Bonding Chapters 8 & 9

Bonds Ionic Covalent Metallic Electrostatic attraction between oppositely charged ions Covalent Shared valence electrons Metallic Pooled electrons around a positive core

Octet Rule Atoms will gain, lose, or share electrons to achieve a full valence shell First energy level is limited to 2 electrons d & f sublevels allow for more than 8

Ionic Bonding Electrostatic attraction due to oppositely charged ions Lattice Energy Energy released when gaseous ions form 1 mole of ionic solid Usually reported as exothermic process Energy required to separate 1 mole of ionic solid into gaseous ions

Lattice Energy Depends directly on magnitude of charges and inversely on size of ions Larger charges means more coulombic attraction Larger ion size means less coulombic attraction

Ionic Compounds Metal and nonmetal High Melting and Boiling Points Hard Crystalline solids Brittle Break when struck (Cleave) Soluble in water Conduct electricity when melted or dissolved Do not conduct as solids

Covalent Bonding Sharing of valence electrons due to equal or near equal electronegativity of 2 nonmetals

Bond Enthalpy Amount of energy required to break a bond Amount of energy released when a bond forms Larger energy = stronger bond ΔH = Σreactants - Σproducts

Bond Polarity Bond type is based on difference in electronegativity between bonding atoms Larger difference = more polar Bond polarity is a continuous spectrum Nonpolar Covalent Polar Covalent Ionic ΔEN

Dipole Moment Measure of Bond Polarity

Molecular Polarity

Molecular Polarity

F Lewis Structures Electron Dot Diagrams Rules for individual atoms Shows valence electrons only Rules for individual atoms Start on any side First two get paired together Next three are separated Fill in as needed F

F- Na+ Lewis Structures Monatomic Ions Positive ions tend to have 0 dots Lost all valence electrons Negative ions tend to have 8 dots Gained enough to fill valence shell Can have brackets around ion with charge outside brackets Na+ F-

Lewis Structures Ionic Compounds NaCl Cl [ ] - Na+

Lewis Structures Covalent Compounds Line for bonding electrons 1 line = single bond 2 lines = double bond 3 lines = triple bond Dots for nonbonding electrons

Lewis Structures Determine number valence electrons Arrange atoms around a central atom with single bonds Usually least electronegative element excluding H Fill in octet for outer atoms Fill in octet for central atom If not enough to fill central atom’s octet, make multiple bonds to central atom

Example CO2 16 valence electrons O C O

Lewis Structures Polyatomic Ions O N O O

Lewis Structures Exceptions to Octet Rule More than octet Less than octet Odd number of electrons

Lewis Structures SF6 F F F S F F F

Lewis Structures BF3 F B F F

Lewis Structures NO N O

Formal Charge Calculated quantity used to determine the validity of Lewis structures For each atom, # of valence electrons – (nonbonding e- + ½ bonding e-) C N H

Formal Charge Best Lewis Structures Have fewest and/or smallest charges Negative charge on most electronegative element

Formal Charge Example -1 S O H +2 S O H -1

Lewis Structures Benzene, C6H6 H C

Lewis Structures Benzene, C6H6 H H C C H C C H C C H H

Resonance Two or more equally valid, but different Lewis Structures Differ only in the position of the electrons Use double headed arrow between resonance structures

Resonance for Benzene Either structure would suggest Double bonds should be shorter than single bonds Electrons are localized between carbon atoms

Resonance for Benzene Experimental evidence shows All bonds are equal in length Length halfway between single and double bond Electrons are “delocalized” Spread around ring equally

VSEPR Theory Valence Shell Electron Pair Repulsion Model accounts for the shape of the molecule Lewis Structures don’t always show proper shape Electron groups repel each other and try to spread as far apart as possible Electron groups (domains) are lone pairs, single bonds, double bonds, and triple bonds

Electron Geometry There are 5 different possible arrangements around a central atom 2 electron groups up to 6 electron groups

Electron Geometry 2 electron groups around a central atom Electron groups on opposite sides Bond Angle of 180° Linear Shape

Electron Geometry 3 electron groups around a central atom Bond Angle of 120° Trigonal Planar

Electron Geometry 4 electron groups around a central atom Tetrahedral Bond Angle of 109.5° Tetrahedral

Electron Geometry 5 electron groups around a central atom 3 equatorial, 2 axial Equatorial angle of 120°, axial angle of 90° Trigonal bipyramidal

Electron Geometry 6 electron groups around a central atom Octahedral Bond angle of 90° Octahedral

Molecular Geometry Molecular Geometry is defined by the position of the bonded atoms in the molecule, not the lone pairs All binary molecules are linear Lone pairs still impact the shape due to their repulsion Same electron geometries may have different molecular geometries

Molecular Geometry Bent 3 electron groups, 2 bonding & 1 lone pair SO2 4 electron groups, 2 bonding & 2 lone pairs H2O

Molecular Geometry Trigonal Pyramidal 4 electron groups, 3 bonding & 1 lone pair NH3

Molecular Geometry Seesaw 5 electron groups, 4 bonding & 1 lone pair SF4

Molecular Geometry T-shaped 5 electron groups, 3 bonding & 2 lone pairs BrF3

Molecular Geometry Linear 5 electron groups, 2 bonding & 3 lone pairs XeF2

Molecular Geometry Square Pyramidal 6 electron groups, 5 bonding & 1 lone pair BrF5

Molecular Geometry Square Planar 6 electron groups, 4 bonding & 2 lone pairs XeF4

Bond Angles Angle between 2 atoms bonded to a central atom Other bonds and/or lone pairs in the molecule affect bond angle

Lone Pairs Electrons in a lone pair take up more space than bonded atoms due to repulsion between the electrons Multiple bonds also take up more space

Larger Molecules For larger molecules, we talk about the geometry around individual atoms

Valence Bond Theory Orbitals must overlap in order for atoms to share electrons and bond together

Internuclear Distance Balancing of attraction between nuclei and electrons with repulsion between electrons Nuclei repel each other if atoms get too close together

Internuclear Distance

Bonding Carbon has an electron configuration of 1s22s22p2 Carbon only has 2 unpaired electrons Carbon should only make 2 bonds

Hybrid Orbitals Orbitals can hybridize to form new equal orbitals Able to form more bonds Carbon mixes 2s and all 2p orbitals to make 4 equal sp3 hybrid orbitals

Hybrid Orbitals

Hybrid Orbitals

Hybrid Orbitals

Sigma and Pi Bonds There are 2 ways that orbitals can overlap to form bonds Sigma bonds Pi bonds

Sigma (σ) Bonds Head to head overlap All single bonds are sigma bonds

Pi (π) Bonds Side to side overlap Multiple bonds have both sigma and pi bonds Double bonds 1 sigma, 1 pi Triple Bonds 1 sigma, 2 pi

Multiple Bonds

Multiple Bonds

Resonance

Resonance

Metallic Bonding Bonding in metals is due to highly mobile valence electrons Delocalized electrons Metal nuclei and inner (core) electrons act as a cation, keeping the valence electrons within the sample Ability of the electrons to move throughout the entire sample give metals their unique properties

Metallic Bonding Metallic Properties Conductivity of heat and electricity Malleable Ductile

Alloys Metallic material that is composed of more than one element More rigid, less malleable and ductile than pure metal Two types Substitutional Interstitial

Substitutional Alloy One metal atom takes the place of original metal atom Brass, Bronze

Interstitial Alloys Small, often nonmetallic, atoms fit in between the metal atoms in a crystal Steel