CHEMICAL BONDING
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Why bond?
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Atoms try to acquire particularly stable electron
configuration (i.e the configuration maximizing binding
energy) for their outer electrons
 form ions or molecules.
Most stable configuration for outer electrons is noble
gas configuration (octet structure)
ns2np6
strategies to achieve more stable configuration:
 give away electrons
 accept electrons
 share electrons
chemical bond is formed when it is energetically
favorable, i.e., when the energy of the bonded atoms is
less than the energies of the separated atoms.
Quantities which play a role in chemical bonding:
 ionization energy (the energy required to remove an
electron from a neutral atom)
 electron affinity (the energy change when a neutral
atom attracts an electron to become a negative ion)
 electronegativity (the ability of an atom in a
molecule to draw bonding electrons to itself)
types of bonds:
 ionic bond
 covalent bond
 metallic bond
IONIC BOND
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atoms achieve noble gas configuration by giving up
or accepting electrons (usually electron transfer
from metal to non-metal)  formation of ion;
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chemical bond is formed due to electrostatic
attraction between two oppositely charged ions;
bond is strong, but becomes quickly weak when ion
is displaced;  materials usually brittle (e.g. glass,
rock, egg shells,..)
compounds formed by ionic bond: e.g. NaCl, CaCl2,..
examples of ions with noble gas configuration:
Noble gas configuration
He
1s2
Ne
[He]2s2 2p6
Ar
[Ne]3s2 3p6
..
…..
ions
H- , Li+ , Be2+
N3- , O2- , F- ,Na+
S2- , Cl- , K+ ,Ca2+
….
METALLIC BOND
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metals have low ionization energy  atoms give up
outer (valence) electron, to be shared by all
metallic lattice = positive ions at fixed positions, in
“sea” or “gas” of mobile electrons;
electron gas has “pressure” due to Pauli principle
electrons mobile  good thermal and electrical
conductivity
positive ions not free to move, but can vibrate
electron not tied down in particular bond  can absorb
and re-emit light over wide frequency range  good
reflector
bond is “elastic” since attraction due to mobile
electrons bond holds even if ions displaced
 metals are malleable
ability of the electrons to spread between the cations
and neutralize their charges  metal ions ions can pack
closely together; closeness of the packing of the atoms
 high densities of metals.
In some sense, piece of metal is like an extremely large
molecule
Covalent Bond
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COVALENT BOND
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well-defined cluster of neighboring atoms share
electrons - form “molecule”
state with shared electrons has lower energy than
individual atoms
“valence electrons” = outer electrons
group number = number of valence electrons
“valence of element” = number of electron pairs shared
to complete octet of electrons
examples: H2 , O2 , N2
 in H2, H atoms share one electron pair;
 n O2, O atoms share two electron pairs;
 in N2, N atoms share three electron pairs;
 single H, O, N (“in statu nascendi”) are much more
reactive than pair
Carbon
2
2
 Carbon's outer electron configuration is4s 4p
 it needs 4 electrons to complete octet
 shares 4 pairs of electrons
 e.g. methane (“swamp gas”)
 possibility of 4 covalent bonds  large variety of
possible compounds with C  organic chemistry
Polar bonds
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POLAR BONDS
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Electron pairs shared between two different atoms
not necessarily shared equally - sharing ratio
depends on electronegativity;
“electronegativity” = ability of atom to attract an
electron to itself;
electronegativity increases from left to right (i.e.
grows with group number) in every row of the
periodic table;
alkali metals (group 1) are least electronegative,
halogens (group 7) are most electronegative;
non-polar bond = bond in which electrons are shared
equally
polar covalent bond = bond in which one of the
atoms exerts greater attraction for the electrons
than the other (has larger electronegativity)
if difference in electronegativity is large enough,
bond becomes ionic bond
GEOMETRY OF BONDS
electron-pair repulsion rule (EPRR):
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Electron pairs surrounding an atom (be they shared
or un-shared with other atom) repel each other and
are directed to be as far apart as possible.
(example: water molecule)
INTER-MOLECULAR FORCES
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electric dipole forces: polar molecules exert
electric dipole forces on each other;
 e.g. water molecule: H(+) attracted to (-) partner
of other molecule (oxygen of other water
molecule, or solute constituent)
 “hydrogen bond”;
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hydrogen bond is reason for water to be liquid
at “normal” temperatures (note e.g. CO2 , CH4
are gases!)
Van der Waals forces:
 fluctuations, shifts of charges within covalent
molecule
 temporary dipole moment
 electrostatic dipole forces;
 charge unbalances small  forces weak (Van der
Waals binding energies are  10-2eV )
 most liquids held together by VdW forces,
 cohesion, surface tension
Water
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water molecule:
 oxygen in water has 4 electron pairs, 2 shared
and 2 unshared;
 EPRR e-pairs point to corners of tetrahedron,
with O atom in center of tetrahedron,
 H atoms sit at two of the corners of the
tetrahedron;
 constituents of water form triangle; angle
between the two lines from O to H is 105o
(109o predicted from this model);
 “Mickey Mouse shape” of water molecule
 O is more electronegative than H  water is
polar molecule, more negative near O atom,
more positive near H atoms.
 polarity of water molecule  good solvent
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Water and Ice
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Why ice is less dense than liquid water
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water is most dense as a liquid and least dense as a
gas; solid water (ice) is less dense than liquid
water, but more dense than a gas;
water vapor (gaseous water) is made up of individual
water molecules that are not hydrogen bonded to
other water molecules;
cooling, pressure  hydrogen bonds form  liquid;
but in liquid, hydrogen bonds keep forming and
breaking up; on average every water molecule in 3.4
H-bonds
at 4oC, water molecules as tightly packed as they
will go
below 4oC , more and more stable hydrogen bonds
form, eventually reaching 4 per molecule; the
hydrogen bonds push the water molecules further
apart, making it less dense ice floats on water