5E - Redox Flashcards
Define:
Oxidation
Reduction
Oxidising agent
Reducing agent
Oxidation = Loss of electrons or increase in oxidation number
Reduction = Gain of electrons or decrease in oxidation number
Oxidising agent = Reagent that accepts electrons from another speies (it itself is reduced)
Reducing agent = Reagent that donates electrons to species being reduced (it itself is oxidised)
Define
Redox reaction
Disproportionation reaction
Redox reaction = Reaction involving oxidation and reduction
Disproportionation reaction = Redox reaction in whcih the same element is both oxidised and reduced
Give an example of a disproportionation of oxygen
Steps to writing a redox equation
- Balance atoms
- Balance charges
For half equations what side do the electrons go on if the atom is being:
- Oxidised
- Reduced
Reduction: Add electrons on LEFT
Oxidation: Add electrons on RIGHT
Steps to writing a redox equation from half equations
- Balance electrons
- Add equations together + cancel electrons
- Cancel any species on both sides
Steps to writing a redox equation from oxidation numbers
- Summarise information given in an equation
- Assign oxidation no’s to identify what has been oxidised and reduced
- Balance ONLY the species that contain the elements that have changed in oxidation no’
- Balance remaining atoms
- Add H2O to balance oxygen
- Add H+ to balance hydrogen
- Check charges are balanced
Define Half cell
Half cell = Contains an element in 2 different oxidation states together with a means of electrical contact (e.g. a piece of metal)
Name the types of half cells you can get and name an example of each
METAL / METAL ION HALF CELLS
- Zinc reacting in CuSO4(aq) as a redox reaction
- Zinc is oxidised to Zinc ions and Copper ions reduced to Copper
METAL ION / METAL ION (ION / ION) HALF CELLS
- Contains the same element in different oxidation states
- Redox equilibrium of Fe3+ + e- ⇌ Fe2+
Draw the half cells and simple electrochemical cell for this reaction:
Zinc reacting in CuSO4(aq) as a redox reaction. Zinc is oxidised to zinc ions and copper ions reduced to copper
Define Salt bridge + example of how to make one
Salt bridge = A concentrated solution of an electrolyte that doesn’t react with the materials in either half cells to complete the circuit by allowing ions to flow between the half cells
e.g. piece of filter paper saturated in KNO3(aq)
Draw the half cell of a metal ion / metal ion for the redox reaction of Fe3+ + e- ⇌ Fe2+
What is the electrode made of and why is it made of such a material?
Electrode = Platinum because it’s inhert
Define Standard electrode potential (E⦵)
Standard electrode potential (E⦵) = The emf of a half cell compared with a standard hydrogen half cell, measured at:
- Temperature of 298 K
- Solution concentrations of 1 moldm-3
- Gas pressure of 100 kPa
What does emf stand for?
emf = electromotive force
Draw a standard hydrogen half cell
State the standard conditions and the electrode used in this half cell
Standard conditions:
- 298 K
- Solutions 1 moldm-3 concentration
- 100 kPa
Electrode used= Platinum
Draw a standard electrode potential diagram for a Zn2+ / Zn half cell to be measured
What special care needs to be taken when measuring standard electrode potentials and why do they need to be taken?
when measuring standard electrode potentials NO current is allowed is flow as this changes the conditions so the system is no longer standard
What do standard electrode potentials tell us?
How likely a half cell is to release or accept electrons:
Electrode potential has a negative sign:
- Half cell is more likely to give up electrons than the hydrogen half cell
Electrode potential has a positive sign:
- Half cell is less likely to give up electrons than the hydrogen half cell
Equation for E⦵(cell) = ?
E⦵(cell) = E⦵(more positive electrode) - E⦵(more negative electrode)
E⦵ value > 0 (feasible reaction)
E⦵ value < 0 (not feasible reaction)
Suggest 2 reasons why, even when such combinations of half cells are correctly connected together, the predicted reaction may not, in fact, take place
- Reaction has a very high Ea so reaction rate is too slow to be obsrved
- Actual conditions may not be standard, so the predictions are no longer accurate
Are oxidation / reduction reactions endothermic / exothermic?
Oxidation = endothermic
Reduction = exothermic
When doing non standard electrode condition calculations what is different about the final answer?
It’s called emf and not E⦵
Using this reaction explain the effects of increasing [Fe2+] and decreasing [Cu2+] on the E value of half cell (emf).
What if you change temperature?
Fe2+(aq) + 2e- ⇌ Fe(s) E⦵ = -0.44 V
Cu2+(aq) + 2e- ⇌ Cu(s) E⦵ = +0.34 V
Fe2+(aq) + 2e- ⇌ Fe(s) E⦵ = -0.44 V
Cu2+(aq) + 2e- ⇌ Cu(s) E⦵ = +0.34 V
[Fe2+] increased above 1 moldm-3
- Equil shifts to favor Fe(s) side
- No’ of electrons produced by this half cell decreases ∴ it’s more likely to accept electrons to form Fe(s)
- emf will be more positive
[Cu2+] decreased below 1 moldm-3
- Equil shifts to favour left side
- Half cell more likely to give up electrons and more likely to accept them
- emf will be less positive
Temperature is explained by le chatelier’s principle.
- Decrease temp → Favours exothermic reaction
- Increase temp → Favours endothermic reaction
Explain the LORA rule
L = Left
O = Oxidises
R = Right
A = Above
How do determine whether a reaction is deasible or not?
For a reaction to be feasible the E value overall MUST be positive (wwhen you add them together)
Question:
Write the redox equation
- Find out that Chromium has been reduced and carbon has been oxidised
- Balance all thigns with chromium and carbon in them
- Add 4H2O to balance out oxygens
- Add 8H+ to balance hydrogens
- Check that the charges balance on both sides (both sides have a +6 charge)
Question:
Write the redox equation
NOTE:
- Only some of the chlorine changes its oxidation number, so need to balance the chlorine when equalising the oxidation number changes.*
- The remainder chlorine will during the final stage*
- Work out that Cl has been oxidised and Mn has been reduced
- Balance species containing Cl and Mn
- Balance the rest of the species
- Check charges balance (both sides = 0)
Name the 3 types of cells that provide electrical power, are they rechargable or not?
Storage cells
- Primary cells - non rechargable (battery)
- Secondary cells - rechargable (battery)
Fuel cells - Fuel + oxygen
Storage cells = finite (unless rechargable)
Fuel cells = Not finite
Define storage cell
Storage cell = Often known as a battery.
- Has a fixed amount of chemicals which are used up as the cell discharges and its energy is used
- Has a specific amount of energy it can supply
- Depending on the cell, it may be able to be recharged
- Primary cell = non rechargable
- Secondary cell = rechargable
Define fuel cell
Fuel cell = Uses a fuel + oxygen which undergo an exothermic reaction
- Reaction arranged so that energy produced is electrical energy
- Fuel can continue to operate as long as furl and oxygen are available
What are primary cells?
How do they work?
Give and example + equations
Primary cells:
- Use only once, non-rechargable
- Electrical energy produced by oxidation and reduction reactions at electrodes
- Most are ALKALINE BATTERIES
ALKALINE BATTERIES:
- Positive electrode = Solid manganese (V) oxide outer layer
- Negative electrode = Powdered zinc core
- solution of KOHaq) is used between them
Electrons move from negative electrode (Zn) → positive electrode (MnO2)
- Zinc - loses electrons and is oxidised because it has the lower E⦵ value
Uses of primary cells (3)
- Conventional batteries
- Smoke alarm
- Power individual calculators
What are secondary cells?
How do they work?
Example + equations
Secondary cells:
- Rechargable nickel - cadmium batteries
- Lead - acid batteries are also a common example
- For the reaction to be reversed and recharge the battery you flip the products and reactants around!
- 2 different types of hydrogen-oxygen fuel cells?
- Are these common?
- Advantages (3)
- Hydrogen-oxygen fuel cells are MOST common type because they don’t produce CO2 during combustion with H2O
- They can be acid or alkaline
Advantages:
- Rechargable / reusable
- Reducing environmental impact (waste materials) and overall long term costs
- Supplies electrical energy constantly
Describe an experiment to measure standard electrode potentials (5)
(Include ones for metal / metal, ion / ion and gas containing half cells)
- Prepare 2 standard half cells:
- For metal / metal ion half cell, the metal ion must have a conc of 1 moldm-3
- For ion / ion half cell, both metal ions present in the solution must have the same conc. Electrode used must be inhert - usually platinum
- For half cell containing gases (e.g. hydrogen half cell), the gas must be at 100kPa pressure, in contact with a solution with an ionic conc of 1 moldm-3. Electrode must be inhert - usually platinum
- Temp must be 298 K for all half cells
- Connect metal electrodes of the half cells to a voltmeter using wires
- Prepare salt bridge by soaking a strip of filter paper into a saturated aqueous solution of KNO3
- Connect the 2 solutions of the half cells with salt bridge
- Record the standard electrode potential from voltmeter
