Electrolysis

Electrolysis is the last of our big concepts in electrochemistry. Unlike te other concepts, we provide electrical energy to force redox to occur (it is not spontaneous).

The key thing to be aware of in an assessment is whether the substance undergoing electrolysis is molten or an aqueous solution. If it is a solution, we need to allow for the electrolysis of water in our observations.

Batteries

Batteries apply all of the redox concepts we have learned so far. We need to apply our understanding of:

• electron transfer
• oxidation numbers
• ion-electron half-equations
• electrode potential calculations

In doing so, we can justify observations, including the voltage generated by a battery.

Spontaneous Reactions

We can use Standard Electrode Potentials to justify why a redox reaction occurs (spontaneous) or does not occur (not spontaneous.

We use the same calculations as we do for electrochemical cells. If the resultant is a positive number, the reaction does occur (spontaneous). If is is less than zero (is negative), then nothing happens.

We can then apply this knowledge to real life situations, such as what happens to a galvanised nail when the zinc coating gets scratched or corroded:

Because the oxidation of zinc, coupled with the reduction of iron, is spontaneous, the compromised section of the nail is protected from further corrosion/damage. Iron ions in the water are actually reduced to iron metal, so "repairing" the scratch/damage.

Electrode Potential Calculations

How do we predict the potential voltage of a cell?

Because we know the standards electrode potential for every half-cell, we can predict the voltage of any combination of half-cells:

Standard Electrode Potentials

We started off looking at the Hydrogen Half-Cell. This is used as a standard electrode to assign "voltages" to every other half-cell.

The convention is to have the hydrogen half-cell as the anode (on the left). However, some half-cells give a negative voltage. This means the hydrogen half-cell is actually being reduced (so is the cathode) and the other half-cell is being oxidised.

We use the values from these cells (with hydrogen) to assign standard electrode potentials to every possible half-cell (and redox pair). We use these to predict which pair would be at the anode and which would be at the cathode. The smaller electrode potential will be oxidised, so will be at the anode.

Cell Diagrams

Today we looked at an application of redox - electrochemical cells. There is a convention in Chemistry for how we draw these, and we explored this:

We draw this cell as:

Zn_(s)/Zn²⁺_(aq)//Cu²⁺_(aq)/Cu_(s)

• States of matter are critically important
• If two solutions (or a solution and a gas) are involved, we need a solid electrode. Usually carbon is used for solutions, and platinum for gases
• Two solutions (or a solution and a gas) are separated by a comma (,) in a cell diagram.
• When we have two solutions (or a gas and a solution), we write them in the order they appear in the half equation

Half Equations

Ion-electron half-equations give us more information about what is happening in a redox reaction:

We will practice this skill, as well as combining half-equations in our practical/double period tomorrow.

Oxidation Numbers

One of the quickest ways to decide if a reaction is redox or not is to see if any species have a change in oxidation number. To do this, we need to know how to assign oxidation numbers to atoms within species.

We then practiced this with some common redox species:

Introduction to Redox

We started today with a familiar experiment - burning magnesium.

We learned that this is an example of a "redox" reaction because:
- electrons are being transferred
- magnesium is having oxygen added to it
- the oxidation numbers of magnesium and oxygen both changed

We have some tasks in Education Perfect to work on over the next week, to solidify and develop our understanding of redox. The next concept we will explore is Oxidation Number.