Bond Enthalpy and Bond Order — Revision Notes

⚡ 30-Second Revision

Bond Enthalpy ($E_b$): — Energy to break 1 mole of bonds (gaseous state). Always positive (endothermic). Unit: kJ/mol.
Bond Order (BO): — Number of bonds between atoms.

- Lewis: 1 (single), 2 (double), 3 (triple). - MO Theory: $\text{BO} = \frac{1}{2}(N_b - N_a)$.

Relationships: — Higher BO $\implies$ Shorter Bond Length $\implies$ Higher $E_b$.
$\Delta H_{rxn}$ from $E_b$: — $\Delta H_{rxn} = \sum E_{b}(\text{bonds broken}) - \sum E_{b}(\text{bonds formed})$.
MO Energy Order:

- $\le 14$ e: $\sigma_{1s} < \sigma_{1s}^* < \sigma_{2s} < \sigma_{2s}^* < (\pi_{2p} = \pi_{2p}) < \sigma_{2p} < (\pi_{2p}^* = \pi_{2p}^*) < \sigma_{2p}^*$. - $> 14$ e: $\sigma_{1s} < \sigma_{1s}^* < \sigma_{2s} < \sigma_{2s}^* < \sigma_{2p} < (\pi_{2p} = \pi_{2p}) < (\pi_{2p}^* = \pi_{2p}^*) < \sigma_{2p}^*$.

Magnetic Properties: — Paramagnetic (unpaired e-), Diamagnetic (all paired e-).

2-Minute Revision

Bond enthalpy is the energy absorbed to break a chemical bond, signifying its strength. It's an average value for polyatomic molecules. Factors like bond order, bond length, atomic size, and electronegativity difference dictate its magnitude; stronger bonds have higher enthalpy.

Bond order quantifies the number of bonds between atoms, directly influencing bond strength and inversely influencing bond length. For simple molecules, it's an integer (1, 2, 3). For resonance structures or using Molecular Orbital (MO) theory, it can be fractional.

MO theory calculates bond order as half the difference between bonding and antibonding electrons ($N_b - N_a$). This theory is crucial for predicting magnetic properties (paramagnetic if unpaired electrons, diamagnetic if all paired) and for comparing the stability and bond lengths of diatomic ions.

Remember the specific MO energy level order for molecules with $\le 14$ electrons versus $> 14$ electrons. Enthalpy changes of reactions can be estimated by summing the energy of bonds broken in reactants and subtracting the energy of bonds formed in products.

5-Minute Revision

To master bond enthalpy and bond order, begin by solidifying definitions. Bond enthalpy is the energy required to break a bond (endothermic, positive value), while bond order is the number of bonds between atoms. These are inversely related to bond length: higher bond order means shorter bond length and stronger bond, hence higher bond enthalpy. For example, a C≡C bond (BO=3) is shorter and stronger than a C=C bond (BO=2), which is shorter and stronger than a C-C bond (BO=1).

For calculating reaction enthalpy ($\Delta H_{rxn}$), use the formula: $\Delta H_{rxn} = \sum E_{b}(\text{bonds broken}) - \sum E_{b}(\text{bonds formed})$. Always draw out the structures to correctly count each type of bond.

For instance, in $\text{H}_2(g) + \text{Cl}_2(g) \to 2\text{HCl}(g)$, you break one H-H and one Cl-Cl bond, and form two H-Cl bonds. If $E_{b}(\text{H}-\text{H}) = 436$, $E_{b}(\text{Cl}-\text{Cl}) = 242$, $E_{b}(\text{H}-\text{Cl}) = 431$, then $\Delta H_{rxn} = (436+242) - (2 \times 431) = -184 \text{ kJ/mol}$.

Molecular Orbital (MO) theory is vital for diatomic species. First, count total electrons. Then, fill them into MOs following the correct energy order: for $\le 14$ electrons, $\pi_{2p}$ are lower than $\sigma_{2p}$; for $> 14$ electrons, $\sigma_{2p}$ is lower than $\pi_{2p}$.

Calculate bond order as $\frac{1}{2}(N_b - N_a)$. For example, $\text{O}_2$ (16 electrons) has MO config $(\sigma_{1s})^2 (\sigma_{1s}^*)^2 (\sigma_{2s})^2 (\sigma_{2s}^*)^2 (\sigma_{2p})^2 (\pi_{2p})^4 (\pi_{2p}^*)^2$.

Here, $N_b=10, N_a=6$, so BO=2. The two unpaired electrons in $\pi_{2p}^*$ make $\text{O}_2$ paramagnetic. Compare $\text{O}_2^+$ (15e, BO=2.5) and $\text{O}_2^-$ (17e, BO=1.5) to understand how adding/removing electrons affects bond order, length, and enthalpy.

Remember, higher bond order means shorter bond length and higher bond enthalpy. Practice these calculations and comparisons rigorously.

Prelims Revision Notes

Bond Enthalpy and Bond Order: NEET Quick Revision

1. Bond Enthalpy ($E_b$ or $\Delta H_{bond}$):

Definition: — Average energy required to break one mole of a specific type of bond in the gaseous state.
Nature: — Always positive (endothermic process). Energy is absorbed to break bonds.
Units: — Kilojoules per mole (kJ/mol).
Bond Dissociation Enthalpy (BDE): — Energy to break a *specific* bond in a *specific* molecule. Differs from average bond enthalpy for polyatomic molecules.
**Factors Affecting $E_b$:**

* Bond Order: Higher BO $\implies$ Higher $E_b$. * Bond Length: Shorter bond length $\implies$ Higher $E_b$. * Atomic Size: Smaller atoms $\implies$ Shorter bonds $\implies$ Higher $E_b$. * Electronegativity Difference: Greater difference $\implies$ More polar bond $\implies$ Stronger bond $\implies$ Higher $E_b$.

**Calculation of $\Delta H_{rxn}$:**

* $\Delta H_{rxn} = \sum E_{b}(\text{bonds broken in reactants}) - \sum E_{b}(\text{bonds formed in products})$ * Remember to multiply bond enthalpies by stoichiometric coefficients and the number of identical bonds within a molecule.

2. Bond Order (BO):

Definition: — Number of chemical bonds between a pair of atoms.
From Lewis Structures:

* Single bond: BO = 1 * Double bond: BO = 2 * Triple bond: BO = 3 * Resonance structures: Fractional BO (e.g., benzene C-C bond order = 1.5).

From Molecular Orbital (MO) Theory (for diatomic species):

* Formula: $\text{Bond Order} = \frac{1}{2} (N_b - N_a)$ * $N_b$: Number of electrons in bonding molecular orbitals. * $N_a$: Number of electrons in antibonding molecular orbitals. * MO Energy Level Order: * **For $\le 14$ electrons (e.

g., $\text{H}_2, \text{N}_2, \text{C}_2, \text{B}_2$):** $\sigma_{1s} < \sigma_{1s}^* < \sigma_{2s} < \sigma_{2s}^* < (\pi_{2p} = \pi_{2p}) < \sigma_{2p} < (\pi_{2p}^* = \pi_{2p}^*) < \sigma_{2p}^*$ * **For $> 14$ electrons (e.

g., $\text{O}_2, \text{F}_2, \text{Ne}_2$):** $\sigma_{1s} < \sigma_{1s}^* < \sigma_{2s} < \sigma_{2s}^* < \sigma_{2p} < (\pi_{2p} = \pi_{2p}) < (\pi_{2p}^* = \pi_{2p}^*) < \sigma_{2p}^*$ * Magnetic Properties: * Paramagnetic: Contains one or more unpaired electrons (e.

g., $\text{O}_2, \text{NO}$). Attracted to a magnetic field. * Diamagnetic: All electrons are paired (e.g., $\text{N}_2, \text{F}_2$). Repelled by a magnetic field.

3. Interrelationship of Bond Parameters:

**Higher Bond Order $\implies$ Shorter Bond Length $\implies$ Higher Bond Enthalpy $\implies$ Greater Stability.**
A bond order of zero implies no stable molecule (e.g., $\text{He}_2$).

Key Strategy: Practice MO configurations and bond order calculations for various diatomic species and their ions. Be meticulous in counting bonds for enthalpy calculations.

Vyyuha Quick Recall

Bond Order Length Enthalpy: Bigger Order, Less Length, More Enthalpy. (Think BOLLE - 'Bigger Order, Less Length, More Enthalpy')