IB MYP 4-5 Chemistry : Types of chemical reaction  Study Notes

Redox
● Definitions
○ Reduction: A reaction that involves the gaining of electrons by one of the atoms involved in a reaction, or two or more chemical species. Oxidation of that element is lowered.

○ Oxidation: Is the loss of electrons during a reaction by a molecule, atom or ion. When oxidation happens the oxidation state of the molecule increases.
○ Reducing Agent: This is an element or compound that loses/donates an electron to another chemical species in a redox chemical reaction.

○ Oxidizing Agent: Is a substance that has the ability to oxidize other substances. In other words, it is the one that gains electrons.

○ The oxidation number is the charge on an element or molecule.
● Oxidation and Reduction can be remembered by the acronym OILRIG. Oxidation is loss of electron, while reduction is the gain of electrons.
● When trying to figure out which elements are oxidized and which are reduced by taking the following example:

•  Let’s break up the example above:

•  Firstly let’s take the Iron (Fe). Before reaction it has an Oxidation number of 0
•  Then let’s look at the \(\mathrm{O}_2\). It is diatomic. It even has an Oxidation number of 0
• Finally let’s take a look at the other side, where we have the element \(\mathrm{Fe}_2 \mathrm{O}_3\)

•  We know this element has a total charge of ” 0 “.
• Oxygen normally have a charge of -2 , meaning that 3 oxygen have a total charge of -6
• Iron has to be positive to cancel it out. So the iron in total has to have a charge of +6 , hence one Fe equals to +3 .
•  From this is can be concluded that oxygen is an oxidation agent, while iron is the reducing agent.
• Half equations an example would be:

\(\begin{aligned} & \mathrm{Cu}(\mathrm{s}) \longrightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{e}^{-} \\ & 2 \mathrm{Ag}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow 2 \mathrm{Ag}(\mathrm{s}) \\ & \mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq})+2 \mathrm{e}^{-} \longrightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})+2 \mathrm{e}^{-}\end{aligned}\)

or

• $\mathrm{Cu}(\mathrm{s})+2 \mathrm{Ag}^{+}(\mathrm{aq}) \longrightarrow \mathrm{Cu}^{2+}(\mathrm{aq})+2 \mathrm{Ag}(\mathrm{s})$
•  For a more detailed look at this can be seen on the Lewis structure level

$
\mathrm{HOOH} \longrightarrow \mathrm{H}_2 \mathrm{O}+1 / 2 \mathrm{O}_2
$

Electrolysis
● Definitions
○ Electrolysis: Is the passing of direct electric current through an ionic substance that is either molten or dissolved in a suitable solvent, producing a chemical reaction at the electrode.
○ Electrolyte: Is a chemical compound that conducts electricity by changing into ions when melted or dissolved into a solution.
○ Anode: Anode is where oxidation takes place
○ Cathode: Where reduction takes place.

○ Corrosion: is the irreversible damage or destruction of material due to a chemical or electrochemical reaction
○ Reactivity series: A series of metals from the most reactive to least.
○ Ore: A natural occurrence of a rock or sediment that contains sufficient minerals with economically important elements. Normally is combined with other elements.

○ In the diagram above you have two examples: Galvanic cells & Electrolytic cells. Galvanic cells are spontaneous and no power is needed. While electrolytic cells require power as the reaction is not spontaneous.
○ Salt bridge is used to allow the current to flow
● Electrolysis cells function due to the difference in charge. Normally one of the metals is more electronegative than the other. Hence, the electrons get attracted to the more electronegative end (the Cathode.) The cathode will normally loose mass as it becomes more soluble. While on the other hand, Cu will increase in mass as it loses electrons.
● Metals are often found in ores hence extracting them has to take place. Electrolysis can be used to extract a more reactive metal from the ore. This can be done through a similar process as above where the metal gets plated.

Electroplating
● Electroplating uses many of the same principles as mentioned above of electrolysis.
○ Firstly you have a solution known as electrolyte. Then in the solution two terminals are placed. They are known as electrodes (they can be anode and cathode). Then when electricity flows through the circuit the electrolysis solution start to split and plate on the cathode creating a thin layer.

● An industrial example of this is copper
○ There is a solution of copper containing compounds such as copper (ll) sulfate. In the solution there is an anode made from impure copper and a cathode made from pure copper.
○ As the reaction takes place copper gets dissolved from the anode as it loses electrons, while the cathode gains mass of the deposited copper.
○ This can be seen through half-life reactions

• Anode: \(\mathrm{Cu}—>\mathrm{Cu}^{2+}+2 e^{-}\)
• Cathode: \(\mathrm{Cu}^{2+}+2 e^{—->} \mathrm{Cu}\)

○ Half-life reactions are equations which show the oxidation and reduction taking place in a reaction.

Voltaic Cell
● Definitions
○ Salt Bridge: The purpose is to stop a reaction from reaching equilibrium too quickly. If salt bridge is not installed, then a high positive and high negative will accumulate on either side causing huge potential difference. Additionally, it helps to complete the circuit.
○ Half Cell: This is half of the normal cell normally consisting of one electrode.

•  As this is a spontaneous reaction, this means that for the flow of electrons no energy is needed. Moreover, the flow of electrons is electricity hence it produces electricity. An example of this would be the Baghdad battery.