Question

Write skeleton equations for these reactions. a. sodium carbonate(s) \(\rightarrow\) sodium oxide(s) \(+\) carbon dioxide \((\mathrm{g})\) b. aluminum(s) \(+\) iodine(s) \(\rightarrow\) aluminum iodide(s) c. iron(II) oxide(s) \(+\) oxygen \((\mathrm{g}) \rightarrow\) iron (III) oxide(s)

Short answer

a. \(Na_2CO_3(s) \rightarrow Na_2O(s) + CO_2(g)\) b. \(Al(s) + I_2(s) \rightarrow AlI_3(s)\) c. \(FeO(s) + O_2(g) \rightarrow Fe_2O_3(s)\)

Step-by-step solution

a. Sodium carbonate decomposes into sodium oxide and carbon dioxide.

The reactants in this reaction are sodium carbonate (Na2CO3), and the products are sodium oxide (Na2O) and carbon dioxide (CO2). Writing the skeleton equation for this reaction, we have: Na2CO3(s) → Na2O(s) + CO2(g)

b. Aluminum reacts with iodine to form aluminum iodide.

The reactants in this reaction are aluminum (Al) and iodine (I2), and the product is aluminum iodide (AlI3). Writing the skeleton equation for this reaction, we have: Al(s) + I2(s) → AlI3(s)

c. Iron(II) oxide reacts with oxygen to form iron(III) oxide.

The reactants in this reaction are iron(II) oxide (FeO) and oxygen (O2), and the product is iron(III) oxide (Fe2O3). Writing the skeleton equation for this reaction, we have: FeO(s) + O2(g) → Fe2O3(s)

Key concepts

Skeleton Equations

When learning about chemical equations, it's important to start with something simple yet fundamental: skeleton equations. These are raw representations of chemical reactions, showing only the formulas of the reactants and products without balancing the quantities. Skeleton equations are like a sketch, providing the basic framework of a chemical process. For example, if you combine sodium carbonate with heat, it decomposes into sodium oxide and carbon dioxide. Showcasing this in a skeleton equation, we write:

• Na2CO3(s) → Na2O(s) + CO2(g)

This equation indicates the substances involved but not the stoichiometry, which means it does not account for how many molecules of one substance react with molecules of another. These initial equations are a crucial step for students to visualize the type of reaction taking place.

Decomposition Reactions

A decomposition reaction is a type of chemical reaction where a single compound breaks down into two or more simpler substances. This is like breaking a complex object into simpler components.

• For example, when sodium carbonate decomposes, it forms sodium oxide and carbon dioxide.

The general form of a decomposition reaction can be represented as:\[ AB ightarrow A + B \]In the case of sodium carbonate (\(\text{Na}_2\text{CO}_3\)), when decomposing, yields sodium oxide (\(\text{Na}_2\text{O}\)) and carbon dioxide (\(\text{CO}_2\)). This type of reaction is often driven by external energy sources like heat, light, or electricity to cause the breakdown.

Synthesis Reactions

Synthesis reactions are quite the opposite of decomposition reactions. Instead of a complex compound breaking into simpler ones, in a synthesis reaction, two or more simple substances combine to form a more complex product. It's like building a new structure from foundational elements.

• A classic example is the reaction of aluminum with iodine to form aluminum iodide: \(\text{Al}(s) + \text{I}_2(s) \rightarrow \text{AlI}_3(s)\)

These reactions follow the general form:\[ A + B \rightarrow AB \]Such reactions are often exothermic, meaning they release energy in the form of heat or light, which can make them quite spectacular to witness in a laboratory setting.

Oxidation Reactions

Oxidation reactions are fundamental to understanding many chemical processes. In chemistry, oxidation refers to the loss of electrons by a molecule, atom, or ion. Often, oxidation reactions involve a substance combining with oxygen.

• An example includes the reaction of iron(II) oxide with oxygen to form iron(III) oxide: \(\text{FeO}(s) + \text{O}_2(g) \rightarrow \text{Fe}_2\text{O}_3(s)\)

Here, iron(II) oxide loses electrons to oxygen, resulting in iron(III) oxide, which is the classic rusting process. These reactions can be identified by the change in oxidation state of the involved elements. They are crucial in various applications, such as metal corrosion and energy production in batteries, highlighting their wide-ranging importance in daily life.