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A delocalized π bond is a π bond in which the electrons are free to move over more than two nuclei.


In a molecule like ethylene, the electrons in the π bond are constrained to the region between the two carbon atoms.


We say that the π electrons are localized.

Even in penta-1,4-diene, the π electrons are still localized.

The #"CH"_2# group between the two π orbitals prevents them from overlapping.

(From iverson.cm.utexas.edu)

However, in buta-1,3-diene, the two orbitals can overlap, and the π electrons are free to spread over all four carbon atoms.

We say that these π electrons are delocalized.

In benzene, the π electrons are delocalized over all six atoms of the ring.


In β-carotene, the π electrons are delocalized over 22 carbon atoms!



WARNING! Long answer! Here's what I get.


Newman projections

In a Newman projection, we are looking down a carbon-carbon bond axis so that the two atoms are one-behind-the-other.

In the Newman projection of propane (below), #"C-1"# is the blue methyl group, #"C-2"# is at the centre of the circle, and #"C-3"# is hidden directly behind it.



Conformations are the different arrangements that the atoms can take by rotating about the #"C-C"# single bond.

There are an infinite number of conformations, but there are two important ones.

In the eclipsed conformation, the substituents on the two carbon atoms are as close to each other as they can get.

The #"H-C-C-H"# and #"CH"_3"-C-C-H"# dihedral angles are 0°.

In the staggered conformation, the groups on the two carbon atoms are as far from each other as they can get, and the #"H-C-C-H"# and #"CH"_3"-C-C-H"# dihedral angles are 60°.

Conformational analysis

The eclipsed conformation is a high-energy conformation because the negatively charged electrons in the #"C-H"# and #"C-CH"_3# bonds repel each other most when the bonds line up.

The staggered conformation is the most stable because the bonds are furthest away from each other and the electron repulsions are minimal.

The energy difference between the two conformations is called torsional strain.

Conformational analysis is the study of the energy changes that occur during the rotations about σ bonds.

In the conformational energy diagram below, we are looking down the #"C1-C2"# bond, and the #"CH"_3# is coming off the back carbon.


We start with the molecule in the eclipsed conformation.

As we rotate the back carbon clockwise, the molecule reaches an energy minimum with a staggered conformation at 60°.

Rotating another 60°, the molecule reaches an energy maximum with an eclipsed conformation.

The pattern repeats twice more as the bond rotates a full 360°.

The energy difference between the maxima and minima is 3.4 kcal/mol (14 kJ/mol).

This represents the total repulsion of three bond pairs, two #"C-H, C-H"# repulsions and a #"C-H, C-CH"_3# repulsion.

We know from the conformational analysis of ethane that one #"C-H, C-H"# repulsion contributes 4 kJ/mol, so a #"C-H, C-CH"_3# repulsion contributes about 6 kJ/mol to the torsional strain.


Here's what I get.



Acetate is the carboxylate ion of acetic acid.


In acetoacetate, an α-hydrogen has been replaced by an aceto or acetyl group, #"CH"_3"CO"#.



Acetyl-CoA is Coenzyme A in which the #"H"# atom in the thiol group has been replaced by an acetyl group.

This is Coenzyme A:


And this is acetyl-CoA:

(From Wikipedia)


Acetoacetyl-CoA is Coenzyme A in which the #"H"# atom in the thiol group has been replaced by an acetoacetyl group, #"CH"_3"COCH"_2"CO"#.

(From Wikipedia)


The high melting point is caused by π-π stacking of the aromatic rings.


In organic chemistry, π–π stacking refers to attractive interactions between the π clouds of aromatic rings.

There are various types of stacking.


Neither benzene nor hexafluorobenzene has a dipole moment.

However, they have strong quadrupole moments, caused by the π clouds above and below the rings.

For example, in benzene, the π clouds are negatively charged and the plane of the ring is positively charged.


The situation is reversed in hexafluorobenzene, because the electronegative fluorine atoms withdraw electron density from the ring.

(From download.e-bookshelf.de)

You can see the charge distribution better in this image:


Both benzene and hexafluorobenzene are destabilized by sandwich stacking, because areas with the same charge are placed next to each other.


However, benzene and hexafluorobenzene are strongly stabilized by sandwich stacking, because areas with opposite charge are placed next to each other.

Theoretical calculations put the stabilization energy at about 20 kJ/mol.

That makes the attractions as strong as many hydrogen bonds and dipole-dipole interactions.

Thus, the strong intermolecular quadrupole attractions cause a 1:1 mixture of benzene and hexafluorobenzene to have a high melting point.

It can be considered as basically spatial crowding.

Steric hindrance is a kinetic factor that limits the ease to which a nucleophile (electron pair donor) can approach an electrophile (electron pair acceptor).

As an example, consider the reaction seen below, which one might hope is #"S"_N2#:


Here, the nucleophile is cyanide (#""^(-):"CN"#) and the electrophile is the central carbon on the alkyl halide (tert-butyl bromide).

We should predict that reaction does not work via an #"S"_N2# mechanism, based on the idea of steric hindrance---the three #"CH"_3# groups are blocking the cyanide from performing its backside-attack.


(LEFT: steric hindrance; RIGHT: reduced steric hindrance)

That reduces the ease to which a successful collision can occur between two reactants, and slows down the first step in a given substitution mechanism (making it the rate-limiting step).

Hence, this reaction proceeds more easily as a first-order mechanism (e.g. #"S"_N1# or #E1#), since the #"Br"^(-)# has the time (during the slow step) to come off on its own.

The rate law for this could then be approximated as first-order based on the slow step, since it dominates the extent of the reaction time:

#r(t) ~~ k_1["Br"^(-)]#

(Depending on the choice of solvent, either #"S"_N1# or #E1# could occur.)


Some examples are: #CO_2#, #O_2#, #CH_3# DNA, non-polar amino acids


Covalent bonds are common in the molecules of living organisms. This bonds are created is by sharing electrons. The more electrons they share the stronger the bond will be.
Non-polar covalent bonds appear between two atoms of the same element or between different elements that equally share electrons.

  • #O_2# and #CO_2#
    #O_2# is non-polar because the electrons are equally shared between the two oxygen atoms. The same case is with #CO_2#.
    Breathe in and out! When you inhale, #O_2# comes in your lungs, and when you exhale, #CO_2# goes out. This process happens because #CO_2# and #O_2# are nonpolar molecules.

  • #CH_3#
    Our intestinal gas/ flatus is produced as a byproduct of bacterial fermentation in our colon.
    Click here for: Why is #CH_4# a non-polar molecule?

  • DNA
    Each strand of the DNA is polar, but each strand with its polarity neutralizes the molecule of the DNA in total.

  • NON-POLAR amino acids #-># their R-groups, white on the picture below, hydrocarbon alkyl groups (alkane branches) or aromatic rings, exception: benzene ring in the amino acid Tyrosine which is polar. Aminoacids are monomers that form proteins.
    Alanine, Cysteine, Glycine, Isoleucine, Leucine, Methionine, Phenylalanine, Proline, Tryptophan, Valine


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