Redox Reactions - SS1 Chemistry Past Questions and Answers - page 2

In redox reactions, oxidising agents and reducing agents are involved in the transfer of electrons.

 

Oxidising Agents:

Oxidising agents are substances that cause oxidation by accepting electrons from other species. They are reducing themselves in the process. Oxidising agents are electron acceptors and are often characterised by their high electronegativity or ability to gain electrons.

 

 

 

Examples of common oxidising agents:

-       Oxygen (O2): Oxygen is a strong oxidising agent and is involved in many combustion reactions.

-       Hydrogen peroxide (H2O2): Hydrogen peroxide is a powerful oxidising agent commonly used as a bleach or disinfectant.

-       Potassium permanganate (KMnO4): Potassium permanganate is a strong oxidising agent used in analytical chemistry and organic synthesis.

 

Reducing Agents:

Reducing agents are substances that cause reduction by donating electrons to other species. They oxidise themselves in the process. Reducing agents are electron donors and are often characterised by their low electronegativity or ability to lose electrons.

 

Examples of common reducing agents:

-       Hydrogen gas (H2): Hydrogen gas is a common reducing agent used in various industrial processes, such as the production of ammonia.

-       Carbon monoxide (CO): Carbon monoxide is a reducing agent involved in metallurgical processes for the extraction of metals from their ores.

-       Sodium borohydride (NaBH4): Sodium borohydride is a mild reducing agent used in organic synthesis and as a reducing agent in some chemical reactions.

 

A redox reaction, also known as an oxidation-reduction reaction, involves the transfer of electrons between species. In these reactions, one species undergoes oxidation (loses electrons) while another species undergoes reduction (gains electrons). This transfer of electrons allows the species involved to achieve a more stable state. Redox reactions are essential in various chemical processes, including energy production, corrosion, and biological reactions.

 

Let's consider the example of the reaction between magnesium (Mg) and hydrochloric acid (HCl):

 

Mg + HCl → MgCl₂ + H₂

 

Step 1: Identify the oxidation states of each element in the reaction.

In this case, Mg is in its elemental form, so its oxidation state is 0. H is typically +1, and Cl is typically -1.

 

Step 2: Determine which species are being oxidised and reduced.

In this reaction, Mg is being oxidised from 0 to +2, while H is being reduced from +1 to 0.

 

Step 3: Balance the atoms other than hydrogen and oxygen.

In this case, we have one magnesium (Mg) atom on the left side and one on the right side, so they are already balanced. The same applies to the chlorine (Cl) atoms.

 

Step 4: Balance oxygen atoms by adding water (H₂O) molecules.

Since there are no oxygen atoms on the left side, we need to add water on the right side. Let's add one water molecule to balance the oxygen atoms:

 

Mg + HCl → MgCl₂ + H₂O

Step 5: Balance hydrogen atoms by adding hydrogen ions (H⁺).

To balance the hydrogen atoms, we need to add two hydrogen ions on the left side:

 

Mg + 2HCl → MgCl₂ + H₂O

 

Step 6: Balance the charges by adding electrons (e⁻).

Since magnesium (Mg) loses two electrons during oxidation, we need to add two electrons on the right side:

 

Mg + 2HCl → MgCl₂ + H₂O + 2e⁻

 

Step 7: Verify that the charges and the number of atoms are balanced.

The equation is now balanced in terms of atoms and charge:

 

Mg + 2HCl → MgCl₂ + H₂O + 2e⁻

 

This balanced equation represents the redox reaction between magnesium and hydrochloric acid.