Transcript Bonding and Structure in PowerPoint ppt format

Bonding and Structure
Octet Rule
• Atoms and ions prefer 8 electrons in the
valence shell. Hence, the unreactivity of
the inert / noble gases.
• If atoms cannot achieve a full octet, they
will settle for a filled or half-filled subshell.
• Atoms can achieve an octet by gaining or
losing electrons (ionic bonding) or sharing
electrons (covalent bonding)
Ionic Bonding
• When metals lose electrons, they acquire
a positive charge. When nonmetals gain
electrons, they acquire a negative charge.
• Likely charges can be inferred from the
periodic table group numbers.
• Atoms lose or gain electrons to achieve
same electron configuration as nearest
noble gas.
• Discussion based on NIST periodic table.
Ions
• Positive ions are smaller than original neutral
atom. Negative ions are bigger than original
neutral atom.
• Positively charged ions are primarily from
groups IA, IIA, IIIA. Maximum charge is the
group number. Transition metals (B-groups)
have multiple possibilities for charges. Due to
proximity of d-electrons to valence shell.
• Negatively charged ions are primarily from
groups VA, VIA, VIIA. Their negative charges
are given by -8 + group number.
Ionic Compounds
• Bonding is electrostatic and strong
• Characterized by ionic crystal lattice structures.
Packing arrangement determined geometrically
by ion size.
• High melting point
• Total positive charge must balance total negative
charge.
• Formulae for ionic compounds represent ratios
of atoms in lattice, not independent molecules.
Nomenclature – ionic compounds
• Positive ions with more than one charge
state have charge included in the name as
a Roman numeral. Otherwise, the names
of atoms with only one charge state omit
the charge in the name.
• Negative ions have only one charge state.
• Necessary to memorize charge states for
common elements. Use periodic table for
clues in case you forget.
Polyatomic ions
• Some ions are covalently bonded units but with
net charge. The deviation from charge neutrality
provides octets for all atoms in the unit. Most of
these ions carry negative charge. They also
form crystals in association with positive ions.
• Overall charge balance between positive and
negative ions must be present.
• Nomenclature – must memorize formulae and
names.
Covalent bonding
• Electrons are shared between atoms to
achieve octets. This forms structured
units with integrity. This is where Lewis
dot structures are useful.
• Simplest – diatomic molecules –
hydrogen, nitrogen, oxygen, halogens
Example
• Draw electron dot structures for the
common diatomic molecules – hydrogen,
nitrogen, chlorine.
Covalent Bonding
• Most covalent bonds are between nonmetals and between metalloids and nonmetals.
• Nomenclature rules – different from ionic
compounds. Structure of name signifies
type of bonding. Actual molecules are
formed.
Example
• Electron dot structures for methane, water,
ammonia, carbon dioxide, hydrogen
sulfide, formaldehyde, methanol, hydrogen
cyanide, ozone, sulfur dioxide, sulfate ion,
carbonate ion, ammonium ion, nitrate ion.
Covalent Bonding
• Sharing is not necessarily equal in a bond.
Electrons shifted to atom with greater
electronegativity. We get a dipole – a
bond with a positive end and a negative
end. When electronegativities are close,
sharing is nearly equal and bond becomes
closer to being non polar.
Molecules and Geometry
• Covalent bonding schemes have geometrical motifs –
fundamental bonding structures underlying what you
see.
• Molecular function depends on molecular structure both
as individual molecules and in aggregate.
• Common structural motifs (not exclusive list): tetrahedral,
flat triangle, linear. Note that some geometries become
truncated because of lone pairs occupying positions.
• Tetrahedral – has single bonds (sp3)
• Flat triangle – often has double bond (sp2)
• Linear – often has triple bond (but not always) (sp).
example
• Tetrahedral motif: methane, carbon
tetrachloride, ammonia, water
example
• Flat triangle: ozone, sulfur dioxide, nitrate
ion, formaldehyde.
example
• Linear: hydrogen cyanide, acetylene
Bond vs. molecular polarity
• Although bonds may be polar, the
molecule itself may not be. Dipole
moments are vector quantities that can be
resolved by trigonometry. If the vector
components cancel, then overall molecule
is nonpolar. Look for a high degree of
symmetry and balance as quick way to
spot nonpolar molecules.
example
• Sketch carbon tetrachloride, chloroform,
methylene chloride, methyl chloride,
methane to show polarity of bonds and of
the molecule.
Intermolecular forces
(among covalent molecules)
• Dispersion – based on fleeting and
instantaneous shifts in electron distribution.
Common in nonpolar and weakly polar
molecules. Weak.
• Polar interactions – among molecules with
permanent polarity. Intermediate.
• Hydrogen bonds – between hydrogen attached
to N,O,F in one molecule and N,O,F in another
molecule. Important biologically. Makes water a
liquid at room temperature while hydrogen
sulfide is a gas.
example
• Identify intermolecular forces in methane,
ammonia, methanol, nitrogen trichloride.