The question asks about the characteristics of formal charges in the most stable Lewis structure. Formal charge is a concept used to determine the most plausible Lewis structure among several possibilities for a given molecule or ion. The stability of a Lewis structure is generally associated with minimizing formal charges on atoms.
A) Closest to zero — The most stable Lewis structure is generally the one where the formal charges on all atoms are closest to zero. This indicates a more even distribution of electrons and less charge separation, leading to greater stability.
The question asks for the specific term that describes the number of covalent bonds an atom can form, guided by the octet rule. We need to understand the definitions of the given options in the context of chemical bonding.
Correct Option: B) Covalency — This term specifically defines the number of covalent bonds an atom can form to achieve a stable electron configuration, often following the octet rule.
The bond order in a molecule or ion exhibiting resonance can be calculated by considering the total number of bonds between specific atoms and dividing it by the number of resonating structures or positions. For the carbonate ion (\(\text{CO}_3^{2-}\)), we need to draw its resonance structures and then apply the formula for average bond order.
D) 1.33 (4/3) is the correct bond order. This value indicates that each C-O bond in the carbonate ion is intermediate between a single bond and a double bond, which is characteristic of resonance.
The question asks to identify the concept that best explains the equal bond lengths in ozone (O₃) when a single Lewis structure cannot. This points to a phenomenon where electron delocalization occurs, leading to averaged properties.
Correct Option: D) Resonance is the concept that explains why a single Lewis structure cannot fully describe the bonding in molecules like ozone, leading to equal bond lengths that are intermediate between single and double bonds. The electrons are delocalized over multiple bonds, creating a resonance hybrid.
The question asks to explain why all four S–O bonds in the sulfate ion (\(\text{SO}_4^{2-}\)) are experimentally equivalent, using Lewis structure concepts. We need to consider how Lewis structures represent bonding and how to account for observed bond equivalency when a single Lewis structure cannot.
A) Resonance — The concept of resonance is used to explain the observed equivalency of bonds in molecules or ions where a single Lewis structure cannot accurately represent the bonding. In the sulfate ion, the experimentally equivalent S–O bonds are a result of electron delocalization over multiple resonance structures, making resonance the best explanation.
The question asks about the general rule for placing an atom at the center when drawing Lewis structures of polyatomic species. This rule is based on the atom's ability to form multiple bonds and its electronegativity.
Correct Option: B) The least electronegative (most electropositive) atom, other than hydrogen
This is the general rule for selecting the central atom in a Lewis structure. The least electronegative atom is typically the one that can form the most bonds and accommodate other atoms around it. Hydrogen is an exception because it can only form one bond, so it is always a terminal atom.