Question

Revise (as necessary) your checklist that can be used to determine if a Lewis structure for a molecule is correct.

Short answer

To check if a Lewis structure for a molecule is correct, one can follow these steps: verify the sum of valence electrons, check the central atom, arrange remaining atoms, ensure the octets are complete, arrange any extra electrons, check individual formal charges and verify the presence of necessary multiple bonds.

Step-by-step solution

Verifying the Sum of Valence Electrons

Check the total number of valence electrons in the molecule. Valence electrons are those present in the outermost shell of an atom which can be involved in the formation of a bond. The total number of valence electrons should be equal to the sum of valence electrons of all individual atoms.

Checking the Central Atom

In general (but not always), the atom with the lowest electronegativity serves as the central atom. Central atoms are typically surrounded by other atoms and electron pairs. It's important to make sure this is represented accurately in the Lewis structure.

Placing the Remaining Atoms

Surround the central atom with the remaining atoms. Draw single bonds between the central atom and surrounding atoms. A single bond represents a pair of shared electrons.

Completing the Octets

Complete the octets of the atoms bonded to the central atom. This involves placing the remaining electrons as lone pairs on the atoms. The atoms (except hydrogen) should now have an octet of electrons.

Arranging Extra Electrons

If there are any leftover electrons after filling the octet of the central atom, place them on the central atom.

Checking Formal Charges

Check the formal charges of each atom. The formal charge is calculated as the number of valence electrons in the free atom, minus the number of lone pair electrons, and minus half the number of bonding electrons. The sum of all formal charges in the molecule should be zero, or match the overall charge of the molecule.

Verifying Multiple Bonds

If the central atom does not have a complete octet, form double or triple bonds as needed. After each modification, repeat the formal charge calculation to verify that the structure with the lowest formal charge is obtained.

Key concepts

Valence Electrons

Valence electrons are crucial in understanding chemical bonding. These are the electrons located in the outermost shell of an atom.
They are the ones that participate in forming chemical bonds, such as the bonds seen in Lewis structures.
Each atom has a specific number of these valence electrons, which corresponds to its group number on the periodic table. For example:

• Hydrogen has 1 valence electron.
• Oxygen has 6 valence electrons.

Calculating the total number of valence electrons in a molecule is the first step in drawing its Lewis structure.
You do this by adding up the valence electrons from each atom in the molecule.
This total helps in accurately arranging the electrons around atoms to depict correct bonding and lone pairs.

Formal Charge

Understanding formal charge is pivotal for ensuring the stability of a Lewis structure. Formal charge is a theoretical charge assigned to an atom in a molecule if all atoms have the same electronegativity.
The formal charge is calculated using the formula: \[ \text{Formal Charge} = (\text{Number of valence electrons}) - (\text{Number of lone pair electrons}) - \frac{1}{2}(\text{Number of bonding electrons}) \]An ideal Lewis structure has the formal charges as close to zero as possible.
This indicates a more stable molecule. When drawing the structure, adjust the arrangement of electrons to minimize formal charges and optimize stability.

• The sum of all formal charges in a neutral molecule should be zero.
• In a charged molecule, the sum should equal the overall charge.

Minimizing formal charges helps predict the most likely and stable structure of a molecule.

Octet Rule

The octet rule is a guiding principle in chemistry, stating that atoms tend to form bonds to have a full set of eight valence electrons, resembling the stable electron configuration of noble gases.
This rule is key in constructing Lewis structures. Most atoms, except for hydrogen and a few others, aim to complete an octet through bonding.

• Hydrogen needs only 2 electrons since it is satisfied with a "duplet," similar to helium.
• Central atoms in molecules are often adjusted by forming multiple bonds or dissociating lone pairs to satisfy the octet rule.

To achieve an octet, atoms will form single, double, or even triple bonds.
The octet rule helps in predicting a molecule's bonding capabilities and structural arrangement.

Electronegativity

Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
It is crucial when determining the central atom in a Lewis structure. Generally, the atom with the lower electronegativity is often central.
This placement allows for others to bond around it effectively, minimizing the formal charge.

• Fluorine is the most electronegative element.
• Generally, electronegativity increases across a period and decreases down a group on the periodic table.

If atoms have closer electronegativity values, they form more covalent character bonds.
Understanding electronegativity can guide decisions about bond formation and molecule polarity.

Single Bonds

In the realm of chemistry, single bonds represent the sharing of one pair of electrons between two atoms.
They are the simplest form of covalent bonds seen in Lewis structures. A single bond is depicted by a line connecting two atoms.
For example, in a water molecule (H₂O), each hydrogen atom is single-bonded to the oxygen atom.

• Single bonds provide a starting point for building molecular structures.
• They usually indicate a sigma bond, which is the strongest type of covalent bond.

However, single bonds might not always satisfy the octet rule by themselves.
In such cases, atoms might form double or triple bonds to complete their octets.
Despite this, single bonds play a pivotal role in maintain simplicity and predicting basic molecular structures.