Structure �and �Bonding-- II
Deepshikha
Department of Chemistry
Bond Length
The bond length refers to the distance between the centers of the nuclei of two bonded atoms in an equilibrium position. The stronger the force of attraction in between the bonding atoms, the smaller is the length of the bond. However, the bigger the atom size, the longer the bond length.
The certain factors upon which the bond length is dependent are
The stronger the force of attraction between the bonding atom, the smaller the bond length. However, the bigger the size of an atom, the longer will be the bond length. Also, it is to be noted that in the case of a covalent bond, the contribution by each atom is referred to as the covalent radius of that atom.
Bond Energy
“The amount of energy required to break one mole of bonds of a particular type so as to separate them into gaseous atoms is called bond dissociation energy or simply bond energy”. Greater is the bond energy, stronger is the bond. Bond energy is usually expressed in kJmol-1.
The bond order is the number of electron pairs shared between two atoms in the formation of the bond. Bond order for C=C and O=O is 2.
The amount of energy required to break a bond is called bond dissociation energy or simply bond energy. Since bond lengths are consistent, bond energies of similar bonds are also consistent.
The bond energy is essentially the average enthalpy change for a gas reaction to break all the similar bonds. For the methane molecule, C(−H)4, 435 kJ is required to break a single C−H bond for a mole of methane, but breaking all four C−H bonds for a mole requires 1662 kJ. Thus the average bond energy is (1662/4) 416 (not 436) kJ/mol.
Factors affecting bond energy
(ii) For the bond between the two similar atoms, greater is the multiplicity of the bond, greater is the bond dissociation energy
.
(iii) Greater the number of lone pairs of electrons present on the bonded atoms, greater is the repulsion between the atoms and hence less is the bond dissociation energy.
(iv) The bond energy increases as the hybrid orbitals have greater amount of s orbital contribution. Thus, bond energy decreases in the following order, sp>sp2>sp3
(v) Greater the electronegativity difference, greater is the bond polarity and hence greater will be the bond strength i.e., bond energy, H-F>H-Cl>HBr>H-I,
(vi) Among halogens Cl – Cl > F – F > Br – Br > I – I, (Decreasing order of bond energy) Resonance increases bond energy.
Bond Angles
Bond angles also contribute to the shape of a molecule. Bond angles are the angles between adjacent lines representing bonds. The bond angle can help differentiate between linear, trigonal planar, tetraheral, trigonal-bipyramidal, and octahedral. The ideal bond angles are the angles that demonstrate the maximum angle where it would minimize repulsion, thus verifying the VSEPR theory.
The following examples make use of this notation, and also illustrate the importance of including non-bonding valence shell electron pairs (colored blue) when viewing such configurations .
The measured bond angles of these compounds (H2O 104.5º & NH3 107.3º) show that they are closer to being tetrahedral than trigonal or linear.
Configuration | Bondingg Partners | Bond Angles | Example |
Tetrahedral | 4 | 109.5º | |
Trigonal | 3 | 120º | |
Linear | 2 | 180º |
The bonding configurations of carbon are easy to remember, since there are only three categories.
The bonding configurations of carbon are easy to remember, since there are only three categories.
The basic goal of resonance structures is to show that molecules can move electrons and charges onto different atoms on the molecule. Resonance generally makes a molecule more stable because the charge (or bond) is now delocalized and not “forced” onto an atom that might not want it. More on this is discussed below!
In one sentence, it is the concept where electrons (bonds) are delocalized over three or more atoms which cannot be depicted with one simple Lewis structure.
RESONANCE
Sometimes, molecules can be represented with more than 1 Lewis structure, where the only difference is the location of pi electrons. Electrons in sigma bonds have a fixed location; these are called localized and never move. Conversely, pi electrons are referred to as delocalized, because they can be easily moved around. Together, these Lewis diagrams are then known as resonance structures or resonance contributors or resonance canonicals. The actual molecule has characteristics of each of the parts, and can represented as a resonance hybrid (which you can think of as a cross-breed). Resonance hybrids are a more accurate way to think about resonance structures, because it is more like what the structure looks like in nature.
Rules for writing of resonating structures
1)The contributing structures should have the same position of atoms.
They should differ only in the position of electrons.
2)The contributing structures should have the same number of unpaired electrons.
3)The contributing structure should have nearly same energy.
4)The contributing structures should have negative charge on the electronegative atom and the positive charge on the electropositive atoms.
5) In a contributing structure, like charges should not be present on adjacent atom while unlike charges should not be widely separated.
The structure of ozone can be written as
Resonance Structures
The Lewis structure of ozone (O3)
1Sum of valence electrons = (6*3) = 18
2. Drawing the bond connectivities:
3. Complete the octets of the atoms bonded to the central atom:
4. Place any leftover electrons (18-16 = 2) on the central atom
Resonating structure of Carbon dioxide
Resonating structure of Carbonate ion
Electromeric Effect can be observed only in organic compounds which contain multiple bonds. It is a temporary effect that arises when the compound is subjected to an attacking reagent.
Electromeric Effect
The instantaneous formation of a dipole in the molecule of an organic compound due to the complete transfer of shared pi electron pairs to one of the atoms under the influence of an attacking reagent is referred to as the Electromeric effect.
This effect can be observed in organic compounds that contain at least one multiple bond. When the atoms participating in this multiple bond come under the influence of an attacking reagent, one pi bonding pair of electrons is completely transferred to one of the two atoms.
The electromeric effect is a temporary effect that remains as long as the attacking reagent is present and exposed to the organic compound. Once this attacking reagent is removed from the system, the molecule that was polarized goes back to its original state.
The electromeric effect is the movement of electrons from one atom to another as a reagent attacks a π bond.
+E EFFECT
If the attacking species is an electrophile, the π electrons are transferred towards the positively charged atom. This is the +E effect.
An example is the protonation of ethene. When the H⁺ comes near the double bond, the bond is polarized towards the proton.
-E EFFECT
If the attacking reagent is a nucleophile, the electrons are transferred away from the attacking reagent and into the π system. This is the –E Effect.
-E Effect
This effect occurs when the electron pair of the pi bond is moved away from the attacking reagent. The attacking reagent attaches itself to the positively charged atom in the molecule, i.e. the atom which lost the electron pair in the transfer.
The -E effect is generally observed when the attacking reagent is a nucleophile and the pi electrons are transferred to the atom which the attacking reagent will not bond with. An example where the -E effect occurs would be the addition of nucleophiles to carbonyl compounds as illustrated below.