Identify the wrong example from the following for the group 14 elements. (a) Element which forms most acidic dioxideCarbon (b) Element which is affected by water - Lead (c) Commonly found in \(+2\) oxidation state - Lead (d) Element used as semiconductor - Silicon.

Dioxides of Group 14 elements exhibit varying degrees of acidity. Acidity refers to a substance's ability to donate protons (H⁺ ions) in an aqueous solution, and in the case of dioxides, this acidity is determined by the ability of the oxygen to accept electrons. Typically, as one descends the group in the periodic table, the acidity of the dioxide decreases.

Starting with carbon dioxide (CO₂), it is a well-known acidic oxide which reacts with water to form carbonic acid (H₂CO₃). Moving down, the dioxides like silicon dioxide (SiO₂) are less acidic and do not readily dissolve in water as silicon has a more metallic character compared to carbon. As we reach the bottom of the group, lead dioxide (PbO₂) shows a marked decrease in acidity and is considered amphoteric - it can react with both acids and bases. Understanding the acidity trends in Group 14 is essential for predicting the chemical behavior of these elements and their oxides.

Lead (Pb), being one of the heavier elements in Group 14, has a more metallic character and shows less reactivity with water compared to its lighter counterparts. Lead does not react with cold or hot water, which can be surprising considering that metals generally tend to react with water to form hydroxides and release hydrogen gas.

The lack of reaction is due to the formation of a protective oxide layer on the lead's surface that prevents further interaction. Because of this inert behavior, lead is commonly found in applications where corrosion resistance is vital, such as in the construction of water pipes, though the use of lead in plumbing has decreased due to its toxicity. By understanding the reactivity of lead with water, students can illustrate the deviations in chemical properties within Group 14.

The common oxidation states of Group 14 elements are +4 and +2, and they are indicative of the group's position in the periodic table straddling metals and nonmetals. For carbon and silicon, the +4 oxidation state is predominant, as seen in carbon dioxide (CO₂) and silicon dioxide (SiO₂). This is because they can form four covalent bonds due to the availability of four valence electrons.

However, as we move to the heavier elements like lead (Pb), the +2 oxidation state becomes more stable. This stability arises from the inert pair effect, a phenomenon where the inner s-electrons are less prone to participate in bonding, leading to the +2 oxidation state being favored. Pb²⁺ is commonly observed in compounds such as lead(II) chloride (PbCl₂). Recognizing these oxidation states is crucial for predicting and understanding the chemical behavior and compounds formation of Group 14 elements.

Semiconductor materials have electrical conductivities that fall between conductors and insulators, making them indispensable in the world of electronics. Silicon (Si), a Group 14 element, is the most widely used semiconductor material due to its stable physical and chemical properties at a wide range of temperatures.

Silicon's ability to conduct electricity can be finely tuned by doping, which is the addition of impurities to create n-type or p-type semiconductors. This manipulation of electrical properties enables the creation of diodes, transistors, and integrated circuits, which are the building blocks of modern electronic devices. Understanding the properties of semiconductor materials like silicon opens a window into the intricate world of electronic engineering and computer chip design.