Transcript Ch. 8 Sections 8.1-8.3 Powerpoint

Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
participate in chemical bonding.
Group
e- configuration
# of valence e-
1A
ns1
1
2A
ns2
2
3A
ns2np1
3
4A
ns2np2
4
5A
ns2np3
5
6A
ns2np4
6
7A
ns2np5
7
•What is a chemical bond?
•Define bonds as the forces that hold groups of atoms
together and make them function as a unit.
•Why do atoms bond?
•A bond will form if the potential energy of the bonded
atoms is lower than that of the separated atoms.
•Two principal classes of bonding forces.
•Covalent bonding: occurs in molecules and involves the
sharing of electrons.
•Ionic bonding: involves the transfer of electrons between
atoms; this produces ions.
•In Figure (a) two hydrogen atoms are being brought
together to form a bond.
•Two unfavorable potential energy interactions: protonproton repulsion and electron-electron repulsion, and one
favorable interaction: proton-electron interaction.
•When will the H2 molecule be favored over the separated
hydrogen atoms?
•Remember nature favors lower potential energy so the
hydrogen atoms will position themselves to achieve the
lowest possible energy.
•The distance where the energy is minimal is called the
bond length.
•Bond length: the distance between the nuclei of the two
atoms connected by a bond or the distance where the total
energy of a diatomic molecule is minimal.
•In ionic bonding the participating atoms are so different
that one or more electrons are transferred to form
oppositely charged ions, when then attract each other.
•In covalent bonding (also called nonpolar covalent
bonding) two identical atoms share electrons equally.
•There are intermediate cases where the atoms are not so
different that electrons are completely transferred but are
different enough that unequal sharing results, forming what
is called a polar covalent bond.
•An example of this type of bond occurs in the hydrogen
fluoride (HF) molecule.
•When molecules of HF are placed in an electric field, the
molecules tend to orient so that the fluoride end is closed to
the positive pole and the hydrogen end closest to the
negative pole (Figure b).
•This suggests the HF molecule has the following charge
distribution.
H-F
δ+ δWhere δ is used to indicate a fractional charge.
•This partial negative and positive charges on the atoms is
called bond polarity.
•Explanation: electrons in the bond are not shared equally.
•The different affinities of atoms for the electrons in a bond
are described by a property called electronegativity.
•Electronegativity: the ability of an atom in a molecule to
attract shared electrons to itself.
•Electronegativity generally increases from left to right
across a period and decreases down a group for the
representative elements.
•Range is from 4.0 for fluorine (most electronegative
element) to 0.7 for cesium.
•For identical atoms (no electronegativity difference), the
electrons in a bond are shared equally, and no polarity
develops (nonpolar covalent).
•For atoms with very different electronegativity values,
electron transfer occurs (ionic bond).
•Intermediate cases give polar covalent bonds with unequal
electron sharing.
•A molecule such as HF which has a center of positive
charge and a center of negative charge is said to be dipolar,
or to have dipole moment.
H-F
δ+ δ•The dipolar character of a molecule is often represented by
an arrow pointing to the negative center charge with the tail
of the arrow indicating the positive center of charge.
•Electrostatic potential diagrams can be used to represent
charge distribution.
•Colors of visible light are used to represent the variation in
charge distribution.
•Red indicates the most electron-rich region of the molecule
and blue indicates the most electron-poor region.
•Any diatomic (two-atom) molecule that has a polar bond
will show a molecular dipole moment.
•In other words, if there is a difference in electronegativity
between the two atoms the molecule is polar.
H - Cl
•But what about polyatomic (more than two atoms)
molecules?
•For example, the oxygen atom in water has a greater
electronegativity value than hydrogen and the charge
distribution is shown in Figure (a) below.
•In Figure (b), when placed in an electric field the water
molecule behaves as if it has two centers of charge.
•Thus water has a dipole moment.
•This type of behavior is observed for the NH3 (ammonia)
molecule also.
•Some molecules have polar bonds but do not have a dipole
moment.
•An example is the CO2 molecule, which is a linear molecule
that has the charge distribution shown below.
•The opposing bond polarities cancel out and the CO2
molecule does not have a dipole moment.
•In other words, it is a nonpolar molecule.