Question

Without using Fig. 8.3, predict which bond in each of the following groups will be the most polar. a. \(\mathrm{C}-\mathrm{F}, \mathrm{Si}-\mathrm{F}, \mathrm{Ge}-\mathrm{F} \quad\) c. \(\mathrm{S}-\mathrm{F}, \mathrm{S}-\mathrm{Cl}, \mathrm{S}-\mathrm{Br}\) b. \(\mathrm{P}-\mathrm{Cl}\) or \(\mathrm{S}-\mathrm{Cl} \quad\) d. \(\mathrm{Ti}-\mathrm{Cl}, \mathrm{Si}-\mathrm{Cl}, \mathrm{Ge}-\mathrm{Cl}\)

Short answer

The most polar bonds in each group are: a. Si-F b. P-Cl c. S-F d. Ti-Cl

Step-by-step solution

Group a: C-F, Si-F, Ge-F

Compare the electronegativity differences for each bond: 1. C-F: electronegativity difference = | 2.55 - 3.98 | = 1.43 2. Si-F: electronegativity difference = | 1.90 - 3.98 | = 2.08 3. Ge-F: electronegativity difference = | 2.01 - 3.98 | = 1.97 The Si-F bond has the greatest electronegativity difference and therefore will be the most polar.

Group b: P-Cl, S-Cl

Compare the electronegativity differences for each bond: 1. P-Cl: electronegativity difference = | 2.19 - 3.16 | = 0.97 2. S-Cl: electronegativity difference = | 2.58 - 3.16 | = 0.58 The P-Cl bond has the greatest electronegativity difference and therefore will be the most polar.

Group c: S-F, S-Cl, S-Br

Compare the electronegativity differences for each bond: 1. S-F: electronegativity difference = | 2.58 - 3.98 | = 1.40 2. S-Cl: electronegativity difference = | 2.58 - 3.16 | = 0.58 3. S-Br: electronegativity difference = | 2.58 - 2.96 | = 0.38 The S-F bond has the greatest electronegativity difference and therefore will be the most polar.

Group d: Ti-Cl, Si-Cl, Ge-Cl

Compare the electronegativity differences for each bond: 1. Ti-Cl: electronegativity difference = | 1.54 - 3.16 | = 1.62 2. Si-Cl: electronegativity difference = | 1.90 - 3.16 | = 1.26 3. Ge-Cl: electronegativity difference = | 2.01 - 3.16 | = 1.15 The Ti-Cl bond has the greatest electronegativity difference and therefore will be the most polar. #Summary of Results# a. Most polar bond: Si-F b. Most polar bond: P-Cl c. Most polar bond: S-F d. Most polar bond: Ti-Cl

Key concepts

Polar Bonds

Polar bonds occur when there is an unequal sharing of electrons between two atoms in a chemical bond. This typically happens in covalent bonds, where atoms share electrons. In polar bonds, one atom attracts the shared electrons more strongly because it has a higher electronegativity, causing a partial negative charge to develop around it, and leaving a partial positive charge on the other atom. This creates a dipole. A classic example is the water molecule (H₂O), where oxygen is more electronegative than hydrogen, resulting in a polar covalent bond. The strength and direction of the dipole depend on the difference in electronegativity between the two atoms involved in the bond.

Electronegativity Differences

Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a bond. When comparing bonds, the difference in electronegativity between the two bonded atoms can tell us how polar the bond will be. A larger difference indicates a more polar bond. For example, in the step-by-step solution provided, the Si-F bond has the largest electronegativity difference of 2.08, making it the most polar bond in its group. Electronegativity differences can categorize bonds as nonpolar covalent (little to no difference), polar covalent (moderate difference), or ionic (large difference, usually above 2.0). These differences help predict the bond's behavior and properties.

Chemical Bonding

Chemical bonds are the attractive forces holding atoms together. They form to lower the potential energy of atoms, thereby creating stable molecules. The main types of chemical bonds are ionic, covalent, and metallic. Covalent bonds, where atoms share electron pairs, can be either polar or nonpolar.

• Ionic bonds occur when electrons are transferred from one atom to another, forming charged ions.
• Polar covalent bonds arise when there is unequal sharing of electrons, as one atom attracts the electron pair more strongly than the other.
• Nonpolar covalent bonds exist when atoms share electrons equally.

Understanding the type of bond helps explain the molecular structure and properties, like melting and boiling points, solubility, and conductivity.

Periodic Table Trends

An understanding of periodic table trends is valuable when predicting electronegativity differences and bond polarity. Electronegativity generally increases across a period (from left to right) and decreases down a group (from top to bottom) on the periodic table.

• Elements like fluorine, oxygen, and chlorine have high electronegativities because they are located towards the top right of the periodic table.
• Alkali metals and alkaline earth metals, found on the left side, have low electronegativities.

These trends arise because as electrons are added across a period, they are attracted more strongly by the increasing nuclear charge due to additional protons. Conversely, going down a group, the addition of electron shells reduces the effective nuclear pull on the valence electrons, resulting in lower electronegativity. Understanding these trends can help in predicting and explaining the properties of elements and their bonds.