Question

Define bond enthalpy. Bond enthalpies of polyatomic molecules are average values. Why?

Short answer

Bond enthalpy is the energy required to break one mole of a bond in a molecule in the gaseous state. Bond enthalpies of polyatomic molecules are considered average values because the same bond types in such molecules can have varied strengths depending on their different surroundings.

Step-by-step solution

Definition of Bond Enthalpy

Bond enthalpy, also known as bond energy, is the amount of energy required to break one mole of a bond in a chemical substance in gaseous state. It is usually expressed in units of kJ/mol. The energy is required to overcome the attractive forces between the atoms.

Bond Enthalpy of Polyatomic Molecules

In polyatomic molecules, the same bond types can exist in many places and the environment surrounding each bond can vary, affecting the bond's strength. Hence, the bond enthalpy values for the same bond types in different compounds, or even different bond types in the same compound, are usually obtained by taking an average over a range of molecules containing that type of bond.

Key concepts

Understanding Chemical Bonds

In the realm of chemistry, chemical bonds are the glue that holds atoms together to form molecules and compounds. These bonds arise from the attraction between the electrons of one atom and the protons in the nucleus of another. The strength and stability of a chemical bond depend on the sharing or transfer of electrons between the atoms involved.

There are three primary types of chemical bonds:

• Covalent bonds, where electrons are shared between atoms,
• Ionic bonds, formed by the transfer of electrons from one atom to another, and
• Metallic bonds, which involve a 'sea' of shared electrons moving among metal atoms.

Each bond type has a characteristic energy associated with it, known as the bond enthalpy. Understanding the concept of bond enthalpy provides insights into the stability of chemical structures and the energy required to transform them during chemical reactions.

Enthalpy Changes in Chemical Reactions

Enthalpy change is a measure of the heat evolved or absorbed in a chemical reaction carried out at constant pressure. It is a crucial concept in thermochemistry that reflects the energy changes associated with the breaking and forming of chemical bonds. When a chemical bond is broken, energy is absorbed, and when a bond forms, energy is released.

The enthalpy change of a reaction is calculated by subtracting the sum of the bond enthalpies of the reactants from the sum of the bond enthalpies of the products. This can be represented by the equation: \[ \Delta H = \sum \text{Bond enthalpies of products} - \sum \text{Bond enthalpies of reactants} \] Reactions resulting in a negative enthalpy change are exothermic, indicating that they release heat, whereas a positive enthalpy change means the reaction is endothermic and absorbs heat from the surroundings.

Polyatomic Molecules and Average Bond Enthalpy

Polyatomic molecules are compounds that contain more than two atoms. These can range from simple molecules like water (H2O) to large, complex proteins. What makes polyatomic molecules particularly interesting is the variety of bond types and bond environments they exhibit as compared to diatomic molecules. The same type of bond—for example, a C-H bond—can be present in various molecules but with different bond enthalpies.

This variation arises because the bond's environment affects its strength. Factors like electronic configuration, molecular geometry, and the presence of other atoms can influence the energy required to break a bond. As a result, scientists use an average bond enthalpy value, derived from an array of molecules containing the bond of interest, to generalize its energy for calculations. This approach simplifies the process of estimating the enthalpy changes in chemical reactions involving polyatomic molecules.