Download Organic Chemistry –III How are chemical bonds formed? Write the

Organic Chemistry –III
1. How are chemical bonds formed? Write the characteristic of ionic,
covalent and corrdinate bonds.
Ans. A chemical bond is an attraction between atoms that allows the formation of chemical
substances that contain two or more atoms. The bond is caused by the electromagnetic
force attraction between opposite charges, either between electrons and nuclei, or as the result of
adipole attraction. The strength of chemical bonds varies considerably; there are "strong bonds"
such as covalent or ionic bonds and "weak bonds" such as dipole-dipole interactions, the London
dispersion force and hydrogen bonding.
Since opposite charges attract via a simple electromagnetic force, the negatively
charged electrons that are orbiting the nucleus and the positively charged protons in the nucleus
attract each other. Also, an electron positioned between two nuclei will be attracted to both of
them. Thus, the most stable configuration of nuclei and electrons is one in which the electrons
spend more time between nuclei, than anywhere else in space. These electrons cause the nuclei
to be attracted to each other, and this attraction results in the bond. However, this assembly
cannot collapse to a size dictated by the volumes of these individual particles. Due to the matter
wave nature of electrons and their smaller mass, they occupy a much larger amount of volume
compared with the nuclei, and this volume occupied by the electrons keeps the atomic nuclei
relatively far apart, as compared with the size of the nuclei themselves.
In general, strong chemical bonding is associated with the sharing or transfer of electrons
between the participating atoms. The atoms inmolecules, crystals, metals and diatomic gases—
indeed most of the physical environment around us— are held together by chemical bonds, which
dictate the structure and the bulk properties of matter.
2. Write the sequence in which electron fill in various energy sub shelf in an
unexcited atom.
3.
What is meant by hybridisation of orbits? Explain sp2 and sp3
hybridization.
Ans. In chemistry, hybridisation (or hybridization) is the concept of mixing atomic orbitals to
form new hybrid orbitals suitable for the qualitative description of atomic bonding properties.
Hybridised orbitals are very useful in the explanation of the shape of molecular
orbitals for molecules. It is an integral part of valence bond theory. Although sometimes taught
together with the valence shell electron-pair repulsion (VSEPR) theory, valence bond and
hybridization are in fact not related to the VSEPR model.[1first developed the hybridisation theory
in order to explain the structure of molecules such as methane (CH4).[2] This concept was
developed for such simple chemical systems, but the approach was later applied more widely,
and today it is considered an effective heuristic for rationalizing the structures of organic
compounds.
For quantitative calculations of electronic structure and molecular properties, hybridisation theory
is not as practical as molecular orbital theory. Problems with hybridisation are especially notable
when the d orbitals are involved in bonding, as in coordination chemistry andorganometallic
chemistry. Although hybridisation schemes in transition metal chemistry can be used, they are not
generally as accurate.
Orbitals are a model representation of the behaviour of electrons within molecules. In the case of
simple hybridisation, this approximation is based on atomic orbitals, similar to those obtained for
the hydrogen atom, the only atom for which an exact analytic solution to itsSchrödinger
equation is known. In heavier atoms, like carbon, nitrogen, and oxygen, the atomic orbitals used
are the 2s and 2p orbitals, similar to excited state orbitals for hydrogen. Hybridised orbitals are
assumed to be mixtures of these atomic orbitals, superimposed on each other in various
proportions. The theory of hybridisation is most applicable under these assumptions. It gives a
simple orbital picture equivalent to Lewis structures. Hybridisation is not required to describe
molecules, but for molecules made up from carbon, nitrogen andoxygen (and to a lesser
extent, sulfur and phosphorus) the hybridisation theory/model makes the description much easier.
The hybridisation theory finds its use mainly in organic chemistry. Its explanation starts with the
way bonding is organized in methane.
4. What gives covalent bonds its strength? Discuss the polar nature of
covalent bonds.
5.
Explain the formation of CH4 and H2c=CH2 .
6. Explain Hund’s rule and Octet Rule.
Ans. In atomic physics, Hund's rules refer to a set of rules formulated by German
physicist Friedrich Hund around 1927, which are used to determine the term symbol that
corresponds to the ground state of a multi-electron atom. In chemistry, rule one is especially
important and is often referred to as simply Hund's Rule.
The three rules are:
1. For a given electron configuration, the term with maximum multiplicity has the lowest
energy. Since multiplicity is equal to
, this is also the term with maximum
. S is the spin angular momentum.
2. For a given multiplicity, the term with the largest value of
where L is the orbital angular momentum.
has the lowest energy,
3. For a given term, in an atom with outermost subshell half-filled or less, the level with the
lowest value of lies lowest in energy. If the outermost shell is more than half-filled, the
level with highest value of is lowest in energy. J is the total angular
momentum,J = L + S.[1][2]
These rules specify in a simple way how the usual energy interactions dictate the ground state
term. The rules assume that the repulsion between the outer electrons is very much greater than
the spin-orbit interaction which is in turn stronger than any other remaining interactions. This is
referred to as the LS coupling regime.
Full shells and subshells do not contribute to the quantum numbers for total S, the total spin
angular momentum and for L, the total orbital angular momentum. It can be shown that for full
orbitals and suborbitals both the residual electrostatic term (repulsion between electrons) and the
spin-orbit interaction can only shift all the energy levels together. Thus when determining the
ordering of energy levels in general only the outer valence electrons need to be considered.
7.
Explain resonance and inductive effect with suitable examples.
8. Explain Electrophilic and Nucleophilic addition reactions with suitable
examples.
Ans. In general, electrophiles are positively charged species that are attracted to an electron rich
centre; but they can also be uncharged species such as a Lewis acid. In chemistry,
an electrophile (literally electron-lover) is a reagent attracted to electrons that participates in a
chemical reaction by accepting an electron pair in order to bond to a nucleophile. Because
electrophiles accept electrons, they are Lewis acids (seeacid-base reaction theories). Most
electrophiles are positively charged, have an atom that carries a partial positive charge, or have
an atom that does not have an octet of electrons.
The electrophiles attack the most electron-populated part of one nucleophile. The electrophiles
frequently seen in the organic syntheses arecations such as H+ and NO+, polarized neutral
molecules such as HCl, alkyl halides, acyl halides, and carbonyl compounds, polarizable neutral
molecules such as Cl2 and Br2, oxidizing agents such as organic peracids, chemical species that
do not satisfy the octet rule such as carbenes and radicals, and some lewis acids such
as BH3 and DIBAL.