R3.2 Definitions

Question 1

Good acronyms

Answer 1

OIL RIG - Oxidation is loss Reduction is gain

Question 2

Oxidation

Answer 2

Loss of electrons (also loss of hydrogen and gain of oxygen but these definitions do no apply to all compounds).
Oxidising agent oxidises another species but is reduced itself.

Question 3

Reduction

Answer 3

Gain of electrons (also gain of hydrogen and loss of oxygen but these definitions do no apply to all compounds).
Reducing agent reduces another species but itself is oxidised.

Question 4

Oxidation State

Answer 4

Regards all compounds as ionic and assigns them a charge.

When calculating it, divide the value by the subscript amount of a molecule present, but not by the co-efficient.

When using this value to balance the electrons, multiply the oxidation state by the subscript and the co-efficient.

Question 5

Calculating oxidation state rules

Answer 5

• Total oxidation states must add to the overall charge.
• Pure elements are zero.
• Maximum oxidation state for anything is its group number.
• More electronegative element has negative charge.

Question 6

Compounds with constant oxidation states

Answer 6

Fluorine = -1
Oxygen = -2 BUT +2 with Fluorine and -1 in peroxides (H2O2).
Chlorine = -1 apart from with O or F
Hydrogen = +1 but -1 when in metal hydrides.
Group 1&2= +1 or +2

MOST ELECTRONEGATIVE ELEMENT TAKES PRIORITY.

Question 7

Oxidation/reduction impact on oxidation state.

Answer 7

Oxidation=increase in oxidation state
Reduction=decrease in oxidation state.

Question 8

Redox Reaction

Answer 8

Both oxidation and reduction occur

Question 9

How to find balanced equation from reactants and products.

Answer 9

1. Write out unbalanced equation with only what is given and find the change in charge for each molecule.
2. Write out half equations and balance their charges by adding the change in charge amount of electrons to the positive side.
3. Balance each side by adding H2O to balance Oxygen and H+ to balance hydrogen.
4. Find Lowest Common Factor of the electrons in each half equation and multiply the equations by that.
5. Put the equations together.

Question 10

Redox Titration

Answer 10

Finding unknown concentration of a substance by finding the equivalence point, which is when it has reacted stoichiometrically by transferring electrons with the titrant (an oxidising (e.g. KMnO4/K2Cr2O7) or reducing agent) - this can be indicated by an indicator or some solutions change colour themselves.

Question 11

Periodic trends with reduction and oxidation

Answer 11

More reactive non-metals are stronger oxidising agents, and more reactive metals are stronger reducing agents.

Question 12

Spectator Ion

Answer 12

A species which does not change in the reaction and can be cancelled out in some representations (e.g. half equations).

Question 13

How to figure out who is reducing/oxidising with two metals/non-metals

Answer 13

The stronger oxidising/reducing agent (more reactive) will ‘win’ and undergo that process, whereas the other species will undergo the other. If you put a more reactive substance in with a less reactive one where the former is in an ion form that prompts it to undergo the opposite, no reaction will occur (e.g. I2 + 2Cl- will not react as the Cl is a stronger oxidising agent and this reaction would force it to be oxidised).

Question 14

Acid and Metal Reaction and when will this react

Answer 14

NOT NEUTRALISATION. Produces hydrogen gas.

Protons (H+ which are in acids) can only be reduced (is an oxidising agent) to H2, meaning it will only react when with a metal that is a stronger reducing agent.

Question 15

What area is more reactive for metals and non metals on periodic table.

Answer 15

Question 16

Voltaic Cell

Answer 16

Converts chemical energy to electrical energy through a spontaneous chemical redox reaction driving electrons around the circuit

Question 17

Electrolytic Cell

Answer 17

Converts electrical energy to chemical energy through an electrical current that reverses the normal direction of chemical change in a non-spontaneous reaction.

Question 18

Anode and Cathode

Answer 18

AN OX AND RED CAT - Anode = Oxidation, Cathode = Reduction

Voltaic - Anode is negative and cathode is positive.

Electrolytic - Anode is positive and cathode is negative.

ELECTRON FLOW FROM ANODE TO CATHODE.

Question 19

What do half cells do initially

Answer 19

The metal will loose cation particles (e.g. Zn^2+) to the electrolyte solution, making this solution more positive, and the metal itself will retain the electrons and will become negatively charged on its surface. Eventually, an equilibrium will be formed as shown until the voltaic cell is completed.

Question 20

What happens at the anode (voltaic)

Answer 20

The electrons transfer from the anode to the cathode and the concentration of cations in the electrolyte build up a positive charge, causing them to flow along the salt bridge to the cathode whilst also receiving anions from the cathode.

Question 21

What happens at the cathode (voltaic)

Answer 21

Electrons come from the anode to the cathode and it becomes negatively charged, attracting the cations from the electrolyte which build up on the metal’s surface. The electrolyte becomes negatively charged, which is offset by the flow of cations from the anode and anions leaving the cathode to the anode along the salt bridge.

Question 22

Cell diagram conventions

Answer 22

• Single vertical line=phase boundary between electrode and solution.
• Double vertical line=salt bridge
• Aqueous solutions placed next to the bridge
• Anode on the left and cathode on the right. - Spectator ions omitted.

Question 23

Advantages and disadvantages of primary general cells.

Answer 23

Advantages:
- Cheap
- Lightweight
- Long shelf life

Disadvantages:
- Single use = more waste.
- Environmental impact.
- Small current delivery

Question 24

How does an acidic hydrogen-oxygen fuel cell work.

Answer 24

Anode - H2 comes in, hits the fuel cell membrane and undergoes oxidation, the e-‘s going through a wire to the cathode and the protons going through the membrane to the cathode.

Cathode - O2 comes in and undergoes reduction, reacting with the H+ and e-‘s and makes H2O.

Energy is harvested as the e- flow from anode to cathode.

Question 25

Advantages and disadvantages of rechargeable/secondary cell

Answer 25

Advantages: - Longer life/more use = less waste. - Can deliver high currents. Disadvantages: - Expensive (apart from led-acid batterys) - Contain toxic heavy metals - Self-discharge even when not in use.

Question 26

Advantages and disadvantages of primary fuel cells.

Answer 26

Advantages: - More efficient than combustion with more efficient chemical energy transformation to useful electrical energy. - No pollution (in hydrogen-oxygen cells) Disadvantages: - Hydrogen is explosive. - Transporting hydrogen is difficult. - Expensive

Question 27

What makes an electrochemical cell reversible?

Answer 27

All products being soluble in the electrolyte, creating an equilibrium with its dissolved ion.

Question 28

Hows does an electrolytic cell work

Answer 28

The electrolyte is a single liquid metal which both electrodes are in, allowing the cation and anion to flow freely. The power supply (little line - negative so e- flow from it to cathode) is connected and the electrons flow towards the cathode where the cations are attracted to the negative charge and are reduced by this, building up the solid metal on the cathode. At the anode, the electrons flow away from it, creating a positive electrode that attracts the anions which loose their electrons and are oxidised, becoming a gas. Electrodes are inert (do not react).

Question 29

How to determine what is at the anode/cathode in electrolytic cells?

Answer 29

It does not follow the same rules as voltaic cells and it can go both ways depending on the flow of electrons (direction of power supply). But generally the anion in the molten metal will be attracted to the positive electrode (the anode=oxidation) and the cation will be attracted to the negative electrode (the cathode=reduction).

Question 30

How to determine what is at the anode/cathode in voltaic cells?

Answer 30

The better reducing agent is shown in the data booklet - the species higher on that list will be the reducing agent and undergo oxidation.

Question 31

Primary Alcohol Oxidation

Answer 31

Becomes aldehyde then carboxylic acid. Requires oxidising agent and heat. Can isolate the aldehyde through distillation (gas rises and is diverted down a chamber to condense) or get the carboxylic acid through reflux (aldehyde gas rises and is cooled and returned back to the reaction mixture). If dichromate is used, it will go from orange to green.

Question 32

Secondary Alcohol and Tertiary

Answer 32

Becomes Ketone. Has water by-product. Requires heat. Tertiary Alcohols do not do oxidise. If dichromate is used, it will go from orange to green.

Question 33

Reduction of carboxylic acids and aldehydes to make primary alcohols. Include conditions.

Answer 33

H- (hydride ion) is the reducing agent, and usually LiAlH₄ or NaBH₄ are added to add H-. Conditions for carboxylic acid: Under heat and in an anhydrous (dry) ether solvent. Conditions for aldehyde: Under heat with NaBH4. The reaction cannot be stopped at the aldehyde phase.

Question 34

Reduction of unsaturated organic compounds.

Answer 34

Addition reaction. The two hydrogen atoms attach simultaneously. Requires a catalyst and heat. Alkynes can also become alkenes or alkanes through this depending on how much H2 is present.

Question 35

Voltaic cell voltmeter reading and Standard Hydrogen Electrode (SHE)

Answer 35

Voltmeter neasures the difference in electrode potential, which exists due to different tendencies to loose electrons (reactivity). Data booklet shows this through electron potential, where the more negative substance is a better reducing agent. Voltage of half cells cannot be measured independently, so they are standardised to a reaction with hydrogen. This hydrogen half-cell is H2(g) ⇌ 2H+(aq) + 2e- (goes both ways depending).

Question 36

Standard Electrode Conditions

Answer 36

* All solutions must have a concentration of 1.0 mol dm-3 * All gases must be at a pressure of 100 kPa * All substances used must be pure * Temperature is 298 K * If the half-cell does not include a solid metal, platinum is used as the electrode. Any half cell under these conditions is known as a standard half cell.

Question 37

Standard Hydrogen Electrode Half-Cell Diagram

Answer 37

Create equilibrium between the reduction of H+ ions and oxidation of H2 which is impacted to what this half cell is attached to.

Question 38

SHE connected with another half cell

Answer 38

Creates a circuit where the voltage generated is the standard electrode potential. Example with hydrogen as reducing agent.

Question 39

Rules for Standard Reduction Value and how to calculate it from half cells values

Answer 39

* All E° values refer to the reduction reaction. * The E° values do not depend on the total number of electrons = no stoich required. * More negative=better reducing agent. Calculated by: E°cell = E°(cathode) - E°(anode)

Question 40

What if half cells do not contain metal?

Answer 40

Platinum will be used as the electrode. Example shown here:

Question 41

E° and spontaneity

Answer 41

Voltaic cell will always run in the direction that gives a positive E° value. If E cell is positive, the reaction is spontaneous. If E° is negative, the reaction is non-spontaneous and the reverse reaction is spontaneous. Non-spontaneous reactions will not occur.

Question 42

Important thing to remember when finding a half-cells standard electrode potential

Answer 42

Some half-cells have multiple electrode potentials in the data booklet depending on the oxidation state change - check this.

Question 43

How to calculate ΔG from E°

Answer 43

ΔG = -nFE° where: n is the number of moles of electrons transferred in the reaction. F is Faraday's constant (charge carried by 1 mol of electrons) Units:ΔG=J, n=mol, F=C mol-1, E°=V=J/C Remember the negative. Convert ΔG to kJ at the end.

Question 44

Spontaneity and electrolytic cells

Answer 44

Electrolytic cells are non-spontaneous, hence require electrolysis.

Question 45

Aqueous Electrolytic Cells

Answer 45

The electrolyte competes with water molecules at both the anode and the cathode to be oxidised/reduced. The outcome is decided through: - The relative E° values of the ions compared to water - the higher (more positive/less negative) value will be reduced and vice versa for oxidation. - The relative concentrations of the ions in the electrolyte - E.g. NaCl at the anode the Cl's electrode potential is only slightly greater so at high concentrations it also reacts. - The nature of the electrodes.

Question 46

What is electroplating and how does it work?

Answer 46

The process of using electrolysis to deposit a layer of a metal on top of another metal or other conductive object. The object to be plated forms the cathode, and the electrolyte contains the metal ions which are to coat the object. The anode can, but does not have to be made from the metal being plated. Used for decoration or protection (e.g. galvanised iron to stop rusting).

Question 47

What hydrogen reactions occur at the anode and cathode for aqueous electrolysis

Answer 47

Cathode - 2H2O + 2e- ⇌ H2(g) + 2OH-(aq) Anode - 4H+ (aq) + O2(g) + 4e- ⇌ 2H2O(l)