Transcript Bonding: General Concepts

Bonding: General Concepts
Chapter 8
Words to know:
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chemical bond
ionic bond
covalent bond
metallic bond
Lewis Symbol
Octet rule
Practice
Classify the following compounds as ionic or
covalent. Justify your answer. Which
compounds contain both types of bonds?
- KBr
- SO2
- H2SO4
- CH3COOH
- Na3PO4
- CaCO3
Lewis Symbol
• Element symbol with valence electrons
written around it as dots
• Elements want to gain, lose, or share electrons
to look like a noble gas (isoelectronic)
Place the following chemical species
into isoelectronic groups:
N3-, K+, Ca2+, O2-, F-, Ne, Br, Kr, Sc3+, Na+, Al3+, Se2-,
Mg2+
Ionic Bonding
• electrons are transferred from an atom with
low electronegativity to one with high EN
• electrostatic attraction between the two
oppositely charged ions
• arranged in a crystal lattice (lattice energy,
DHL, is energy required to completely separate
solid ionic compound into gaseous ions)
Lattice Energies?
• Lattice energy increases when the atoms are
smaller or have a higher charge (exchanging
more electrons)
• higher lattice energy means the ionic
compound is more strongly bonded
• high lattice energies also explains why ionic
compounds are brittle and hard
In each of the following pairs of
compounds, identify the one with the
higher lattice energy
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KCl, CaS
LiF, NaCl
Fe2O3, MnO2
CaO, CaCl2
Ions
• Representative elements follow the “hill of
oxidation numbers” when in ionic compounds
• transition metals (including lead & tin) are a
little weird:
– Their valence electrons are the highest filled s
sublevel and occasionally 1 or more of their d
electrons
Transition Metal Ions
• Iron
– Fe: [Ar] 4s23d6
– Fe2+: [Ar] 3d6
– Fe3+: [Ar] 3d5
• Lead
– Pb: [Xe] 6s24f145d106p2
– Pb2+: [Xe] 6s24f145d10
– Pb4+: [Xe] 4f145d10
Write the electron configurations for Cr3+
and Sn4+
Covalent Bonding
• 2+ atoms sharing electrons
• Lewis structures show shared and lone pairs of
electrons
• polar or non-polar determined by difference
in EN values for the 2 bonding elements
– 0-0.4 = NPC
– .4-1.0 = PC
– 1.0-2.0 = really PC
– >2.0 = ionic
Dipoles
• With polar covalent bonds, there is a dipole
(one end of the bond hogs the “shared”
electrons a little more than the other)
• symbolized with d+ and d- and are not whole
number charges
Drawing Lewis Structures
1. Add up the total valence electrons of all bonding
atoms
2. Use one pair of electrons to bond each outer
atom to the central atom (usually the least EN or
the one present in the least abundance)
3. Complete octets around all of the outer atoms
4. Place any remaining electrons around the
central atom
5. If there aren’t enough electrons to give the
central atom an octet, make multiple bonds
Practice
Write the Lewis structures for each of the
following compounds:
• NO3• CO2
• PCl5
• NO3
Resonance Structures
• used when 2+ Lewis structures are equally
good representations of the bonds
• actual structure is kind of an average of all the
possibilities
• examples include ozone & nitrate ion
Exceptions to Octet Rule
• When there is an odd # of electrons, one atom
will only have 7 electrons around it. (NO, NO2)
• When the compound has a group 2 or 3
element as the central atom, the number of
electrons around the central atom will be
twice the group number
• If central element is big, it can have an
expanded octet (PCl5)
Covalent Bond Strength
• Multiple covalent bonds are stronger than
single covalent bonds (they are also shorter)
• Higher the number of electrons shared, the
stronger the bond