The correct order of acidic strength is: (a) \(\mathrm{Cl}_{2} \mathrm{O}_{7}>\mathrm{SO}_{2}>\mathrm{P}_{4} \mathrm{O}_{10}\) (b) \(\mathrm{CO}_{2}>\mathrm{N}_{2} \mathrm{O}_{5}>\mathrm{SO}_{3}\) (c) \(\mathrm{Na}_{2} \mathrm{O}>\mathrm{MgO}>\mathrm{Al}_{2} \mathrm{O}_{3}\) (d) \(\mathrm{K}_{2} \mathrm{O}>\mathrm{CaO}>\mathrm{MgO}\)

Understanding the acid-base properties of oxides can be tricky, but focusing on two key factors—electronegativity and oxidation state—can make it much simpler. Electronegativity refers to an atom's ability to attract and hold on to electrons. Generally, the higher an element's electronegativity, the more acidic its oxide.

Oxidation state, on the other hand, indicates the degree of oxidation of an atom within a compound. When it comes to non-metal oxides, a higher oxidation state usually equates to a greater acidic character. That's why in our example, \(\mathrm{Cl}_2\mathrm{O}_7\), with chlorine at a high oxidation state, is more acidic than \(\mathrm{SO}_2\) and \(\mathrm{P}_4\mathrm{O}_{10}\). This is a fundamental concept in chemistry and can help explain many reactions and properties of different compounds.

Oxides can be broadly classified into acidic, basic, and amphoteric, based on their reactions with acids and bases. Acidic oxides tend to react with bases to form salts and water, while basic oxides react with acids in a similar fashion. The acid-base character of an oxide is closely related to the position of its element on the periodic table.

For example, oxides of metals tend to be basic, whereas oxides of non-metals are typically acidic. This is directly linked to the metal's tendency to lose electrons and form cations, and the non-metal's propensity to gain electrons and form anions. By understanding this underlying chemistry, it becomes clearer why certain oxides behave the way they do and how they can be used in various applications.

The periodic table isn't just a chart of elements; it's a map that shows trends in properties, including acidity and basicity of oxides. If we move from left to right across a period, we generally see an increase in the acidic strength of oxides. This is because elements become more electronegative, which leads to more acidic oxides.

Down a group, basicity typically increases. Metal oxides become more ionic, which usually translates into a stronger basic character. That's why when comparing \(\mathrm{Na}_2\mathrm{O}\), \(\mathrm{MgO}\), and \(\mathrm{Al}_2\mathrm{O}_3\), we see an increase in basicity from aluminum oxide to sodium oxide. By remembering these trends, predicting the nature of an oxide becomes an easier task.

Some oxides don't neatly fit into the categories of acidic or basic—they're amphoteric, meaning they can react both as acids and bases. This dual nature makes them fascinating and versatile. Aluminum oxide, \(\mathrm{Al}_2\mathrm{O}_3\), is one such amphoteric oxide.

Amphoteric oxides are often found among metalloids or elements adjacent to the metalloid border in the periodic table. They provide an excellent opportunity for chemical reactions, as they can interact with a wide range of other substances, making them invaluable in both industrial processes and teaching about chemical reactivity.

• Can react with acids to form salts and water
• Can also react with bases to form salts and water
• Amphoteric oxides straddle the line between metals and non-metals in the periodic table

Understanding amphoteric behavior deepens our appreciation for the complexities of chemical properties.