Question

Write a Lewis formula for each of the following. (a) \(\mathrm{HClO}_{2}\) (c) \(\mathrm{HClO}_{4}\) (e) \(\mathrm{ClO}_{3}{ }^{-}\) (b) \(\mathrm{HClO}_{3}\) (d) \(\mathrm{BrF}_{3}\)

Short answer

Lewis Structures for the given compounds are drawn out as follows: (a) HClO2: Cl is single-bonded to each oxygen atom and to H, with 2 lone pairs of electrons remaining on Cl. One oxygen has a double bond with Cl. (b) HClO3: Cl is single-bonded to each oxygen atom and to H, with 1 lone pair of electrons remaining on Cl. One of the oxygens has a double bond with Cl. (c) HClO4: Cl is single-bonded to each oxygen atom and to H, with no lone pairs of electrons remaining on Cl. Two of the oxygens have a double bond with Cl. (d) BrF3: Br is single-bonded to each F atom, with 2 lone pairs of electrons remaining on Br. (e) ClO3-: Cl is single-bonded to each oxygen atom, with 1 lone pair of electrons remaining on Cl, and one of the oxygens carries the extra electron from the negative charge of the ion. The negative charge is shared among the three Oxygen atoms.

Step-by-step solution

Count the total number of valence electrons

The first step to writing a Lewis structure is to count how many valence electrons are present in the molecule or ion. Using the periodic table, each Group 1 element has 1 valence electron, Group 2 has 2, and so forth. However, note that after Group 2, things jump over to Group 13 which has 3, Group 14 has 4, Group 15 has 5, Group 16 has 6, Group 17 has 7, and Group 18 has 8. For polyatomic ions, add one electron for each negative charge and subtract one electron for each positive charge.

Start by arranging the atoms

Next, we want to sketch out a rough structure of the molecule/ion. Hydrogen and Halogens (except when it is a central atom like in interhalogens) are always terminal, so they get placed on the outside, while the rest of the atoms are connected usually by single bonds which you can signify with a straight line. These establish the 'skeleton' of the molecule.

Distribute the remaining electrons

After forming the bonds, the remaining electrons should be placed on the structure. Start adding the remaining electrons as lone pairs to the terminal atoms first, then to the central atom(s) if any are left over. Remember that each bond represents 2 electrons. Performing this step should yield correct Lewis structures for HClO2, HClO3, HClO4, and BrF3.

Check Formal Charges

Formal charge is a tool used to keep track of electrons during bond formation. A correct Lewis structure will typically be the one where the formal charges on each atom are closest to zero, and the sum of the formal charges equals the overall charge on the molecule or ion. To calculate the formal charge, use the formula: Formal Charge = [# of valence electrons on atom] - [non-bonded electrons + number of bonds]. Adjust the structure as necessary based on this check.

Draw Resonance Structures if necessary

Resonance structures are multiple Lewis structures that can be drawn for a molecule or ion. For ClO3-, we can draw three equivalent resonance structures showing that the extra charge is delocalized over the three Oxygen atoms. To represent resonance, a double-headed arrow is used and brackets are used for charged species. After this step, we should have correct Lewis structures for all compounds.

Key concepts

Valence Electrons

Valence electrons are the electrons in the outermost shell of an atom and are crucial in determining how atoms bond with each other. In a Lewis structure, these electrons are represented as dots surrounding the atoms. The number of valence electrons for an element is typically the same as its group number in the periodic table. For example, elements in Group 1 have 1 valence electron, while elements in Group 17 have 7 valence electrons. When determining the number of valence electrons in ions, it's important to adjust for the ion's charge by adding an electron for each negative charge or subtracting one for each positive charge.

Understanding valence electrons is vital as they are used to form bonds and lone pairs in Lewis structures. For instance, hydrogen has 1 valence electron and is often found with another electron to complete its duet, while elements in Group 17 often seek one additional electron to complete their octet.

Formal Charge Calculation

Formal charge is a concept used to assign charges to individual atoms within a molecule or ion. The calculation helps in determining the most stable Lewis structure among possible structures. The formula for calculating the formal charge on an atom is:
\[ \text{Formal Charge} = (\text{Number of valence electrons in the free atom}) - (\text{Number of non-bonded electrons}) - (\text{Number of bonds connected to the atom}) \]
The goal is to have formal charges that are as close to zero as possible, which often corresponds to the most stable structure. Calculating formal charges is crucial for molecules or ions that have more than one valid Lewis structure because it helps to identify the most favored structure based on the distribution of electrons.

To practice formal charge calculations, students should look at each atom in a molecule and apply the formula. Small mistakes in counting electrons or bonds can lead to incorrect formal charges, so attention to detail is paramount.

Resonance Structures

Resonance structures are alternative ways of drawing the Lewis structure for a molecule or ion, reflecting the delocalization of electrons within the molecule. These structures are particularly important when a single Lewis structure cannot fully explain the bonding in a molecule. For example, in the ion \( \mathrm{ClO}_{3}^{-} \), the negative charge can be associated with any of the three oxygen atoms, leading to three possible resonance structures where the double bond and the negative charge location rotate among the oxygen atoms.

In resonance structures, electrons are not static but are depicted as being spread out or shared between different atoms. This concept is represented by drawing multiple structures and connecting them with a double-headed arrow, indicating that the actual structure is a hybrid of the depicted resonance forms. When multiple resonance structures exist, they contribute to the overall stability of the molecule, and none of the individual resonance structures are viewed as the sole correct structure.