lewis structure hc2-

2 min read 27-02-2025

The acetylide ion, HC₂⁻, presents a fascinating example of how to apply Lewis structure rules to a molecule with a negative charge and a triple bond. Understanding its structure is key to comprehending its reactivity and properties. This article will guide you through drawing the Lewis structure step-by-step.

Step 1: Count Valence Electrons

First, we need to determine the total number of valence electrons.

Hydrogen (H): 1 valence electron
Carbon (C): 4 valence electrons each (2 carbons = 8 electrons)
Negative Charge (-): 1 additional electron

Total Valence Electrons: 1 + 8 + 1 = 10 electrons

Step 2: Identify the Central Atom

Carbon is less electronegative than hydrogen, making it the central atom. We arrange the two carbon atoms together.

Step 3: Form Single Bonds

Connect the carbon atoms with a single bond (2 electrons). This uses 2 of our 10 valence electrons, leaving us with 8. Connect each carbon to the hydrogen with a single bond, using another 2 electrons (leaving 6).

Step 4: Complete Octet Rule (Where Possible)

We distribute the remaining 6 electrons around the carbon atoms to satisfy the octet rule (8 electrons around each carbon atom, where possible). However, we'll quickly see this isn't enough.

Step 5: Form Multiple Bonds

To satisfy the octet rule for both carbon atoms, we need to form a triple bond between the two carbon atoms. This involves moving two electron pairs from lone pairs on the carbons to create three bonds between them. This uses a total of 6 electrons. The final Lewis structure now shows a triple bond between the two carbons. One carbon has a single bond to the hydrogen atom.

Step 6: Formal Charges

Let's check the formal charges of each atom:

Hydrogen: 1 valence electron - 1 bond = 0 formal charge
Left Carbon (bonded to Hydrogen): 4 valence electrons - 1 bond - 2 lone electrons = +1 formal charge
Right Carbon: 4 valence electrons - 3 bonds - 0 lone electrons = -1 formal charge

The overall charge of the molecule (-1) is the sum of these formal charges (+1 + (-1) = 0). While the individual formal charges are non-zero, this is a valid structure for the ion.

The Final Lewis Structure of HC₂⁻

The final Lewis structure of the acetylide ion, HC₂⁻, shows a triple bond between the two carbon atoms, with a single bond between one of the carbons and the hydrogen atom. One carbon atom will have a +1 formal charge, and the other carbon will have a -1 formal charge.

     H-C≡C⁻

HC₂⁻: Properties and Importance

The linear structure and the presence of the triple bond contribute to the acetylide ion's properties. It's a strong nucleophile, readily reacting with electrophiles due to the negative charge. This makes it important in organic chemistry and as an intermediate in various reactions. It's also a component in some metal acetylides, which find applications in materials science.

Further Considerations

While this Lewis structure is a useful representation, more sophisticated models like molecular orbital theory provide a more complete description of the bonding in HC₂⁻. However, the Lewis structure serves as an excellent starting point for understanding its basic structure and reactivity.