Problem 23 - Chemistry Principles and Practice

Question

Balance the following reactions, and specify which species is oxidized and which is reduced. (a) \(\mathrm{Na}+\mathrm{FeCl}_{3} \rightarrow \mathrm{Fe}+\mathrm{NaCl}\) (b) \(\mathrm{SnCl}_{2}+\mathrm{FeCl}_{3} \rightarrow \mathrm{SnCl}_{4}+\mathrm{FeCl}_{2}\) (c) \(\mathrm{CO}+\mathrm{Cr}_{2} \mathrm{O}_{3} \rightarrow \mathrm{Cr}+\mathrm{CO}_{2}\)

Step-by-Step Solution

Answer

(a) Na is oxidized, Fe is reduced; (b) Sn is oxidized, Fe is reduced; (c) C in CO is oxidized, Cr is reduced.

Step 1: Identify the type of reaction and changes in oxidation states

For each reaction, identify the species that undergo changes in oxidation states to spot oxidation and reduction processes.

Step 2: Balance the first reaction (a)

In reaction (a) \(\mathrm{Na} + \mathrm{FeCl}_{3} \rightarrow \mathrm{Fe} + \mathrm{NaCl}\), oxidation occurs as \(\mathrm{Na}\) goes from 0 to +1 and reduction occurs as \(\mathrm{Fe}\) goes from +3 to 0. The balanced equation is: 3 \(\mathrm{Na} + \mathrm{FeCl}_{3} \rightarrow \mathrm{Fe} + 3\mathrm{NaCl}\). \(\mathrm{Na}\) is oxidized and \(\mathrm{Fe}\) is reduced.

Step 3: Balance the second reaction (b)

In reaction (b) \(\mathrm{SnCl}_{2} + \mathrm{FeCl}_{3} \rightarrow \mathrm{SnCl}_{4} + \mathrm{FeCl}_{2}\), \(\mathrm{Sn}\) is oxidized from +2 to +4, and \(\mathrm{Fe}\) is reduced from +3 to +2. The balanced equation is: \(\mathrm{SnCl}_{2} + 2\mathrm{FeCl}_{3} \rightarrow \mathrm{SnCl}_{4} + 2\mathrm{FeCl}_{2}\). \(\mathrm{Sn}\) is oxidized, and \(\mathrm{Fe}\) is reduced.

Step 4: Balance the third reaction (c)

In reaction (c) \(\mathrm{CO} + \mathrm{Cr}_{2}\mathrm{O}_{3} \rightarrow \mathrm{Cr} + \mathrm{CO}_{2}\), \(\mathrm{C}\) in \(\mathrm{CO}\) oxidizes from +2 to +4, while \(\mathrm{Cr}\) reduces from +3 to 0. The balanced equation is: \(3\mathrm{CO} + \mathrm{Cr}_{2}\mathrm{O}_{3} \rightarrow 2\mathrm{Cr} + 3\mathrm{CO}_{2}\). \(\mathrm{C}\) in \(\mathrm{CO}\) is oxidized, and \(\mathrm{Cr}\) is reduced.

Key Concepts

Oxidation States

Oxidation states, also known as oxidation numbers, are a way to keep track of electrons in atoms during a chemical reaction. They are useful in identifying how electrons are transferred or shared among different elements in a compound. Assigning oxidation states helps us to determine which elements are oxidized and which are reduced in reactions.

Here's how you can assign oxidation states:

• The oxidation state of an element in its pure form (like \(\mathrm{Na}\) or \(\mathrm{Fe}\) in reaction a) is always 0.
• In ions, the oxidation state is equal to the charge of the ion. For example, \(\mathrm{Cl}^-\) in a compound typically has an oxidation state of -1.
• For molecules, the sum of the oxidation states in a neutral molecule must be zero, and in an ion, it must equal the ion's overall charge.

By knowing these rules, it becomes easier to identify changes in oxidation states and determine which elements are oxidized or reduced in chemical reactions.

Redox Reactions

Redox reactions are a type of chemical reaction where oxidation and reduction occur simultaneously. These processes involve the transfer of electrons between chemical species. Understanding redox reactions is crucial for balancing the overall chemical equation and predicting the products formed in a reaction.

A key feature of redox reactions is the reciprocal change in oxidation states. One element's oxidation state increases (indicating oxidation), while another decreases (indicating reduction). For example:

• In reaction (a), \(\mathrm{Na}\) is oxidized from an oxidation state of 0 to +1, as it loses electrons.
• In the same reaction, \(\mathrm{Fe}\) is reduced from an oxidation state of +3 to 0, as it gains electrons.

Recognizing the transfer of electrons allows for deeper insights into the relationships and transformations that occur in chemical reactions.

Oxidation and Reduction Process

Oxidation and reduction processes are fundamental to understanding how redox reactions operate. Oxidation involves a loss of electrons, while reduction involves a gain of electrons.

To remember these processes, think of the acronym OIL RIG - Oxidation Is Loss, Reduction Is Gain. This can help identify electron transfer between reactants. In chemical reactions, it's common to see:

• Elements that lose electrons (and thus increase in oxidation number) are said to be oxidized. For instance, in reaction (b), \(\mathrm{Sn}\) loses electrons, increasing its oxidation number from +2 to +4.
• Elements that gain electrons (and thus decrease in oxidation number) are said to be reduced. In the same reaction, \(\mathrm{Fe}\) gains electrons, decreasing its oxidation number from +3 to +2.

Recognizing who "gives" or "receives" electrons in a reaction helps in deciding which species is the oxidizing agent and which is the reducing agent. Understanding these processes enables one to correctly balance chemical equations and predict reaction outcomes.