Exam 4 Practice Problems – Step-by-Step Guidance

Q3. For the following molecule, which representation correctly represents the bond polarity using delta notation?

Background

Topic: Bond Polarity and Delta Notation

This question tests your understanding of how to represent bond polarity in molecules using the delta notation ($\delta^+$ and $\delta^-$), which indicates partial positive and partial negative charges due to differences in electronegativity.

Key Terms:

• Bond Polarity: A measure of how electrons are shared between atoms in a bond. If one atom is more electronegative, it attracts electrons more strongly, creating a partial charge.

• Delta Notation ($\delta^+$, $\delta^-$): Used to show partial positive and negative charges in polar covalent bonds.

• Electronegativity: The tendency of an atom to attract electrons in a bond.

Step-by-Step Guidance

1. Identify the atoms in the molecule and their relative electronegativities. For example, in HF, fluorine is much more electronegative than hydrogen.

2. Determine which atom will have a partial negative charge ($\delta^-$) and which will have a partial positive charge ($\delta^+$). The more electronegative atom (F) will attract electrons more strongly.

3. Draw the bond polarity using delta notation: place $\delta^-$ on the more electronegative atom and $\delta^+$ on the less electronegative atom.

4. Check the Lewis structure for lone pairs and the correct placement of charges.

Try solving on your own before revealing the answer!

Final Answer: The correct representation is the image showing H with $\delta^+$ and F with $\delta^-$.

Fluorine is more electronegative, so it attracts electrons more strongly, resulting in a partial negative charge ($\delta^-$) on F and a partial positive charge ($\delta^+$) on H.