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Why form bonds?
1. To get to a noble gas configuration (octet rule)– Ionic bonds (metal + nonmetal)• Metals want to lose their valence electrons• Nonmetals want to gain electrons
– Covalent bonds (nonmetals)• Both nonmetals want to share to reach a noble
gas configuration
2. To get to the lowest possible energy
Exceptions to the Octet Rule
• Hydrogen and helium want only 2
• Beryllium (metalloid) wants only 4
• Boron wants only 6
• 3rd energy level or below could be expanded if necessary– bonds using electrons in d orbitals
Reminders
• Electronegativity– How tightly an atom holds onto
electrons in a bond
– A measure of the attraction of an atom to another atom’s electrons in a bond
Reminders
• Ionic bond– Transfer of electrons
• Covalent bond– Sharing of electrons (a covalent bond is
always TWO shared electrons)
• Polarity – Unequal distribution of charge
Types of Bonds
• Based on the DIFFERENCE in electronegativity– System by Linus Pauling (1932)
– The more electronegative atom pulls the electrons in a bond
Types of Bonds
• Pauling’s system – Set fluorine as 4.0
– Differences in electronegativity• < 0.4 = nonpolar covalent• 0.4 – 2.0 = polar covalent• > 2.0 = ionic
Types of Bonds
• What it really means…– Same atom bonded to itself =
NONPOLAR
– Two different nonmetals bonded together = POLAR• The larger the difference, the more polar
the bond
– Metal and nonmetal = IONIC
More Ionic Bonding
• Coulomb’s Law– Describes the energy of interaction (attraction)
between a pair of ions
– Also called “lattice energy”
– C = Q1 x Q2 Q = charge of ion
r r = distance between centers of two
ions
More Ionic Bonding
• Coulomb’s Law– Examples:• NaCl vs. CaS – which has stronger
attraction / lattice energy?
• CaF2 vs. BaBr2 – which has stronger attraction/lattice energy?
More Ionic Bonding• Coulomb’s Law– Implications• A negative value means there is attraction
between the pair of ions, and you are getting lower energy by putting them together (good)• A positive value means there is repulsion,
and it’s endothermic (no bond formed)
• Charge larger influence than radius mathematically
More Ionic Bonding
• Ionic character– Measure of how ionic something is
– Can use locations on periodic table to predict• Shown – Which one is most ionic? NaF LiBr KBr NaI
More Covalent Bonding
• Two nonmetals sharing electrons
• Two different things bonded together is a polar bond
• Overall molecule may be polar or nonpolar based on symmetry
Bond Length Diagrams
• http://phet.colorado.edu/en/simulation/atomic-interactions
Bond Length Diagrams
• Size of dip = bond energy– Larger molecules = more energy– Double/triple bonds = more energy
More Covalent Bonding
• Nonpolar = symmetrical, equal pull in every direction
• Polar = not symmetrical, greater pull in one direction
• Polar and nonpolar don’t mix
More Covalent Bonding• Types of covalent bonds (rubber bands)
– Single bond• One shared pair of electrons• Sigma bond• Longest bond length• Lowest bond energy
– Double bond• Two shared pairs of electrons• Sigma bond + pi bond• Shorter bond length than single• Higher bond energy than single (less stable)
– Triple bond• Three shared pairs of electrons• Sigma bond + 2 pi bonds• Shortest bond length• Highest bond energy
More Covalent Bonding
• LOTS more to come with drawing and what the shape of each of these molecules is
• Stay tuned….
Lewis Structures
• G.N. Lewis (1916)
– Used lines and dots to represent bonding and valence electrons
– A line was a covalent bond (2 e-)– A dot represents each valence electron
Lewis Structures
• Atoms– The number of dots is equal to the
number of valence electrons in an atom– Fill each side of the symbol before
pairing any electrons (except H and He)
Examples:
Lewis Structures
• Ionic Compounds1. Draw symbol for each atom2. Transfer electrons until each metal has
0 valence electrons and each nonmetal has 8 valence electrons
3. Add multiple atoms of each type if necessary