Bond parameters are the measurable characteristics of chemical bonds that describe their strength, stability, and geometry. It helps in predicting the properties and reactivity of molecules. These parameters are essential for building a initial knowledge of chemical bonding.
1.Bond Length
Definition: Bond length is the average distance between the nuclei of two bonded atoms. Bond length is expressed in Angstrom or in picometer.
The covalent radius is measured approximately as the radius of an atom’s core which is in contact with the core of an adjacent atom in a bonded situation. The covalent radius is half of the distance between two similar atom joined by covalent bond in the same molecule.
The van der Waals’ radius represents the overall size of the atom which includes its valence shell in a non-bonded situation. The van der waals’ radius is half of the distance between two similar atoms in separate molecules in a solid.
Factors Affecting Bond Length:
Atomic Size: Larger atoms form longer bonds because their electron clouds are more spread out.
Bond Order: Higher bond order (e.g., triple bond) means shorter bond length due to increased electron density between nuclei.
Electronegativity: A higher difference in electronegativity can shorten bond length as the bond becomes stronger.
Examples:
In diatomic molecules like H2, bond length is ~74 pm (picometers).
Triple bonds like C ≡ C are shorter (~120 pm) compared to single bonds like C−C (~154 pm).
2. Bond Angle
Definition: Bond angle is the angle formed between two adjacent bonds around a central atom. It defines the shape and geometry of molecules. Bond angle gives an idea of the distribution of orbitals in three dimensional apace around the central atom.
Factors Affecting Bond Angle:
Lone Pairs: Lone pairs exert greater repulsion than bonded pairs, decreasing bond angles.
Hybridization: The type of hybrid orbitals influences bond angles. For example, sp3 hybridization in methane gives a bond angle of 109.5∘.
Electronegativity: Greater electronegativity differences can slightly alter bond angles.
Examples:
Water (H2O) has a bond angle of 104.5∘ due to lone pair-bond pair repulsion.
Carbon dioxide (CO2) has a bond angle of 180∘, indicating a linear geometry.
3. Bond Energy
Definition: Bond energy (or bond enthalpy) is the amount of energy required to break one mole of a specific bond in a gaseous molecule. It expressed in kj mol-1.
Factors Affecting Bond Energy:
Bond Order: Bonds with higher bond orders (e.g., triple bonds) require more energy to break.
Bond Length: Shorter bonds are stronger and have higher bond energy.
Molecular Environment: The surrounding atoms and functional groups can influence bond strength.
Examples:
The bond energy of H − H in H2 is ~436 kJ/mol.
Triple bonds like N ≡ N have higher bond energies (~945 kJ/mol) compared to single bonds.
4. Bond Order
Definition: Bond order is the number of chemical bonds between a pair of atoms. It can be determined using Molecular Orbital Theory or Lewis structures.
Isoelectronic molecules and ions have identical bond orders. For example F2 and O2- have bond order one, CO and NO+ have the bond order three.
A general correlation useful for understanding the stabilities of molecules is that : with increase in bond order, bond enthalpy increases and bond length decreases.
Formula (for diatomic molecules):
Number of bonding electrons−Number of antibonding electrons
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Bond Order = 2
Significance:
A bond order of 1 indicates a single bond, 2 indicates a double bond, and so on.
Fractional bond orders indicate resonance or delocalized bonding, as seen in O3 (ozone).
Examples:
O2 has a bond order of 2, indicating a double bond.
N2 has a bond order of 3, making it very stable.
Dipole Moment
Definition: Dipole moment is a measure of the polarity of a bond. It is the product of charge separation and the distance between charges. Dipole moment = electric charge x distance of separation. Dipole moment measured in ‘debye’ unit (D). 1D = 1018 esu cm = 3.33 x 10-30 coulomb meter. Dipole moment is vector quantity.
Formula:
μ = q × d
Where q is the charge, and d is the distance.
Factors Influencing Dipole Moment:
Electronegativity: Greater differences result in higher dipole moments.
Molecular Shape: Symmetrical molecules may have zero dipole moments despite having polar bonds (e.g., CO2).
Examples:
Water (H2O) has a dipole moment of ~1.85 D, making it highly polar.
Carbon tetrachloride (CCl4) has zero dipole moment due to its symmetrical geometry.
Fajan’s rule :- According to Fajan’s rules, magnitude of covalent character in an ionic bond depends upon the extent of polarization caused by cation.
Key points
Bond parameters like bond length, bond angle, bond energy, bond order, and dipole moment provide valuable insights into molecular structure and behavior. They are interconnected and depending on the nature of bonding and molecular environment. Smaller the size of cation, larger is its polarizing power.
Among two cations of similar size, the polarising power of cation with noble gas configuration is larger than cation with noble gas configuration , polarising power of Ag+ is more than K+. Larger the anion more will be its polarisibility..
Bond parameters are measurable characteristics of chemical bonds that include bond length, bond angle, bond energy, bond order, and dipole moment. These parameters provide approach into the bond’s strength, stability, and geometry.
Bond length is the average distance between the nuclei of two bonded atoms. It depends on:
Atomic size: Larger atoms result in longer bonds.
Bond order: Higher bond order (e.g., triple bonds) results in shorter bonds.
Electronegativity: Greater differences can slightly shorten bond length due to increased attraction.
Bond angle determines the spatial arrangement of atoms around a central atom, defining the molecular geometry (e.g., linear, trigonal planar, tetrahedral). It is influenced by factors like lone pair repulsion, hybridization, and the electronegativity of surrounding atoms.
Bond energy is the energy required to break one mole of a bond in a gaseous molecule. It indicates bond strength; stronger bonds (shorter bonds or bonds with higher bond order) have higher bond energies.
Bond order is the number of bonds between two atoms. It can be calculated using Molecular Orbital Theory as:
Number of bonding electrons−Number of antibonding electrons
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Bond Order = 2
For simple molecules, it can also be determined using Lewis structures.
Dipole moment is a measure of the polarity of a bond, calculated as the product of charge and distance between charges μ = q × d. A higher dipole moment indicates greater bond polarity, while symmetrical molecules may have zero dipole moment despite polar bonds.
Bond parameters influence each other:
Shorter bond lengths typically mean higher bond energy.
Molecules with greater bond angles may experience less repulsion, leading to stability.
Polar bonds (dipole moment) can affect bond strength and molecular geometry.
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