Draw the Lewis structure and determine the formal charge on each atom in (a) \(\mathrm{NO}^{+} ;\)(b) \(\mathrm{N}_{2} ;\) (c) \(\mathrm{CO} ;\) (d) \(\mathrm{C}_{2}^{2-}\); (c) \(\mathrm{CN}^{-}\).

Understanding the concept of formal charge is key to mastering chemical structures and reactions. A formal charge is a hypothetical charge on an atom in a molecule or polyatomic ion. We calculate it assuming that electrons in chemical bonds are shared equally between atoms, regardless of actual electronegativity differences.

The formula to calculate formal charge is: Formal charge = (Valence electrons) - (1/2 Bonding electrons) - (Non-bonding electrons).

Here's a breakdown of the terms:

• Valence electrons are the electrons in the outermost shell of an atom that can participate in chemical bonding.
• Bonding electrons refers to the electrons that are shared between atoms in bonds. In the formula, they're divided by two because we allocate half to each atom in the bond.
• Non-bonding electrons, also known as lone pairs, are the valence electrons not involved in bonding.

To ensure accuracy in structures such as \(\text{NO}^{+}\), or \(\text{N}_2\), one must carefully follow these calculations. A correct formal charge is critical for understanding the stability and reactivity of molecules. Incorrect formal charges can lead to a misunderstanding of the molecule's behavior in chemical reactions, hence the importance of meticulous calculation.

Valence electrons play a central role in Lewis structures and chemical bonding. They are the electrons located in the outermost shell of an atom and are key in determining how an atom can interact with others.

During the creation of Lewis structures, as demonstrated in our textbook solutions for molecules like \(\text{CO}\), and ions such as \(\text{C}_2^{2-}\), it's essential to count the correct number of valence electrons for each atom concerned. This count includes adding electrons for negative charges and subtracting for positive charges.

The number of valence electrons dictates the atom's bonding capabilities since they can be used to form bonds with other atoms, either by sharing electrons to form covalent bonds, donating or accepting electrons in ionic bonds, or by remaining as non-bonded lone pairs, which can affect the shape and polarity of the molecule.

To simplify, remember that elements in the same group on the periodic table have the same number of valence electrons, which allows for predictions about their chemical behavior and bonding patterns. Noting the wrong number of valence electrons can lead to incorrect Lewis structures and misinterpretation of a molecule's properties.

Chemical bonding is the force that holds atoms together in a molecule or compound. There are several types of chemical bonds, each with distinct properties and implications for the structure and stability of compounds.

The three main types of chemical bonds are:

• Covalent bonds, which involve the sharing of electron pairs between atoms. These are typically formed between non-metal atoms.
• Ionic bonds, which occur when one atom donates electrons to another, resulting in positive and negative ions that attract each other. These usually occur between metals and non-metals.
• Metallic bonds, found in metals where electrons are free to move around multiple nuclei, giving rise to properties like conductivity.

In context with the given exercise involving \(\text{CN}^{-}\) for instance, understanding chemical bonding helps one to determine how likely atoms are to share or transfer electrons.

Sketching Lewis structures, as performed in the exercise, is fundamental to visualizing these bonds. It requires knowledge of valence electrons and formal charges to predict molecular geometry, bond angles, and potential reactivity. Remember to distribute electrons carefully and use double or triple bonds if necessary when atoms need more electrons to complete their octet, as such bonding can greatly affect the molecule's overall properties and interactions.