chapter 8 bonding general concepts. types of bonding ionic bonding –occurs when atoms gain or lose...
TRANSCRIPT
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
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
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
Resonance
• Species where equivalent Lewis structures exist
• Electron density is spread out evenly between resonant bonds– Delocalized – Spread out
• Often present in polyatomic ions
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
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
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
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)
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)
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
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