Question

Pass a piece of paper around the group and ask each group member in turn to perform the next step in the process of determining a correct l.ewis structure (including formal charges on all atoms and resonance structures, if appropriate) for the following molecules and ions: \(\mathrm{N}_{2} \mathrm{H}_{4}, \mathrm{CCl}_{4}, \mathrm{CO}_{3}^{2-},\) and \(\mathrm{NH}_{4}^{+}\)

Short answer

\text{For each molecule/ion: 1. Count total valence electrons. 2. Draw skeleton structure. 3. Distribute electrons to satisfy octets. 4. Adjust with multiple bonds if necessary. 5. Calculate formal charges. 6. Draw resonance structures if needed. 7. Finalize the Lewis structure.}

Step-by-step solution

Calculate total number of valence electrons

For each molecule or ion, count the total valence electrons available for bonding. For anions, add electrons equal to the charge. For cations, subtract electrons equal to the charge.

Draw a skeleton structure

Sketch a skeleton structure for the molecule or ion with atoms connected by single bonds. Put the least electronegative atom at the center (except hydrogen, which can only form one bond).

Distribute electrons

Starting from the skeleton, distribute the electrons around the atoms to fill their octets (or duet for hydrogen). Place remaining electrons as lone pairs.

Determine the need for multiple bonds

If there are any atoms that do not have an octet and there are no remaining electrons, form double or triple bonds as needed by sharing electron pairs.

Calculate formal charges

Calculate the formal charges for all atoms using the formula: Formal Charge = (Number of valence electrons in the free atom) - (Number of lone-pair electrons) - (1/2 the number of bonding electrons).

Draw resonance structures if applicable

If there's more than one way to arrange the electrons that results in the same charge distribution, draw all valid resonance structures.

Check and finalize the structure

Make sure that the final structure has the correct number of electrons, follows the octet rule (or duet for hydrogen), and has the lowest possible formal charges, preferring negative charges on more electronegative atoms.

Key concepts

Valence Electrons

Imagine valence electrons as the hands of an atom, reaching out to hold hands with other atoms. These are the electrons found in the outermost shell of an atom and play a key role in the formation of chemical bonds. In the context of Lewis structures, they serve as the primary building blocks. Like people holding hands in a circle, atoms use their valence electrons to connect with neighboring atoms, forming molecules.

Each element has a characteristic number of valence electrons based on its position in the periodic table. Group 1 elements have one, while Group 18 have eight, except for helium, which has two. During the construction of Lewis structures, counting the valence electrons tells us how many connections—or 'hands'—an atom can offer in chemical bonding.

Formal Charge Calculation

Think of formal charge as a way to keep score of the electrons given, taken, or shared by an atom within a molecule. To calculate the formal charge, we use the formula: \[ \text{Formal Charge} = (\text{Number of valence electrons in the free atom}) - (\text{Number of lone-pair electrons}) - \left(\frac{1}{2} \text{the number of bonding electrons}\right) \]
Imagine each atom starts with a certain amount of cash (their valence electrons) and they 'pay' or 'receive' part of it through bonds (like transactions). The goal is to have each atom's 'account' reflect the arrangement where they have spent or received electrons in a way that reflects their 'original wealth', which ideally should result in a formal charge of zero. When we draft Lewis structures, if we find an atom 'owing' or 'owning' too much—that is, having a high positive or negative formal charge—we'll want to revisit our 'electron economy' and see if we can redistribute the 'wealth' more evenly.

Resonance Structures

Resonance structures are to molecules what different poses are to a dancer: they represent different arrangements of the same ensemble. While the atoms remain in the same place, electrons can be arranged in different ways to show the same total charge distribution. It's like having several choreographies that the dancer knows, each showcasing different steps but all fitting the same piece of music. Drawing all possible valid resonance structures for a given molecule or ion helps to visualize the full range of 'dances' that the electrons can perform, demonstrating the molecule's 'flexibility'.

It is crucial to note that real molecules exist in a state that is a hybrid mix of their resonance structures. They don't 'flip' between forms, rather, the actual molecular structure is more stable than any individual resonance structure would suggest, much like a dancer's fluid motion that is smoother than any individual pose.

Octet Rule

The octet rule is the chemical equivalent of the saying, 'good things come in eights'. It is based on the observation that atoms tend to bond in such a way that they have a full set of eight valence electrons, resulting in a noble gas electron configuration. This rule is a general guideline for building Lewis structures, much like following a building code when constructing a house. Atoms bond by either sharing, losing, or gaining electrons to achieve a 'complete' outer shell, which is often eight electrons—though hydrogen and helium are content with just a pair.

However, it's important to remember that the octet rule isn't absolute—some atoms like those in the third period and beyond can accommodate more than eight electrons due to their larger orbitals. Ensuring compliance with the octet rule (or duet rule for hydrogen) is how we give structure to the molecular 'home', providing stability and balance within a molecule or ion.