remember last time we were talking about
metallic bonding metals
solids
now as we have seen before we have
defined the metallic bonding attraction
between the electron cloud which moves
freely in the solid metal and the
positively charged metal ions we simply
imagine a metal as if the positively
charged metal ions are put into an
electron see now these electrons II or
electron cloud you can call it are
actually the valence electrons off the
metal atom so a net attraction between
the this electron cloud and the
positively charged metal atom simply
holds the metal together and we call
that metallic bonding now again as we
said last time this happens because the
metals do not have enough electron
valence electrons to make the electron
configuration as the noble gas electron
configuration by Charon making bonds
like for bonds or with the surrounding
metal atoms therefore they get together
to form an electron called electron C
and the positively charged metal ions
there that's how we simply consider and
imagine the metal atom but as I said
last time there are some problems with
this model because when you consider for
example the transition metal atoms 3 d 4
D first transition metal atoms second
and so on the metal bond strength which
is indicated by the melting points of
the metals for example one property
increases as the number of electrons D
electron increases then when we reach to
the middle it becomes maximum the
strongest metal metallic bonding
actually and then as the number of
electron increases
the metallic bond decreases again
actually as the number of electrons
increases that will provide more
abundant with the other metals and
should make the metallic bond and
stronger but it
that doesn't happen actually as I said
experimentally so there is another model
that we can explain the metallic
properties and the electronic structure
of the metals electronic structure of
the metals and this can be consider
again as based on the molecular orbital
theory now how let me try to explain now
remember again we usually talked about
the molecular orbital energy diagrams or
the electronic structure of the
molecules based on by molecular diatomic
molecules right made of two atoms and we
combined the atomic orbitals to form the
molecular orbitals now remember again
that the number of molecular orbitals
are always equal to the number of atomic
orbitals that we linearly combine if
they combine to we get two molecular
orbitals and one of those is the
abundant lower energy the other one is
anti-bonding again higher energy and
then the electrons valence electrons
actually goes into these molecular
orbitals first we start filling them the
lowest energy bonding and the
antibonding and we also define a simply
bond order as the number of electrons in
the bonded molecular orbital minus the
number of electrons an anti-bonding
divided by two and this actually gave us
informations about the bond strength
between those two atoms actually which
form the molecule now again we said the
molecular orbitals are very similar to
the atomic orbitals but they cover all
over the molecule so the electron which
is in a molecular orbital means that
move around all atoms in the molecule in
the molecule free to move because it
orbitals is that big
and again molecular orbital's just takes
two electrons just the same as the
atomic orbitals we can put two electrons
in one molecular orbital and they should
be in spin paired opposite and spin now
let's simply consider the lithium
molecule like we have done before so
when we consider in lithium molecule
remember the lithium three electron
configuration is 1s2 2s1 it has 1s
electron as valence electrons so again
these two s orbitals actually from one
lithium and from the other lithium
combine to form bonding molecular
orbitals and antibonding molecular
orbitals and the bonding molecular
orbitals is Sigma type remember
anti-bonding is also Sigma type and
since we only have two electrons one
from each lithium valence electrons that
go into the bonding molecular orbital so
the lithium in fact the electronic
structure if you write down the
electronic structure that will be like
Sigma 1s two electrons there and then we
have the core electrons which is
involves this one s but forget about
that just take the valence electrons now
they are simply indicated as here this
is the bonding molecular orbital okay
this is the bonding molecular orbital
this is the shape again an electron goes
here in this bonding molecular orbital
they move all around the molecule which
is made off from two lithium atoms in
this case and this is the anti-bonding
molecular orbital where we have a nodal
plane sort of between the two atoms and
the electrons are more there as we said
again more electron is in abundant
molecular orbital in a molecule means
that the more stable the molecule
because the bond order will increase now
what happens if I don't if I take for
example for lithium atoms right now
these four lithium atoms can be again as
follow we can think about this as one
lithium and other lithium and another
lithium and another lithium here all
same distances away from each other or
you can think about this as one lithium
here one leave him there they're very
the same distance again and in this case
again if we consider the molecular
orbitals we will have four atomic
orbitals to combine these are actually
2's orbitals right and when they combine
they will form for molecular orbitals
now where are these molecular orbitals
these molecular orbiters will be one
will be the abundant the lowest again
and now the second one will be little
higher
over here and then than the third one
and the fourth one and so on right like
this now if I simply increase the number
of lithium atoms and then go to the
lithium atom infinite number of lithium
atoms this can be again in three
dimensions now again you can think about
the molecular orbitals that will form
how many molecular orbitals will form
there's going to be actually if they're
infinite as the number of lithium atom
goes to infinite that means that the
number of molecular orbitals will will
go to the infinite two now what does
that mean that means half of those
molecular orbiters will be bonding half
will be anti-bonding but when you have
that much molecular orbital actually
what really happens that the energy
differences like between the molecular
orbitals becomes almost zero right
almost zero so that means there is
molecular orbital here molecular orbital
very close to that molecular orbital
here here here infinite almost number of
molecular orbitals now again when you
have those infinite molecular orbiters
and Delirium infinite case again as I
should say the number of n goes to
infinity that probably put it that way
can't have infinite but that's like
we're talking about a piece of
lithium metal a piece of just solid
material right there in you can you can
say the end goes to infinite their
number of lithium atoms now there's
going to be like two n that means twice
as much as the number of lithium atoms
with the molecular orbiters here bonding
and anti-bonding but they're very close
is energy you can't really separate
those so you fill those electrons what
happens it will fill up to the middle
right here because when we combine n
lithium atoms you will form in molecular
orbitals but you will have actually n
electrons since you put two electrons in
one lip molecular orbital that means
half of the molecular orbitals will be
full half will be empty now this series
of molecular orbitals this is the
highest energy this is the lowest energy
again
this is again in metal metallic solid
actually and you feel the electrons
there is this is called bend right this
series of molecular orbitals which are
very very closely spaced between this
energy and that energy is called bend
again
now the electrical conductivity can
easily be explained here because when
simply these electrons here filled with
the molecular orbitals and filled
molecular orbitals electrons absorb
energy from the light or from other
sources they will be excited to the
higher molecular orbitals so I will have
molecular orbitals here which are
partially filled now those partially
filled molecular orbitals that means one
electron in one molecular orbital will
function to carry the electrical charge
through the metal lithium that means
that the electron in those molecular
orbitals can move from one end of the
metal to the other man and and that will
carry the electrical charge right that's
how we explain that now on the other
side again with
Bend approach the more electrons we have
in the bonding molecular orbitals and
less electron in the anti-bonding
molecular orbiters you can think about
this these are sort of anti bonding then
the more stable will be the metal and
higher will be the melting point this is
simple lithium atom case right now what
happens if I have transition metals
which have D orbitals S orbitals and P
orbitals as valence orbitals for example
3d transitions we have four s orbitals
lower energy here three D orbitals there
and then 4s orbital there actually
remember the energy of these 4s and
three D changes as we fill out that but
think about this again as follow we have
four s + 3 d + 4 P now we can think
about exactly the same way we thought
about the lithium metal now all the four
s orbitals from each metal atom will
combine to form and molecular orbitals
as n goes to infinity they will simply
spend an energy range and that will be
from here to here and this is called the
s Bend this is again the series of
molecular orbitals which are very
closely spaced energy wise and they're
made up from four s linear combinations
of the four s orbitals like that then
the D orbitals five of those D orbitals
from each metal atom they will also
linearly combine to form molecular
orbitals and since the end goes to
infinite these molecular orbitals which
simply obtain from the linear
combinations of the D orbitals will have
a band like this and this is simply
called 3d band this is called for air
spend now then the p orbitals here
they will also combine linearly to form
molecular orbitals exactly the same
fashion as the 4s and 3b and they will
simply have a band
here we call that P bet what does that
bank means that means series of
molecular orbitals very large number of
molecular orbitals spaced very very
closely and energy wise now once you
have these separate bands s Bank P band
and D band in this energy scale actually
the relative energies with respect to
each other as shown here then we will
take the electrons this is actually the
nickel right twenty-eight atomic number
and we take these electrons here and
then we start filling them from the
lowest energy molecular orbital which is
in the S band and then continuing over
here when we reach to the here then we
fuel SMD together like this okay and
then when we reach here the P we simply
start fill-in s P and D this looks like
you know pouring water here and as the
water reaches to this region then the
water goes over here and as it reaches
to that level it goes to here here and
here
that means we fill out the electrons in
these bands like this this is the
electronic structure of the nickel atom
you have three bands now again here we
have empty molecular orbitals here we
have empty molecular orbitals and here
also we have empty molecular orbitals
any electron excited to these level
actually is simply will provide the
electrical conductivity actually this
part fill part of the as we will see
bends are called the valence band and
the empty part is called the conduction
band again that means the electrons in
the conduction band are the ones that
conducts carry the electrical charge and
make the metal atom conductor okay this
is the approach again molecular orbital
approach to explain the metallic
properties and it does better than the
other electron and electron C and the
positively charged metal ions okay now
the next we will talk about the ionic
solids now ionic solids remember are the
solids made of positive and negatively
charged ions
and they are crystalline structures and
again the distances very regular pattern
they have a legislative irregular
pattern and they have a unit cell just
like other crystalline solids now the
easiest way to imagine the ionic solids
would be as follow now in this case we
have positive and negative charges and
remember again the attractions of
positive and negative charges results
the ionic bent and that is the one keeps
the solid ionic solid together and they
are very hard there are properties of
those do not conduct the electricity and
so on but we can imagine the structure
of an ionic lattice as follow usually
one of the ion is bigger than the other
and that is usually the negative
negatively charged ion right you can
imagine that the larger ion or
negatively charged ion will form a
close-packed lattice that means
hexagonal lattice exactly with a
hexagonal unit cell or cubic close
packing which have the face and the
cubic unit cell and then the other iron
smaller ion usually the positively
charged ion will fill into the holes now
remember what kind of holes we had in
the close back in structures we had
remember last time I try to show you
that the holes between four metal atoms
we call that tetrahedral holes and the
holes that is between six metal atoms
and that is called the octahedral holes
and the octahedral holes are always
larger than the tetrahedral holes and we
also said that the number of tetrahedral
holes is equal to two times the number
of metal metal atoms that form that
lattice and the number of octahedral
holes are equal to the number of metal
atoms as a matter of fact so this is how
we can imagine that now with this for
example
let's see if I concede here
we can consider for example the sodium
chloride structure like over here like
this let's consider this sodium chloride
structure now sodium chloride you have
sodium plus ions and you have chloride
ions this is definitely much larger than
the sodium chloride ion is much larger
sodium are smaller so we can think about
that the chloride ion forms the cubic
close packing lattice which have
actually phase centered cubic
no yes face centered cubic unit cell
considering the chloride ions chloride
ions and then the sodium atoms fill into
the octahedral holes fill to the
octahedral holes
now remember number of octahedral holes
always equal to the number of atoms that
form the lattice close packed lattice
that means in this case all of the
octahedral holes will be full in a
sodium chloride right all of the
octahedral was full because the number
of octahedral holes equal to the number
of atoms that form the letters in this
case the number of atoms actually ions
which is the chloride ions is the one
that formed lattice so that simply means
that one chloride and one sodium ratio
comes out for the sodium chloride
structure now then if you consider the
zinc sulfide again again the zinc
sulfide we have zinc two plus and the
sulfide 2 minus ions here
here again the sulfide is much bigger
than the zinc two plus ion so the zinc
actually sulfide ion actually will form
the closed pack structure and the zinc 2
plus cation will fill in the detector he
drew holes right now this way again you
see in nature try to
use the space with the maximum
efficiency to get the highest density
right so this way you know the the empty
spaces between the atoms are also
partially fueled and that is the
structure of the zinc sulfide this is
called zinc blende actually zinc sulfide
now let's think about this way let's
assume that the sodium chloride
structure chlorides forming the cubic
close pack structure that means that the
unit cell with respect to the chloride
ion is face centered cubic so how many
unit cell belongs to it lets ability
this way how many chloride ions belongs
to one unit cell only chloride ion how
many four because it is face centered
cubic with respect to the chloride
remember the face centered cubic has
four atoms or four atoms make one face
centered cubic now let me ask you this
how many octahedral holes will be
present in one unit cell for exactly
because we said that the number of
updater laws equals the number of atoms
that form the letters for chloride for
Octavia laws so now when you have all
the octahedral holes are full with
sodium ion that means in one unit cell
how many sodium ion for so the unit cell
of sodium chloride contains four
chloride and for sodium ion right but
remember the ionic compounds do not have
molecules so the formula of those
compounds are not called molecular
formula but they are called the unit
formula so the actually the sodium
chloride lattice will be represented as
sodium for chloride for but the simplest
form of that one stood in one chloride
that is the formula of the sodium
chloride now in the case of zinc
sulphide actually how many tetrahedral
holes it's difficult to see it here but
there is actually tetrahedral holes
there
how many tetrahedral holes will be exact
zinc sulphide if the sulphide forms the
face centered cubic unit cell eight
right because the number of tetrahedral
holes is two times as the number of
spheres that we knew packing so eight
but in this case these eight actually
tetrahedral holes in this unit cell in
this cube are at this corner that corner
not exactly corner but some somewhere
inside over here here here here there
and then also form or below there eight
but in the zinc sulphide only half of
those tetrahedral holes are full with
zinc two plus iron the other half is
empty why because the formula actually
the ratio of the zinc two plus ion to
the sulfide has to be one right from the
formula so only half of that is full and
with the zinc sulfide now the question
here is what determines where the
smaller cation will go octahedral or
tetrahedral for example in a case of the
sodium chloride structure we said the
sodium plus goes into the octahedral
holes fills the octahedral holes but the
zinc sulphide goes in a tetrahedral
holes and the factor which determines
that is the size ratio of positive and
negatively charged ions if actually R +
/ r minus that means radius of the
positively charged smaller to the radius
of the negatively charged is about equal
to or less than I should probably say
two point
two-five then the iron will go into the
tetrahedral holes right now when it goes
into the tetrahedral holes what really
happens in that case it doesn't leave
really disturb the close-packed
structure it doesn't push away the other
atoms for example chloride ions if it is
exactly equal to that but in the case of
sodium this ratio is bigger than that
what does that means that means that in
this structure over here if I ever find
my okay in this structure over here that
means that chloride ions I'm not
touching each other
remember in close packed structure all
the spheres has to touch the each other
right why doesn't touch each other
because when the sodium goes into the
octahedral holes that ratio is a little
bit bigger than that and what that
happen that makes that push the chloride
ions a little bit away from each other
right but still you know pretty good
efficiency of packing there considering
that the empty holes are fueled by the
sodium plus ion again now if the ratio
of a plus a minus is greater than or
about somewhere about like 0.41
somewhere about there I can't remember
exactly for - I think this is somewhere
now if the this ratio is bigger than
that then it goes into the tetrahedral
holes no sorry
octahedral goes octahedrons okay I guess
we have simply confused here with the
sodium chloride and zinc sulphide
structure now this is for tetrahedral
hole this is for octahedral hole so that
means that this ratio is for the zinc
sulphide that is zinc over sulfide ratio
actually bigger than that and that's
sulfides are not touching each other
because that ratio is bigger now this is
for the sodium chloride cases because
the ratio is bigger than that and the
sodium goes into the octahedral holes
here zinc 2 plus goes into T 2 8 rolls
zinc 2 plus is smaller than sodium plus
actually right and the sulfide is big -
2 - sulfide is pretty big bigger than
the chloride and so on so this is how we
can imagine the ionic structures ionic
solid actually and we can calculate and
the number of ions that belong to one
unit cell let's for example take that
for sodium chloride now this is the
sodium chloride now sodium chloride
structure is as again as follow chloride
forms the face centered unit cell cubic
unit cell and the sodium goes into the
octahedral holes but now the octahedral
holes are here one is in the middle
right there in the middle of this unit
cell face centered cubic unit cell and
now the other one is here in the middle
of edge and this is in the middle of
edge and so on all on the middle of
edges so if I calculate the number of
chloride in sodium chloride unit cell
let's say unit cell so again I have
chloride cases that say chloride - iron
1 over 8 times 8 form the corner this is
one from there and the again 1 over 2
times 6 middle of faces this is equal to
3 so we have simply four chloride ions
if I simply take now the sodium eye and
calculate the number of sodium ion per
unit cell that would be again as follow
one will be from the center that belongs
to this unit cell which is right in the
middle of here there in octahedral holes
one sodium octahedral hole Center and
then we have simply one at the middle of
edges now the edge will be shared by how
many unit cell how many coop in three
dimension edge like this how many no
four right because this edge will be
shared by unit cell here by unit cell
here and up here and up there right four
so I have simply from the edges I will
have 1 over 4 and how many edge we have
12 right
and that times 12 and that makes three
sodium these are the edges here so if I
simply add them up I have 4 so again you
see in one unit cell sodium chloride I
have 4 chlorides and for sodium atom
actually if I would like to represent
the formula of a unit cell and sodium
chloride structure will be sodium for
chloride for right like this now zinc
sulphide cases we can also calculate
that but tetrahedral holes in this case
are all inside of this and there are 4
sulfides and 4 actually zinc 2 plus and
the formula again 1 zinc and 1 sulfide
ok
now you see the ratio of the positive to
the negative charges now this is the
sodium chloride and the ratio is 0.6 T 9
so it has to be what
not to disturb the close packing
structure this has to be point four
right but it is much higher than point
four actually what does that means the
chlorides are pushed away and this one
is point 52 now that is also much larger
than point two twenty five right and
that means that the sulfides are also
pushed away by the zinc two-plus iron
when it fills into the tetrahedral holes
you can think about that now how about
this one
now this one is the cesium chloride
structure and that is not a closed pack
structure again with respect to the
chlorides or with with respect to the
cesium's it is different actually so the
lattice cesium chloride lattice will not
be a closed pack lattice but these two
can be considered as a closed pack rat
lattice with respect to the negatively
charged ions as a matter of fact now
here again we have a cube of chloride
the chlorides are on the corner of a
cube simple cubic unit cell which refers
to the chloride ion cesium is in the
center now cesium plus actually is much
larger than the sodium and and this this
this hole which is called cubic hole but
that's not present in the close packed
structures but in cesium chloride
structure we say the chloride ion forms
the simple cubic lattice and the cesium
two plus ion fills in two cubic holes
cubicles in that case cubicles between
eight Spears here right in the middle of
that okay now the next is we talked
about the again metallic solids and X is
the molecular solids we talk about ionic
solids now the molecular solids
molecular solids remember the attraction
forces between the molecules in the
solids are London forces actually
hydrogen bonding dipole-dipole type
forces and they are usually
weaker as we said before and that simply
means that the molecular solids are
usually soft and the melting points are
lower than the ionic solids for example
now you see again in the molecular
solids the lattice points or the units
actually are molecules right molecules
so the attractions are the as I said
intermolecular forces that we have
talked about before and they are pretty
weak but we should simply consider the
shape of the molecules particularly when
we account for the intermolecular forces
compare the intermolecular forces for
example if you look at this benzene
toluene and phenol now this is nonpolar
molecule so the only force that
effective here holding that is London
forces right or dispersion forces same
things and this one is also nonpolar
again not polar tahleen
very similar here also dispersion forces
now this one here is phenol we have
dispersion forces we have hydrogen
bonding and we also have dipole-dipole
this is also polar molecule so when you
compare the mountain points five benzene
and 95 toluene and 43 phenol much higher
so this simply shows the intermolecular
force is much stronger in phenol and
then benzene and toluene but if we
consider only the dispersion forces
toluene which have higher molecular
weight than benzene should have a
stronger dispersion forces right
remember the special forces increases
with increase in molecular weight now
the reason here is because the packing
of the Tholian
in the solid is not the same as benzene
so the
distances between two toluene molecules
is much larger than the distances
between two benzene
so the intermolecular forces which are
effective in short distances is more
effective in benzene than Tolley in for
example when it comes to the liquid you
see again in the boiling point of
liquids benzene is 80 Tolley in 11 and
phenol is 182 because here the molecules
are simply moving around than directly
we can compare the intermolecular forces
with the molecular weight of the liquid
in fact where liquid molecules
now our covalent Network solids is
another type of solids and we said that
the covalent Network solids are hold
together by covalent bonds covalent
bonds and these covalent bonds are very
very strong and therefore the solids has
very very high melting point and very
strong again and the example is one
example is diamond now in diamonds is
carbon but it is structure is like this
and every carbon is bonded to four other
carbons around with a covalent bond
that's a very very strong bond and
actually the diamond is known as the
hardest solids hardest solid and and the
melting point of the diamond diamond is
about 3500 very very high and that means
that the attraction forces holding the
solid together are very very strong they
are actually covalent bonds again now
graphite is another Network solids again
we have the graphite layer actually the
carbon layers here these are all
hexagons and all of these carbons are
sp2 hybridization and we have a covalent
bond within the plane here one layer of
graphite
but between the layers we have actually
dispersion forces dipoles and other
intermolecular forces as a matter of
fact this is another network solids now
the silicon dioxide the glass actually
the crystalline silicon dioxide is
called quartz quartz is crystalline
normal glass is not crystalline but
amorphous that is also covalent Network
solids because the all atoms in three
dimensional lattice are bonded to each
other by covalent bond or force of
attractions are actually covalent bond
very very strong now uh let's look at
the semiconductors now remember we said
that the elements can be class elements
in the periodic table can be classified
as metals non-metals and metalloids and
the metals have high electrical
conductivity actually their conductor a
good electrical conductor heat
conductors and the nonmetals are not
conductors they do not conduct
electricity they do not conduct heat
again but there are also in the periodic
table remember sort of elements between
the metal region and nonmetal region and
those are called metalloids so the
properties of those are between the
metals and intermediate between metals
and nonmetals that simply means that
their electrical conductance is not zero
but as not as high as metals at all but
they're still conductors and they are
called semiconductors we call them
actually semiconductors now let's let's
look at the semiconductors now the
element semiconductors are actually
silicon germanium tin that is for a
group
where we have the diamond like structure
diamond like structure that means every
silicon for example is bonded to other
four silicones and every germanium
actually has bonded to four and so on
now when we consider the electronic
structure of semiconductors for example
we can think of as the diamond or just
the silicon like we can think about that
as again diamonds in a chronic structure
or the silicon again now the bent theory
still applies to those semiconductor
substances again as follow now take for
example think about the carbon and then
think about the electronic structure of
diamond now the carbon has two electrons
in S orbitals and four carbon this is
two as for silicon it is 3 S 4 s it goes
down like this and then it has P orbit
at P electron electrons in AP orbital
two of those four valence electrons and
four valence electrons again will make
for bond here like this and then we have
the network solids covalent Network
solids now this s orbitals from each
carbon will combine to form molecular
orbitals like this bonding molecular
orbitals and antibiotic molecular
orbitals these P orbitals three of those
will combine to form bonding and
anti-bonding over here but as the number
of carbon atom increases then the
molecular orbitals here increases
and that means that they will be placed
very close in energy like this just like
a band in the metals remember any metals
so we have actually series of molecular
orbitals here in this band and what are
these actually these are some of them
are s resulted from the linear
combinations of s orbitals and some are
resulted from the linear combination of
P orbitals on the carbon and
correspondent if we have simply here
like n molecular orbitals as the number
of carbon atoms goes to infinite in fact
then we also have n anti bonding
molecular orbitals here now if you
simply take the valence electrons and
start filling them out here you will
fill out all of the bonding molecular
orbitals and this will be the empty
non-bonding molecular which is empty so
we have a bent here which is full of
electrons that means every molecular
orbital I can accommodate two electrons
there made full and here there is no
electrons now this mo diagram or bent
structure will apply for carbon silicon
germanium and tin they are all in the
same group they have four valence
electrons now this energy difference
here between the filled molecular
orbitals filled bent which we call it
valence band and the empty molecular
orbitals and that is the empty band this
one is called conduction band this one
is called valence band right now the
energy gap here for carbon it is
let me give you the exact values yes for
carbon carbon actually diamond it is 5.5
now diamond and let's call this eg
energy difference between valence band
and conduction band diamond it is 5
point 5 electron volt for silicon it is
1 point 1 1 electron volt for germanium
and it is point 67 electron volts and
for tin SM tin
this is 0.08 electron volt and for lead
let is in the last one their LED and
this is actually 0 okay now the energy
that corresponds the energy of the
visible light that means the energy
carried by photons visible photons but
remember this is going to be like one
mole molar wise or no in this case
single photons energy of the visible
light photons is about 3 electron volt
less than 3 electron volt right and that
means that actually you can excite the
electron from the valence band by
electromagnetic radiation if you give
enough energy energy that means if the
energy of the photons is equal to
the eg the bandgap here which is called
the bandgap now for the carbon it is not
possible to excite electron from the
valence band to the conduction band by
visible light right because the energy
of the visible light photon is not
enough for that is less than that but
for the silicon in fact you can excite
the electron from the valence band to
the conduction band by the visible light
actually and when you excite the
electron from the valence band to the
conduction band you create partially
filled molecular orbitals right
partially filled molecular orbitals
now those partially filled molecular
orbitals will provide the electrical
conductivity to that substance actually
the electrical conductivity will be
obtained by the partially filled or due
to the actually due to the partially
filled molecular orbitals that means one
molecular object contains one electron
now how many electrons not much but
there enough to give some conductivity
to the silicon right not as much as
metal much less than that but not as
much as non metal nonmetal case is
almost zero now if you go to the
germanium you see the band gap is
smaller that means there will be better
electrical conductance here tin is much
better team actually almost like metal
the lead is metal because there's no
energy gap between the full valence band
and the conduction band there now
actually the more electron goes there
that means the more will be the
electrical conductance of the substance
that semiconductors so this is how the
semiconductor become conductors actually
right it's exciting the electron from
the valence band to the conduction band
and having that partially filled band
provides the electrical conductivity
why's the electrical conductivity but
now here is the point the main
difference between the semiconductors
and the metals semiconductor if you
increase the temperature of the
semiconductor electrical conductivity
increases but if you increase the metal
temperature electrical conductivity of
metal decreases and how as follows this
is explained as follow for the
semiconductor increase in the
temperature means that we increase the
number of electrons that are excited in
the conduction band because that is also
temperature energy actually like higher
temperature and that means more
electrons in the conduction band means
higher electrical conductivity so
increasing the temperature increases the
electric conductivity but for the metal
actually in fact since when we have zero
gap there increase in the temperature
increases the vibrational motions of the
metal in the lattice and that increase
in the vibrational motion sometimes
breaks the molecular orbital which goes
from one end to the other end because
you change the distances between the
metal metal atoms and that decreases the
electrical conductivity of the metal so
we say that the metal conductivity
decreases with increase in temperature
but semiconductor conductivity increases
with increase in temperature okay
question so let's have a break and come
back after 10 minutes let's