Pre AP Chapter 6 Notes
A chemical bond is a mutual electrical attraction
between the nuclei and valence electrons of
different atoms, and binds those atoms together.
There are three types of bonding: ionic, covalent,
and metallic
There are three types of bonding: ionic, covalent,
and metallic
1.Ionic bonding results from the transfer of
valence electrons from one atom to another.
There are three types of bonding: ionic, covalent,
and metallic
1.Ionic bonding results from the transfer of
valence electrons from one atom to another.
2.Metallic bonding results from a sharing of very
weakly held valence electrons
There are three types of bonding: ionic, covalent,
and metallic
1.Ionic bonding results from the transfer of
valence electrons from one atom to another.
2.Metallic bonding results from a sharing of very
weakly held valence electrons
3.Covalent bonding results from the sharing of
valence electrons, whether equally or
unequally.
The type of bonding between two atoms can be
estimated by calculating the difference in the
elements’ electronegativity.
The type of bonding between two atoms can be
estimated by calculating the difference in the
elements’ electronegativity. To determine the
type of bond, if the electronegativity difference
is:
0-.3
then the bond is nonpolar covalent
>.3 - 1.7
then the bond is polar covalent
>1.7
then the bond is ionic
Nature favors chemical bonding because most
atoms are at lower potential energy when
bonded to other atoms than when they are
independent particles.
Nature favors chemical bonding because most
atoms are at lower potential energy when
bonded to other atoms than when they are
independent particles. The bond length is the
average distance between two bonded atoms.
Nature favors chemical bonding because most
atoms are at lower potential energy when
bonded to other atoms than when they are
independent particles. The bond length is the
average distance between two bonded atoms.
Bond energy is the energy required to break a
chemical bond and form neutral, isolated
atoms. The unit for bond energy is in kilojoules
per mole (kJ/mol).
An ionic compound is composed of positive and
negative ions (cations and anions) that are
combined in the right ratio to that the total of the
positive and negative ions are zero.
Most ionic compounds are crystalline solids. A
crystal is a three-dimensional network of
positive and negative ions.
Most ionic compounds are crystalline solids. A
crystal is a three-dimensional network of
positive and negative ions. A formula unit is the
simplest collection of atoms from which the
ionic compound’s formula can be established.
In an ionic crystal, ions minimize their potential
energy by combining in an orderly arrangement
known as a crystal lattice.
Most ionic compounds are crystalline solids. A
crystal is a three-dimensional network of
positive and negative ions. A formula unit is the
simplest collection of atoms from which the
ionic compound’s formula can be established.
In an ionic crystal, ions minimize their potential
energy by combining in an orderly arrangement
known as a crystal lattice. Lattice energy is the
energy released when one mole of an ionic
crystalline compound is formed from gaseous
ions.
The forces of attraction between compounds and
molecules are much weaker than the forces
inside holding the compound or molecule
together. In addition, the forces holding ionic
compounds together are much stronger than
the forces holding covalent molecules together.
This accounts for the high melting and boiling
points of ionic compounds.
Generally ionic compounds do not vaporize at
room temperature, and they are hard but brittle.
They are poor conductors of electricity in their
solid state but are good conductors in their
liquid state. Most ionic compounds are easily
dissolved and will also be good conductors in
their aqueous form.
A polyatomic ion (radical) is a charged group of
covalently bonded atoms. The charge of a
polyatomic ion results from an overall excess of
electrons (negative charge) or a shortage of
electrons (positive charge). Polyatomic ions will
form ionic bonds with other charged elements
or other polyatomic ions to form a neutral
compound.
Metallic bonding accounts from the unique
properties of metals. Their high conductivity is
due to the highly mobile valence electrons
(often referred to as a sea of electrons).
Metallic bonding accounts from the unique
properties of metals. Their high conductivity is
due to the highly mobile valence electrons
(often referred to as a sea of electrons). These
electrons roam freely throughout the atom and
are called delocalized.
Metallic bonding accounts from the unique
properties of metals. Their high conductivity is
due to the highly mobile valence electrons
(often referred to as a sea of electrons). These
electrons roam freely throughout the atom and
are called delocalized. This also accounts for
the luster of metals because the electrons
absorb light, jump to a higher energy level, and
then fall back down to ground level while
emitting the energy in the form of light.
Metallic bonding accounts from the unique
properties of metals. Their high conductivity is
due to the highly mobile valence electrons
(often referred to as a sea of electrons). These
electrons roam freely throughout the atom and
are called delocalized. This also accounts for
the luster of metals because the electrons
absorb light, jump to a higher energy level, and
then fall back down to ground level while
emitting the energy in the form of light. The
ability of the metal nuclei to “ride over” the sea
of electrons helps account for the ductility and
malleability of metals.
The metallic bond strength varies with the nuclear
charge of the metal and the number of
electrons in the sea. Both of these factors are
reflected in the heat of vaporization which is the
amount of heat required to vaporize one mole
of metal.
Covalent bonding results from the sharing of
electrons in pairs between two nuclei. If two
atoms share the electron pair equally, the bond
is nonpolar covalent. If the two atoms have an
unequal sharing it is a polar covalent bond.
Again, the type of boding is determined by the
difference in the electronegativities of the
atoms.
To indicate the positive end of a bond, “δ+” is
used and for the negative end of the bond, “δ-”
is used. For example, in the bond between H
and Cl, the bond is polar covalent because
hydrogen’s electronegativity is 2.1 and
chlorine’s is 3.0, so the difference is 0.9. We
would show this as:
δ+
H - Cl
δ-
A molecule is a neutral group of covalently
bonded atoms. It is capable of existing on its
own (does not have a crystal lattice). A
chemical formula or molecular formula indicates
the relative numbers of atoms of each kind in a
chemical compound by using symbols for the
elements and subscripts. A molecule of only
two elements is called a diatomic molecule.
NO2
H2 O
H2O2
N2O5
To draw compounds and molecules,
the valence electrons act as a
guide using the Octet Rule.
Elements form compounds and
molecules with gaining, losing, or
sharing electrons to have an octet
of valence electrons (eight), filling
up the s and p sublevels to give a
noble gas configuration and
stability.
For ionic compounds show the electrons being
lost from the metal and gained by the nonmetal.
Place a bracket around the newly formed
cations and anions with the charge outside.
Examples:
..
...
..
.
.
.
.
.
+
2+
- Ba2+ [ . I . ].
.
.
.
K [ O
]
K
and
[
I
]
..
..
..
Notice it is the negative nonmetal that gets all the
dots!
For covalent molecules, there is a simple method
that can be used to determine how many bonds
and how many lone pairs should be around
each atom. This method is the W - A = S
method
W = the total number of electrons all the atoms
want
(each atoms wants 8, but hydrogen only wants 2)
A = the total number of electrons all atoms have
available
(count the dots off the periodic table, add for each charge and subtract for each + charge)
S = the total number of electrons that need to be
shared
(divide by 2 for the number of bonds)
Examples:
SO2
PO33C2H2
Limitations to the octet rule (cases in which the
octet rule may not apply):
1.Most covalent compounds of Be
2.Most covalent compounds of Group 3A (13)
3.Compounds in which the central atom must
share more than 8 valence electrons to
accommodate all the attached atoms (like SF6)
4.Compounds containing d or f metals
5.Species with odd numbers of electrons (like NO
with A = 11 electrons)
Unshared electron pairs are called lone pairs and
belong to one atom only. Lewis structures are
these structures in which atomic symbols
represent nuclei, dashes represent shared
pairs, and dots represent lone pairs.
Unshared electron pairs are called lone pairs and
belong to one atom only. Lewis structures are
these structures in which atomic symbols
represent nuclei, dashes represent shared
pairs, and dots represent lone pairs. Structural
formulas indicate the kind, number and
arrangement, and bonds of the atoms, but not
the unshared lone pairs.
A single bond is produced by sharing one pair of
electrons between two atoms. A double bond is
produced by sharing two pairs of electrons
between two atoms. A triple bond is produced
by sharing three pairs of electrons between two
atoms.
Double and triple bonds are referred to as
multiple bonds. Resonance can occur with
multiple bonds, and refers to situations where
more than one structure can be drawn.
Example:
SeO2
In real life molecules form three dimensional
structures that are not as simple as the two
dimensional structures above. Thus it takes a
new theory - Valence Shell, Electron Pair
Repulsion, or VSEPR to tell us how to draw the
structures three dimensionally.
In real life molecules form three dimensional
structures that are not as simple as the two
dimensional structures above. Thus it takes a
new theory - Valence Shell, Electron Pair
Repulsion, or VSEPR to tell us how to draw the
structures three dimensionally. This theory
states that the repulsion between the sets of
valence level electrons surrounding the atom
causes these sets to be arranged as far apart
as is possible.
In real life molecules form three dimensional
structures that are not as simple as the two
dimensional structures above. Thus it takes a
new theory - Valence Shell, Electron Pair
Repulsion, or VSEPR to tell us how to draw the
structures three dimensionally. This theory
states that the repulsion between the sets of
valence level electrons surrounding the atom
causes these sets to be arranged as far apart
as is possible. This could lead to molecular
polarity where there is an uneven distribution of
molecular charge.
There are several forces that act between
compounds and molecules to hold them
together (not referring to the ones inside the
compounds and molecules). These forces are
called Intermolecular forces (as opposed to
Intramolecular forces).
There are three types of intermolecular forces:
1.Hydrogen bonding - the attraction of a hydrogen
on one polar molecule to the lone pair of a
highly electronegative atom (N, O, F) of a
second polar molecule.
Click here for a demo of hydrogen bonding in water.
2.Dipole-dipole - a dipole is created by equal but
opposite charges that are separated by the
short distance of a bond. The positive end of
one dipole molecule will be attracted to the
negative end of a second dipole molecule, and
vise versa.
3.London dispersion forces - resulting from the
constant motion of electrons and the creation of
instantaneous dipoles.
Covalent Network
Intramolecular
Forces
Ionic
Metallic
Intermolecular
Forces
Hydrogen bonding –
polar with hydrogen
Dipole-dipole –
polar only
London –
nonpolar