2SO4 to the conical flask to provide the H+ ions required for the reduction of MnO4-
4. The reaction is self-indicating. As the reaction is happening, the solution in the conical flask becomes DECOLOURISED. The end point is judged by the occurance of the FIRST PERMANENENT PINK COLOUR (indicating excess of MnO4- ions) and measure the volume lost from the burette
5. Repeat until you obtain concordant tirtres (within O.10cm3 of one another)
KMnO4 is a DEEP PURPLE colour, so its difficult to read measurements from the bottom of the meniscus. So, in this titration, burette readings are read from the TOP OF THE MENISCUS (the titre will be the same as if you were measuring from the bottom as measuring CHANGE - will be the same, as long as you measure it from the top, both times) - 10Aqueous, acidified potassium dichromate redox titrations (Cr2O7 2-)They can be used as an oxidising agent, when supplied with H2SO4/HCl (source of H+ ions)
It can be used to find the concentration of reducing agents + all the other calculations done with KMnO4
Cr2O7 2- + 14H+ + 6e- ---> 2Cr3+ + 7H2O
(reduced from +12 to +6) - Cr2/2Cr - 11Iodide/thiosulfate redox titrations equationsThe thiosulfate ion S2O3 2- is a REDUCING AGENT (becomes oxidised itself), and is used to find the concentration of aqueous iodine to IODIDE ions, which is a REDUCTION reaction
2S2O3 2- ---> S4O6 2- + 2e- (oxidation)
I2 + 2e- ---> 2I- (reduction)
Overall = 2S2O3 2- + I2 -> S4O6 2- + 2I- - 12Iodide/thiosufate redox titrations procedure1. Na2S2O3 is added from a burette, into the concical flask containing aqueous I2.
2. When the solution in the conical flask starts to become a pale-straw colour, it is nearing the end-point, so STARCH SUSPENSION is added
3. Adding starch to the iodine in the conical flask forms a DEEP PURPLE colour
4. The endpoint of the titration is when this deep purple colour DISAPPERS (decolourises) - 13Why is starch added near the end-point and not before?If added too early on in the reaction, the starch will start to BREAK DOWN!
- 14How do you convert from moldm-3 to gdm-3?Multiply the value in moldm-3 by the Mr!
- 15Name four purposes of I2/S2O3 2- titrations?1. To determine the concentration of I2
2. To determine the percentage purity of iron in a sample
3. To work out the ClO- content in bleach
4. To work out the Cu2+ content in copper compounds/alloys - 16Working out the ClO- content in bleach using an iodine/thiosulfate redox titrationActive ingredient in household bleach = ClO- ions (chlorate (I) ions)
NaClO usually supplies the ClO- ions
You can calculate the CONCENTRATION of ClO- ions in bleach, in moldm-3
1. ClO- ions first react with I- and H+ ions, to form I2:
ClO- + 2I- + 2H+ ---> I2 + Cl- + H2O
2. The I2 formed then reacts with S2O3 2- ions in the titration. The moles of I2 can be calculated and, using the first equation, the concentration of ClO- can be calculated using these moles and a fixed volume that you reacted initially
REMEMBER TO SCALE UP YOUR QUANTITIES!! - 17Working out the Cu2+ content in copper compounds/alloys using an iodine/thiosulfate reactionFor copper SALTS (compounds) = the Cu2+ ions are produced by dissolving it in WATER (insoluble ones dissolve in acids)
For copper ALLOYS = the Cu2+ ions are obtained by dissolving it in CONCENTRATED NITRIC ACID, followed by NEUTRALISATION
These ions then react with I- to form I2 (then used in a titration with S2O3 2- ions) - 18What do voltaic electrochemical cells do?They convert CHEMICAL energy to ELECTRICAL energy, using a redox reaction.
Voltaic electrochemical cells are in modern cells and batteries
These reactions are REDOX reactions as the electrical energy results from the MOVEMENT OF ELECTRONS, from one half cell to another - 19What is a half-cell?A half-cell contains the chemical species that are present in a redox HALF EQUATION.
When a METAL ROD is immersed in a solution of its own ions, a VOLTAGE is created between the metal rod and its ionic solution. This is the half-cell's ELECTRODE POTENTIAL (measured in volts) - 20How are voltaic electrochemical cells constructed?They are made by connecting two different HALF CELLS, allowing the flow of electrons between the half cells.
The chemicals in the two half cells must be kept APART. If they were to be together, the electrons would flow uncontrollably and heat energy would then be released instead. - 21What is the format of half-cell electrode potential equationsThey are all represented as REDUCTION equations (as, when measuring standard cell potentials, you are always measuring reduction)
If you are given two half cells' equations and are told they make up an electrochemical cell, the MORE NEGATIVE one will be OXISIDED (so need to reverse the equation) and the MORE POSITIVE one will be reduced, so leave the equation as it is.
When constructing electrochemical cells, the species being oxidised goes on the LEFT and the species being reduced goes on the RIGHT - 22Determining reduction/oxidationMetals can be oxidised/reduced, depending on their chemical environments
In a cell with two metal/metal ion half cells connected, the MORE REACTIVE metal RELEASES electrons, and becomes OXIDISED. The electrode with the less reactive metal gains electrons, and is reduced. - 23What are the 2 types of half-cells?1. Metal/metal ion half cells
2. Metal ion/metal ion half cells - 24Describe what comprises a metal/metal ion half cellWhen the cells are connected, oxidation occurs in the cell supplying the electrons and reduction occurs in the other, to accept them. Each half cell's element will change oxidation state. The DIRECTION of the electron flow depends on the relative TENDENCIES of each electrode to release electrons.
- 25Describe what comprises a metal ion/metal ion half cellThey contain an aqueous solution of the same element, in TWO DIFFERENT OXIDATION STATES eg. Fe2+ and Fe3+. This means there is no solid metal to transport the electrons in/out the half cell, so an INERT METAL PLATINUM ELECTRODE IS USED, to transfer the electrons
- 26What is the purpose of a salt bridge?The cell electrodes are connected by WIRES, whilst the solutions are connected by a SALT BRIDGE
Salt bridges allow the FLOW OF IONS between the two solutions. This MAINTAINS THE CHARGES of the two solutions.
The salt bridge is made up of a concentrated solution that won't react with either of the two solutions in the two half cells it is connecting eg. filter paper soaked in KNO3 - 27Define the standard electrode potentialThis is the emf of a half cell, when connected to a standard hydrogen electrode half cell, under standard conditions (solutions at 1 moldm-1, gases at a pressure of 100kPa, temperature of 298K)
- 28How are standard electrode potential values of metals obtained?They are obtained EXPERIMENTALLY, using a STANDARD HYDROGEN ELECTRODE.
- 29Describe what makes up a standard hydrogen electrode- Hydrogen gas
- A solution of aqueous H+ ions
- A platinum electrode is used (as no solid metal there - a gas/metal ion half cell)
Standard hydrogen electrodes are used because the standard electrode potential of a standard hydrogen electrode is Ov
If, when measuring the standard electrode potential, you get a POSITIVE value, the species is more likely to GAIN electrons and become REDUCESD, vice-versa - 30Describe how to set out the conventional representation of electrochemical cellsSalt bridge = 2 vertical lines
Separate different ions in a metal ion/metal ion half cell with COMMAS
A SINGLE vertical line separates metal electrodes and their ionic solutions
Include state symbols!
Platinum electrode = Pt (s) - 31Describe how to calculate a standard cell potentialStandard cell potential = emf of most positive half-cell - emf of least positive half cell
eg. 0.73 - (-0.43) = +1.16v (include if +ve or -ve)
Essentially = emf rhs - emf lhs
As, in cells, the half cell being reduced goes on the right and the one being oxidised goes on the left - 32How can emf values be used to predict feasibilty? What is the threshold for feasible reactions?Emf values can be used to predict the feasibility of redox and cell reactions.
In THEORY, any redox reaction with a POSITIVE emf value will work.
In PRACTICE, reactions will proceed if the emf value is greater than 0.4V - 33What determines a spontaneous reaction?Any reaction with a positive emf will be spontaneous, but only if it is FEASIBLE and other factors, such as reaction kinetics, have been deemed to not hinder the reaction proceeding
- 34Predicting what is oxidised/reduced from electrode potentialsAs we know a reaction isn't feasible unless the emf is +ve (or, greater than O.4v), we can use electrode potential data to find out what the oxidising and reducing agents are.
This proves that equations with a more positive emf are stronger OXIDISING AGENTS and the equations with a more negative emf are stronger REDUCING AGENTS
The more positive the emf, the greater the tendency for the reaction to proceed FORWARDS
The more negative the emf, the greater the tendency for the reaction to proceed in the REVERSE direction - 35How are emf values related to Gibb's Free Energy?Gibb's Free Energy considers standard enthalpy changes and entropy changes to determine feasibility
But, it can still connect back to emf values
The feasibility of redox reactions can be considered in terms of Gibb's Free Energy and emf values. They are connected mathematically
Gibbs Free Energy = -n x F x emf
n = mol of e- transferred in reaction
F = the Faraday constant - 96,000 c mol-1)
If the cell emf is positive, Gibb's Free Energy will be NEGATIVE (so, feasible)
The smaller the emf of the cell, the LESS NEGATIVE the value of Gibb's Free Energy will become - 36What are the limitations of predictions using emf values?Electrode potentials indicate THERMODYNAMIC FEASIBILITIES, but don't consider reaction kinetics
- A reaction rate may be very SLOW (isn't considered) as ACTIVATION ENERGIES aren't considered. Large activation energies result in very SLOW rates
Some cells, therefore, may have a positive emf value so can be seen to be feasible initially and should happen, but the reaction may not be SPONTANEOUS and may not take place, if the activation energy is too high - 37Standard electrode potential equilibria and Le Chatelier's PrincipleElectrode potentials are calculated at STANDARD CONDITIONS, so any changes in:
- Concentration
- Temperature
- Pressure
affect the emf values
Only need to know effects of concentration and temperature in detail
For changes in pressure, look at the MOLES of each side - will go in side of least/most moles - depends - 38How do changes in concentration affect equilibria? (if not at 1 moldm-3)Increasing the concentration of the REACTANTS over 1 moldm-3 means the equilibrium shifts to the RIGHT, and so the emf value of the PRODUCT becomes MORE POSITIVE
Decreasing the concentration of the REACTANTS to under 1 moldm-3 means the equilibrium shifts to the LEFT, so the emf value of the product becomes LESS POSITIVE (less product) - this decreases the emf value. - 39How do changes in temperature affect equilibria?Most reactions are EXOTHERMIC, so increasing the temperature would then DECREASE the cell potential, as the equilibrium would move to the left as this is the endothermic direction - more likely oxidised, then)
- 40What are the 3 types of electochemical cells?1. Primary cells
2. Secondary cells
3. Fuel cells - 41Primary cellsThese are NON-RECHARGEBALE
This is because the reactions are non-reversible. Electrical energy produced by oxidation and reduction at the electrodes can't be reversed.
The chemicals are eventually used up, voltage falls, battery goes flat and the cells are then discarded/recycled
APPLICATIONS:
- Usually in wall clocks and smoke detectors (low current so long storage)
- Most are ALKALINE, made with Zn and MnO2 and a KOH alkaline electrolyte - 42Secondary cellsThese are RECHARGEABLE as the cell reaction producing electrical energy can be REVERSED during recharging - the chemicals in the cell are regenerated and can be used again. Used eg. in car batteries (storage cells)
EXAMPLES
- Lead-acid batteries (car batteries)
- Nickel-cadmium batteries (radios, torches etc.)
- Lithium-ion cells (cameras, laptops, phones) - lithium-ion = risk of toxicity + fire, but low costs - 43Fuel cells, using the example of hydrogen-fuel cellsA fuel cell is an electrochemical cell, using the energy from the reaction of a fuel with oxygen, to create a voltage.
Most common fuel used = hydrogen
The fuel and oxygen flow into the fuel cell each side, and the product flows OUT
Fuel cells DON'T have to be recharged and can operate CONTINUOUSLY, as long as the fuel and oxygen are constantly supplied into the cell
H2 GIVES the electrons, at one electrode, and oxygen GAINS these electrons, at the other electrode.
The GASES are bubbled through 2 porous platinum coated electrodes.
Hydrogen fuel cells can have either an ACID or an ALKALINE electrolyte
The acid/alkaline cells would have different redox systems/equations, but the emf of both cells is 1.23V
Both have the same overall equation of:
2H2 (g) + O2 (g) -> 2H2O (l) - 44ACIDIC hydrogen-fuel cellsH+ ions diffuse through porous membrane
ANODE (oxidation): H2 ---> 2H+ + 2e-
CATHODE (reduction): 4H+ + 4e- + O2 ---> 2H2O - 45ALKALINE hydrogen fuel cellsAn alkaline electrolyte eg. KOH is used. OH- ions diffuse through porous electrode, from anode to cathode
ANODE (oxidation): 2H2 + 4OH- ---> 4H2O + 4e-
CATHODE (reduction): 2H2O + 4e- + O2 ---> 4OH- - 46Advantages of hydrogen-fuel cells- Only product is harmless water (no pollutants)
- More efficient than conventional energy production methods
- Operates continuously when fuel is available
- Less complex than conventional methods - 47Disadvantages of hydrogen-fuel cells- Expensive to extract/store/transport hydrogen (need a large fuel tank with thermal insulation + pressure withstanding)
- CO2 can be released during HYDROGEN PRODUCTION
- Risks of leaks during hydrogen storage (can lead to fire + explosion) - 48Oxidation equation of a standard hydrogen electrode?H2 ---> 2H+ + 2e-
- 49What is the pH of the solution in a standard hydrogen half-cell?0
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- A level Chemistry Fichas sobre Redox and Electrode Potentials Flashcards, creado por Rhea Sekhon el 26/12/2021.
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- Rhea Sekhon
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