Lewis Structure Lab
• For this you need to know covalent bonds, Lewis
dots, electronegativity, geometric shapes, duet &
octet, single/double/triple bonds, etc…...
• I can’t assume you have had all these, so bear with
the comprehensive review
• gather round the campfire, this will take a while...
everything we see around us - natural or man• essentially
made, cmpds or elements - has atoms bonded with atoms
or soft? solid, liquid, or gas? these properties are
• hard
determined by bonding
• molecular structures play the major role in chm rxns
all around us, and…
• and the way they’re bonded determines structure
• how do most nonmetals bond together?
• when these two critters get close, the e of one is
-
attracted to the p of the other;
• the e (or p) of one is repelled by the e (or p) of the other
-
-
• if they are
most comfy
at a
distance
where their
e- shells
overlap - ba
da bing! we gots
us a
bond!
• mini-summary
• this is a covalent bond because outer shell
(valence) electrons are shared by both atoms
• bigger-than-mini summary
• ionic bonding (transferring e s) and covalent
-
bonding (sharing e-s) are the extremes…
• b/t these 2 extremes are…
• polar covalent bonds!
• here one atom is not strong enough to strip the
electron off the other, but it can hog it
• the hog gets a ∂- to show
it is almost (not quite) a
negative ion; the loser
gets a ∂+
• but how much do these
critters like electrons???
electronegativity
• when different nonmetals
form a covalent bond, they
almost never share the
electrons equally ➙
electronegativity
• = the relative ability of an atom
in a molecule to attract electrons
to itself
• here are electronegativity values; see a general pattern?
• F rules the en table! and the numbers generally decrease as
you move away from him
• the polarity of a bond depends on the two battling for the
electrons
• in fact, the difference in electronegativity is the
Big Deal here:
• great differences in en = more ionic sort of bonds
• small differences = polar covalent bonds
• no difference = (nonpolar) covalent
•
but don’t think it’s so black and white; it’s more like this
bond polarity and dipole moments
• if the molecule actually ends up having a partial positive
side and a partial negative side we say it has a dipole
moment
• an arrow tells us
where the
negative side is
• it always happens in diatomic
molecules with unequal sharing;
it can also happen in
polyatomic molecules
• like water:
•
this unequal distribution
of charge in a water
molecule is
unbelievably
important to life
• it can surround and dissolve ions (more on that later)
• they can attach to each other and stick (more later)
• which is why it is a liquid and a solid on earth
• here two polar bonds cancel each other
out to give a nonpolar molecule (red O’s
pull away electrons from the weaker C)
• do you see that this
molecule’s polar
bonds do not get
cancelled out, so the
molecule overall is
polar???
lewis structures
• remember: bonding involves the valence electrons of
atoms!!!
• the valence are transferred in ionic bonding and
shared in covalent
• the Lewis structure is just a simple
representation to show how the valence e-s are
distributed around a molecule
• it was developed 100
years ago by G.N.
Lewis
• chemists had figured
out by then that the
valence electrons are
used to make
everyone look noble
• he used dots to show
valence e-s, sorta
like…
•
this!
•
see the Big Pattern?
• we will skip ionic Lewis structures and go straight to ….
• covalent critters
• they still have to
look noble and they
share until all are
happy
• let’s start easy…
• H’s with their lowly single valence only need one more to be
dressed like helium, so they agree to share:
• see how we represent it with the dots?
• notice each H “sees” two electrons
• its desire for 2 total e s (to
-
look like He) is called the
duet rule
• so whenever H is involved
make sure it ends up with
2 electron dots!
NMs want
• the second-row
2 6
an s p look;
they want a total of
eight!
• called the octet rule
• they will share two per bond like this:
• notice they both “see” 8 e s
-
• a shared pair is called a bonding pair
• not shared? = lone pair or
unshared pair
• noble gases are not invited to these soirees
• ready for LD structures for bigger molecules?
(answer: yes)
• first, the rules…
example
• draw the lewis structure for water
1. collect all the valence electrons
H+H+O=1+1+6=8
2. give each bond 2
H:O:H
3. give the leftovers to the ones that need them
..
..
H:O:H
or with lines for bonds H–O–H
..
..
• what’s all this then? why is the LD structure all crooked???
• all well! all the requirements are fulfilled!
H sees 2, O sees 8
• there are several correct versions of most lewis
structures; remember, it’s just telling us where the
electrons are not how the molecule is structured
example
•
write the lewis structures for NH3 and CH4
lewis structures of molecules with
multiple bonds
• what happens if you run out of electrons before everything
has its 8?
• e.g. CO : C + O + O = 4 + 6 + 6 = 16
• start with O–C–O and you have 12 left
• but each O will need 6 and then you’ve run out!
• what about poor carbon???… :(
2
• how can you get carbon its 8?
• the O’s agree to share their lone pairs
• and we make the single bond into a double
bond!
• [think symmetrically!]
• you can form this:
•you may have written this:
•this is a triple bond! the sharing of 3 pairs
•notice that any of these is correct according to the rules, but
experiment shows the symmetric one to be the best
•when >1 structure can be drawn for a cmpd, that is
resonance
• what about lewis structures for ions? no problem!
• if it is a neg ion just add electrons to cover the
charge
• if positive, subtract from total to cover the
charge, like…
• write the lewis structure for CN
• C + N +1 = 4 + 5 +1 = 10
• C–N is the base
• add electrons so everyone is happy
• but it must be written [:C:::N:]• this would not have worked w/o the extra
-
electron
small exceptions to the octet rule
• some atoms are too small to get a full eight; some
are so big they can accommodate more
• e.g. little boron is
happy with 6 or 8
• Be can make cmpd
with only 4 e-s
molecular structure
• so far: where the electrons go
• now: how they look in 3D
• called molecular structure
(or geometric structure)
• first the basic shapes
• then figuring out which of those shapes molecules take…
• this molecule has a “bent” shape with
bond angles of ~105˚
• some molecules have a bond angle of 180˚ - called
linear structure
• there’s a third shape called trigonal planar
with 120˚ bond angles
• one of the most important, though, is
the tetrahedral
critter has four things sticking out
• this
of it in four symmetric directions
• the magic bond angle is 109.5˚
what determines what
• but
the shape will be?…
molecular structure:
the vsepr model
• the importance of molecular structure in life is high level!
• one tiny change on a giant biomolecule can render it useless
• here we learn how to predict the basic structures of molecules
based on those electron pairs (shared and unshared)
•
we’ll use a simple but effective intro to molecular 3D structure
called VSEPR
•
= valence
•
the name says it all; we’re talking about shapes of molecules due
to:
model
shell electron pair repulsion
1. pairs of e-s (shared & lone) in the valence shell and…
2. their dislike of each other…
• we want the pairs of e s to get as far away from each other as
-
they can
• we’ll start with an easy one, and an exception to the octet
rule, BeCl2
• the Lewis structure
is like this:
• concentrate on the central atom, Be
• those shared pairs don’t like each other and will move as far
away as they can…
• the farthest two pair can get is 180˚
• this forces BeCl into a linear molecule
2
• now BF (another exception to the octet rule)
• has a central atom surrounded by three
3
shared pairs
• how far away can they get?
• 120˚!
• [note: Lewis structure nothing like geometry
of real thing; was never meant to be]
• this molecule, shaped like a triangle and flat (aka
planar) is called trigonal
planar
• what about 4 pairs of electrons?
• how do we build CH , a central atom with 4
4
electron pairs around it?
• this is real new and counterintuitive
people when seeing the
• most
Lewis structure think it must be
a cross
the electron pairs can get
• but
farther than 90˚!
can actually get 109.5˚
• they
away
• a central atom with four electron pairs forms
something called a tetrahedron
• important distinction alert!
• we separate all the electron pairs, shared or alone, using
vsepr, but!
• we are most concerned about where the atoms are as
a result
• we will name the molecular structure based on
where the atoms end up
• watch…
•
first the rules…
1. draw the Lewis structure!
2. count the electron pairs (bonded and unshared) and arrange
them so they’re as far away from each other as they can get
3. determine where the atoms go
4. name the thing from where the atoms are!
now, some examples…
• predict the structure for NH3
1. get the Lewis structure
2. see that there
are four electron
pairs around N;
they would be
placed like this
• put the atoms in:
that bad boy, remembering when you name it
• name
just worry about where the atoms are
are tetrahedral;
• pairs
but molecule is
trigonal pyramid
trigonal planar;
• [not
the lone pair has
pushed the H’s down
into a pyramid]
• now water
1. get the Lewis
structure
2. see that there
are four electron
pairs around O;
they would be
placed like this
• put the atoms in:
• name that bad boy, remembering when you name it
just worry about where the atoms are
• pairs are tetrahedral;
but molecule is
bent
(aka V-shaped)
• [not linear! the lone
pairs pushed the
H’s over]
examples (molecular)
• Cℓ O
2
molecule is bent
• NF
3
molecule is trigonal pyramid
examples (molecular)
• Cℓ O
2
molecule is bent
• NF
3
molecule is trigonal pyramid
examples (molecular)
• H Se
2
• CℓO
4
-
examples (molecular)
• H Se
2
molecule is bent (V-shaped)
• CℓO
4
-
molecule is tetrahedral
examples (just molecules)
NH4+
H2S
BeF2
examples (just molecules)
NH4+
molecule is tetrahedral
H2S
molecule is bent
BeF2
molecule is linear
molecular structure: molecules with
double bonds
• vsepr works with double bonds, too!
• we see in experiments that CO takes on a linear look
2
• it’s as if the double bonds of CO
being repulsive)!
2
act like single bonds (as far as
helps look at single and double bonds as electron clouds
• ifofit repulsion
single!
double!
double pair are restricted to an area just like a single pair
• the
(they’re just a wee fatter)
•
it happens here, too, with ozone
.. .. ..
:O–O=O:
..
•
experiment shows that ozone is shaped like this:
•
do you see why it is bent and not
linear?
• look at the center oxygen
.. ..
..
:O–O=O:
..
• it has a lone pair, a single bond, and a double bond!
• this behaves
as if there were
just three areas
of electrons
• thus, 120˚
• conclusion: when using the vsepr to predict the
geometry of a molecule treat the double
or triple bond as if it were a single
• but wait! there’s more!
• some of the bigger atoms can hold more than 8!!!!
8.13
• when there are 5 areas coming
out, we get a trigonal
bipyramid
8.13
• but even these can get
weird
• a lone pair can get this
guy called a seesaw
• see the orbital
geometry is still trigonal
pyramid?
8.13
• two e- pairs can
force a T-shape
8.13
• three lone pairs can
force the bonded atoms
in opposite directions =
linear
• see the orbital
geometry is still trigonal
pyramid?
8.13
-
Three possible arrangements of the electron pairs in the I3 ion.
Which is best???
8.13
SIXth and Last
6 areas of electrons
= octahedral
8.13
Molecular structure of PCl6-
8.13
• one lone pair
forces this...
8.13
two lone pairs
this
seems the lone
pairs want to be
as far away as
they can...
8.13
Possible electron-pair arrangements for XeF4.
Which is probably best???
8.13
• a summary of it all...
8.13
8.13
8.13
• What are these shapes?
8.13
• SN = stearic number
• how many electron pairs, bonded and alone, are
coming out of the central atom
• OG - orbital geometry
• the geometry including everything, bonded and
non-bonded
• MG - molecular geometry
• geometry of only the bonded pairs
•
all our predictions will be based on these:
8.13
one more thing
• remember the s’s and p’s and d’s? they never took
these bizarre angles!!!!!
• how did these shapes come about based on
those atomic orbitals?
• they hybridized to make new orbitals!
9.1
• when they share one pair = single bond = sigma bond
• when they share a second pair = pi bond
• a third pair = just another pi bond
• single = sigma
• double = sigma + pi
• triple = sigma + pi + pi
9.1
• let’s start easy:
one s orbital and three p’s get together to
make four sp3 orbitals...
9.1
9.1
• an s and two p’s can get together to make
three sp2 hybrid orbitals
• s + 2 p’s = 3 sp
2
9.1
9.1
• one s and a p get together to form two sp
hybrid orbitals
• s + p = 2 sp
9.1
9.1
• how about double bonds and triple bonds?
• some orbitals hybridize, some don’t
• e.g...
9.1
9.1
• what about the exceptions to the octet rule, like
PCl5?
• bring in some d’s!!!
• s + 3 p’s + d = 5 sp d (five areas = trig bipyramid)
• s + 3 p’s + 2 d’s = 6 sp d (six areas, octahedral)
3
3 2
9.1