Question

Which statement(s) is/are true about bond enthalpy? (a) Energy is required to break a bond. (b) \(\Delta H\) for the formation of a bond is always a negative number. (c) Bond enthalpy is defined only for bonds broken or formed in the gaseous state. (d) Because the presence of \(\pi\) bonds does not influence the geometry of a molecule, the presence of \(\pi\) bonds does not affect the value of the bond enthalpy between two atoms either. (e) The bond enthalpy for a double bond between atoms \(A\) and \(B\) is twice that for a single bond between atoms \(\mathrm{A}\) and \(\mathrm{B}\).

Short answer

a. Energy is required to break a bond. b. ΔH for the formation of a bond is always a negative number. c. Bond enthalpy is defined only for bonds broken or formed in the gaseous state. Answer: a, b, c

Step-by-step solution

Statement (a)

Energy is required to break a bond. This is true. Bond enthalpy is the amount of energy required to break a bond and convert the participating atoms into individual gaseous atoms. Since breaking a bond requires energy to overcome the electrostatic attractions between atoms, this statement is true.

Statement (b)

\(\Delta H\) for the formation of a bond is always a negative number. This is true. When two gaseous atoms form a bond to create a stable molecule, energy is released as the electrostatic attractions between atoms are satisfied. This energy released is negative since it is being lost by the atoms. Therefore, the enthalpy change (\(\Delta H\)) for bond formation is a negative number.

Statement (c)

Bond enthalpy is defined only for bonds broken or formed in the gaseous state. This is true. Bond enthalpy is defined as the change in enthalpy due to the breaking or formation of a bond in the gaseous state. It is measured using gaseous state particles because it becomes easier to isolate the bond formation or breaking since gas-phase systems have no intermolecular forces.

Statement (d)

Because the presence of \(\pi\) bonds does not influence the geometry of a molecule, the presence of \(\pi\) bonds does not affect the value of the bond enthalpy between two atoms either. This statement is false. The presence of \(\pi\) bonds does influence bond enthalpy between two atoms because a \(\pi\) bond strengthens the overall bond and requires more energy to break than a single sigma bond. The statement is false because it incorrectly associates molecular geometry with bond enthalpy.

Statement (e)

The bond enthalpy for a double bond between atoms \(A\) and \(B\) is twice that of a single bond between atoms \(\mathrm{A}\) and \(\mathrm{B}\). This statement is false. While it is correct that a double bond between atoms \(A\) and \(B\) is stronger than a single bond between these atoms, it is not true that the bond enthalpy is exactly twice as high. The bond enthalpy of a double bond is greater than that of a single bond, but not necessarily twice as much. This is due to factors such as resonance and altered electron distribution that affect bond strength in double bonds.

Key concepts

Chemical Bond Energy

Understanding the energy associated with chemical bonds is fundamental in chemistry. Chemical bond energy, also known as bond enthalpy, is a measure of the strength of a chemical bond.

Imagine chemical bonds as the glue holding atoms together in a molecule. This glue isn't free; it requires energy to both establish and disrupt the bond between atoms. So when a bond is formed, energy is released, and conversely, when a bond is broken, energy must be supplied. The energy we're talking about here is the bond enthalpy.

For example, when we're considering the formation of a water molecule from hydrogen and oxygen gasses, energy is released as the H-O bonds are formed. We can consider this release of energy as a sign of the stability gained when atoms form a bond. On the other hand, if we want to break the H-O bond in a water molecule to release the hydrogen and oxygen atoms, we'd have to input a specific amount of energy that equals the bond enthalpy.

To assess a molecule's reactivity and determine reaction dynamics, one must consider the chemical bond energies involved. Moreover, as students solve problems related to bond enthalpies, understanding that the bond energy values provided in data books are average values and that actual energy changes can vary slightly depending on the specific molecular environment is a key insight for accurate analysis.

Enthalpy Change

Enthalpy change, denoted as \( \Delta H \), plays a pivotal role in thermodynamics and chemistry, as it represents the heat change in a reaction at constant pressure. It's an extensive property, meaning it depends on the system's scale or amount of matter.

The sign of \( \Delta H \) tells us if a process is exothermic (releases heat, \( \Delta H < 0 \)) or endothermic (absorbs heat, \( \Delta H > 0 \)). When you see a negative \( \Delta H \) value, like in the formation of a chemical bond, it signifies that energy is being released to the surroundings - a scenario common in the establishment of stable molecular structures.

Positive \( \Delta H \) values are just as informative. They highlight reactions that require energy input to proceed, such as the breaking of chemical bonds. In the classroom, students often calculate \( \Delta H \) for reactions using bond enthalpies, reinforcing the concept that reactions favoring the release of energy (exothermic) are generally more spontaneous than those that absorb energy (endothermic).

Improving textbook exercises involves not just memorizing that bond formation releases energy and is associated with a negative \( \Delta H \) but also understanding why this happens from a subatomic perspective. It’s the movement of electrons to lower energy states and the resulting stability that provides the exothermic characteristics of bond formation.

Molecular Geometry

The shape of a molecule, known as its molecular geometry, has an immense impact on the molecule's physical properties and reactivity. Molecular geometry is determined by the number of electron pairs, both bonding and lone pairs, surrounding the central atom.

For instance, water (\(H_2O\)) has a bent shape because of the two bond pairs and two lone pairs of electrons. These electron pairs repel each other, adopting a geometry that minimizes repulsion, which in the case of water, results in a roughly 104.5-degree bond angle.

The VSEPR (Valence Shell Electron Pair Repulsion) theory serves as a reliable method to predict molecular shape. Understanding this can help explain the physical properties such as polarity and intermolecular forces which dictate the state of matter, boiling and melting points, and solubility.

Molecular geometry doesn't directly influence the bond enthalpy between two atoms, but it does impact the overall stability and reactivity of molecules. Compounds with different shapes have different types of stress on their bonds and thus could require different amounts of energy to react. Students must grasp the significance of geometric arrangements to correctly predict reaction mechanisms and product formations, moving beyond simple rote learning to a solid understanding of how molecular shapes influence chemical behavior.