Bonding General Concepts Chapter 8 Bonding General Concepts

Types of Bonding Ionic Bonding Covalent Bonding Metallic Bonding Occurs when atoms gain or lose electrons to become ions Very Strong Attractions Covalent Bonding Occurs when atom share electrons Metallic Bonding Occurs when metal atoms allow a “sea of electrons” to be shared

Examples Ionic Covalent Metallic Sodium chloride, Lithium Sulfate, Iron (II) Chloride Covalent Carbon dioxide, Octane, Ethanol Metallic Aluminum, Copper, Bronze

But How Do I Know Ionic Covalent Metallic Metals and Nonmetals Metals and Polyatomic Ions Polyatomic Ions and Polyatomic Ions Covalent Nonmetals Metallic Metals

Electronegativity The ability of an atom in a molecule to attract shared electrons to itself. Developed by Linus Pauling Values range 0.7 to 4.0 Fluorine = 4.0 Francium and Cesium = 0.7 What is the periodic trend? Top to Bottom – Decrease Left to Right Increase

H----F Electron Sharing Electrons are not always shared evenly in covalent bonds. Called Polar Covalent Bonds Example of HF H----F

But How Do I Know Revisited Ionic Between atoms with a large difference in electronegativity Nonpolar Covalent Between atoms with no difference in electronegativity Polar Covalent Between atoms with a medium difference in electronegativity

Ion Size Ions are not the same size as their parent atom Positive Ions are smaller than parent Negative Ions are larger For a group of isoelectronic ions the most positive ion is the smallest

Example Place the following ions in order of decreasing size Na+, K+, Rb+, Cs+ Cs+> Rb+> K+ > Na+ Se-2, Br-, Rb+, Sr+2 Se-2> Br-> Rb+> Sr+2

Ionic Compound Formation The formation of ionic compounds from their elements is an exothermic process Several energy aspects that must considered

Energy Considerations What must happen for the reaction Na(s) + 1/2Cl2(g)  NaCl(s) We need to get Na+ and Cl- ions Sublimation of Na Ionization of Na (Ionization Energy) Breaking Cl2 bond (Bond Energy) Ionizing Cl (Electron Affinity) Combination of Na+ and Cl- (Lattice Energy) Sum is ΔHfº = Energy Change

Example #42 p. 404 Find ΔHfº Mg(s) + F2(g)  MgF2(s) Lattice Energy = -3916 kJ/mol Sublimation of Mg = 150 kJ/mol First Ionization Energy = 735 kJ/mol Second Ionization Energy = 1445 kJ/mol Bond Energy = 154 kJ/mol Electron Affinity = -328 kJ/mol

Lattice Energy Comparisons The lattice energy for two sets of ions can be compared with a form of Coulomb’s Law K is a constant (don’t worry about it) Q1 and Q2 are the charges of the ions R is the distance between the centers of the ions The LE will be neg. if the charges are opposite

Example Compare the lattice energies of Sodium Fluoride and Magnesium Oxide Sodium Fluoride will have a smaller LE because of the smaller charges

Homework p. 403 #’s 21,22,23,30,35abce, 40,41,46

Lewis Structures A method for determining the arrangement of bonds in covalent species Similar to dot structures but shows all bonds present

How 2 Draw Determine the number of valence electrons Determine the central atom. Usually the single atom, or the one in the middle of the formula Place other atoms around the middle and bond Complete octets with remaining electrons If each atom does not have 8 electrons Multiple Bonds my be necessary

Bond Types Single Bonds = 2 electrons Double Bonds = 4 electrons Weakest and longest bonds Double Bonds = 4 electrons In the middle Triple Bonds = 6 electrons Strongest and shortest There are not quadruple bonds!

General Rules Hydrogen will only form single bonds Halogens usually only form 1 bond. Why? 7 valence electrons Oxygen will have 2 bonds and often forms multiple bonds Carbon likes to form chains

Example – Draw the Lewis Structure for CBr4

Example – Draw the Lewis Structure for NF3

Example – Draw the Lewis Structure for O2

Example – Draw the Lewis Structure for CS2

Example – Draw the Lewis Structure for BeF2

Example – Draw the Lewis Structure for C6H14

Polyatomic Ions Atoms that are covalently bonded together and have a charge Lewis structure rules Negatively charged add electrons Positively charged subtract electrons Place Lewis structure in brackets when you are finished

Example – Draw the Lewis Structure for NO+

Resonance Species where equivalent Lewis structures exist Electron density is spread out evenly between resonant bonds Delocalized – Spread out Often present in polyatomic ions

Example – Draw the Lewis Structure for CO3-2

Example – Draw the Lewis Structure for NO2-

Example – Draw the Lewis Structure for AsF5

Homework p. 405 #’s 61,63,65,72

Formal Charge Difference between the number of valence electrons on free atom and the valence electrons in a species FC=Valence Electrons on free atom – valence electrons on the species Atoms desire lowest formal charge possible Negative formal charge should reside with most electronegative element

Example Use formal charge to compare the molecules

Molecular Geometry Lewis Structures do not show us the shape of molecules Use VSEPR Theory Valence Shell Electron Pair Repulsion Theory Electron Groups want to be as far apart as possible in molecules 1 Electron Group = Single, Double or Triple Bond or Lone Pair of Electrons Lone Pair Decrease the Bond Angle

2 Electron Groups Name Linear Bond Angle 180

3 Electron Groups Name Trigonal Planar Bond Angle 120

3 Electron Groups 1 Lone Pair Name Bent Bond Angle <120

4 Electron Groups Name Tetrahedral Bond Angle 109.5

4 Electron Groups 1 Lone Pair Name Trigonal Pyramidal Bond Angle 107

4 Electron Groups 2 Lone Pairs Name Bent Bond Angle 104.5

5 Electron Groups Name Trigonal Bipyramidal Bond Angle 90 &120

5 Electron Groups 1 Lone Pair Name SeeSaw or Distorted Tetrahedral

5 Electron Groups 2 Lone Pairs Name T Shape

5 Electron Groups 3 Lone Pairs Name Linear

6 Electron Groups Name Octrahedral Bond Angle 90

6 Electron Groups 1 Lone Pair Name Square Pyramidal

6 Electron Groups 2 Lone Pairs Name Square Planar

6 Electron Groups 3 Lone Pairs Name T Shape

6 Electron Groups 4 Lone Pairs Name Linear

Molecular Polarity A polar molecule is one that has a partially positive and partially negative side Molecules are Always nonpolar if they are one of the 5 base shapes w/ the same atom at the ends Molecules are Always polar their bond dipoles do not cancel out Molecules are polar if they do not have the same atoms at the end

Bond Dipole Example the Bond Dipole for CO2 and CH2O

Predict whether the molecule is polar or nonpolar

Homework P. 406 #’s 73,77,82,86 explain,92

Bonding Carbon forms four bonds with Hydrogen but, how! Carbon [He] 2s2 2p2 There are only 2 electrons to share Something more must have to happen!

Hybridization Mixing of different energy orbitals to form new bonding orbitals In CH4 Carbon needs to blend 1 s orbital and 3 p orbitals to be able to bond Called sp3 hybridization 4 electron groups gives sp3 hybridization

What is a Bond? σ (Think of it as a single bond) A bond is the overlap of orbitals Two hybrid orbitals, a hybrid and a nonhybrid, or two nonhybrid First bond to form is called a sigma bond σ (Think of it as a single bond)

Draw C2H6 in terms of orbitals How are the H’s aligned?

sp2 Hybridization Blending of 1s and 2p orbitals Used for 3 electron group geometry There is still 1 unhybridized p orbital left over Runs perpendicular to hybrid orbitals Unhybridized p is used for double bond Called a pi bond (π)

Draw C2F4 in terms of orbitals How are the F’s aligned? How many sigma bonds? Pi bonds?

sp Hybridization Blending of 1 s and 1 p orbital Used for 2 electron group geometry There is still 2 unhybridized p orbitals left over Run at 90 degrees of each other

Draw HCN in terms of orbitals How many sigma bonds? Pi bonds?

Expanded Octets Some atoms can expand their octets by utilizing unused d orbitals Must be in period 3 or greater 5 electron groups uses 1 d orbtital dsp3 hybridized 6 electron groups uses 2 d orbitals d2sd3 hybridized

Draw PCl5 in terms of orbitals

Draw SF6 in terms of orbitals

Bond Order The number of bonds between two atoms Ex. H2 is 1 O2 is 2 N2 is 3

Homework P. 441 #’s 11-15,22,24,28d