it's professor Dave, let's talk about molecular orbitals so we know about
so we know aboutatomic orbitals and we learned that an atom can take its atomic orbitals and
hybridize them into molecular orbitals. these hybridized orbitals have shapes
that resemble combinations of the atomic orbitals and these are the orbitals that
are used to make covalent bonds which are simply the overlapping of such
orbitals. let's start with something simple, a molecule of H2. each hydrogen has one
electron that sits in a 1s orbital so to form a sigma bond the atomic orbitals
overlap and form one molecular orbital which both electrons will reside in. the
hydrogen molecule will sit at a lower energy than the two lone hydrogen atoms
because of the additional electrostatic interaction between each electron and
opposite nucleus. this is why hydrogen spontaneously forms diatomic molecules
but the number of orbitals must be conserved so when two atomic orbitals
come together to form molecular orbitals they will form one bonding orbital and
one antibonding orbital. electrons in the covalent bond occupy the lower
energy bonding orbital since any orbital, atomic or molecular, can hold two
electrons. but if one electron were to become excited it could jump from the
sigma or bonding orbital to the sigma star or antibonding orbital. this is
where hybridization comes into play. if an atom is participating in multiple
bonds it will utilize hybrid orbitals to form them. for example if carbon is making four
bonds as it tends to do it will take it's 2s orbital and all of its 2p
orbitals to create four sp3 hybridized molecular orbitals. in order for carbon
to hybridize its orbitals it must first promote one of the 2s electrons up
to the vacant 2p orbital. then it can hybridize them leaving four molecular
orbitals
that are identical in energy or degenerate. each of them has one electron
leaving room for another that will be provided by the other atom participating
in each covalent bond. every energy level has one s orbital, three p's and five d's and
the number of bonds an atom can participate in determines the number of
atomic orbitals that will need to be hybridized to form the necessary
molecular orbitals to generate the bonds. an sp3 hybridized carbon is typically
participating in four sigma bonds but carbon can also be sp2 or sp3 hybridized
let's look at an sp2 carbon. if it is sp2 hybridized it has just three electron
domains which can happen if it has a double bonds to something.
for example each of these carbons is sp2 hybridized because of the double bond
present. what this means for molecular orbitals is that the carbon will promote
one electron and then take it's 2s orbital and just two of the 2p
orbitals to generate the sp2 hybridized orbitals. that leaves one p orbital
unhybridized. it is the overlap of these unhybridized orbitals each containing one
electron that generates a pi bond which is what the second bond in this double
bond is. the p orbitals extend in perpendicular fashion from the plane of
the molecule and they overlap to make the pi bond. the same goes for C2H2
which has a triple bond between the carbons. each carbon has two electron
domains which means it will be sp hybridized. after promoting one electron
it needs to only utilize the 2s orbital and one 2p orbital to get the sp
hybridized orbitals. that leaves two unhybridized 2p orbitals, each with one
electron. in this case they would generate two pi bonds, one from the
overlap of these p orbitals and another one from the other set of p orbitals
that will extend perpendicularly from the first pair. what you end up with is a
linear molecule
with this propeller-like orientation of the unhybridized to 2p orbitals which
form two pi bonds beyond the initial sigma bond, which is the overlap of the
sp orbitals. triple bond. sometimes we have to generate orbital diagrams for a
molecule. to do this we show each atom on the sides with its respective atomic
orbitals and the molecular orbitals go in between. here each hydrogen has one
electron in the 1s orbital so they both go to fill up the bonding orbital
generating a molecule of H2, we can use these diagrams to calculate the bond
order of the molecule which is given by the number of electrons in bonding
orbitals minus the number of electrons in antibonding orbitals divided by two
here that gives us a bond order of one which is why H2 has a single covalent
bond. now look at 02. each atom displays its electron configuration and every
electron will go into a molecular orbital. there's the 1s atomic
orbitals which are both full that go to make the 1s sigma and 1s sigma
star which are also both full. then the 2s orbitals also both full which go
to make the 2s sigma and 2s sigma star again both full. next we have the
2p orbitals. each oxygen atom has four 2p electrons. there are six 2p orbitals total
3 per atom, so we need six 2p molecular orbitals. those will be in order of
increasing energy, the sigma, pi, pi star and sigma star. remember the starred ones
are antibonding orbitals. as we take the four electrons from each oxygen atom and
place them in the molecular orbitals bottom to top
we end up with full sigma and pi orbitals as well as half full pi stars
calculating bond order we have (10-6)/2 which gives us 2. that's why atoms
in an oxygen molecule have a double bond between them. we can predict the covalent
bonding behavior of any
two atoms using orbital diagrams like this one
let's look at nitrogen. here the bond order calculation turns out to be 3
which is why atoms in a nitrogen molecule have a triple bond between them
let's check comprehension
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