Problem 109 - Chemistry: An Atoms-Focused Approach

Question

Hydrogen isocyanide (HNC) has the same elemental composition as hydrogen cyanide (HCN), but the H in HNC is bonded to the nitrogen atom. Draw a Lewis structure for HNC, and assign formal charges to each atom. How do the formal charges on the atoms differ in the Lewis structures for HCN and HNC?

Step-by-Step Solution

Answer

Question: Compare the formal charges on the atoms in the Lewis structures of Hydrogen Isocyanide (HNC) and Hydrogen Cyanide (HCN). Answer: In HNC, the formal charges are Hydrogen (0), Nitrogen (+1), and Carbon (-1). In HCN, the formal charges are Hydrogen (0), Carbon (+1), and Nitrogen (-1). The formal charges on the Hydrogen atoms in both structures are the same (zero), but the formal charges on the Nitrogen and Carbon atoms are flipped between the two structures.

Step 1: Count the total number of valence electrons

To draw the Lewis structure for HNC, the first thing we need to do is determine the total number of valence electrons. Hydrogen has 1 valence electron, Nitrogen has 5 valence electrons, and Carbon has 4 valence electrons. So, the total number of valence electrons in HNC is 1 + 5 + 4 = 10 electrons.

Step 2: Arrange the atoms and connect them with single bonds

In HNC, the Hydrogen atom is bonded to the Nitrogen atom. Since Nitrogen is more electronegative than Carbon, it will be placed in the middle. So, the arrangement of atoms can be written as H-N-C. Now we'll connect the atoms with single bonds. A single bond contains 2 electrons, so there are 10 - 2(2) = 6 electrons remaining.

Step 3: Distribute the remaining electrons

We now have 6 electrons left. We will place these electrons around the atoms to complete their octet (except for Hydrogen, which only needs 2 electrons to complete its valence shell). We'll add 2 more electrons to Nitrogen and 4 electrons to Carbon, which will complete their octets. There are no remaining electrons to add, so we have the final Lewis structure of HNC. The Lewis structure for HNC is: H - N = C Now let's calculate the formal charges for each atom.

Step 4: Assign formal charges to each atom

The formula for calculating the formal charge of an atom is as follows: Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - 1/2(Bonding Electrons) Let's calculate the formal charges for H, N, and C: 1. Hydrogen: There are 1 valence electron, 0 non-bonding electrons, and 2 bonding electrons (in the H-N bond). Formal Charge of H = 1 - 0 - 1/2(2) = 0 2. Nitrogen: There are 5 valence electrons, 2 non-bonding electrons, and 6 bonding electrons (in the H-N bond and N-C bond). Formal Charge of N = 5 - 2 - 1/2(6) = +1 3. Carbon: There are 4 valence electrons, 0 non-bonding electrons, and 8 bonding electrons (in the N-C bond and the triple bond to itself). Formal Charge of C = 4 - 0 - 1/2(8) = -1 The Lewis structure for HCN is: H - C ≡ N The formal charges for HCN are as follows: Hydrogen (0), Carbon (+1), and Nitrogen (-1).

Step 5: Compare formal charges

Now, we can compare the formal charges on the atoms in the Lewis structures of HNC and HCN. For HNC: Hydrogen (0), Nitrogen (+1), and Carbon (-1) For HCN: Hydrogen (0), Carbon (+1), and Nitrogen (-1) The formal charges on the Hydrogen atoms in both structures are the same (zero). However, the formal charges on the Nitrogen and Carbon atoms are flipped between the two structures. In HNC, Nitrogen carries a positive charge, Carbon carries a negative charge, while in HCN, Carbon carries a positive charge, and Nitrogen carries a negative charge.

Key Concepts

Formal Charge

Understanding formal charge is crucial when analyzing Lewis structures, as it helps us predict molecule stability and reactivity. The formal charge of an atom in a molecule can be calculated using the formula:
Formal Charge = (Valence Electrons) - (Non-bonding Electrons) - \(\frac{1}{2}\) (Bonding Electrons).
The idea is to simulate how each electron is distributed in a particular bond. If the formal charge is zero or close to zero, it indicates that the atom is in its most stable state. For example, in the exercise, we find:

• For Hydrogen in HNC and HCN, the charge is zero, reflecting hydrogen's stability with one bond.
• In HNC, Nitrogen holds a formal charge of +1 due to an uneven share of electrons.
• Conversely, Carbon in HNC has a formal charge of -1, indicating it possesses more electrons compared to its isolated state.

When comparing HNC to HCN, notice how the charges reverse for Nitrogen and Carbon, emphasizing their different electron-sharing configurations in the bonds.

Valence Electrons

Valence electrons play a pivotal role in chemistry, especially in constructing Lewis structures. These are the electrons available in the outermost shell of an atom. They are instrumental in forming chemical bonds.
For the molecule HNC, calculating valence electrons requires us to account for each atom's contribution:

• Hydrogen adds 1 valence electron.
• Nitrogen contributes 5 valence electrons.
• Carbon provides 4 valence electrons.

Collectively, HNC has a total of 10 valence electrons for bond formation and to be distributed throughout the molecule for stability. In drawing the Lewis structure, we use these electrons to create bonds and satisfy the octet rule for Carbon and Nitrogen (with Hydrogen only requiring 2 electrons). These defined bond allocations also help us determine formal charges by identifying electrons shared or owned by each atom.

Chemical Bonding

Chemical bonding refers to the interaction between atoms, leading to the formation of molecules. There are different types of bonds, including covalent (sharing of electrons) and ionic (transfer of electrons). In Lewis structures, these bonds are illustrated by lines.
In HNC, the chemical bond arrangement is depicted as:

• The bond between Hydrogen and Nitrogen represents a single covalent bond, sharing one pair of electrons.
• Between Nitrogen and Carbon, a triple bond is formed, depicting a stronger interaction with three shared electron pairs.

The distinction in bond types (single vs. triple) indicates different strengths and lengths of bonds. In chemistry, stronger bonds like the triple bond in HNC generally have higher energy and are shorter. Understanding these concepts allows us to predict molecular shape and interaction dynamics. Bonding not only defines molecular structure and arrangement but also influences properties such as boiling and melting points, solubility, and reactivity of molecules.