Chapter 6: Bonding… Chemical Bonding  Describe covalent, ionic and metallic bonds  Classify bond type by electronegative difference  Explain why atoms form bonds  Draw Lewis structures  Classify covalent and ionic compounds by physical properties  Predict molecular geometry using VSEPR theory

A. Vocabulary (Talkin’ the talk)  Chemical Bond  attractive force between atoms or ions that binds them together as a unit  bonds form in order to…  decrease potential energy (PE)  increase stability

A. Vocabulary CHEMICAL FORMULA Molecular Formula Unit IONICCOVALENT CO 2 NaCl (True molecules)

A. Vocabulary ION – charged atom or group of atoms Positive Charge Negative Charge SO 4 -2 Cl - Mg +2 NH 4 + CATION ANION

B. Basics of Ionic and Covalent Bonds Na Cl + - IONIC Electrons are: Transferred!

Ionic Bonding - Crystal Lattice B. Ionic and Covalent traits

Cl COVALENT Electrons are: Shared!

Covalent Bonding - True Molecules B. Ionic and Covalent traits Diatomic Molecule

B. Ionic and Covalent traits

C. Bond Polarity O Electronegativity Trend (p. 151) | Increases up and to the right.

C. Bond Polarity  Electronegativity 1.Attraction an atom has for electrons. 2.In covalent molecules: a)higher e - neg atom   - b)lower e - neg atom   +

C. Bond Polarity 3)Most bonds are a blend of ionic and covalent characteristics. 4)Difference in electronegativity determines bond type Difference in electronegativity 100% 50% 5% 0% % Ionic Character Non polar Covalent Polar Covalent Ionic

 Nonpolar Covalent Bond  e - are shared equally  symmetrical e - density  usually identical atoms C. Bond Polarity

++ --  Polar Covalent Bond  e - are shared unequally  asymmetrical e - density*  results in partial charges (dipole)

 Nonpolar  Polar  Ionic View Bonding Animations.Bonding Animations C. Bond Polarity

Elements Electronegativity difference Bond type More negative atom H & S Cs & S S & Cl C & H O & O 0.38 Polar S 1.79 Ionic S 0.58 Polar Cl 0.35 Polar C 0 Non-polar none H = 2.20 Cs = 0.79 S = 2.58 Cl = 3.16 C = 2.55 O = 3.44

balanced attraction & repulsion increased repulsion attraction vs. repulsion D. Bond Formation  Potential Energy Diagram

no interaction attraction vs. repulsion increased attraction D. Bond Formation  Potential Energy Diagram

Bond Energy D. Bond Formation Bond Length  Bond Energy  Energy required to break a bond

E. Lewis Structures  Electron Dot Diagrams  show valence e - as dots  distribute dots like arrows in an orbital diagram*  4 sides = 1 s-orbital, 3 p-orbitals* X Na Mg Al C N O F Ne

E. Lewis Structures  Octet Rule  Most atoms form bonds in order to obtain 8 valence e -  Full energy level stability ~ Noble Gases Ne

E. Lewis Structures Structural Formula Lewis Structure H C C H H What’s the difference?

E. Lewis Structures ElementValence e-BondsOctet Rule Hydrogen Group 2 Group 3 (B) Carbon Nitrogen Oxygen Halogens  For covalent bonds (rare) 6 (boron)

E. Lewis Structures  Examples Na Cl → Na Cl +1  Covalent – show sharing of e-  Ionic – show transfer of e-

E. Lewis Structures Ca Cl → Ca Cl +2

E. Lewis Structures  Nonpolar Covalent - no charges  Polar Covalent - partial charges F F → F F H O → H H O H δ-δ- δ+δ+

E. Lewis Structures Isomers – Same formula different structure. C 3 H 7 OH

IONIC COVALENT Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties e - are transferred from metal to nonmetal high yes (solution or liquid) yes e - are shared between two nonmetals low no usually not Melting Point crystal lattice true molecules F. Ionic and Covalent traits Physical State solid liquid or gas mostly odorous, Make true molecules Make ‘formula units’

G. VSEPR Theory  Valence Shell Electron Pair Repulsion Theory  Electron pairs orient themselves in order to minimize repulsive forces.

G. VSEPR Theory  Types of e - Pairs  Bonding pairs - form bonds  Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!

F. VSEPR Theory  Use a chart to determine shape

H. Common Shapes 2 Bonded to central atom 0 lone pairs LINEAR Be H H BeH 2

3 Bonded to central atom 0 lone pairs TRIGONAL PLANAR H. Common Shapes B F F F BF 3

4 bonds to central atom 0 lone pair TETRAHEDRAL H. Common Shapes C H H H H CH 4

3 bonded to central atom 1 lone pair TRIGONAL PYRAMIDAL NH 3 H. Common Shapes N H H H

2 bonded to central atom 2 lone BENT H2OH2O H. Common Shapes O H H

Metallic Bonding - “Electron Sea” I. Metallic Bonds

“electron sea” METALLIC Bond Formation Type of Structure Solubility in Water Electrical Conductivity Other Properties Melting Point I. Metallic Bonds (on pg 5) Physical State e - are delocalized among metal atoms very high yes (any form) no malleable, ductile, lustrous solid

I. Metallic Bonding  Malleability - Hammer into sheets  Ductility - Draw into thin wire

J. Intermolecular Forces...that's all well and good, but what holds different molecules together? When H 2 O is in its solid state, what is holding it like that? When CO 2 is in its solid state, what is holding it like that?

J. Intermolecular Forces Dipoles - polar covalent molecules where there is an asymmetrical electron density across the entire molecule (not just in the bonds) Dipole-dipole interaction occur when dipoles line up and attract each other Consider H 2 O, NH 3, CH 4, and CO 2, given the shape of the molecule, will one end have a different charge than the other?

J. Intermolecular Forces

K. Hydrogen bonds Hydrogen bonds - super dipole-dipole interactions between δ+ on hydrogen and δ- charges on other parts of a molecule

K. Hydrogen bonds - Ice

L. London Dispersion forces What about non-polar molecules and atoms, like the noble gases? How do they form things like liquid helium and solid oxygen without dipole interactions? London dispersion forces are the attractive forces between temporary dipoles created by random motion of electrons The more e - involved, the stronger the London dispersion forces

L. London Dispersion forces of Helium