How Bond Dissociation Enthalpy Affects Chemical Reactions
Most of the chemical conversion is built upon the formation and dissociation of bonds. Chemical bonds are created when electrostatic forces between atoms combine. Chemical bonds can only be broken with the application of energy. The bond dissociation energy is the amount of energy needed to dissolve a chemical bond. Understanding how to calculate reaction enthalpies from bond dissociation enthalpy does not automatically translate into an understanding of bond breaking energies or even of energies and interactions. This article gives a clear view on the bond dissociation enthalpy with examples of halogens.
What is Bond Dissociation Enthalpy?
Bond Dissociation Enthalpy is the term utilised to describe the quantity of energy necessary in an endothermic reaction to dissociate a chemical bond and generate 2 distinct atoms, each having one of the original mutual pair's electrons. Whenever a bond is dissolved through homolytic dissociation, bond dissociation enthalpy is the typical shift in enthalpy.
The concept of bond-dissociation enthalpy, which is linked to the phrase bond-dissociation energy and is frequently utilised interchangeably, is equivalent to that of bond-dissociation energy. While some researchers utilise the phrase "bond-dissociation enthalpy" to apply to the enthalpy shift at 298 K (DH0298), others employ the word "bond-dissociation energy" (D0) to relate to the enthalpy shift at 0 K.
Features of Bond Dissociation Enthalpy
The following are a few crucial aspects of the bond dissociation enthalpy idea:
It refers to the quantity of energy required to supply a chemical connection between two molecules.
It is a way to determine how strong a chemical bonding is.
It is equivalent to the bond energy level in particular for diatomic compounds.
According to some, the bond involving silicon and fluorine has the highest bond dissociation enthalpy.
The bond dissociation energies of covalent bonds involving atoms or molecules are referred to as weak.
Example of Bond Dissociation Enthalpy
The reaction's enthalpy, that is by design the bond dissociation enthalpy of the molecule AB, DH0298 (AB), equals to the energy needed to dissociate a homolytic bond at 298 K. The following gives an instance of the bond dissociation enthalpy of a diatomic molecule.
The first rule of thermodynamics states that thermal energy equal to the bond dissociation enthalpy is produced if the radicals A and B join once more to create the molecule AB. These concepts make it easy to calculate the energy of the system of a variety of straightforward but significant processes including the transfer of a single bond. Through deducting, the energy needed to dissociate the previous bond from the energy acquired from the newly created bond is accomplished.
Bond Dissociation Enthalpy of Halogens
As previously stated, in the particular instance of diatomic molecules like halogens, the bond dissociation enthalpy of halogens is equivalent to the bond energy. This happens since bond energy is the median of every one of the bond dissociation enthalpies of halogen bonds of the identical form in a molecule.
As the size of the atom rises, the bond dissociation enthalpy of the halogen group falls down the hierarchy. Nevertheless, because of the inter-electronic repulsion occurring in the small atom of fluorine, fluorine's bond dissociation enthalpy is lower than that of chlorine and bromine. Hence, the order of bond dissociation enthalpy of halogen decreases as
Cl2 > Br2 > F2 > I2
This implies the decreasing bond dissociation enthalpy of halogens.
Halogen acids strength
The compounds formed while halogens react with hydrogen gas are recognised as halogen hydracids or halogen acids holding the formula HX. The 4 hydracids are referred to as: HF - Hydrofluoric acid, HCl - Hydrochloric acid, HBr - Hydrobromic acid, and HI - Hydroiodic acid.
Halogen acid's bond dissociation energy and level of ionisation in the aqueous phase both influence the acid strength. Overall, when bond dissociation energy rises, the degree of dissociation correspondingly reduces. This implies increasing bond dissociation enthalpy, by decreasing acidic strength.
Since the order of bond dissociation enthalpy of halogen acids are HI< HBr<HCl<HF. Thus, the above hydracids increasing acidic order is represented as,
HF< HCl< HBr< HI
Examples of Strong and Weak Chemical Bonds
The strongest single bonds, as per bond dissociation enthalpy research, are SiF bonds. The bond dissociation enthalpy for H3SiF is over 50% stronger than the bond dissociation enthalpy for CH3F. Much more bond dissociation enthalpy is present for SiF3F. These findings have the effect of generating silicon fluorides in a variety of processes, including glassware etching, deportation in chemical synthesising, and volcano releases. Due to the significant variation in electronegativity involving silicon and fluorine, respectively, ionic and covalent bonding significantly contributes to the bond's overall ability. This is why the bond is so strong. Diacetylene is one of the strongest carbon single bonds, connecting two sp-hybridised carbon atoms.
In contrast, it is unclear where an extremely weak covalent bond ends and an intermolecular interaction begins. The weakest bonds with significant covalent characteristics are those formed by Lewis acid-base combinations involving components of transition metal and noble gases.
Conclusion
Therefore, it can be concluded that the bond-dissociation enthalpy (DH°) is the measure of the capacity of a diatomic molecule's chemical bond. The bond-dissociation enthalpy varies from the bond energy with the exception of diatomic molecules.
Key Features
Bond-dissociation enthalpy, average bond energy, and bond strength are other names for the enthalpy of bond formation.
Extra energy is required to break down chemical bonds, owing to the thermodynamic preference of chemical bonds. As a result, bond enthalpy estimates are often positive and comprise units.
A stronger bond signifies a higher bond enthalpy, and a more force is needed to dissolve it.
FAQs on Bond Dissociation Enthalpy Explained with Examples and Applications
1. What is bond dissociation enthalpy?
Bond dissociation enthalpy is the energy required to break one mole of a specific covalent bond in a gaseous molecule, resulting in gaseous atoms or radicals. It is a direct measure of the strength of a chemical bond and is typically expressed in kilojoules per mole (kJ/mol). A higher value indicates a stronger bond.
2. Can you provide a common example of bond dissociation enthalpy?
A classic example is the breaking of the H-H bond in a hydrogen molecule. The bond dissociation enthalpy for H₂(g) is +436 kJ/mol. This means that 436 kilojoules of energy are needed to break all the bonds in one mole of gaseous hydrogen, forming two moles of gaseous hydrogen atoms: H₂(g) → 2H(g).
3. How does bond dissociation enthalpy differ from average bond enthalpy?
The key difference lies in specificity. Bond dissociation enthalpy refers to the energy needed to break a single, specific bond in a particular molecule (e.g., the first O-H bond in water). Average bond enthalpy, however, is the mean energy required to break a particular type of bond (like an O-H bond) averaged across various molecules or within a polyatomic molecule where multiple such bonds exist.
4. What are the main applications of knowing bond dissociation enthalpy values?
Understanding bond dissociation enthalpy is crucial for several applications in chemistry, including:
Calculating Reaction Enthalpies: It allows for the estimation of the overall enthalpy change (ΔH) for a chemical reaction.
Predicting Reaction Mechanisms: Knowing which bonds are weaker (have lower BDE) helps predict which bonds are likely to break first in a reaction.
Assessing Chemical Stability: Molecules with high bond dissociation enthalpies are generally more stable and less reactive.
5. What factors determine the magnitude of bond dissociation enthalpy?
Several factors influence bond strength and thus the bond dissociation enthalpy:
Atomic Size: Smaller atoms form shorter and stronger bonds, leading to a higher BDE.
Bond Order: The BDE increases with bond order (i.e., triple bond > double bond > single bond).
Lone Pair Repulsion: Repulsion between lone pairs on adjacent small atoms (like in F-F or O-O) can weaken the bond and lower the BDE.
Hybridisation: A higher percentage of s-character in the hybrid orbitals results in a stronger bond and higher BDE.
6. Why is the bond dissociation enthalpy of F₂ an exception in the halogen group?
The bond dissociation enthalpy trend for halogens is Cl₂ > Br₂ > F₂ > I₂. Fluorine (F₂) is an exception to the general trend of decreasing bond strength down the group. This is because the fluorine atom is very small and has three lone pairs of electrons. The high electron density in a small space leads to significant lone pair-lone pair repulsion between the two F atoms, which weakens the F-F covalent bond and lowers its dissociation enthalpy compared to chlorine.
7. How does bond dissociation enthalpy help explain the relative acidic strength of HF and HCl?
In the gaseous phase, the H-F bond has a much higher bond dissociation enthalpy (565 kJ/mol) than the H-Cl bond (427 kJ/mol), making it a significantly stronger bond. This high bond strength makes it much more difficult for the H-F bond to break and release a proton (H⁺). Although other factors like hydration enthalpy play a role in aqueous solutions, the very high stability of the H-F bond is a primary reason why HF is a weak acid, whereas HCl, with its weaker bond, is a strong acid.
8. Which diatomic molecule has the highest bond dissociation enthalpy for a single bond?
For a single covalent bond, the diatomic hydrogen molecule (H₂) has the highest bond dissociation enthalpy, valued at approximately 436 kJ/mol. The strength of this bond is attributed to the very effective overlap of the 1s orbitals of the small hydrogen atoms, resulting in a short and very stable bond.
