Chemistry I | © Ray Lovegrove; contributors: Reema Lodha, Antonio Codina, Madhusree Bathini, Ray Lovegrove, Rajesh Rathod et al. Source: PEOI |
During the chapter 3 we built up an idea of what atoms are like. We know that they are made up of
tiny sub-atomic particles - electrons, protons and neutrons. We also understand that specific atoms
of a given element will always have a given number of protons and electrons but may differ in terms
of the number of neutrons that they have. Chemistry, for the main part is not concerned with what
atoms are, but very concerned with what atoms do , in particular how individual atoms join together
to form more complex substances. It is the atomic
electrons that determine
the way that atom combines with other atoms to form
complex compounds. This will be the only topic of this chapter and it is essential that we
understand it well, as it is the key to much of the chemistry that follows.
Atomic Structure and Bonding
In 1914 W Kossel found that by reacting elements of group 1A and 2A with elements of groups VIA &
VIIA, produced salt like
compounds. He concluded that the ions achieved a noble gas configuration and that electrostatic
forces hold these ions.
In 1916, the American chemist Gilbert Newton Lewis proposed that chemical bonds are formed between
atoms because electrons from
the atoms interact with each other. Lewis had observed that many elements are most stable when they
contain eight electrons in
their valence shell. He suggested that atoms with fewer than eight valence electrons bond together
to share electrons and
complete their valence shells.
While some of Lewis' predictions have since been proven incorrect (he suggested that electrons
occupy cube-shaped orbitals),
his work established the basis of what is known today about chemical bonding.
In 1919 Irving
Langmuir extended Lewis's ideas
and introduces concept of the covalent or homopolar bonding. The success of Schrodinger’s quantum
wave theory in 1926, dealing
with the atomic structure, directed the entire attempt to understand the covalent bond towards this
new approach. W. Heitler
and F. London threw new light on the nature of
covalent bond. They introduced number of concepts that are now of general use in chemical theory.
These concepts are used along
with the Lewis’s theory to give more insight into chemical bonding.
Filling the orbitals
Hunds Rule
The atom is a central, positively charged nucleus surrounded by extra-nuclear electrons.
These extra-nuclear electrons
are in such a number as to achieve electrical neutrality and are grouped in several
shells - K, L, M, N... These shells are divided into subshells - orbitals called: s, p, d, f..
and the electrons occupy each of these orbitals (subshells) in pairs. The number
of electrons that a shell can contain increases with the distance of the shell to the nucleus. The number of electrons in a
n is the
quantum shell number,
3ii
i.e for K (first shell n is 1) = 2, L = 8, M = 18, N = 32 etc.
Fig. 2 (Electrons in green dots).
Sodium (Na) is a alkali metal (group IA element) with atomic number 11 in
Periodic table
Elements.
It has 11 protons in the nucleus and to achieve electrical neutrality it has 11
electrons in
the K, L, M shells. These
electrons are arranged as, 2 (K shell), 8 (L shell), 1 (M shell) i.e. 2. 8. 1
(11 electrons) Fig. 1.
Calcium (Ca) is a
alkali earth metal (group II element) with atomic number 20. It has 20 protons in
the nucleus and to achieve electrical
neutrality it has 20 electrons in the K, L, M, N shells. These
electrons are arranged as, 2 (K shell), 8 (L shell), 8 (M shell), 2 (N shell)
i.e. 2. 8. 8. 2 (20 electrons) Fig. 3.
We already mentioned that these shells have sub-shells or more accurately termed orbitals; s, p, d or f. These orbitals are occupied by paired electrons in an opposite spin. The s orbital can occupy 2 such electrons (e–), p orbital 6e–, d orbital 10e–, f orbital 14e–. And the orbital electrons (in superscript) represented as follows:- 1s2 (K shell); 2s 2, 2p6 (L shell); 3s2, 3p6, 3d10 (M shell); 4s2, 4p6, 4d10, 4f14 (N shell). The order in which the orbitals s, p, d, f... become filled with electrons, is governed by their energy level, i.e. lower energy orbital filled first;
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Fig. 1 |
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Fig. 3 |
When the required number of electrons (equal to the Atomic Number) has been accounted for, the electrons which occupy the outer quantum shell (valency shell) are regarded as the valence electrons. The valence electrons are most exposed to another atoms and will participate in the formation of chemical bonds.
The octet rule
When the elements combine to form chemical compounds, the electrons in the
outermost shell or valency shell may be transferred
from one
atom to another or there may be a mutual sharing of the electrons i.e. atoms are held together
by the exchange or sharing of
electrons, to achieve a stable configuration or ‘octet’ (G. Lewis, 1916).
Bonds are formed by sharing pairs of electrons between two nuclei, so that, each
achieves a noble gas (Group VIII/O)
configuration. Noble gases have complete filled orbitals (ns2 np6) that confer extra stability. Atoms are
therefore considered
to combine together in such a manner as to
achieve this stable Octet configuration
(octet rule). Atoms achieve this stable state chemically in three ways: Electrovalent (Ionic) Bonding,
Covalent Bonding, and Dative (co-ordinate, semi-polar, semi-ionic) Covalent Bonding.
Valency and Bonding
In 1914 W Kossel found that by reacting elements of group 1A and 2A with elements of groups VIA & VIIA, produced salt like
compounds. He concluded that the ions achieved a noble gas configuration and that these ions are held by electrostatic forces.
In Electrovalent (Ionic) bonding the valency or the combining power of an atom is equal to the number of charges on Cation
(positive ion) or Anion (negative ion). Monovalent include; Li, Na, K, Rb, Cs, F, Cl, Br, I and
Divalent; Mg, Ca, Sr, Ba, O, S, Se, Te.
Table showing six periodic elements and the electron(s) in outer valence KLM shell Fig. 4 |
Positive Valency
In the Group I alkali metal atom, loss of an electron destroys the electrical neutrality (free state)
of the atom and it
becomes a positive ion (+1 valency). The element Sodium (Na) has 11 protons and 11 electrons
in its free state. It has one
electron in the outer M shell, that it loses to obtain noble gas configuration and become stable. On losing the electron, it
will have + 1 charge
(Cation) and the valency
is +1 Fig. 5
Negative Valency
In the Group VIIA halogen atom there is one less electron (in outermost shell) than the noble gas (Group VIII/O) structure, so
the capture of an
electron by a halogen atom will give the eight electrons of
the noble gas octet structure.
As before, the electrical neutrality (free state) is destroyed, and the atom became a negative
ion (Anion). The element
Chlorine (Cl) has 17 protons and 17 electrons in free state. It has 7 electrons in the outer M shell,
one less, to fit the
stable octet configuration. On gaining the electron, it will have a –1 charge (Anion) and the
valency will be –1 Fig. 6
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Electrovalent (Ionic) Bonding
The electrovalency of an element is the number of electrons its atom must gain or lose to
attain a stable electron structure.
In ionic (electrovalent) bonding, electrons are completely transferred from one atom to another. Ionic
bonds are a type of
electrostatic bond between two atoms and are formed from the mutual attraction of ions. Atoms that
have charges are called
ions.
Ionic bonds can be explained using the theory of electron orbitals. Every atom has a number of electron
orbitals equal to or
greater than one. Each orbital has a maximum electron limit, after which a new orbital is created.
The number of orbitals in
the elements ranges from one for hydrogen or helium, to six, seven, or eight for the larger molecules with
atomic numbers
greater than uranium. Atoms are "most stable" when their outermost orbitals are filled. As the atoms
approach each other, each nucleus begins to attract the electron held by the other nucleus.
Eventually, the
electron clouds overlap and fuse into one "molecular orbital". Like an atomic orbital, a molecular orbital is
most stable when
filled by a pair of electrons. This shared orbital acts as a "chemical bond" between the two atoms.
In the process of either losing or gaining negatively charged electrons, the reacting atoms form ions.
The oppositely charged
ions are attracted to each other by electrostatic forces, which are the basis of the ionic bond.
Typical participants in an ionic bond are a metal and a non-metal, such as sodium and chlorine. We
revisit our example of the
reaction of sodium with chlorine that creates sodium chloride, or common table salt.
We have the positively charged Sodium ion and a negatively charged Chlorine ion. These attract each
other because opposite
charges attract.
Elements like the Alkali metals (Group IA) or halogens (Group VIIA) whose atoms lose or gain
only one electron, have a valency of 1. As it is only the electrons in the outermost shell which are affected,
A single electrovalent bond is formed by the sharing
of two electrons. Sodium (Na) reacts
energetically with Chlorine to form
the very stable salt Sodium Chloride.The essential chemical action is the capture of an electron
from each sodium atom by each
chlorine atom. After they have reacted the atoms are now oppositely charged ions in close
association, and powerful
electrostatic forces of attraction hold them in contact
Fig. 7.
Ions combined together this way are
said to be linked by an Ionic
bond, and compounds formed from this type of binding are called Electrovalent or Ionic.
A crystal of NaCl is a closed packed
system.
The Alkali Earth metals in Group IIA have an electrovalency of 2, because two electrons must be
removed to expose the noble gas
structure in the underlying shell. Example shown in Fig. 8 is Magnesium (Mg).
Properties of Ionic or Electrovalent Compounds. At ordinary temperatures they invariably exist as
crystalline solids although individual particles
such as Na+ Cl- may exist in the vapour state.
The crystal consists of an ‘infinite assembly’ of ions,
that is, a large and indefinite number of oppositely charged ions joined together in a
regular manner. Electrovalent compounds have high melting points and
boiling points owing to the strong
attraction between the oppositely charged ions. In the fused state they are good conductors of electricity and are
electrolysed by a current. This is explained
by the breakdown of the crystal into ‘free’ ions, which are attracted to,
and discharged at, the
electrode of opposite sign. They often dissolve readily in water, but are only sparingly soluble in organic
solvents like ether and benzene. When oppositely charged ions react together in aqueous solution
they do so almost instantaneously.
Covalent Bonding
Covalent bonding occurs because the atoms in the compound have a similar tendency for electrons (generally to gain electrons).
This most commonly occurs when two nonmetals bond together. Because both of the nonmetals will want to gain electrons, the
elements involved will share electrons in an effort to fill their valence shells.
According to modern atomic theory, atoms have electrons circling them in shells called orbitals. Each orbital has a maximum
number of electrons, and each atom "wants" to max out its electrons in each orbital. The bonds are formed between twp atoms
each of its electron orbitals seeking a similar number of other electrons to max out their electron shells. When the atoms are
brought together, their electron shells intermingle, and create something called "molecular orbitals," where electrons wander
freely between both atoms and orbit the nuclei of both.
Depending on the number of shared electron pairs, a covalent bond is characterized as a single bond, a double bond, triple
bond, etc. Some of the metals with the highest melting points, molybdenum and rhenium, have quadruple bonds. Quintuple covalent
bonds and sextuple covalent bonds are quite rare, and we have good reason to believe that nothing on the periodic table can go
beyond a sextuple bond.
Single Covalent Bond
A good example of a covalent bond is that which occurs between two hydrogen atoms Fig. 8a. Atoms of hydrogen (H) have
one
valence
electron in their first electron shell. Since the capacity of this shell is two electrons, each hydrogen atom will "want" to
pick up a second electron. In an effort to pick up a second electron, hydrogen atoms will react with nearby hydrogen (H) atoms
to form the compound H2. Because the hydrogen compound is a combination of equally matched atoms, the atoms will
share each
other's single electron, forming one covalent bond. In this way, both atoms share the stability of a full valence shell.
In methane molecule (CH4)
Fig. 8b, the Hydrogen (H) requires one electron to
form the helium electron (duplet) structure and carbon needs four more electron to form the neon (octet)
electron structure. Therefore
by sharing
the electrons of four hydrogen atoms with each electron of the carbon atom, all
the atoms gain the
noble gas (octet) electron structure. A pair of electrons shared in this way constitutes a
covalent bond.
Multiple Covalent Bonds
For example Fig. 9, oxygen (which has six valence electrons) needs two electrons to complete its valence shell. When
two
oxygen atoms
form the compound O2, they share two pairs of electrons, forming two covalent bonds.
In Fig. 10* only the electrons in the outer valency shells are shown in these examples since electrons in the inner
completely filled shell, are not involved in bond formation. (The outer shell electrons in the atoms can also be reprensented
by dots and crosses instead of colours). These examples shows the formation of single (Cl2), double (O2),
triple (N2) covalent bonds. Covalent linkages of the above type are non-polar because each atom contributes the same
number of electrons to the shared electron pairs and no separation of charges arises (other examples; NH3, BF3
).
A Shorthand Method for Drawing Covalent Bonds
Polar and Nonpolar Covalent Bonding
Whenever two atoms of the same element bond together, a nonpolar bond is formed. The H2 molecule is a good example
of the nonpolar bond. Because both atoms in the H2 molecule have an equal attraction (or affinity) for electrons,
the bonding electrons are equally shared by the two atoms and a nonpolar covalent bond is formed
A polar bond is formed when electrons are unequally shared between two atoms. Polar covalent bonding occurs because one atom
has a stronger affinity for electrons than the other (yet not enough to pull the electrons away completely and form an ion). In
a polar covalent bond, the bonding electrons will spend a greater amount of time around the atom that has the stronger affinity
for electrons. A good example of a polar covalent bond is the hydrogen-oxygen bond in the water molecule.
Water molecules contain two hydrogen atoms bonded to one oxygen atom. Oxygen, with six valence electrons, needs two additional
electrons to complete its valence shell. Each hydrogen contains one electron. Thus oxygen shares the electrons from two
hydrogen atoms to complete its own valence shell, and in return shares two of its own electrons with each hydrogen, completing
the H valence shells. Fig. 10c
Dative (Co-ordinate) Covalent Bonding Fig. 12 Metallic Bonding Conductivity of electricity Conductivity of heat Ductility (being able to draw out
metallic structures into wires) Metallic sonority (metals have a unique
'clanging' or ringing
sound if you strike them with something hard)
Structures
Diamond Fig. 13 Fig. 14
Intermolecular Bonding
Fig. 15 Fig. 16
The way in which charges on a water molecule
are distributed is
important, as it shows how hydrogen bonds are
formed.
The hydrogen atoms have their electrons
involved in covalent bonds with
oxygen. This leaves the hydrogen nuclei
exposed, giving a small
positive area of charge to this area of the molecule (δ+).
The oxygen
atom of the water molecule has all of the electrons in the molecule
gathered around it, giving this part of the
molecule a small negative
charge (δ-).
Molecular orbital and Bonding
Hybridisation
Boron atom has electronic confirguration 1s2 2s2 2p1 (one more
electron than beryllium atom). Two of the three p orbitals in the second quantum shell are
unoccupied, one of the empty p orbitals, giving two p electrons. Hybridisation occurs between the
remaining s orbital and the two p orbitals, and three equivalent sp2 hybrid orbitals are
produced. The axes of these lie in the same plane at 120 degrees to each other Fig. 17.
The Carbon atom again has one more electron than the boron atom. Its electron configuration is 1s2, 2s2,
2px1, 2py1 (K 2, L 4), the 2px and 2py signify that two different p orbitals in the second quantum are
occupied by single electrons. ‘Excitation’ of the carbon atom results in the carbon atom results in one of the 2s electrons
being 'promoted' to the empty 2pz orbital, so that now all three p orbitals contain one electron. Hybridization takes place
between the s orbital and the three p orbitals, giving four equal sp3 hybrid orbitals Fig. 18. These are
directed
outwards from the carbon atom so that they form the tetrahedral angle of 109 degree 28’minutes with each other. Thus a molecule
of methane or carbon tetrachloride has the shape of a regular tetrahedron, the carbon atom being at the center and the hydrogen
or chlorine atoms at the four corners Fig.19 e.g.
The four Hydrogen atoms in methane CH4 is considered as being bonded by the overlap of Hydrogen 1s orbitals with the
half-filled orbitals of the Carbon atoms because of the s orbital being non-directional and p orbitals are directional, the
four bonds (right angles to each other) in CH4would not be equivalent in direction. Also abnormality arises between
experiment and theory of these four bonds, as they would neither be equivalent in direction (due to s orbital being
non-directional & p orbital are at right angles to each other) nor be of equivalent strength (as s orbitals have lower energy
than p orbitals). It has been found experimentally that 4 C––– H bonds in CH4 to be equivalent in direction and
reactivity and the orbitals that constitute this must also be equivalent.
Maximum overlapping will occur if the orbitals have the same symmetry with respect to the C — C axis, that is, if all six
atomic nuclei lie in the same plane. Overlapping then takes place both above and below the plane. [Your opinion is important to us. If you have a comment, correction or question pertaining
to this chapter please send it to
comments@peoi.org
.]
In order to maintain overall electrical neutrality (free state), doubly charged ions are always
associated with two single
charged ions of opposite sign or another double charged ion of opposite sign.
Other e.g. MgO (Mg2+ + O2-
).
When there is sharing of electrons between atoms, the bonds are called covalent or homopolar bonds. Valency of the atom is the
number of electron pairs shared with the other nuclei.
Transferring electrons is not the only way that atoms can combine, covalent bonding
occurs when two (or more) elements share electrons.
Most common type of chemical bonding is single covalent bonding, where one pair of valence electrons is shared by the two
atoms.
Hydrogen molecule [H2] Fig. 8a
Fig
8b
For every pair of electrons shared between two atoms, a single covalent bond is formed. Some atoms can share multiple pairs of
electrons, forming multiple covalent bonds.
Oxygen molecule [O2] Fig. 9
The multiple bond in N2 molecule, is formed by, each of the two atoms in the molecule of
nitrogen completing its octet by
sharing three of the five electrons belonging to the other atom. Two electrons with opposite spins
constitute the ordinary
covalent bond, and as the six shared electrons in the nitrogen molecule form three such pairs the
result is three covalent
bonds joining two atoms. Acetylene (C2H2) also has triple bonds.
Double bonds are formed in similar way; e.gs. ethylene (C2H4), carbon dioxide (CO2). Double
bonds are more common than triple bonds.
The main contribution that Lewis made to structural theory was the suggestion that a pair of
electrons could be shared between
two atoms to form a bond while still contributing to the stable octet of each (or duplet in the
case of the H atom).
Fig. 10
The problem with this "solar system" diagram of the atoms is that sharing two pairs of electrons gets even more confusing in
molecules having triple or quadruple covalent bonding. Thus a method to simplify a diagram of the molecule was devised. It only
shows the valence electrons as dots. It is called the electron dot notation or Lewis dot notation.
Lewis dot structure is shorthand to represent the valence electrons of an atom. Valence electrons are those that are in the
outer orbit or shell of an atom. The structures are written as the element symbol surrounded by dots that represent the valence
electrons.Fig 10a
Fig.10a
Lewis structures for H2 and O2 Fig.10b
There are, in fact, two subtypes of covalent bonds.
Because the valence electrons in the water molecule spend more time around the oxygen atom than the hydrogen atoms, the oxygen
end of the molecule develops a partial negative charge (because of the negative charge on the electrons). For the same reason,
the hydrogen end of the molecule develops a partial positive charge. Ions are not formed; however, the molecule develops a
partial electrical charge across it called dipole Fig. 16.
Water molecule [H2O] Fig. 10c
Properties of Covalent Compounds.
Molecules with a lone pair of electrons like those of Methanol (CH3OH),
Ammonia (NH3
), Oxygen (O2) or Water (H2O), frequently behave as
electron donors and use the lone pair to form a
covalent
bond with an
electron acceptor. Molecules which behave as electron acceptor have available vacant
orbitals and gain extra stability from the
additional
electrons. Boron fluoride is an electron acceptor and forms a stable additional compound with
Methanol
Fig. 11.
As a result of an atom losing a share of two electrons, it is left with a partial positive charge
and the electron acceptor
that gains the extra share of electrons, acquires an equal partial negative charge. So that associated
with the co-ordinate
bond, there is also a weak ionic bond and the co-ordinate bond,
could equally be represented as A
б+—>B
б- (the symbol implies a partial charge on the atom).
Because of
this it has
been called a Semi-polar bond. The term dative bond is also used. Co-ordinate bond is represented
by an arrow sign with or
without partial charges.
Another example is Ammonium Chloride (NH4Cl), which contains
electrovalent and simple covalent
linkages as well.
Covalently bound to the N atom in the NH4+ ion in this way, the H+
ion from hydrogen chloride is
indistinguishable from the other H atoms, and the positive charge is spread symmetrically over the
whole structure giving it
the characteristics of a unipositive ion. This then combines ionically with negative Cl-
ion in the usual way,
ultimately producing an ionic solid
Fig. 12.
The properties of a particular salt are conditioned by the size and charge of the cation.
The alkali metals form unipositive
ions, which have large ionic radii. The NH4 radical has a single positive charge
and a large radius, this explains
why there are similarities between NH4salts and the salts of alkali metals.
We can use the idea of ionic and covalent
bonding to explain all of the
compounds formed between metals and non-metals, non-metals and other
non-metals and elements in which more than one atom is usually
present ( H2, O2, N2 etc.).
However, most of the elements in the periodic
table are metals and we
need to look at how these bond together. Metal
atoms, as you will
remember, all have very few electrons in their
outer shells.
When metal atoms get together, the outer
electrons
(depending on the element concerned, this
will be one, two or three) are given away,
changing the atoms to
cations. All the electrons form a negative
charge 'cloud' or 'sea'
surrounding
the cations. The negative sea of electrons
acts like a glue which holds
the cations together.
All the properties of metals are a direct
result of metallic bonding:
We will look more closely at the structure of metals in chapter six.
It would be unwise to consider the bonding between atoms as the only
way that we can understand them. Chemists use the word
bonding to think
about what is happening on a very small scale. If we wish to
think on a bigger scale that considers how
substances behave,
chemists use the word Structure. Table below shows Structure, type of bonding and examples.
The structure of diamond is an excellent
example of how closely
the properties of a substance are related to its structure.
Diamond is
made up of carbon atoms, each forming four strong covalent bonds
with its neighbours. As a covalent bond is
strong,
four covalent bonds (in a three dimensional
model at 109° to one
another) make for a very strong structure indeed.
Diamond is what chemists call a 'giant'
covalent structure, namely that a single diamond is one structure, made
up perhaps of millions of carbon atoms.
The unique optical properties of
diamonds are due to the fact that the atoms
are arranged in a very
clear pattern. Diamonds, apart from being 'a girl's best friend,' have
many uses where a hard and strong substance
is vital, such as a
dentist's drill. It is possible to make
artificial diamonds - albeit
with some
difficulty - but these are reserved for industrial use as their
appearance makes them unacceptable for jewellery
Fig. 13.
Graphite
Like diamond, graphite is a giant covalent
molecule composed entirely
of
carbon atoms. Unlike diamond, graphite is not
particularly strong and
does not transmit light; in fact it is a
rather dull, grey solid.
In the graphite structure, carbon atoms are
joined to each other by three,
strong covalent bonds, forming a layered
structure
Fig. 14.
If you think of the
carbon atom you will remember that it has four electrons in its outer
shell, the fourth electron from each atom being free to move around.
Electrons that move will of course carry a
charge, and that is why
graphite is unusual among non-metals (but not
unique) in that
it conducts electricity.
The free electrons in graphite behave in a
similar way to
the electrons in a metal (see above).
So far, we have restricted our discussions to the bonding that takes
place between atoms - ionic, covalent and metallic bonding.
There are a number of weaker intermolecular
attractions that
occur between molecules.
These are called "intermolecular bonds" and
can be defined as
a weak attractive force between molecules,
caused by a small difference
in electrostatic charge. While these bonds
are very easily broken,
they do have a great influence on the physical
properties of a
substance, such as its melting and boiling point.
The most common of these intermolecular
attractions are
called VAN DER WAALS forces and
exist when small areas of negative
(δ-)
and small areas of positive charge (δ+) form
on the surface of the atom for a very short time.
Even inert atoms like helium have momentary
small areas of charge on
their surface due to the movement of electrons;
this very small attractive force enables
helium to be liquified at very low
temperatures (-269°C or 4K). An element that
had no attractive
forces between atoms would remain a gas even at absolute zero
(-273°C or 0K) - such particles do not exist.
A very important form of intermolecular
bonding is the HYDROGEN BOND of
which we will have more to say in later chapters.
Hydrogen bonding
occurs in molecules where hydrogen is bonded
covalently to one of a
few other elements, most importantly oxygen and nitrogen.
These elements have a much stronger pull on
the electrons in a bond than
the hydrogen atom, and are said to have a
greater
ELECTRONEGATIVITY
3vii.
Electronegativity is a measure of how good
atoms are at attracting
electrons; a numerical value is given to each element which is nothing
more than an index, a way of comparing it with the other elements.
Electronegativity values have no units (you
generally don't need
to refer to the actual
values for electronegativity. Electronegativity increases across the period and decreases
down the group e.g the value
increases from Li (1.0) to F (4.0) and decreases from Li (1.0) to Cs (0.7) Fr (0.7). Metals
have lower values
and non metals - especially those in groups
VA (15) - VIA (16) and VIIA (17) of the
periodic table - have high values). If a
covalent bond is formed between
two
atoms of differing electronegativity, the
electrons are unevenly
distributed and the bond is described as
POLAR.
A bond between two atoms of similar
electronegativity is
not going to have an unequal distribution of
charge and the bond is called NON-POLAR.
In the diagram below
Fig. 15, all the bonds
are polar except the H-C bond.
In the diagram
Fig. 16
we
can see how hydrogen bonds form between the water molecules
( hydrogen bonds are shown by a dotted line
----------).
It is hydrogen bonding that accounts for the
very unusual behaviour of
water (see Chapter 10).
As the atoms are brought together, the 1s atomic orbital of hydrogen overlap and coalesce to form molecular orbital, the
contour of which defines the region, which both electrons (provided that they are of opposite spin) can be considered to occupy
jointly, Figure below.
In chlorine molecule half-filled 3p orbital are involved. Stability of the Hydrogen molecule is due to the fact that electrons
are within larger region of the molecular orbital, although greatest electron density is found in the region between the two
nuclei i.e. shared electron pair.
In dative co-ordinate covalent bonding, the bond is formed by the overlapping of an atomic orbital occupied by an electron pair
with a suitable vacant orbital.
The electronic configuration of Carbon [C] is 1s2, 2s2, 2px1, 2py1 (K 2, L 4), in
which px & py are half filled, and C is divalent. It is concluded that the quadrivalent nature of carbon is due to one of the
2s electrons being promoted to the vacant 2pz orbital thus giving four half-filled orbital.
When an atom is in an ‘excited’ state through absorption of energy, unpairing of its electrons may occur. If there is an
unoccupied orbital in a higher energy sub-level one of the unpaired electrons is ‘promoted’ to the vacancy and both electrons
become available for bond formation. Beryllium atom in its ground state has an electron configuration 1s2 2s2
. When beryllium combines with chlorine the 2s electrons become unpaired and one is ‘promoted’ to one of the three empty
orbitals in the 2p sublevel. One might expect that the beryllium atom would combine with two chlorine atoms by means of an s-p
bond and p-p bond. This would result in two bonds of different strengths, a p-p bond being stronger than an s-p bond because of
greater overlapping of the charge clouds. Actually the two bonds in Beryllium Chloride (Be Cl2) are equal in
strength. This is because of a process known as hybridisation of atomic orbitals. The wave systems of the two-beryllium
electrons interact to give two similar orbitals, with axes at 180 degrees to each other. The beryllium chloride molecule is
thus linear (Cl—Be—Cl). Since the new orbitals are derived from one s orbital and one p orbital we describe them as ‘sp hybrid’
orbitals. The advantage of hybridisation of atomic orbitals is that the hybrid orbitals are concentrated in particular
directions, so that there is greater overlapping of charge clouds, producing stronger bonds. . Similar sp hybridisation of
orbitals occurs in the mercury atom when it combines with two chlorine atoms. Mercury (II) Chloride also has a linear molecule
(Cl—Hg—Cl).
Fig. 17
Fig. 18
Fig. 19
The abnormality is solved by modifying the direction and energies of the orbitals by mathematical solution i.e. hybridisation
of one s orbital and the three p orbitals (sp3 hybridisation) leads to replacement of these orbitals with four new
hybrid orbitals of equal energy which are directed towards corners of the bonds in the methane molecule, result from the
overlap of equivalent hybrid in orbit with hydrogen 1s orbital to form experimentally verified tetrahedral molecule [image
methane]. Other modes of hybridisation of C atomic orbitals are possible.
In formation of the ethene molecule (C2H4) H2C: :CH2 which contains a
‘double’ bond, hybridisation of
atomic orbitals of the ‘exited’ carbon atom takes place differently. ‘After promotion’ of one of the 2s electrons to the 2p
sub-level the remaining s orbital combines with only two of the p orbitals (sp2 hybridisation), so that each carbon
atom has three sp2 hybrid orbitals and one unchanged p orbital Fig. 20. As with the boron atom, the axes of
the sp2
hybrid
orbitals lie in the same plane and are at 120 degrees at each other. Two of the hybrid orbitals are used in bonding with two
hydrogen atoms, and the third in establishing a single bond with the other carbon atom. The axis of the remaining p orbital is
at right angles to the axis of the hybrid orbitals, and the second bond between the carbon atoms is formed by overlapping of
the two p orbitals or charge clouds.
Fig. 20
Fig. 21
Two types of bonds involved in hybridisation.
1] A Sigma bond is formed due to maximum overlapping between two nuclei (internuclear axis)- called ‘collinear overlapping’
i.e. covalent bond formed by collinear overlapping of atomic orbitals is called a Sigma (σ) bond. Examples of sigma bonds
with collinear overlapping are H—H, O—H, C—H in methane (all sigma bonds).
2] A bond formed by ‘collateral’ overlapping of p orbitals is known as a pi (π) bond. Stronger of two bonds is Sigma bond.
Ethene cannot spin or twist because of the pi bond Fig. 21.
Prior: Atoms
Modified: 2014
Next: The Periodic Table