Lewis Structures

Unit 8: Covalent Bonding

1Covalent Bonds (2 nonmetals)atoms share e to get a full valence shellC1s2 2s2 2p2F1s2 2s2 2p5

*Both need 8 v.e for a full outer shell (octet rule)!*o4 valence e-7 valence e-oxooCxxxxxxFBonding Review2Draw the Lewis dot structure for the following elements (write e- config first):SiOPBArBr

1s2 2s2 2p6 3s2 3p31s2 2s2 2p41s2 2s2 2p11s2 2s2 2p6 3s2 3p61s2 2s2 2p6 3s2 3p6 4s2 3d10 4p51s2 2s2 2p6 3s2 3p2 4 valence e-6 valence e-5 valence e-3 valence e-8 valence e-7 valence e-1 2 3 4 5 6 7 8 HHeLi Be B C N O F NeNa Mg Al Si P S Cl ArK Ca Se Br KrRb Sr Te I XeCs Ba Notice any trends?TRANSITION METALSThe group # corresponds to the # of valence e 4CFFFFLets bond two F atoms togetherEach F has 7 v. e and each needs 1 more e FFFFNow lets bond C and F atoms together

CFFFFcarbon tetrafluoride (CF4)F2Lewis Structures: 2D Structures

NH3CH2OCO2SO2CH4

6Sum the # of valence electrons from all atomsAnions: add e (CO32- : add 2 e ) Cation: subtract e (NH4+: minus 1 e )Predict the arrangement of the atoms Usually the first element is in the center (often C, never H)Make a single bond (2 e) between each pair of atomsArrange remaining e to satisfy octets (8 e around each) Place electrons in pairs (lone pairs) Too few? Form multiple bonds between atoms: double bond (4 e ) and triple bond (6 e )Check your structure! All electrons have been used All atoms have 8e-

Exceptions: Remember that H only needs 2e !

Drawing Lewis Structures7Lewis Structure PracticeCH4

H2O

NF3

HBr

OF2

HCN

NO3-

CO32-

Draw a Lewis Structure for the following compounds:

NHC

NHC

Lewis Structure TrendsHere are some useful trendsC groupForms a combo of 4 bonds and no LP (Lone Pairs) i.e. CO2N groupForms a combo of 3 bonds and 1 LP i.e. NH3O groupForms a combo of 2 bonds and 2 LP i.e. CH2OF group (halogens)Forms 1 bond and 3 LP i.e. OF2Note that these are NOT always true!

10CarboniteCarbonate?CO32-CO22-

Resonance StructuresResonance structures differ only in the position of the electronsThe actual structure is a hybrid (average) of the resonance structuresTechnically NOT two single bonds and one double bondAll 3 Oxygen atoms share the double bond3 equal bonds (somewhere between a double and single)Arrow formalism: curved arrows show electron movement

Show resonanceShow movement of e-12Lewis Structures and Formal ChargeFor each atom in a moleculeFormal charge= # valence e # lone e # bonds or = # valence e dots linesSum of formal charges = overall charge on molecule/ionWaterH2O

HydroniumH3O+Formal ChargesH: 1 - 1 = 0O: 6 - 4 - 2 = 0Total charge = 0H: 1 - 1 = 0O: 6 - 2 - 3 = +1Total charge = +1

13Predicting Molecular Shape: VSEPR(Valence Shell Electron Pair Repulsion)Electrons repel each otherThe molecule adopts a 3-D shape to keep the electrons (lone pairs and bonded e-) as far apart as possibleDifferent arrangements of bonds/lone pairs result in different shapesShapes depend on # of bonds/lone pairs (things) and LP around the central atom

Selected Shapes and Geometries using VSEPRThings15Carbon Dioxide: CO2

Two things (bonds or lone pairs) Linear geometry 0 LP Linear Shape 180o Bond angleCOOLewis Structure

CHHOFormaldehyde: CH2O Three things Trigonal planar geometry 0 LP Trigonal planar shape 120 bond anglesLewis Structure

Sulfur Dioxide: SO2 Three things Trigonal planar geometry 1 LP Bent shape 120 bond anglesLewis StructureSOO

BAAA

Methane: CH4Lewis Structure

Four things (bonds/LP) Tetrahedral geometry 0 LP Tetrahedral shape 109.5o bond angles

Ammonia: NH3

Lewis Structure Four things (bonds/LP) Tetrahedral geometry 1 LP Trigonal pyramid shape 107o bond angles

Water: H2O

Lewis Structure 4 things (bonds/LP) Tetrahedral Geometry 2 LP Bent Shape 104.5o bond angle

Hydrogen Chloride: HCl

Four things (bonds/LP) Tetrahedral geometry 3 LP Linear Shape No Bond angleLewis StructureClH

ClA special noteFor any molecule having only two atoms e.g. N2, CO, O2, Cl2, HBr, etc.

Geometry = Linear Shape = Linear Bond Angle(s)? = None It is much like geometrywhat is formed by connecting two points?a line.

NNHBrClClOO

You will need to commit these to memory!Things24VSEPR Practice (w/o aid of yellow sheet)CO2G:S:Angle:

ClO2-G:S:Angle:

NO2-G:S:Angle:

CH3COO-G:S:Angle:

PBr3G:S:Angle:

AsO43-G:S:Angle:

Electronegativity and Bond TypeThe electronegativity difference between two elements helps predict what kind of bond they will form.Bond type

CovalentPolar covalentIonicDefinitione- are evenly sharede- are unevenly sharede- are exchanged (gained or lost)

Electronegativity difference

0.40.5 1.8> 1.826Electronegativity difference

0.4

0.5 1.8> 1.8Practice with Bond TypesBond type

CovalentPolar covalentIonicH2.1Li1.0Be1.5B2.0C2.5N3.0O3.5F4.0Na0.9Mg1.2Al1.5Si1.8P2.1S2.5Cl3.0Br2.8I2.5K0.8Ca1.0Sample BondsNaClCl-ClC-OC-HElectronegativity Difference3.0 0.9 = 2.13.0 3.0 = 03.5 2.5 = 1.02.5 2.1 = 0.4Bond Type?IonicCovalentPolar covalent Covalent27Dipole Moments and PolarityArrow points toward partially - end Occurs in polar covalent bonds Uneven distribution of e- Atoms become partially chargedPartially + chargedend

-

+

28Polarity ExamplesHCNCO2CO32-CH2OSO2

CH4CH3FC3H8CONH3Check molecule for dipole moments (polar bonds)When determining overall polarity, an imbalanced structure will likely be polar (at least partially)Even with polar bonds, a balanced structure is non-polar overall Any structure with lone pairs on the central atom is automatically polar!Try these with your neighborsPolarNon-polarNon-polarPolarPolarNon-polarPolarNon-polarPolarPolarIntermolecular ForcesIon-ion (ionic bonds)Ion-dipoleDipole-dipole (and H-Bond)London dispersion forces

+ ++ +

+30Intermolecular Forces (IMFs)Intramolecular Forces = bonding within a moleculee.g. ionic, covalent, polar covalent bondsIntermolecular Forces = interactions between two moleculesIntercity v. Intracity v. Innercity

Intermolecular Forces are ALL weaker than Intramolecular bonds31IMFs: Ion-Ion Force+-++Similar Charges RepelOpposite Charges AttractNa+ClNa+Na+Cl ClAttractive and repulsive forces between two separate ions.32HHHIMFs: Ion-Dipole ForceNa+ClHCl++-

+The interaction between an ion and another molecule that has a dipole moment. (polar covalent)ONa+

Lewis Structure-

+

+

+

-

33IMFs: Dipole-Dipole Force

The interaction between two separate molecules, each having a dipole moment. (polar covalent)HClHClHClHCl

HCl = Stomach Acid-

+

-

+

34IMFs: Hydrogen BondingA specific type of dipole-dipole interaction between an H bond donor and an H bond acceptor.

HHOHHO

H bond donor: an H bonded to N, O, or FH bond acceptor: any lone pair of e35IMFs: London Dispersion ForcesInvolves an instantaneous dipole. This dipole will induce dipoles in other molecules.HHProbable? YesPossible? YesProbable? YesPossible? YesProbable? YesPossible? YesProbable? NO!Possible? YES!+-Why instantaneous?This dipole will only remain for an instant! The electrons will quickly move to another part of the molecule!HH-+HHHHHHHHHHHHInstantaneous = WEAKEST!All molecules will exhibit LDF mass, LDF36IMF ReviewIon-IonIon-DipoleHydrogen BondingDipole-DipoleLondon Dispersion Forces (LDFs)a.k.a. van der Waals ForcesSTRONGEST Weakest *Remember: These are all weaker than actual bonds (ionic, covalent, etc.). These are just attractions.Involves an ion (+ and charged)Involves a dipole (polar molecule)Involves a non-polar molecule37

CHHOFormaldehyde: CH2OTrigonal planar geometry120 bond anglesPolar C=O bond = Net dipole momentIMF= Dipole-DipoleLewis Structure

IMF Practice38

Methane: CH4Lewis StructureTetrahedral geometry109o bond angles

Covalent bonds = No net dipoleIMF = London dispersion forces

39

Ammonia: NH3

Lewis StructurePolar Bonds, Lone pairs = Dipole1 H bond acceptor (LP), 3 H bond donors (N-H)IMF = Hydrogen Bonding

Trigonal pyramid107o bond angles

40Carbon Dioxide: CO2

Linear geometry 180o Bond angleCOOLewis StructureC=O bond is polar, butDipoles cancel = No net dipoleIMF = London Dispersion Forces!

41

BentWater: H2OPolar bonds, lone pairs = Net dipole2 H bond acceptor (LP), 2 H bond donors (O-H) IMF = Hydrogen bonding

Lewis Structure

42Orbital ReviewEach atom contains electrons in atomic orbitalsEach orbital can hold a max of 2 electronsOrbitals are in energy levelsInner energy levels = core e- (kernel e-)Outermost energy level = valence e-Valence e- are used for bondingIonic bonding: transfer of those e-Covalent bonding: sharing of those e-1s2s2pE Core e- Valence e-

43Orbitals and BondingCovalent bonds are formed by sharing two e- Sharing occurs in overlapping atomic orbitals1s1s:E1sExample: H2 (H-H)Use 1s orbitals for bondingExample: H2OFrom VSEPR: bent, 104.5 angle between H atomsUse two 2p orbitals for bonding?E2s2p90How can we explain this bonding and match VSEPR?

2p2p1s1s44BondingSingle bonds (one sigma bond) Overlap of s orbitals on bond axisDouble bonds (one sigma + one pi bond) Sharing of electrons between p orbitals perpendicular to the bonding atomsTermed pi or bonds

2p

2pBond Axis of bondOne bondO=O or O245Hybrid Orbitals2H: 1sH2C: 2s, 2pC: two spspC: 2s, 2px, 2pyC: three sp2sp2C: 2s, 2px, 2py, 2pzC: four sp3sp3LINEARTRIGONAL PLANARTETRAHEDRAL++++====46

Hybrid Orbitals: Tetrahedral1 s and 3 p orbitals on O mix to form 4 identical sp3 hybrid orbitals (tetrahedron)Two sp3 orbitals contain O lone pairsTwo form bonds with H 1s orbitalsMix starting atomic orbitals together to form hybrid orbitalsLone pairsBondswith HHOH

Hybrid Orbitals: Trigonal Planar1 s and 2 p orbitals on C mix to form 3 identical sp2 hybrid orbitals (trigonal planar).One p orbital remains unhybridized.sp2 orbitals form bonds with O and H atomsLeftover p orbital forms bond with O bonds bond

Hybrid Orbitals: Linear Bonding

1 s and 1 p orbital on C mix to form 2 identical sp hybrid orbitals (linear)Two p orbitals remain unhybridized.Each sp orbital forms a bond with OEach leftover p orbital forms a bond with O bonds bonds

Determining HybridizationTo determine an atoms hybridization, count things around the atomThings = bonded atoms and lone pairsTwo things: sp hybrid orbitalsThree things: sp2 hybrid orbitalsFour things: sp3 hybrid orbitals

4 thingssp3 hybridization

MethaneAcetonitrilespsp3