Write the complete Lewis structure for (a) zinc cyanide; (b) potassium tetrafluoroborate; (c) barium peroxide (the peroxide ion is \(\mathrm{O}_{2}^{2}\) ).

Valence electrons are the electrons located in the outermost shell of an atom. These electrons are paramount when it comes to chemical bonding because they're the ones that participate in forming bonds with other atoms. For instance, in the case of zinc cyanide (Zn(CN)2), zinc (Zn) is a transition metal that normally adopts a +2 charge, indicating it has given up two valence electrons.

These valence electrons are then accepted by the cyanide ions (CN-), which consist of carbon (C) and nitrogen (N) atoms. Carbon and nitrogen both have their own valence electrons that they utilize to form triple bonds with each other in CN-, where three pairs of electrons are shared. Since zinc loses its valence electrons, its Lewis structure generally does not show these electrons, unlike the detailed structures you'd draw for nonmetallic elements and their compounds.

Understanding valence electrons is crucial for predicting how atoms will bond and what molecular structures will look like. This understanding lays the foundation for exploring more complicated concepts in chemistry, such as the bonding in ionic compounds like potassium tetrafluoroborate (KBF4) and coordination compounds, including the intriguing peroxide ion.

A coordinate bond, also known as a dative covalent bond, occurs when both electrons in the bond are contributed by one of the two atoms involved in the bond. In our example of potassium tetrafluoroborate (KBF4), the boron (B) atom forms a coordinate bond with one of the four fluorine (F) atoms. This bond is special because the fluorine atom donates both electrons that create the bond with boron.

The Lewis structure of KBF4 shows single bonds between the boron and three fluorine atoms, which is typical covalent bonding where each atom contributes one electron to the bond. The fourth bond is the coordinate bond where the boron atom, lacking a sufficient number of electrons to form conventional bonds with all four fluorine atoms, benefits from the generous fluorine that offers its lone pair to establish the bond.

This type of bonding is integral in forming complex ions and molecules, where an atom that has a lone pair can donate it to an electron-deficient atom. Coordinate bonds are significant in biochemistry and inorganics, often resulting in the formation of stable compounds. They offer a fine example of the nuanced ways that atoms come together to form the myriad substances around us.

The peroxide ion is an intriguing species with the formula \( \mathrm{O}_2^{2-} \). This ion features two oxygen atoms that are covalently bonded to each other with a single bond. In the context of barium peroxide (BaO2), this ion takes on a -2 charge. Each oxygen atom has six valence electrons to start with and shares one pair to form the bond between them. Additionally, they share another pair of electrons, giving rise to what's known as the peroxide linkage (-O-O-).

The presence of the peroxide ion delineates a special type of oxide, distinctly different from regular oxides by having a single bond between the oxygen atoms rather than the more common double bonds found in, for example, carbon dioxide (CO2). The Lewis structure of the peroxide ion shows each oxygen atom surrounded by three lone pairs of electrons in addition to the shared pairs, illustrating that despite the negative charges, the ion is stable.

Understanding the peroxide ion's Lewis structure is not only important in basic chemistry but also in practical applications where peroxides serve as oxidizing agents, in disinfection, bleaching, and as initiators in polymerization reactions. Overlooking the unique nature of this bond can lead to misunderstanding the reactivity and properties of substances that include the peroxide ion.