Drawing Lewis Structures Find the total # of valence electrons from all atoms Anion: add e- (NO3- : add 1 e-) Cation: subtract e- (NH4+: minus 1 e-) Predict the arrangement of the atoms Place least electronegative atom in center Usually written first in the formula (often C, never H) Make a single bond (2 e-) between each pair of atoms Use remaining electrons to satisfy octets (8 e- around each) Place electrons around atoms in pairs (lone pairs) b. Try forming multiple bonds between atoms: double bond (4 e-) and triple bond (6 e-) Check your structure by verifying that all e- have been used and that all atoms have a satisfied octet Exceptions: Remember that H only needs 2e- for its “octet”

Lewis Structure Trends Certain trends can be used to help predict the bonding behavior of particular atoms C group Forms combination of 4 bonds (4 singles, 2 doubles, etc.) and no LP (i.e. CH4) N group Forms combo of 3 bonds and 1 LP (i.e. NH3) O group Forms combo of 2 bonds and 2 LP (i.e. H2O) F group (halogens) Forms 1 bond and has 3 LP (i.e. HCl) Note that these are NOT always true!

Lewis Structures and Formal Charge Formal charge = (# valence e-) – (# lone e- + ½ # bonded e-) or = (# valence e-) – (dots + bonds) Sum of formal charges = overall charge on molecule/ion Formal Charges Water H2O H: 1 – 1 = 0 O: 6 – (4 + 2) = 0 Total charge = 0 Hydronium H3O+ H: 1 – 1 = 0 O: 6 – (2 + 3) = +1 Total charge = +1

Resonance structures differ only in the position of the electrons Show resonance Show movement of e- The actual structure is a hybrid (average) of the resonance structures All three oxygen’s share part of the double bond In reality, there are three bonds of equal length (in between a double and single bond) Arrow formalism: curved arrows show electron movement

Predicting Molecular Shape: VSEPR (Valence Shell Electron Pair Repulsion) Electrons repel each other The molecule will adopt a 3-D shape that keeps the electrons (lone pairs and bonded e-) as far apart as possible Different arrangements of bonds/lone pairs result in different shapes Using Lewis structures, the shapes will depend on the total number of bonds/lone pairs (“things”) and how many lone pairs are around the central atom

Carbon Dioxide: CO2 Lewis Structure Two “things” (bonds or lone pairs) Linear geometry 0 LP → Linear Shape 180o Bond angle C O

Formaldehyde: CH2O Lewis Structure Three “things” Trigonal planar geometry 0 LP → Trigonal planar shape 120° bond angles

Sulfur Dioxide: SO2 Lewis Structure Three “things” B A Lewis Structure Three “things” Trigonal planar geometry 1 LP → Bent shape 120° bond angles

Methane: CH4 Lewis 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 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 3 LP → Linear Shape 180o Bond angle Cl H

Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al Si 1.8 P S Cl Br 2.8 I K 0.8 Ca Electronegativity difference ≤ 0.4  0.5 – 1.8 > 1.8 Bond type Covalent  Polar covalent Ionic Sample Bonds NaCl Cl-Cl C-O C-H Electronegativities 3.0 – 0.9 = 2.1 3.0 – 3.0 = 0 3.5 – 2.5 = 1.0 2.5 – 2.1 = 0.4 Bond Type Ionic Covalent Polar covalent

Dipole Moment Occurs in polar covalent bonds where e- are unequally shared One atom will attract more e- density than the other One atom becomes partially – charged, and the other partially + charged “+” charged end Arrow points toward “-” end δ+ δ-

Carbon Dioxide CO2 Lewis Structure Linear geometry 180o Bond angle O is more electronegative than C e- more attracted to O in each bond Creates polar bonds (unequal sharing of e-) O atoms pulling e- equally in opposite directions, so no net polarity or dipole formed Intermolecular forces: London dispersion forces

Trigonal planar geometry Formaldehyde CH2O Lewis Structure C H O Trigonal planar geometry 120° bond angles Polar C=O bond: Net dipole moment Intermolecular forces: dipole-dipole

Methane CH4 Lewis Structure Non-polar covalent bonds No net dipole Intermolecular forces: London dispersion forces Tetrahedral geometry 109o bond angles

Ammonia: NH3 Lewis Structure Trigonal pyramid shape 107o bond angles No Polar Bonds Shape is unbalanced e- around LP, creating dipole

Intermolecular forces: Water: H2O Lewis Structure Polar bonds Net dipole moment Intermolecular forces: Dipole-dipole H-bonding Bent

Intermolecular Forces Ion-ion (ionic bonds) Ion-dipole Dipole-dipole Hydrogen bonding London dispersion forces − + + − − + + −

Bonding and Electrons Each atom contains electrons in atomic orbitals Each orbital can hold a max of 2 electrons E 1s 2s 2p  Valence electrons  Core electrons Atomic orbitals are contained in energy levels (1, 2 etc.) Inner energy levels = core electrons (a.k.a kernel e-) Outermost energy level = valence electrons Valence electrons are used for bonding Ionic bonding: transfer of electrons Covalent bonding: sharing of electrons

Atomic Orbitals and Bonding Bonds between atoms are formed by electron pairs in overlapping atomic orbitals E 1s 1s 1s Example: H2 (H-H) Use 1s orbitals for bonding : Example: H2O From VSEPR: bent, 104.5° angle between H atoms Use two 2p orbitals for bonding? E 2s 2p 2p 1s How do we explain the structure predicted by VSEPR using atomic orbitals? 90°

Bonding Single bonds Double bonds Overlap of bonding orbitals on bond axis Termed “sigma” or σ bonds Double bonds Sharing of electrons between 2 p orbitals perpendicular to the bonding atoms Termed “pi” or π bonds 2p 2p Bond Axis of σ bond One π bond

Hybrid Orbitals Mix starting atomic orbitals together to form hybrid orbitals Lone pairs Bonds with H One s and three p orbitals on O mix to form four identical sp3 hybrid orbitals that point towards the corners of a tetrahedron Two sp3 orbitals contain O lone pairs Two form  bonds with H 1s orbitals

Hybrid Orbitals for Trigonal Planar Bonding  bonds One s and two p orbitals on C mix to form three identical sp2 hybrid orbitals pointing towards the corners of a triangle. One p orbital remains unhybridized. sp2 orbitals form  bonds with O and H atoms Leftover p orbital forms  bond with O

Hybrid Orbitals for Linear Bonding  bonds  bonds One s and one p orbital on C mix to form two identical sp hybrid orbitals pointing 180° away from each other Two p orbitals remain unhybridized. Each sp orbital forms a  bond with O Each leftover p orbital forms a  bond with O

Determining Hybridization To determine an atom’s hybridization, count “things” around the atom “Things” = bonded atoms and lone pairs Two things: sp hybrid orbitals Three things: sp2 hybrid orbitals Four things: sp3 hybrid orbitals Methane Acetonitrile sp sp3 4 “things” sp3 hybridization