Question

For each of the following bonds, draw a figure indicating the direction of the bond dipole, including which end of the bond is positive and which is negative. a. $\mathrm{H}-\mathrm{C}$ b. $\mathrm{N}⟶\mathrm{O}$ c. $\mathrm{N}-\mathrm{S}$ d. $\mathrm{N}-\mathrm{C}$

Short answer

a. $\mathrm{H}⟶\mathrm{C}$ (H is the positive end, C is the negative end) b. $\mathrm{N}⟶\mathrm{O}$ (N is the positive end, O is the negative end) c. $\mathrm{S}⟶\mathrm{N}$ (S is the positive end, N is the negative end) d. $\mathrm{C}⟶\mathrm{N}$ (C is the positive end, N is the negative end)

Step-by-step solution

Determine the electronegativities of the atoms involved in each bond

To draw the bond dipole, we first need to know the electronegativity of each atom involved. Here are the approximate electronegativity values (in Pauling scale) of the elements involved: - Hydrogen (H): 2.2 - Carbon (C): 2.5 - Nitrogen (N): 3.0 - Oxygen (O): 3.5 - Sulfur (S): 2.5 Step 2: Determine the bond dipole direction

Identify the direction of the bond dipole based on the electronegativity difference

Using the electronegativity values, we can identify the direction of the bond dipole for each bond. The direction will point from the less electronegative atom (positive end) to the more electronegative atom (negative end). a. $\mathrm{H}-\mathrm{C}$ Since H has a lower electronegativity than C (2.2 vs 2.5), the bond dipole points from H to C. The positive end of the bond is at the H atom, and the negative end of the bond is at the C atom. b. $\mathrm{N}⟶\mathrm{O}$ Here, N has a lower electronegativity than O (3.0 vs 3.5). The bond dipole goes from N to O. The positive end of the bond is at the N atom, and the negative end of the bond is at the O atom. c. $\mathrm{N}-\mathrm{S}$ In this case, N has a higher electronegativity than S (3.0 vs 2.5). The bond dipole points from S to N. The positive end of the bond is at the S atom, and the negative end of the bond is at the N atom. d. $\mathrm{N}-\mathrm{C}$ For the N-C bond, N has a higher electronegativity than C (3.0 vs 2.5). The bond dipole goes from C to N. The positive end of the bond is at the C atom, and the negative end of the bond is at the N atom. Step 3: Draw figures indicating bond dipoles

Draw figures with bond dipoles, including positive and negative ends

Based on the bond dipole directions identified in step 2, we can now draw the figures: a. $\mathrm{H}⟶\mathrm{C}$ (H is the positive end, C is the negative end) b. $\mathrm{N}⟶\mathrm{O}$ (N is the positive end, O is the negative end) c. $\mathrm{S}⟶\mathrm{N}$ (S is the positive end, N is the negative end) d. $\mathrm{C}⟶\mathrm{N}$ (C is the positive end, N is the negative end)

Key concepts

Electronegativity

Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond. Imagine each atom wanting electrons to join their side, and some atoms are better at insisting than others. This is all determined on the Pauling scale, which gives numerical values to these tendencies.

In general, higher electronegativity values mean the atom is more capable of drawing electrons towards itself. For instance, in the case of the elements discussed in the exercise, oxygen with an electronegativity of 3.5 can attract electrons more powerfully than nitrogen with 3.0.

• Hydrogen (H) has an electronegativity of 2.2
• Carbon (C) stands at 2.5
• Nitrogen (N) measures 3.0
• Oxygen (O) is at the top with 3.5
• Sulfur (S) equals carbon at 2.5

These values are key in determining how electrons are distributed between atoms in a molecule, hence affecting the bond polarity.

Dipole Direction

The dipole direction in a chemical bond indicates how the electron density is skewed between two bonded atoms. It points from the atom with lower electronegativity toward the one with higher electronegativity. This sets up a 'positive' end and a 'negative' end within the bond.

Think of it as a game of tug-of-war, where electrons are the rope. The stronger the team (more electronegative atom), the closer the rope gets to them. For a bond like H-C, the dipole direction goes from hydrogen to carbon because carbon's higher electronegativity means it pulls electrons closer. This results in hydrogen being the positive end, and carbon the negative end.

To summarize:

• Direction is from weaker (less electronegative) to stronger (more electronegative)
• The end with more electron density is negatively charged
• The other end becomes positively charged

These dipole directions help define the partial charges within a molecule, essential for predicting chemical behavior.

Chemical Bonds

Chemical bonds are the connections holding atoms together in molecules and come in various forms like covalent, ionic, and metallic bonds. In our focus on polar covalent bonds, such as those in the exercise, atoms share electrons, but not equally due to differences in electronegativity.

When a bond is described as polar, it means there is a noticeable difference in electronegativity between the two atoms, leading to a dipole moment. The electrons are shared but tend to hang out more with the atom having higher electronegativity.

Some key points include:

• Polar covalent bonds are formed between atoms with different electronegativities
• Non-polar covalent bonds have nearly equal sharing of electrons (very similar electronegativities)
• High electronegativity differs so much that electrons get completely transferred in ionic bonds

Understanding chemical bonds and the concept of polarity is crucial because they affect properties such as solubility, melting points, and electrical conductivity of substances.