Electrolysis - Chapter 8

Electrolysis

Uses electricity to break down a compound into its simple elements

Uses electrical energy to form chemical energy (reverse of Galvanic Cells)

How Electrolysis Works

The electrolyte is compound which is to be broken apart

The battery initiates electron movement

As the anions and cations break apart to the electrodes, distinct compounds are formed

Visual Explanations

What Electrodes Are Used?

Inert Electrodes are used (carbon, platinum, graphite)

So that the electrodes themselves do not interact with the reaction

Difference between spontaneous and non-spontaneous

Spontaneous doesn’t require external input for a reaction

Non-spontaneous requires external power stimulus for the reaction

Why do products need to be kept apart in electrolytic cell

If the products contact one another, they can form back to being reactants

When there are multiple reactants in the cell

Electrodes themselves, water and reactanst could all be reactants

Strongest oxidising agent (one with highest emf value - at cathode)

Strongest reducing agent (one with lowest emf value - at anode)

Example

Inert Vs Reactive Electrodes

Inert electrodes - don’t participate in reaction

Reactive electrodes - are reactants in reaction (use electrochemical series to identify possible reactants)

Galvanic Vs Electrolytic

Why Anode In Electrolytic Cell Is Positive

Electrons are being stripped/deprived contunally due to the power supply - hence turns into a positive region

“electron deficient”

Why Anode In Galvanic Cell Is Negative

Electrons are being continually formed and are accumulated, despite the electron flow occuring

“Electrons are born there”

Why Are There Gases As Products

1) Gases form in voltaic (galvanic) cells when oxidation or reduction involves ions or molecules that become neutral gas molecules after gaining or losing electrons.

2) Gases form in voltaic cells when oxidation or reduction converts aqueous ions into neutral molecules, which are unstable in solution and therefore evolve as gases.

Down Cell

An electrolysis cell that seperates NaCl into Na and Cl in commercial quantities

Is done through molten NaCl

Down Cell (Image)

Why Water IS Not In Down Cell

Water is a stronger oxidizing agent than Na ions - would create H₂ gas instead of Na (l) at the anode

Na ions also react very violently with H₂0

Why Is Iron Inert As An Electrode (Unclear)

Electrons are continually pumped through the electrode, preventing it from oxidizing

Why Are Aqueous Electrolytes More Preferable to Molten Electrolytes

It is more cost-effective

Energy to keep molten electrolyte is very high - keeping it in aqeous state requires less energy and cost

Membrane Cell

Used to create hydrogen gas, chlorine gas and NaOH

Membrane Cell (Visual Representation)

Benefits Of Membrane Cell

The NaOH is seperated from the NaCl

The aqeuous electrolyte reduces the need for intensive heat and, reduces costs on energy

Why Metals in cathode don’t corrode

Oxidation = corrosion

In cathode, excess of electrons are being fed in, preventing the metal oxidising and loosing electrons - there is a continual supply of electrons

Electrolytic Production Of Aluminium

Done through Hall-Heroult cell, and doen with molten aluminium

Visual Of Hall-Heroult Cell

Why is Cryolite (Na₃Al₂O₃) used

It is used as a solvent to pure alumiunm, to reduce the melting point

Allows to conduct electricity better

Reduces overall melting point (2072 - 1000 degrees)

Prevents excessive anode wear, as lower temperature extends lifespan

Equation at Anode (Graphite) - 2 Step Process

Two Step Process At Anode

The oxygen is oxidised to oxygen gas

The oxygen gas then immediately reacts with the carbon anode, producing oxygen gas

Reaction at cathode

Is lined with graphite

Overall Reaction

Why this current model is not ethical

Intensive energy requirement

Constant replacements of the graphite anode (is reactive - participates in the reaction) - very expensive

Releases CO₂

Ways to shift to another model

Using inert electrodes (platinum, gold) - not feasible on a global scale

Alcoa and Rio Tinto Idealogy - using proprietayr electrodes - makes O₂ gas instead of CO₂

Can reduce costs, improve aluminum production and be done sustainably

Common Features & Reasoning In Electrolytic Cells

Electroplating

Anotehr metal (thin layer) is cast upon another surface to prevent corrosion or oxidation of the metal underneath

Can have a stable metal or a strong reductant, used as sacrificial protection

How Does It Work

Cathode is the item needing to be plated - negative electrode

From the anode, ions are flowing out, keeping the electrolyte’s concentration of ions balanced

The ions are attracted to the negative electrode and react with the electrons to form a solid element

They then coat around the cathode (the item needing to be plated)

Visual Representation

Example with tin (at anode)

Example with tin (at cathode)

Faraday’s First Law

The mass of metal produced at the cathode is directly proportional to the amount of electric charge is passed through (Q - Coulombs)

Relationship for charge

Q (Coloumbs) = current (amps) * time (seconds)

Faraday’s Second Law

Though Mass is directly proportional to Q, it is done at different rates per different metals

The no. of moles of electrons required to create the mass is measured

Requiring less electrons = more yield, (sharper gradient)

e.g. Ag only needs 1 mol of electrons - has sharpest gradient

Charge of 1 mol of electrons

1 mol of e^- = 96,500C (coloumbs)

Relationship between Q & C (2nd Law)

Q = n(e^-) * F

e.g. Copper:

Cu⁺² + 2e^- ——→ Cu_(s) = 2×96,500 → 193,000 Coloumbs