Covalent Bonds (2 nonmetals) …atoms share e– to get a full valence shell C1s 2 2s 2 2p 2 F1s 2 2s 2 2p 5 *Both need 8 v.e – for a full outer shell (octet rule)!* o 4 valence e- 7 valence e- o x o o C x x x x x x F Bonding Review

Draw the Lewis dot structure for the following elements (write e- config first): Si O P B Ar Br 1s 2 2s 2 2p 6 3s 2 3p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 1 1s 2 2s 2 2p 6 3s 2 3p 6 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 1s 2 2s 2 2p 6 3s 2 3p 2 4 valence e- 6 valence e- 5 valence e- 3 valence e- 8 valence e- 7 valence e-

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar K Ca Se Br Kr Rb Sr Te I Xe Cs Ba Notice any trends…? TRANSITION METALS The group # corresponds to the # of valence e –

C F F F F Let’s bond two F atoms together… Each F has 7 v.e. and each needs 1 more e- F F F F Now let’s bond C and F atoms together… C F F F F carbon tetrafluoride (CF 4 ) F2F2

Lewis Structures: 2D Structures NH 3 CH 2 O CO 2 SO 2 CH 4

1.Sum the # of valence electrons from all atoms Anions: add e – (CO 3 2- : add 2 e – ) Cation: subtract e – (NH 4 + : minus 1 e – ) 2.Predict the arrangement of the atoms Usually the first element is in the center (often C, never H) 3.Make a single bond (2 e – ) between each pair of atoms 4.Arrange 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 – ) 5.Check your structure! All electrons have been used All atoms have 8e- Exceptions: Remember that H only needs 2e – ! Drawing Lewis Structures

Lewis Structure Practice CH 4 H 2 O NF 3 HBr OF 2 HCN NO 3 - CO 3 2- Draw a Lewis Structure for the following compounds: NHC

NHC

Lewis Structure Trends Here are some useful trends… C group Forms a combo of 4 bonds and no LP (Lone Pairs) e.g. CO 2 N group Forms a combo of 3 bonds and 1 LP e.g. NH 3 O group Forms a combo of 2 bonds and 2 LP e.g. CH 2 O F group (halogens) Forms 1 bond and 3 LP e.g. OF 2 Note that these are NOT always true!

Carbonite Carbonate? CO 3 2- CO 2 2-

Resonance Structures Resonance structures differ only in the position of the electrons The actual structure is a hybrid (average) of the resonance structures Technically NOT two single bonds and one double bond All 3 Oxygen atoms share the double bond 3 equal bonds (somewhere between a double and single) Arrow formalism: curved arrows show electron movement Show resonanceShow movement of e -

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

Selected Shapes and Geometries using VSEPR “Things”

Carbon Dioxide: CO 2 Two “things” (bonds or lone pairs) Linear geometry 0 LP → Linear Shape 180 o Bond angle COO Lewis Structure

C H H O Formaldehyde: CH 2 O Three “things” Trigonal planar geometry 0 LP → Trigonal planar shape 120° bond angles Lewis Structure

Sulfur Dioxide: SO 2 Three “things” Trigonal planar geometry 1 LP → Bent shape 120° bond angles Lewis Structure S O O B A A A

Methane: CH 4 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 0 LP → Tetrahedral shape o bond angles

Ammonia: NH 3 Lewis Structure Four “things” (bonds/LP) Tetrahedral geometry 1 LP → Trigonal pyramid shape 107 o bond angles

Water: H 2 O Lewis Structure 4 “things” (bonds/LP) Tetrahedral Geometry 2 LP → Bent Shape o bond angle

Hydrogen Chloride: HCl Four “things” (bonds/LP) Tetrahedral geometry 3 LP → Linear Shape No Bond angle Lewis Structure ClH

A special note… For any molecule having only two atoms… e.g. N 2, CO, O 2, Cl 2, HBr, etc. Geometry = Linear Shape = Linear Bond Angle(s)? = None It is much like geometry… what is formed by connecting two points? …a line. NN HBr Cl OO

You will need to commit these to memory! “Things”

VSEPR Practice (w/o aid of yellow sheet) CO 2 G: S: Angle: ClO 2 - G: S: Angle: NO 2 - G: S: Angle: CH 3 COO - G: S: Angle: PBr 3 G: S: Angle: AsO 4 3- G: S: Angle:

Electronegativity and Bond Type The electronegativity difference between two elements helps predict what kind of bond they will form. Bond type Covalent  Polar covalent  Ionic Definition e- are evenly shared e- are unevenly shared e- are exchanged (gained or lost) Electronegativity difference ≤ 0.4  0.5 – 1.8  > 1.8

Electronegativity difference ≤ 0.4  0.5 – 1.8  > 1.8 Practice with Bond Types Bond type Covalent  Polar covalent  Ionic 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 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 Br 2.8 I 2.5 K 0.8 Ca 1.0 Sample Bonds NaCl Cl-Cl C-O C-H Electronegativity Difference 3.0 – 0.9 = – 3.0 = – 2.5 = – 2.1 = 0.4 Bond Type? Ionic Covalent Polar covalent Covalent

Dipole Moments and Polarity Arrow points toward partially “-” end Occurs in polar covalent bonds Uneven distribution of e - Atoms become partially charged Partially “+” charged end δ-δ- δ+δ+

Polarity Examples HCN CO 2 CO 3 2- CH 2 O SO 2 CH 4 CH 3 F C 3 H 8 CO NH 3 1.Check molecule for dipole moments (polar bonds) 2.When determining overall polarity, an imbalanced structure will likely be polar (at least partially) 3.Even with polar bonds, a balanced structure is non-polar overall 4. Any structure with lone pairs on the central atom is automatically polar! Try these with your neighbors… Polar Non-polar Polar Non-polar Polar Non-polar Polar

Intermolecular Forces (IMF’s) Intramolecular Forces = bonding within a molecule e.g. ionic, covalent, polar covalent bonds Intermolecular Forces = interactions between two molecules …Intercity v. Intracity v. Innercity Intermolecular Forces are ALL weaker than Intramolecular bonds

IMF’s: Ion-Ion Force Similar Charges Repel Opposite Charges Attract Na + Cl – Na + Cl – Attractive and repulsive forces between two separate ions.

H H H IMF’s: Ion-Dipole Force Na + Cl H δ+δ+ + δ-δ- + The interaction between an ion and another molecule that has a dipole moment. (polar covalent) O Na + Lewis Structure δ-δ- δ+δ+ δ+δ+ δ+δ+ δ-δ-

IMF’s: Dipole-Dipole Force The interaction between two separate molecules, each having a dipole moment. (polar covalent) H Cl H H H HCl = Stomach Acid δ-δ- δ+δ+ δ-δ- δ+δ+

IMF’s: Hydrogen Bonding A specific type of dipole-dipole interaction between an H bond donor and an H bond acceptor. H H O H H O H bond donor: an H bonded to N, O, or F H bond acceptor: any lone pair of e –

IMF’s: London Dispersion Forces Involves an instantaneous dipole. This dipole will induce dipoles in other molecules. H H Probable? Yes Possible? Yes Probable? Yes Possible? Yes Probable? Yes Possible? Yes Probable? NO! Possible? YES! δ+δ+ δ-δ- Why instantaneous? This dipole will only remain for an instant! The electrons will quickly move to another part of the molecule! H H δ-δ- δ+δ+ H H H H H H H H H H H H Instantaneous = WEAKEST! All molecules will exhibit LDF ↑ mass,↑ LDF

IMF Review Ion-Ion Ion-Dipole Hydrogen Bonding Dipole-Dipole London Dispersion Forces (LDFs) a.k.a. van der Waals Forces STRONGEST 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 molecule

C H H O Formaldehyde: CH 2 O Trigonal planar geometry 120° bond angles Polar C=O bond = Net dipole moment IMF= Dipole-Dipole Lewis Structure IMF Practice

Methane: CH 4 Lewis Structure Tetrahedral geometry 109 o bond angles Covalent bonds = No net dipole IMF = London dispersion forces

Ammonia: NH 3 Lewis Structure Polar Bonds, Lone pairs = Dipole 1 H bond acceptor (LP), 3 H bond donors (N-H) IMF = Hydrogen Bonding Trigonal pyramid 107 o bond angles

Carbon Dioxide: CO 2 Linear geometry 180 o Bond angle COO Lewis Structure C=O bond is polar, but… Dipoles cancel = No net dipole IMF = London Dispersion Forces!

Bent Water: H 2 O Polar bonds, lone pairs = Net dipole 2 H bond acceptor (LP), 2 H bond donors (O-H) IMF = Hydrogen bonding Lewis Structure