Download Covalent bond - general description 1H: 1s : F • + • F : → : F : F

Covalent bond - general description
There are many compounds consisting of non-metallic atoms, e.g. CO2, H2O, HCl etc.
Atoms in them have similar electronegativities indicating that their atoms have similar
tendency (ability) to take or give electrons. Electronegativity measures the ability of an
atom to pull the electrons toward itself. So nonmetals will share their electrons rather than
transferring them because neither atom can take the electron from the other. For this
reason the electron transfer does not happen, so the only chance to gain a stabile noble
gas configuration is to share all or some of the valence electrons. Bonding in which
electrons are shared by two atoms is called covalent bond.
Let’s examine the formation of a hydrogen molecule from two hydrogen atoms.
Since the electronic configuration of a hydrogen atom is 1s1, it needs one electron to attain
the stable configuration of helium. So, both hydrogen atoms in the hydrogen molecule in
attempt to reach stability need one electron. The electronegativity of hydrogen is 2.2, but
the electronegativity difference is 0, because both atoms are identical, none of them can
take an electron from the other. Therefore, the two hydrogen atoms will share their
electrons and a covalent bond is formed. Formation of covalent bond in hydrogen
molecule can be represented in three ways as we said previously:
@ orbital representation:
1s1
1
1H: 1s
1H:
q
@ electron dot representation (Lewis symbol):
@ line representation;
bonding pair
r
H• + •H →
H : H
H−H
Covalent bonds can be classified into three groups: nonpolar, polar and coordinate
covalent bonds.
Non-polar covalent bonds exist in homoatomic molecules such as H2, O2, N2, Cl2, P4,
S8 or in molecules in which atoms have very similar electronegativities, like in CH4 where
the electronegativity difference is 0,2 or in PH3 (∆Ev = 0,01). Both atoms involved in the
bond pull the bonding electrons with equal strength, due to the bonding electrons are
shared equally which means that electrons are not displaced toward either atom. The
centres of positive and negative charge coincide; each is found at a point of midway
between the atomic nuclei. The bond describes above is non-polar covalent bond.
Let’s consider another example of nonpolar covalent bond. The fluorine molecule,
F2, consists of two fluorine atoms joined together. The electron configuration of fluorine is
1s22s22p5 and its orbital representation is:
Two atoms join together to increase their number
of valence electrons to eight. When their half-filled 2pz
orbital overlaps, a bond is formed. As a result, each
9F:
fluorine atom completes its octet and together they form
1s
2s 2px 2py 2pz
the stable fluorine molecule.
:
:
: :
:
:
:F•+ •F: → :F : F:
each F atom
has 3 lone or
non-bonding
pairs
:
σ
:
σ
Unpaired electrons of
both fluorine atoms
form a pair of e– and
this type of a bonding
is called sigma bond
and it denotes by the
Greek letter sigma - σ.
bonding pair
in F2 molecule
σ-bond
Depend on the number of shared electron pairs between atoms, covalent bonds
may be single or multiple (double and triple). This is called bond order. The term bond
order describes whether a covalent bond is single, double or triple. The higher the bond
order - that is, the more electrons present - the more tightly the atoms are held together. If
two atoms share one electron pair, the chemical bond is single, like in fluorine molecule,
F2. Similar, In order to attempt stability, an oxygen atom needs two more electrons so the
Lewis structure for oxygen molecule is
In nitrogen molecule, atoms needs to gain
three electrons to reach stability and so it forms a triple bond to one atom of nitrogen more
(:N ≡ N:). In case of formation of a molecule of oxygen or nitrogen beside sigma there is
another type of covalent bonding called pi-bond and it designates by the Greek letter pi - π.
In oxygen’s molecule each atom has two nonbonding pairs and two unpaired electrons.
Similar, in molecule of nitrogen, both atoms shared three electron pairs and so in its
molecule there is triple bond, in which one is a sigma and the other two are pi-bonds.
σ
π
π
The triple covalent bond in N2 is a very strong bond that is difficult to break in chemical
reactions. The unusual strength of this bond makes gaseous nitrogen quite inert. As a
result, N2 coexists with O2 in the atmosphere but no reaction takes place among them. The
lack of reacting of N2(g) toward N2(g) is an essential condition for life on Earth. The
inertness of N2(g) also makes it difficult to synthesize nitrogen compounds artificially.
Polar covalent bonds exist in heteroatomic molecules such as HCl, H2O, NH3 or
PCl3. A covalent bond in which bonding electrons are not shared equally between two
atoms is polar covalent bond. Bonding atoms differ a lot in electronegativities, so that the
electronegativity difference must be more than 0,4 but less than 1,8(2). In such a bond,
electrons are displaced toward the more non-metallic element.
We will examine the formation of a molecule of hydrogen chloride, HCl. In this
molecule, chlorine (electronegativity value of 3.16) pulls electrons more strongly than the
hydrogen atom does (electronegativity value of 2.1). The centre of negative charge lies
closer to the chlorine nucleus. We say that there is a separation of charges and such bond
is polar. The polarity of HCl molecule can be described by Lewis formula in which the
bond pair shifts closer to the Cl atom than H-atom.
The δ+ signifies that the centre of positive charge is displaced toward the H nucleus
whereas δ– signifies that the centre of positive charge is displaced toward the Cl nucleus.
To express the separation of charges we use the terms partial positive charge on H and
negative charge on Cl.
The Lewis dot structures (diagrams) for water and ammonia (NH3) are shown below:
The ionic character of
a bond increases with
increasing electronegativity
difference between the
bonding elements (see the
diagram left).
Orbital Overlapping Concept /Overlap of atomic orbitals/
Orbital overlap was an idea first proposed by Linus Pauling to explain the molecular bond
angles observed through experimentation and is the basis for the concept of orbital
hybridisation. s-orbitals are spherical and have no directionality in space while p-orbitals
are oriented 90° to one another. A theory was needed therefore to explain why molecules
such as water, ammonia (NH3) or methane (CH4), instead of 90° had observed bond
angles of 105°, 107° and 109,5° respectively. According to the Pauling's theory, bonds are
created by overlapping of atomic orbitals of adjacent atoms.
In this section we will explain more precisely the orbital overlap concept. According
to this concept, the formation of a covalent bond between two atoms results by pairing of
electrons present in the valence shell having opposite spins. When two atoms approach
each other, the final configuration depends on the spin of the two electrons. If the spin of
the two electrons is parallel, the two atoms remain separated. If the spin of the two
electrons is antiparallel the two atomic orbitals combine to form a molecular orbital
(molecule). In the formation of a molecule, there is a minimum energy state when two
atoms are so near that their atomic orbitals undergo partial interpenetration. This partial
merging of atomic orbitals is called overlapping of atomic orbitals which results in the
pairing of electrons. The extent of overlap decides the strength of a covalent bond. In
general, greater the overlap the stronger is the bond formed between two atoms.
The simplest chemical bond is that formed between two hydrogen atoms. Each
hydrogen atom has one electron. As the two atoms approach each other, the nucleus of
one atom attracts the electron of the other. The electrons of both atoms and the two nuclei
repel each other due to their similar charges. But in a moment both the attractive and
repulsive forces cancel out and the overlapping between the atoms stops. Eventually the
two orbitals overlap, becoming a single i.e. molecular orbital containing two electrons. This
orbital encompasses space around both nuclei.
Although the electrons may be in any part of
this orbital, we can predict that they are most
likely to be in the space between the nuclei,
shielding one nucleus from the other and
being attracted by both. In the resulting
molecule, both atoms have two electrons
and a filled outer (valence) shell. These
shared electrons form a bond between the
two atoms. This chemical bond is a covalent
bond, a pair of electrons shared between two
atoms. When this bond forms, energy is
released. This release of energy shows that
the molecule of hydrogen is more stable than
the separate atoms.
There is a distinction made between the types of bonds which form as orbitals overlap.
The two most important types are sigma bonds (σ) and pi bonds (π). One way in which
these bonds differ is in their location with respect to the internuclear axis (a line which
connects the nuclei of two bonded atoms). In a sigma bond, orbital overlap is always along
the internuclear axis so the bond is centered directly between the two nuclei. σ bond is
formed by head-to-head overlap of two s-orbitals (σs-s); two p-orbitals (σp-p) or s- and porbital (σs-p). All single bonds between two atoms are σ-bonds. Pi bonds can only be
formed after a sigma bond has already been formed. Therefore a double bond contains
one σ and one π bond and a triple contains one σ and two π bonds.
s-orbital
σs-s; σs-p
and
σp-p
p-orbital
σ−bond
s-orbital
σ−bond
p-orbital
π- bond is formed by side by side overlapping of two parallel p orbitals. Because π bond
forms from sideways overlap of orbitals, the electron cloud lies above and below the plane
formed by σ bond. Pi bond is weaker than sigma bond. To form a pi bond, two atoms must
form sigma bond first.
p-orbital
p-orbital
π−bond
two separate electron clouds
π−bond
Here are some examples of molecules formed by overlap of atomic orbitals.
– formation of σs-p bond in a molecule of HCl as a result of head-to-head overlap of half-filed A.O.
– formation of σp-p bond in a molecule of F2 as a result of head-to-head overlap of half-filed A.O.
– bonding in a molecule of H2S (Note: this is not fully correct description; H2S is a bent molecule
and the bond angle is 109° but can not be of 90° like the picture indicates for).
At the end of this section, introduce the properties that ionic and covalent
compounds are displayed (notice the differences).
Property
Ionic Compounds
Covalent Compounds
Elements
metal - nonmetal
nonmetal - nonmetal
Phase (at STP)
solid (in crystal lattice)
solid, liquid or gas
Hardness
hard and brittle (salt)
brittle and weak (sugar)
or soft and waxy (butter)
Melting Points &
Boiling Points
high
low
Solubility
mostly soluble in water
solubility varies widely
Electrical
Conductivity
solid - nonconductor,
liquid or aqueous solution conductor
insulators