Problem 74 - Chemistry The Central Science In Si Units

Question

We can define average bond enthalpies and bond lengths for ionic bonds, just like we have for covalent bonds. Which ionic bond is predicted to have the smaller bond enthalpy, \(\mathrm{Li}-\mathrm{F}\) or \(\mathrm{Cs}-\mathrm{F} ?\)

Step-by-Step Solution

Answer

Cs-F has the smaller bond enthalpy due to larger ion size.

Step 1: Understand Bond Enthalpy

Bond enthalpy refers to the amount of energy required to break one mole of a bond in a gaseous state. In general, stronger bonds have higher bond enthalpies.

Step 2: Consider Ionic Bond Characteristics

Ionic bonds form between atoms with significantly different electronegativities, involving an electron transfer that results in attraction between ions. Factors like ion size significantly influence bond enthalpy in ionic bonds.

Step 3: Compare Ion Sizes

Lithium (Li) is a smaller ion than Cesium (Cs). Since the force of attraction in ionic bonds is inversely proportional to the distance between ions, and smaller ions have closer distances, Li-F will have a stronger attraction than Cs-F.

Step 4: Predict Bond Enthalpy Based on Ion Size

Given Li is much smaller compared to Cs, the 3F bond in Li-F is closer than in Cs-F. Hence, the bond enthalpy of Li-F is higher due to the stronger ionic bond, while Cs-F has a smaller bond enthalpy.

Key Concepts

Bond Enthalpy

Bond enthalpy is a measure of the energy needed to break a bond between two atoms in a gaseous state. It's a crucial concept when studying chemical bonding. The higher the bond enthalpy, the more energy is required to break the bond, indicating a stronger bond. In ionic compounds, bond enthalpy can tell us a lot about the stability and strength of the bond formed between ions.

In ionic bonds, the bond enthalpy is influenced by several factors, including the size of the ions involved and their charge. Since ionic bonds form due to the electrostatic attraction between oppositely charged ions, a strong attraction results in high bond enthalpy. The concept is essential for predicting chemical behavior, especially reactions where bonds are broken, and new ones are formed.

• Stronger bonds = Higher bond enthalpy
• Weaker bonds = Lower bond enthalpy

Ion Size

Ion size plays a significant role in ionic bond strength and hence bond enthalpy. Smaller ions, like lithium (Li), will generally have a stronger ionic bond than larger ions such as cesium (Cs). This is because the force of attraction between oppositely charged ions is stronger when the ions are closer together. Hence, the bond formed will be stronger, requiring more energy to break.

The size of the ion affects the distance between the ions in an ionic compound. For example, Li is much smaller than Cs. As the size of the ion increases, the distance between the ions also increases, weakening the electrostatic force of attraction. This results in a lower bond enthalpy.

• Smaller ions = Stronger bond = Higher bond enthalpy
• Larger ions = Weaker bond = Lower bond enthalpy

Electronegativity

Electronegativity is a measure of how strongly an atom can attract and hold onto electrons. In the context of ionic bonds, it helps explain why certain atoms can form ionic bonds together.

An atom with a high electronegativity, such as fluorine (F), tends to attract electrons strongly, which can result in electron transfer from a less electronegative atom, such as lithium (Li). This transfer of electrons produces ions that are oppositely charged, thus forming an ionic bond due to electrostatic attractions.

• High electronegativity = Strong tendency to form ionic bonds
• Low electronegativity = Weaker tendency to form ionic bonds

Gaseous State

The gaseous state is an important condition for measuring bond enthalpy as it ensures that the energy measured is under standardized conditions with no interference from other bonding forces.

When substances are in their gaseous state, the individual atoms or ions are not influenced by any intermolecular forces, making it easier to measure the true energy needed to break the bond. This is crucial for accurate comparisons of bond enthalpies across different chemicals.

• Gaseous state ensures standardized measurements
• No interference from intermolecular forces