Transcript Covalent Bonds!

Covalent Bonds!
Yeah!
 Elements with high electronegativities (non-
metals) will not give up electrons. Bonds are
not formed by a transfer of electrons, they
are formed by sharing electrons.
 Molecules are neutral groups of covalently
bonded atoms
 A diatomic molecule is two atoms of the
same element covalently bonded together
Weird, huh
Molecular Compounds
 Molecular compounds tend to have lower
melting points than ionic compounds
 Many of them are either gases or liquids at
room temp.
 Some molecules can conduct electricity but
most don’t.
 Polyatomic ions are covalently bonded atoms
with a charge.
Why!?
Why!?
Why!?
Octet Rule…. again
 Atoms what to attain the electron
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configuration of a noble gas, (8 electrons in
the outer shell)
Nonmetal will share from 1-3 electrons in
order to achieve eight.
Single covalent bonds, two shared electrons
Double, four shared. Triple, 6 shared
Each shared pair makes a bond
More sharing
 Some electrons will not be involved in the
bonding process and are called an unshaired
pair. Single electrons are always bonded.
 A dashed line represents a bond, multiple
dashes, multiple bonds.
 Some molecules are exceptions to the octet
rule, multiple bonds make up for this, NO2
 Chemical symbols with dashes represent a
structural formula, compared to a chemical
formula which is just symbols and subscripts
Sophia in 30 years?
Must have been a rough life
Polyatomic ions
 Covalently bonded atoms with a charge,
several of them
 Many ionic compounds end either “ate” or
“ite”
 Many of them are coordinate covalent
compounds.
 Coordinate covalent compounds are
compounds where one atom donates both
bonding electrons. NH3 and NH4 for
example
Resonance
 Resonance structures are
different electron dot
configurations for the
same molecule
 Ozone, for example can be
drawn 2 different ways.
Bond Dissociation Energy
 The energy required to break a covalent
bond.
 A large bond dissociation energy corresponds
to a strong covalent bond.
 Single bond is weaker than a double weaker
than a triple.
 Some single bonds can be stronger than
other single bonds.
Molecular Orbitals
 At0ms have atomic orbitals. When atoms
bond together, it is theorized that these
orbitals overlap to form molecular orbitals, or
a combination of the two atomic orbitals.
 A sigma bond forms when two orbitals are
symmetrical around the two nuclei or the axis
between them , s or p orbitals for example
 A pi bond forms when p orbitals overlap side
by side, electrons are found above and below
the bonding axis.
What does that mean?!?
Molecular Orbitals
 Pi bonds overlap less than sigma bonds and
are weaker than sigma bonds
 This is one of several theories to explain the
principles behind atomic bonding, how it
occurs, and the shapes that result.
VSEPR
 Valence Shell Electron Pair Repulsion Theory,
notice it says theory. Another way to try to
explain molecular bonding.
 According to this theory, valence shell
electron pairs repel each other in order to
stay as far apart as possible.
 This accounts for bonding electrons and
unbonded pairs.
VSEPR
 Shapes include Linear triatomic, trigonal
planar, bent triatomic, pyramidal, and many
others. The shape depends on the number of
atoms, bonds, and unbonded electrons.
Why you wear sunscreen
Hybridization
 Long story short, different orbitals in the
same atom form one hybrid orbital in that
atom
 Methane, CH4 for example, Carbon has an
outer configuration of 2s2 3p2 It has to bond
with four hydrogens, but there are only 2
unpaired electrons. One electron comes up
from the s orbital to the p orbital to make it
2s1 3p3 and now we have four single electrons
to bond with hydrogen and an sp hybrid
orbital
Polar Bonds
 Covalently bonded atoms become polar
when one atom has a higher electronegativity
than the other. (usually, just more electrons)
 A polar covalent bond is one where atoms are
shared unequally. One side of the molecule
develops a positive charge and the other side
develops a negative charge due to the
imbalance of electrons
Polar
 Polar covalent bonds form polar molecules
 Polar bonds can cancel each other out if they
are in the same plane and linear, CO2 for
example
 Polar molecules are attracted to each other
by opposite charges.
If you are watching from the
ski lodge below, you might
think about moving
Intermolecular forces
 Molecules are attracted to each other by a
variety of ways called intermolecular forces.
 Intermolecular forces are weaker than atomic
forces such as covalent or ionic bonds.
 The two weakest forces are collectively called
Van der Waals Forces. They are dipole and
dispersion.
 Dipole is the same as polar, the negative end
of one molecule is attracted to the positive
end of another
More intermolecular
 After dipole are dispersion forces , the
weakest of all intermolecular forces.
 Dispersion is due to the movement of
electrons and is slightly stronger with more
electrons present.
 Hydrogen bonds, the strongest, occur
between molecules that due to their polarity,
share a hydrogen, same as polar or dipole but
with a hydrogen in the middle
Hydrogen Bonds
 They are the strongest and account for a lot
of important properties in water and
biological processes.
Covalent Bonding is kinda
hairy
Seriously