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...
essentially everything we see around us - natural or manmade, cmpds or elements - has atoms bonded with atoms hard or soft? solid, liquid, or gas? these properties are 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???
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 electronegativity
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!
the second-row NMs want an s 2 p 6 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
lewis structures of molecules with multiple bonds what happens if you run out of electrons before everything has its 8? e.g. CO 2 : 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??? :(
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 this critter has four things sticking out of it in four symmetric directions the magic bond angle is 109.5 but what determines what 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 shell electron pair repulsion model the name says it all; we re talking about shapes of molecules due to: 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, BeCl 2 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 2 into a linear molecule
now BF (another exception to the octet rule) 3 has a central atom surrounded by three 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 4, a central atom with 4 electron pairs around it? this is real new and counterintuitive
most people when seeing the Lewis structure think it must be a cross but the electron pairs can get farther than 90! they can actually get 109.5 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 NH 3 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:
name that bad boy, remembering when you name it just worry about where the atoms are pairs are tetrahedral; but molecule is trigonal pyramid [not trigonal planar; 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]
molecular structure: molecules with double bonds vsepr works with double bonds, too! we see in experiments that CO 2 takes on a linear look it s as if the double bonds of CO 2 being repulsive)! act like single bonds (as far as
if it helps look at single and double bonds as electron clouds of repulsion single! double! the double pair are restricted to an area just like a single pair (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 XeF 4. Which is probably best??? 8.13
a summary of it all... 8.13
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What are these shapes? 8.13
SN = stearic number how many electron pairs, bonded (count single, double, and triple as one) 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 sp 3 orbitals... 9.1
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an s and two p s can get together to make three sp 2 hybrid orbitals s + 2 p s = 3 sp2 9.1
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one s and a p get together to form two sp hybrid orbitals s + p = 2 sp 9.1
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how about double bonds and triple bonds? some orbitals hybridize, some don t e.g... 9.1
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what about the exceptions to the octet rule, like PCl5? bring in some d s!!! s + 3 p s + d = 5 sp3 d (five areas = trig bipyramid) s + 3 p s + 2 d s = 6 sp3 d 2 (six areas, octahedral) 9.1