Chapter 8

Bonding General Concepts

Types of Bonding

• Ionic 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– Sodium chloride, Lithium Sulfate, Iron (II)

Chloride

• Covalent– Carbon dioxide, Octane, Ethanol

• Metallic– Aluminum, Copper, Bronze

But How Do I Know

• Ionic– 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

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 ΔHf

º = Energy Change

• Example#42 p. 404

Find ΔHfº

Mg(s) + F2(g) MgF2(s)

Lattice Energy = -3916 kJ/molSublimation of Mg = 150 kJ/molFirst Ionization Energy = 735 kJ/molSecond Ionization Energy = 1445 kJ/molBond Energy = 154 kJ/molElectron 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

)( 21

r

QQkrgyLatticeEne

• 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– 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

• Polar

• Nonpolar

• Polar

• Nonpolar

• Polar

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?

• 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