Bond Energy (bond enthalpy)
When a bond is formed between two atoms, energy is released. The same amount of energy is absorbed when the bond is broken to form neutral atoms. So, the bond energy is the general amount of energy needed to break all bonds of a specific type in one mole of the substance.
It is identified experimentally, by measuring the heat involved in a chain reaction. It is also called bond enthalpy, as it is a procedure of enthalpy change at 298 K. The enthalpy changes in splitting a particle into its part atoms are called, enthalpy of atomization. The bond energy is given in kj mol-1 which is the energy required to break an Avogadro’s number (6.02 x 1023) of bonds. It is also released when an Avogadro’s number of bonds are formed.
It may be noted that energies of several bonds are greater than those of single bonds. But a double bond is not two times as strong as a single bond or a triple bond is not thrice as strong as a single bond. It suggests that s- bond is more powerful than a p-bond. Likewise, a polar covalent bond is more powerful than a non-polar covalent bond.
Ionic Character and Bond Energy
Bond energy is a measure of the strength of a bond. The strength of a bond depends on the following aspects.
- Electronegativity difference of bonded atoms
- (ii) Sizes of the atoms
- (iii) Bond length
Let us think about, first the part played by electronegativity difference. Take a look at the bond energies of H-X type of substances, where X= F, Cl, Br, I. The table shows that electrons are not equally shared in between the bonded atoms i.e., HX. As the halogen atom is more electronegative, the bonded set is more drawn in towards the X atom and consequently, polarity establishes. This triggers additional attracting force for bonding.
From the difference between speculative bond energies and those determined by presuming equivalent sharing, it is possible to approximate relative electronegativities. The comparison of these values shows that the discrepancy is the greatest for HF and the least for HI.
Let us determine, the increase in the strength of the H-Cl bond, due to the ionic character present in it. The H-H bond energy is 436 kJ mol-1. It indicates 436 kJ of heat is needed to break the Avogadro’s number of H2molecules into individual atoms. Hence, bond energy per bond is 72.42 x 10-23kJ.
This is obtained by dividing 436 by 6.02 x 1023. As the bonding electron set is similarly shared in between the two H atoms, we can presume that each bonded H-atom contributes half of the bond energy i.e., 36.21×10-23kJ. Similarly, the bond energy for Cl2 is 240 kJ mol-1. Therefore, each Cl-atom should contribute 19.93 x 10-23 kJ to any bond, where sharing of an electron set is equivalent.
Let us, now consider, the bond in HCl. This bond is polar; however, we think about the electron set to be similarly shared. On adding up the bond energy contributions of H-atom and Cl-atom, we anticipate the bond energy of H-Cl to be 56.14 x 10-23kJ per particle which is the sum of 36.21×10-23kJ and 19.93 x10-23kJ.
For Avogadro’s number of HCl molecules, the calculated bond energy is 337.96 kJmol-1 which is acquired by increasing 56.14 x 10 -23 with 6.02 x 10-23. The experimentally found bond energy for HCl is 431 kJmol-1. The observed bond energy is significantly greater than the computed value which means a more steady H-Cl bond. This stability is because of the ionic character present in the molecule. The decreasing polarity from HF to HI shows a pattern towards equivalent sharing of electrons which is consistent with decreasing electronegativity from F to I.
The bonds with greater bond energy values have shorter bond lengths. The bond energies of C-to-C bonds being in the order C ≡ C >C=C>C-C. Their bond lengths remain in the reverse order i.e.,C -C > C=C > C ≡ C.
The distance between the nuclei of two atoms forming a covalent bond is called the bond length.
between the positive and negative centers (r
The bond lengths are experimentally determined by physical methods. The techniques might be electron diffraction, X-ray diffraction, or spectral studies. The covalent bond length in between 2 atoms is frequently but not constantly independent of the nature of the particles. For instance, in the majority of the aliphatic hydrocarbons, the C-C bond length is really near to 154 pm. The C-C bond length is also found to be the very same in diamonds.
The covalent radii for different elements are almost addictive in nature. The single bond covalent radius of carbon is 77 pm which is half of the C-C bond length (154 pm). Likewise, the covalent radius of Cl is 99 pm i.e., one-half of the Cl-Cl bond length (198 pm). So, the bond length of the C-Cl bond will be 77 + 99= 176 pm.
With an increase in electronegativity difference between the bonded atoms, the bond ends up being shortened. For example, Si-F bond length in SiF4 is found to be 154-159 pm, whereas the addition of their covalent radii (Si= 117 pm and F= 64 pm) give Si-F bond length to be equal to 181 pm. The calculated values are often higher due to electronegativity differences. The ionic character leads to reducing of the bond length due to the force of attraction in between the polar ends.
Furthermore, the hybridization plan included, likewise discusses the reduction of bonds due to the predominant participation of s-orbitals. Because the 2s-orbital of carbon has a smaller mean radius than the 2p-orbitals, it would be expected that the greater the s character in the hybrid orbitals utilized, the shorter will be the bond distance. Hence, the C-C bond lengths are 154,133 and 120 pm for ethane, ethene, and ethyne, respectively where s orbital contribution increases from sp3 to sp. Further, p-bonding likewise minimizes the internuclear bond distance. The bond length increases, as we move from top to bottom in group IV-A of the periodic table.
Hence, Si-Si bond length is more than C-C bond length in group IV-A and P-P bond length is a lot more than N-N bond length in group V-A. As the atomic radii increase in a group (N to P or C to Si), the effect of the effective nuclear charge decreases on electrons. As a result, the bond length will increase.
In the periodic table, reduction of bond lengths occurs from delegated right in, a period. This can be attributed to the pull by a nuclear charge with the same value of the principal quantum number. For that reason, C-C bond length is greater than N-N bond length.
In heteronuclear molecules, e.g., HCl where the bonded atoms are of different elements, the molecule ends up being polar due to the electronegativity difference. Partial positive and negative charges become separated on the bonded atoms. The separation of these charges on the particle is called a dipole and the molecule is stated to have a dipole moment. The dipole moment is a vector quantity, which has a magnitude along with a direction.
” The dipole moment (m) is defined as the product of the electric charge (q) and the distance between the positive and negative centers (r):”
µ= q x r
The dipole moments of simple heteronuclear diatomic molecules like HF, HCl, HBr. HI, CO, NO, and so on are directed from electropositive ends to electronegative ends.
The dipole moments are determined in Debye (D) systems. Let us think about a theoretical molecule (A *– B-), or a unit negative charge separated from a unit positive charge by distance r = 100 pm (1 Å) The dipole moment of such a molecule can be calculated by increasing the distance 100pm to charge of one electron or proton is 1.6 orx10-19C m= (1.6022×10-19C) x(100×10-12m) = 1.6022×10-29 mC.
Another unit of dipole moment is Debye. The equivalence of Debye and mC is 1 D = 3.336×10-30mC. So, the dipole moment of the, above system in Debye units is
Dipole moment provides two kinds of information about the molecular structure:
- (i) Percentage ionic character of a bond
- (ii) Angles between the bonds or the geometry of molecules
(1) Percentage ionic character of a bond
From the experimentally identified dipole moments, the percentage ionic character in a bond can be calculated. For this function, we ought to understand the real dipole moments of the particle and real bond, length. The dipole moment of 100% ionic compound is represented as µionic.
(2) Bond Angles or the Geometry of Molecules
We can understand this element by taking some important examples.
The dipole moment of water is 1.85 D which eliminated its linear structure. The estimations reveal that water has an angular structure with a bond angle of 104.5 ° in between the two O-H bonds. A linear water particle (H-O-H) would have absolutely no dipole moment. Similarly, the triatomic particles H2S or SO2 and so on are also bent like H2O.CO has a dipole moment while CO2 does not have any.
The factor is that CO2 has a linear structure, where the dipoles being equivalent and opposite, counteract each other’s effect. Likewise, CS2 has no dipole moment. Symmetrical triangular planar molecules of BF3, AlCl3, and completely tetrahedral particles like CH4, SiH4, CCl4 also have absolutely no dipole moments. This is all due to the cancellation of individual bond moments.