Module #10: Covalent Bonding “A” students work

Module #10: Covalent Bonding “A” students work
(without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall 402 Office Hours W – F 2-3 pm Module #10: Covalent Bonding

G3. Scientific Knowledge is Referential G4. Watch out for Red Herrings
FITCH Rules G1: Suzuki is Success G2. Slow me down G3. Scientific Knowledge is Referential G4. Watch out for Red Herrings G5. Chemists are Lazy C1. It’s all about charge C2. Everybody wants to “be like Mike” (grp.18) C3. Size Matters C4. Still Waters Run Deep C5. Alpha Dogs eat first General Chemistry

Properties and Measurements
Property Unit Reference State Size m size of earth Volume cm3 m Weight gram mass of 1 cm3 water at specified Temp (and Pressure) Temperature oC, K boiling, freezing of water (specified Pressure) x10-24g amu (mass of 1C-12 atom)/12 quantity mole atomic mass of an element in grams Pressure atm, mm Hg earth’s atmosphere at sea level Energy: Thermal BTU 1 lb water 1 oF calorie 1 g water 1 oC Kinetic J 2kg mass moving at 1m/s Energy, of electrons energy of electron in a vacuum Electronegativity F

Bonding = sharing –electrons between repulsive + nuclei
Lewis Dot structures help us visualize sharing of electrons Octets Double and triple bonds Resonance structures and No Clean Socks Formal Charge to help distinguish between alternatives Violations of the Octet Rule 2 electrons >8 electrons Using electrons to predict the SHAPE of the molecules VESPR Effect of unpaired electrons on the central atom on molecular shape Effect of Multiple bonds How to deal with “no central atom” Bond polarity Effect of electronegativity difference between atoms in bond Effect of molecular shape How to symbolize bond polarity Discrepancies between Electron Orbital Diagrams and VESPR Introducing orbital hybridization, s, sp, sp2, sp3 Using orbital hybridization to visualize resonance or smearing out of electron density in the molecule.

Covalent bonding Patterns in abundance suggest a. periodicity
b. preferred electronic configuration of elements Leading to the Rule: “Everybody wants to be “Like Mike” a. Ions: Groups 16 and 17 gain electrons; Groups 1 and 2 lose b. Other atoms share electrons to have eight electrons = COVALENT BONDING

Covalent Bonding – getting to a noble gas electron configuration by
sharing electrons Bring two elements close together When very close the positive nuclei repel each other Repulsion of two hydrogen atoms with their Proton core + +

+ + + + Repulsion of two hydrogen atoms with their proton core e e e e
Repulsion is high close where Protons see each other Atoms which are far apart Do not even see each other There is no energy, repulsive Or attractive between the two Repulsion is low where Electrons shield nucleus, and where Electrons can be stabilized by both Positive charges + e Repulsive energy + e + e Electrons are the jelly and peanut butter between the slices of bread (protons) Attractive energy

(without solutions manual) ~ 10 problems/night.
“A” students work (without solutions manual) ~ 10 problems/night. Alanah Fitch Flanner Hall 402 Office Hours W – F 2-3 pm

Lewis dot structure (electron dot
Structure or diagram) are diagrams that Show the bonding between atoms of A molecule based on shared “valence” shell (outer shell) electrons and shows the Presence of any “lone pair” of electrons That may exist in the covalently bonded Molecule. Gilbert Newton Lewis ; Caltech Physical Chemist Covalent Valence = outermost shell electrons of an atom shared Latin: valere – to be strong

When the two hydrogen atoms are together, the electron configuration
Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? The “inner” shell electrons do not Show in this diagram

Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? Only the “outer-most” or valence shell electrons Show in this Lewis Dot Structure How many valence electrons?: = last number in group

Looks like? He When a hydrogen atom and a fluorine atom share electrons, the Electron configuration on fluorine looks like? The shared pair of Electrons = covalent bond The unshared pairs of electrons are “regions of high Charge density”

When a hydrogen atom and an oxygen atom share valence electrons plus an
Extra electron, the electron configuration on hydrogen and oxygen look like? Valence electrons on oxygen? Valence electrons on hydrogen? Invoking Rule: Chemists are Lazy the diagram above is too tedious to write out all the time make shared electrons (bond) a line Lewis dot structure for hydroxide The single electron pair shaired between the two bonded atoms Is called a single bond It is drawn as a line.

When two hydrogen atoms and an oxygen atom share valence electrons, the
electron configuration on hydrogen and oxygen look like? Two shared electron pairs = Two single bonds Two shared electron pairs

electron configuration on hydrogen and oxygen look like? When three hydrogen atoms and a nitrogen atom share valence electrons, the electron configuration on hydrogen and nitrogen look like?

electron configuration on hydrogen and oxygen look like? When three hydrogen atoms and a nitrogen atom share valence electrons, the electron configuration on hydrogen and nitrogen look like? Valence shell of nitrogen? Three pairs of shared electrons = three single bonds

When four hydrogen atoms and two carbon atoms share valence electrons, the
electron configuration on hydrogen and carbon look like? Valence shell of carbon? Two electron pairs shared is a Double bond

When two hydrogen atoms and two carbon atoms share valence electrons, the
electron configuration on hydrogen and carbon look like? Three electron pairs shared is a Triple bond

Rules for Writing Lewis Dot Structures
Count the number of valence electrons (last number of group) of all atoms a. For an anion add the appropriate extra number of electrons b. For a cation subtract the appropriate extra number of electrons Draw a molecular skeleton, joining by single bonds to the central atom. a. The central is usually the atom written first in the formula (N in NH4+, S in SO2, and C in CCl4). b. The terminal atoms are usually H, O. c. Halogens are always terminal atoms. Determine the number of valence electrons still available for distribution after subtracting two electrons for each single bond. Determine the number of electrons required to complete the octet a. H gets only two electrons b. Other exceptions to be noted below 5. Fill in the region required for the octet. Make up deficit of electrons by creating double bonds a. C, N, O, S H can only have one bond because it can share only one Electron. Poor H. Halogens have lots of electrons but really do not like to share. Greedy halogens All they want is one more to make up the Mike configuration

Draw Lewis structures of
Hypochlorite ion Methyl alcohol N2 SO2 Valence shell electrons? O 6 +Cl 7 +Negative charge 1 Total electrons 14 -1Single bond -2 12 -2(6 electrons for O,Cl) 12 remaining 0 Hypochlorite? Hypo – smallest number of oxygens OCl- Skeleton

Hypochlorite ion Methyl alcohol, CH3OH N2 SO2 Skeleton Carbon is first in formula Hydrogen is always terminal Valence shell electrons? O 6 +C 4 +4(H) 4 Negative charge 0 Total electrons 14 -5single bonds -10 remaining 4 -octet for oxygen -4 remaining 0 Octets Carbon has its octet Hydrogen has its duet Oxygen requires 4 more electrons

Hypochlorite ion Methyl alcohol, CH3OH N2 SO2 Skeleton Octets Each nitrogen requires 6 more Valence shell electrons? 2N 10 Negative charge 0 Total electrons 10 -1single bond -2 Remaining 8 Octet completion -12 Difference -4 -3 single bonds -6 Remaining 4 We are short 4 electrons for the octet, The only way to get extra ones is to Share four more electrons = triple Bond. Place the remaining 4 electrons equally On the two equal nitrogens

Skeleton, First atom in formula is central
Draw Lewis structures of Hypochlorite ion Methyl alcohol, CH3OH N2 SO2 Octets We are short 2 electrons for the octet, The only way to get extra ones is to Share two more electrons = double bond. Valence shell electrons? 2O 12 +1S 6 Negative charge 0 Total electrons 18 -2(single bonds) -4 Remaining electrons 14 Octet for S -4 2(Octet for each O) -12 Deficit? -2 -3(single bonds) -6 12 Place the remaining 12 electrons to fill octets

We got this Lewis dot structure
No reason not to write instead Which would lead to Is there any reason for us to Presume one of these is correct And not the other? No Grammar: double-headed arrow is used to separate resonance structures

Remember our Low Charge Density; Spectator Polyatomic Anions?
No Clean Socks NO3- N 5 3(O) 18 Charge 1 Total 24 -single bonds -6 Remaining 18 Octets (6x3 O +2 for N) -20 Deficit of 1 electron pair -2 Bonds -8 16 Charge is distributed over All three of the resonance Forms = one big fat marshmallow

Resonance The “real” molecule is none of the three nitrates we drew but something intermediate to the three. Resonance can be “assumed” or “predicted” when there are equally plausible Lewis dot structures. Resonance forms differ only in the distribution of electrons and not in the arrangement of atoms.

Write three resonance forms for SO3
Valence electrons 4(6) 24 Sulfur central atom Three single bonds to the sulfur -3(2) -6 Remaining electrons 18 2 electrons to complete S octet -2 3(6) electrons to complete O octets = Deficit of two electrons = double bond -2 Four single bonds to the sulfur -3(2) -8 Remaining electrons 16

Formal Charge helps determine the correct Lewis Dot Structure
X=number of valence e- in the free atom (last number of group) Y = number of unshared e- owned by the atom in the Lewis structure Z = number of bonding e- shared by the atom in the Lewis structure The correct Lewis dot structure is generally the one in which The formal charges are as close to zero as possible Negative charge is located on the more electronegative atom

X=number of valence e- in the free atom (last number of group)
Y = number of unshared e- owned by the atom in the Lewis structure Z = number of bonding e- shared by the atom in the Lewis structure Which is correct?

The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas
In same fashion: not everybody can share enough Electrons to make up a perfect octet These guys will have 1, 2, and 3 bonds only

The unpaired electron means The compound is paramagnetic
The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas In same fashion: not everybody can share enough Electrons to make up a perfect octet This guy may end up “holding the bag” Having an unpaired electron Because he is Not really Strong enough To always Get the lion’s Share of the Electrons in A covalent bond, Particularly With oxygen The unpaired electron means The compound is paramagnetic The presence of these unpaired electrons on these gases Gives rise to the many atmospheric reactions involved In ozone destruction and formation of smog. Para = paramour = love = similar orientation Dia – diatribe = against = opposite orientation

A singlet electron is also called a “free radical”
Aging Hydroxy radical hydroxide Cellular membrane damage

After hydroxyl radical cleavage and denaturing gel electrophoresis, the gel patterns differ. The unfolded RNA molecule is cleaved uniformly, giving rise to a homogeneous ladder of bands on the gel (left, bottom). The gel pattern for the folded RNA (right, bottom), in contrast, shows several region where strand cleavage is inhibited, corresponding to the sites of low solvent accessibility in the folded structure (right, top). TD Tullius, JA Greenbaum, Curr Opin Chem Biol 2005, 9:127–134 (Figure 1) Commentary from grove.ufl.edu/~dmorgan/Articles/DER/Journal%20Club.ppt

The Octet Rule is a form of Rule #C2- Everybody wants to be like a Noble gas
Some guys can take on more electrons because they Make use of their d orbitals 3p these guys Have d orbitals That allow them To have more Than 8 electrons 4s 3d

Draw the Lewis structure of XeF4
8+4(7) =36 electrons 4bonds = 8 electrons Remainder = 28 electrons Octets: 4(6) for F = 24 Remainder to Xe = 4

Draw the Lewis structure of XeF2
8+2(7) =22 electrons 2bonds = 4 electrons Remainder = 18 electrons Octets: 2(6) for F = 12 Remainder to Xe = 6

Can now predict SHAPE of molecule
determines: how two molecules orient themselves for reaction together. Can they dock? And actually do work together in three dimensional space?

Valence Shell Electron Repulsion Pair
Rule C1 = It’s all about charge Eel d Valence electron pairs surrounding an atom repel one another. Consequently, the orbitals containing those electron pairs are oriented to be as far apart as possible

Geometries of AX2-AX6 molecules
Repulsion of valence shell electrons in Bonds and on F push F apart to a 180o orientation

What orientation would put electrons as far apart as possible? 120o degrees apart in a “circle”

CH4 Valence shell electrons? +C 4 +4(H) 4 Negative charge 0 Total electrons 8 -4single bonds -8 remaining 0 Skeleton Carbon is first in formula Hydrogen is always terminal Octets Carbon has its octet Hydrogen has its duet

X X X A X How can four bonds be organized in 3-D space to
be farthest apart? X X X A X

A pyramid is a space figure with a square base and 4 triangle-shaped sides.
(5 “faces”) 5 2 3 4 1 A tetrahedron is a space figure and 4 triangle shaped faces. 2 3 4 Dictionary: a four-sided solid; a Triangular pyramid 1

(5 “faces”) 2 3 4 Dictionary: Square base and sloping Sides rising to an apex 5 1 A tetrahedron is a space figure and 4 triangle shaped faces. 2 3 Dictionary: a four-sided solid; a Triangular pyramid 4 1

(5 “faces”) Dictionary 1: Square base and sloping Sides rising to an apex 2 5 3 4 Dictionary 2: A solid figure with a polygon base. The surface, or lateral faces, are triangles having a common vertex. In a regular pyramid the base is a regular polygon and the lateral faces are congruent triangles 1 A tetrahedron is a space figure and 4 triangle shaped faces. 2 3 4 Dictionary: a four-sided solid; a Triangular pyramid SIGNIFICANT AMBIGUITY In nomenclature!!!! 1

Why tetrahedron and not this orientation?
(Square planar) Calculate the repulsion experience By atom X1 with charge -1 X1 ,1- 90 o 180 o X4,1- A X2,1- Assume Bond distance=1 X3,1- Since our example has all Q the same

Calculation suggests that the electrostatic charge repulsion
Why tetrahedron and not this orientation? X1,1- 90 o d14 d12 1 X4,1- A X2,1- d13 X3,1- Calculation suggests that the electrostatic charge repulsion Energy is proportional to for a “square planar” orientation Of four identically charged atoms

Square planar orientation= 1.914
Compare to tetrahedron A X1 X4 60o 1 d14 A 4 2 3 This means: atom 1 experiences less charge Repulsion from 2, 3, and 4 when tetrahedral

Triangular bipyramid Triangular pyramid A=Central atom X= atoms How can five bonds be arranged in space to be as far apart as possible?

octahedron SF6 How can these six guys best position themselves away from each other?

AX2E Some of the molecules we constructed using Lewis Dot structures
had UNSHARED PAIRS of electrons on the CENTRAL ATOM What effect will this have on the geometry?. AX2E Unshared electron pairs orient themselves pretty much the same as single bonds. The observed molecular geometry (invisible electrons) is very different

AX2E This geometry, with respect to electron pairs and bonds,
is triangular planar (three guys trying to get out of each others way) But one of the “terminal atoms” is missing so the molecular geometry differs from triangular planar <120 Actual degrees observed is slightly less than 120o because unshared electron pair expands Molecular Geometry is “bent”

AX3 AX2E 120 o Bent Molecular Geometry Triangular planar Triangular

AX3E This geometry, with respect to electron pairs and bonds,
is tetrahedral (four guys trying to get out of each others way) But one of the “terminal atoms” is missing so the molecular geometry differs from tetrahedral Triangular Pyramid Explains why amines like ammonia can “steal” a proton From water – high charge density from the lone pair

Effect of lone pairs on substituents (non-central atoms)

Effect of F is NOT by geometry of it’s lone pairs
BUT By it’s electronegativity which pulls electrons along the bond, lowers Density of electrons in the bondings area Allows N lone pair to expand Compressing the angle between F atoms F F 107.2o F 102.3o

AX2E2 What happens when we have Two Unbonded electron pairs on the
Central atom, A? AX2E2 This geometry, with respect to electron pairs and bonds, is tetrahedral (four guys trying to get out of each others way) But two of the “terminal atoms” are missing so the molecular geometry differs from tetrahedral This shape is “bent”

AX2E AX2E2 Both are bent, but the angle is different.
Depends upon the number of valence shell electron pairs AX2E AX2E2

AX4 AX3E AX2E2 Molecular Geometry Tetrahedron Triangular pyramid
Bent AX2E2

XeF2 ClF3 Triangular bipyramid Triangular pyramid Triangular bipyramid
5 ELECTRON PAIRS Molecular Orientation Triangular bipyramid Triangular pyramid Triangular bipyramid Triangular bipyramid 180o Linear AX2E3 XeF2 Triangular bipyramid AX3E2 90o 180o ClF3 T-shape Why put the E at equator?

1 2 Electrostatic repulsion is sum of all near neighbor repulsions
Comparing where the non-bonded electron pair will go Variations: Axial versus Equatorial orientation equatorial axial 90o S E S S 120o E S E S S E 1 2 E = non-bonding Electron pair S = Shared (bonded) electron pair Electrostatic repulsion is sum of all near neighbor repulsions

If we put the E at the axial orientation they
minimize E-E repulsion; increase E-S repulsion If we put E at the equatorial orientation E-E repulsion exists, but we decrease the E-S repulsion E E Minimizes impact Of E on S

XeF2 ClF3 Triangular bipyramid Triangular pyramid Triangular bipyramid
5 ELECTRON PAIRS Molecular Orientation Triangular bipyramid Triangular pyramid Triangular bipyramid 180o Linear AX2E3 XeF2 Triangular bipyramid AX3E2 90o 180o ClF3 T-shape

SF6 XeF2 AX4E2 90o 180o Square planar octahedron octahedron octahedron
Molecular Orientation octahedron SF6 octahedron Square pyramidal octahedron AX4E2 90o 180o Square planar XeF2

Anticancer Drug Why electron configuration is important:
controls shape of molecule dictates 3D interaction of molecules Anticancer Drug Square planar lets it slide into the DNA grove

MULTIPLE Bonds Has no effect on geometry Multiple bond acts as a single bond Compare BF3 and SO3 Triangular Planar

Compare Molecular Geometries
for BeF2 and CO2 1. We already did BeF2

for BeF2 and CO2 1. We already did BeF2 Skeleton Valence shell electrons for CO2? Carbon is first in formula= central atom +C 4 +2(O) 4 Negative charge 0 Total electrons 16 -2single bonds -4 remaining 12 e required for octets -16 deficit = multiple bonds -4 -4bonds -8 Remaining 8 Octet for C = 0 Octet for each O = 4

for BeF2 and CO2 Both are linear

AX3 Figure out geometry with NO CENTRAL ATOM
Consider each carbon separately AX3 Geometry around the carbon = Triangular Planar

AX2 Figure out geometry with NO CENTRAL ATOM
Consider each carbon separately AX2 Geometry around the carbon =Linear

Tetrahedral Trigonal Tetrahedral Planar
Number of Electron Domains Electron-domain geometry Predicted Bond Angles Tetrahedral Trigonal Tetrahedral Planar 109.5o 120o o

Bond Polarity Polarity – distribution of electrons in the bond Depends upon the difference in electronegativity of bonded atoms If two atoms in the bond are identical =nonpolar Otherwise all bonds are polar

Decreasing size H–H H–C H–F E.N. = 2.2-2.2=0 nonpolar
Electronegativity is a measure of Positive Nuclear charge density experienced by Electron on another atom , scaled to a maxium of 4 H–H H–C H–F E.N. = =0 nonpolar E.N. = = 0.3 slightly polar E.N. = 4-2.2=1.8 strongly polar

Molecular Polarity Depends on 1. bond polarity 2. molecular shape
a. diatomic molecules are linear molecule polar if atoms differ b. Polyatomic molecules can have polar bonds and still be non-polar H–Cl Cl–Cl polar nonpolar Polar molecules line up in an electric field

CCl4 Non-polar molecule Polar molecule
Arrow indicates the direction in which electrons are biased - the negative pole CCl4 F  Be  F Bonds are polar E.N. = 4-1.6=2.4 Bonds are polar E.N. = =1.3 Vectors cancel each other Non-polar molecule Net charge direction Polar molecule Bonds are polar E.N. = =0.7 No net charge direction Non polar molecule

Very polar bonds (F=4; Be=1.6)
Non-polar polar Changing Central atom so That a lone Electron pair Bends the molecule Means Vectors Don’t cancel Changing One atom Means Vectors Don’t cancel polar Non polar Very polar bonds (Cl = 3.2; C=2.5)

shell electrons pushes F apart to a 180o orientation VSEPR AX2 Linear
The “fly in the ointment” Repulsion of valence shell electrons pushes F apart to a 180o orientation VSEPR AX2 Linear VSEPR model suggests that once Be bonds to F the orbitals are “equivalent” and therefore are equidistant from each other. But the electron orbital diagram suggests otherwise that Be has paired electrons and would not make bonds at all.

Be has no unpaired electrons available for bonding
Orbital diagram of isolated atoms F and Be Be promotes 1 “2s” electron to a “2p” orbital = “sp” hybridization Orbital diagram of isolated F and hybridized Be atoms 4. The new “sp” electrons engage in bonding with the unpaired electrons on F sp Shared electrons Orbital diagram of F and Be in BeF2

Formation of Hybrid Atomic Orbitals
S + p orbitals = 2 “sp” hybridized orbitals = The number of hybridized orbitals formed = number of atomic orbitals mixed Energies of hybridized orbitals intermediate to the atomic orbitals mixed

+ What does the bond made from the atomic “sp” and
“p” atomic orbitals look like? sp + Sigma bond Single bond

VSEPR AX3 Triangular Planar sp2
Orbital diagram of isolated atoms F and B 1. Boron needs to have three of its electrons shared with the three F electrons to create the three single bonds. B mixes 2 “s2” and 1 “2p” electrons using 1 s and 2 “p” orbitals 3 sp2 atomic orbitals with 1 e each formed 4. Electrons in these orbitals are shared with unpaired electrons on F to create single bonds (blue) sp2

What does an sp2 orbital look like?

C needs four energetically equivalent bonds
VSEPR AX4 Tetrahedral C needs four energetically equivalent bonds C mixes 1 “2s” and 3 “2p” atomic orbitals 4 sp3 atomic orbitals formed sp3 orbitals used to create bonds Orbital diagram of isolated atoms H and C sp3

sp3 orbital on C + s orbital on H Form a sigma bond

So far we have considered
sp VSEPR AX2 Linear 2 orbitals VSEPR AX3 Triangular Planar sp2 3 orbitals VSEPR AX4 Tetrahedral sp3 4 orbitals What about the guys with expanded octets?

VSEPR AX5 Triangular bipyramid these guys 3p Have d orbitals
That allow them To have more Than 8 electrons 4s 3d

VSEPR AX5 triangular bipyramid sp3d

VSEPR Nothing new here – same as we got with VSEPR AX2 AX3 AX4 AX5 AX6
Note to myself from 2006 – F does not hybridize – has one unpaired electron Forming one bond

Hybridize to Create Equivalent Orbitals on One atom for Bonding with another atom Hybridized Orbitals From two diff atoms combine to create Bonding orbitals Atomic Orbitals 0kJ 180o 180o -kJ

Double and Triple Bonds
Whenever we have a “single” bond we can assume that it has the sigma shape, resulting from hybridization between atomic orbitals Sigma bond Single bond For double and triple bonds, we do not need to create more equivalent bonds which can be moved as far apart as predicted by Valence Shell Electron Pair Repulsion. We need to simply create additional bonds within the shape predicted by VSEPR Pi bond Double bond around single bond

AX3 What does ethylene, C2H4 look Like? Geometry around
the carbon = Triangular Planar

AX2 What does acetylene, C2H2, look like?
Geometry around the carbon =Linear

Example problem: What is the hybridization of nitrogen in
a) NO3-, b) NH2Cl, c) N2, d) and N2O? Make the Lewis dot structure Count sigma bonds on central atom Count the unbonded electron pairs on central atom Sum the 2 (n) Number of hybridized orbitals =n Cmpd σ Bonds Unbonded electron pairs on N n Hybridizationon N NO3- 3 sp2 NH2Cl 1 4 sp3 N2 2 sp Cmpd σ Bonds Unbonded electron pairs on N n Hybridizationon N NO3- 3 sp2 NH2Cl 1 4 sp3 N2 2 sp N2O a N2O b Cmpd σ Bonds Unbonded electron pairs on N n Hybridizationon N NO3- 3 sp2 NH2Cl 1 4 sp3 Cmpd σ Bonds Unbonded electron pairs on N n Hybridizationon N NO3- 3 sp2 NH2Cl 1 4 sp3 N2 2 sp N2O a N2O b Cmpd σ Bonds Unbonded electron pairs on N n NO3- 3 Cmpd σ Bonds NO3- 3 Cmpd σ Bonds Unbonded electron pairs on N NO3- 3 Cmpd σ Bonds Unbonded electron pairs on N n Hybridizationon N NO3- 3 sp2

RESONANCE

Benzene, C6H6, is a very common compound
Usually don’t show Hydrogens or Carbons P orbitals Sigma bonds Benzene resonance Pi (double) bonds Charge delocalization

Nitrate: no clean socks
BIG KEY POINT!!!!!! Delocalization/Resonance Structures Nitrate: no clean socks Nitrate is a is not charge dense Poor Nitrate 1 charge Low Charge density Large radius

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