5E - Redox

Question 1

Define:

Oxidation

Reduction

Oxidising agent

Reducing agent

Answer 1

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)

Question 2

Define

Redox reaction

Disproportionation reaction

Answer 2

Redox reaction = Reaction involving oxidation and reduction

Disproportionation reaction = Redox reaction in whcih the same element is both oxidised and reduced

Question 3

Give an example of a disproportionation of oxygen

Answer 3

Question 4

Steps to writing a redox equation

Answer 4

1. Balance atoms
2. Balance charges

Question 5

For half equations what side do the electrons go on if the atom is being:

• Oxidised
• Reduced

Answer 5

Reduction: Add electrons on LEFT

Oxidation: Add electrons on RIGHT

Question 6

Steps to writing a redox equation from half equations

Answer 6

1. Balance electrons
2. Add equations together + cancel electrons
3. Cancel any species on both sides

Question 7

Steps to writing a redox equation from oxidation numbers

Answer 7

1. Summarise information given in an equation
2. Assign oxidation no’s to identify what has been oxidised and reduced
3. Balance ONLY the species that contain the elements that have changed in oxidation no’
4. Balance remaining atoms
   
   • Add H₂O to balance oxygen
   • Add H⁺ to balance hydrogen
5. Check charges are balanced

Question 8

Define Half cell

Answer 8

Half cell = Contains an element in 2 different oxidation states together with a means of electrical contact (e.g. a piece of metal)

Question 9

Name the types of half cells you can get and name an example of each

Answer 9

METAL / METAL ION HALF CELLS

• Zinc reacting in CuSO_4(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⁺

Question 10

Draw the half cells and simple electrochemical cell for this reaction:

Zinc reacting in CuSO_4(aq) as a redox reaction. Zinc is oxidised to zinc ions and copper ions reduced to copper

Answer 10

Question 11

Define Salt bridge + example of how to make one

Answer 11

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)

Question 12

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?

Answer 12

Electrode = Platinum because it’s inhert

Question 13

Define Standard electrode potential (E^⦵)

Answer 13

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

Question 14

What does emf stand for?

Answer 14

emf = electromotive force

Question 15

Draw a standard hydrogen half cell

State the standard conditions and the electrode used in this half cell

Answer 15

Standard conditions:

• 298 K
• Solutions 1 moldm^-3 concentration
• 100 kPa

Electrode used= Platinum

Question 16

Draw a standard electrode potential diagram for a Zn²⁺ / Zn half cell to be measured

Answer 16

Question 17

What special care needs to be taken when measuring standard electrode potentials and why do they need to be taken?

Answer 17

when measuring standard electrode potentials NO current is allowed is flow as this changes the conditions so the system is no longer standard

Question 18

What do standard electrode potentials tell us?

Answer 18

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

Question 19

Equation for E^⦵(cell) = ?

Answer 19

E^⦵(cell) = E^⦵(more positive electrode) - E^⦵(more negative electrode)

E^⦵ value > 0 (feasible reaction)

E^⦵ value < 0 (not feasible reaction)

Question 20

Suggest 2 reasons why, even when such combinations of half cells are correctly connected together, the predicted reaction may not, in fact, take place

Answer 20

• Reaction has a very high E_a so reaction rate is too slow to be obsrved
• Actual conditions may not be standard, so the predictions are no longer accurate

Question 21

Are oxidation / reduction reactions endothermic / exothermic?

Answer 21

Oxidation = endothermic

Reduction = exothermic

Question 22

When doing non standard electrode condition calculations what is different about the final answer?

Answer 22

It’s called emf and not E^⦵

Question 23

Using this reaction explain the effects of increasing [Fe^2+] 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

Answer 23

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

Question 24

Explain the LORA rule

Answer 24

L = Left

O = Oxidises

R = Right

A = Above

Question 25

How do determine whether a reaction is deasible or not?

Answer 25

For a reaction to be feasible the E value overall MUST be positive (wwhen you add them together)

Question 26

Question: Write the redox equation

Answer 26

1. Find out that Chromium has been reduced and carbon has been oxidised 2. Balance all thigns with chromium and carbon in them 3. Add 4H₂O to balance out oxygens 4. Add 8H⁺ to balance hydrogens 5. Check that the charges balance on both sides (both sides have a +6 charge)

Question 27

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*

Answer 27

1. Work out that Cl has been oxidised and Mn has been reduced 2. Balance species containing Cl and Mn 3. Balance the rest of the species 4. Check charges balance (both sides = 0)

Question 28

Name the 3 types of cells that provide electrical power, are they rechargable or not?

Answer 28

Storage cells * Primary cells - non rechargable (battery) * Secondary cells - rechargable (battery) Fuel cells - Fuel + oxygen Storage cells = finite (unless rechargable) Fuel cells = Not finite

Question 29

Define **storage cell**

Answer 29

**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

Question 30

Define **fuel cell**

Answer 30

**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

Question 31

What are **primary cells**? How do they work? Give and example + equations

Answer 31

_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 KOH_aq) 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

Question 32

Uses of primary cells (3)

Answer 32

* Conventional batteries * Smoke alarm * Power individual calculators

Question 33

What are **secondary cells**? How do they work? Example + equations

Answer 33

_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!

Question 34

* 2 different types of hydrogen-oxygen fuel cells? * Are these common? * Advantages (3)

Answer 34

* 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

Question 35

Describe an experiment to measure standard electrode potentials (5) (Include ones for metal / metal, ion / ion and gas containing half cells)

Answer 35

1. 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 2. Connect metal electrodes of the half cells to a voltmeter using wires 3. Prepare salt bridge by soaking a strip of filter paper into a saturated aqueous solution of KNO₃ 4. Connect the 2 solutions of the half cells with salt bridge 5. Record the standard electrode potential from voltmeter