ib chem redox

Oxidising agent

• An oxidising agent is a substance that oxidises another atom or ion by causing it to lose electrons

• An oxidising agent itself gets reduced – gains electrons

• Therefore, the oxidation number of the oxidising agent decreases

what makes the cell rechareabel

• Reverse the redox reactions that occurred during discharge

• Restore the original chemicals at each electrode

“The products do not disperse/dissolve/migrate away from the electrodes (as gas).”

draw a diagram of apparatus to measure the standard electrode potenial of zinc

2 agents for reducing carbonyl compounds

1. LiAlH₄, in anhydrous conditions, commonly dry ether, followed by the addition of aqueous acid

   • This is the stronger of these reducing agents and can reduce carboxylic acids

2. NaBH₄, in aqueous or alcoholic solutions

   • This is the less hazardous of these reducing agents but it cannot reduce carboxylic acids

   Both of these reagents produce the nucleophilic hydride ion, H^– 

electrolyte

substanced conducting electricity as liquid/aq and is chemically decomposed in the process. when its a solid it is an insulator

Reducing agent

• A reducing agent is a substance that reduces another atom or ion by causing it to gain electrons

• A reducing agent itself gets oxidised – loses/donates electrons

• Therefore, the oxidation number of the reducing agent increases

battery

moves electyrons toward hwere they ddont wanna go - move electrons to the cathode

[o] meaning

heat with acidified potassiummdichromate

H+/Cr2O7 2-

how to oxidize primary alcohol to get an aldehyde NOT CARBOX

distillation

• aldehyde have no hydrogen bonds so lower boiling point

charge carriers

in wires - electrons

in electrolyte - ions

why do we use a salt bridge

Let’s say we didn’t have a salt bridge:

• Mg is losing electrons → Mg²⁺ builds up in the solution → solution becomes increasingly positive

• Ag⁺ is gaining electrons → forming solid Ag → Ag⁺ concentration decreases, leaving anions behind → solution becomes increasingly negative

This charge imbalance would:

• Create a voltage buildup,

• Stop the reaction because the solutions resist further electron flow.

• The salt bridge is filled with a neutral salt solution, like KCl or KNO₃, and allows ions (not electrons!) to move:

  • Anions (e.g., NO₃⁻) flow toward the anode (Mg) to balance the buildup of Mg²⁺

  • Cations (e.g., K⁺) flow toward the cathode (Ag) to balance the loss of Ag⁺

  This keeps both sides electrically neutral, so the redox reaction can keep going.

cell notation

Anode (oxidation) | solution (highest ox state) || solution (highest ox state) | Cathode (reduction)

FUEL CELL

electrochemical cell in which a fuel donates electrons at one electrode and oxygen gains electrons at the other electrode

• as fuel enters cell it becomes oxidised, setting up a potential difference (voltage) in the cell

H-O fuel cell components

• A reaction chamber with separate inlets for hydrogen and oxygen gas

• An outlet for the product - water

• An electrolyte of aqueous sodium hydroxide 

• A semi-permeable membrane that separates the hydrogen and oxygen gases

  The half equations are

2H₂ (g) + 4OH^– (aq)  →  4H₂O (l) +  4e^–                     E^θ = -0.83 V 

O₂ (g) +  2H₂O  +  4e^– →  4OH^– (aq)                      E^θ = +0.40 V 

• The overall reaction is found by combining the two half equations and cancelling the common terms:

2H₂ (g) + 4OH^– (aq) + O₂ (g) +  2H₂O   +  4e^– →   4H₂O (l) +  4e^– + 4OH^– (aq)

2H₂ (g) + O₂ (g)  →   2H₂O (l)           E^θ = +1.23 V

benefits of fuel cells (4)

• only product is water → no pollution

• The reaction is the same as hydrogen combusting in oxygen, but since the reaction takes place at room temperature without combustion, all the bond energy is converted into electrical energy instead of heat and light

• No NOₓ (nitrogen oxide) emissions cuz it not in high temp

• Water produced can be used on spacecraft)

1 difference between fuel and other cells

the cell operates continuously as long as there is a supply of hydrogen and oxygen

• The energy is not stored in the cell

limitations of fuel cells

• Hydrogen is flammable → safety risk

• Storage is difficult → needs heavy, costly containers

• Currently made from crude oil → non-renewable

• hydrogen has High energy per gram but low energy per volume → bulky storage

voltaic vs electrolytic cells

• Voltaic cells generate electricity from chemical reactions

  • This is a spontaneous reaction which drives electrons around a circuit

• Electrolytic cells drive chemical reactions using electrical energy

  • An electric current reverses the normal directions of chemical change and this is non-spontaneous

secondary (rechargeable) cell

employ chemical reactions which can be reversed by applying a voltage greater than E°cell, causing electrons to push in the opposite direction

• lead-acid

• NiCad

• Lithium-ion

lead acid battery

electrolytic cell

• n electrolysis, the substance that the current passes through and splits up is called the electrolyte

• The electrolyte contains positive and negative ions

• Negative ions move to the anode and lose electrons - this is oxidation

• Positive ions move to the cathode and gain electrons - this is reduction

• Since metals are always cations and non-metal anions, it is easy to predict the products of electrolysis of molten salts:

  • Metals will always be formed at the cathode and non-metals at the anode

distillation and reflux (reboiling) of alcohol

• To  produce an aldehyde from a primary alcohol the reaction mixture must be heated

• The aldehyde product has a lower boiling point than the alcohol ( since it has lost the H-bonding) so it can be distilled off as soon as it forms 

FORMATION OF ALDEHYDE FROM CARBOXYLIC ACID

• You have to use LiAlH₄ refluxed in dry ether, followed by dilute acid

• This reaction cannot be stopped at the aldehyde because the LiAlH₄ is too powerful

• To form an aldehyde from a carboxylic acid, you have to reduce the carboxylic acid down to a primary alcohol and then oxidise it back up to the aldehyde

draw, using a diagram, essential components of the electrolytic cell

power supply, connecting wires, pos/neg electrode, and labelled electrolyte eg. NaCl

state one example showing economic importance of electrolysis

electroplating / purifying a metal

standard hydrogen electrode

• Standard Hydrogen Electrode (SHE) = reference electrode at E° = 0.00 V

• Involves hydrogen gas, acid solution (to be reduced, often a HCL), and an inert platinum electrode. MUST BE 1 MOL OF H

• Balances: 2H⁺ + 2e⁻ ⇌ H₂

• Used to measure standard electrode potentials of other half-cells by connecting and using a high-resistance voltmeter.

connecting another half cell to the SHE allows us to measure that half cells standard EP relative to SHE’s 0v

• If the half-cell pushes electrons toward the SHE → its E° is negative.

• If the half-cell pulls electrons from the SHE → its E° is positive.

how to find EMF (electromotive force)

EMF = E°(reduction) - E°(oxidation)

pls dont multiply them if u balance lol they are standardized

aqueous hydrolysis tip

Tables usually show standard reduction potentials (E°), so:

• A higher (more positive) E° → easier to reduce (not oxidize).

• A lower (more negative) E° → harder to reduce → easier to oxidize!

✅ So:

• Species with lower reduction potential (more negative E°) are harder to reduce and thus easier to oxidize.

• More negative E° → more likely to be oxidize

why can polyatomic ions not oxidize

polyatomic ions central atom is already at its maximum oxidation number so cant lose more electrons. water will win and get oxidized at anode

what to do at cathode vs anode at a an aqueous hydrolysis reation

• At the cathode:

  • Always reduce the species with more positive E° (no difference if electrode is active or inert).

• At the anode:

  • Inert electrode → oxidize water (if anion like SO₄²⁻ can't oxidize).

  • Active electrode → electrode itself can dissolve (like copper oxidizing into Cu²⁺).