2 Na(s) + Cl2(g) ---> 2 Na1+Cl1-(s)
The sodium cation, Na1+, is formed when a sodium atom loses an electron.
Na --------> Na1+ + e1-
[Ne] 3s1
[Ne]
The chloride anion, Cl1-, is formed when a chlorine atom gains an electron.
Cl + e1- --------->
Cl1-
[Ne] 3s23p5
[Ne] 3s23p6
= [Ar]
Note that when the ions are formed in the previous equations, both ions adopt a noble gas electron configuration.
A covalent bond
is a chemical bond formed by the sharing of a pair of electrons between
atoms. The bonding in hydrogen
chloride, HCl, is
covalent bonding. More detail will be given later in the notes
on covalent bonding.
|
|
. |
. C . . |
. N : .. |
. O : .. |
: F : .. |
: Ne : .. |
H . + H . -----> H:H {":" represents a covalent bond, that is, a shared pair of electrons}
A Lewis structure
is a convention used to represent the covalent bonding in a molecule.
Dots or lines are used to represent
valence electrons
(electrons available for bonding). The Lewis structure for the hydrogen
molecule is as follows.
H:H or H-H
A bonding pair is an electron pair shared between two atoms.
A nonbonding pair
(lone pair, unshared pair) is a pair of electrons that remains on one atom
and is not shared.
In the Lewis structure for the chlorine molecule, Cl2, there is a single bonding pair located between the Cl atoms and a total of six (6) lone pairs distributed equally around the Cl atoms. Each Cl atom is surrounded by three (3) lone pairs and shares one lone pair. This makes a total of eight (8) valence electrons surrounding each Cl atom. (See octet rule.) |
. .
. .
: Cl - Cl : . . . . |
A coordinate covalent
bond is a bond formed when both electrons in a bond are donated by
one atom.
A + :B ---> A:B
A coordinate covalent
bond is not essentially different from other covalent bonds; it involves
the sharing of a pair
of electrons between two atoms.
The octet rule
is the tendency of atoms in molecules to have eight electrons in their
valence shell.
A single bond
consists of one pair of shared electrons between two atoms.
{e.g. H2} |
H-H |
A double bond
consists of two pairs of shared electrons between two atoms.
{e.g. O2} |
. . .
.
O=O . . . . |
A triple bond
consists of three pairs of shared electrons between two atoms.
{e.g. N2} |
: N=N : |
X = I.E. - E.A./2
Linus Pauling's scale
is based upon bond energies. Fluorine is assigned a value of 4.0 and Li
a value of 1.0.
The general trend
is that electronegativy increases left to right across a period and bottom
to top in a group.
The absolute value
of the difference in electronegativity of two bonded atoms gives a rough
measure
of the polarity to
be expected in a bond. When this difference is small, the bond is
nonpolar. When it
is large, the bond
is polar (or, if the difference is extra large, perhaps ionic).
POLAR COVALENT BONDS
A polar covalent
bond is a covalent bond in which the bonding electrons spend more time
closer to one atom
than the other.
The majority of covalent bonds are polar.
The bond between hydrogen
and chlorine in HCl is a polar bond. The chlorine is more electronegative
and therefore attracts
the bonding electrons to itself (i.e. electrons are not equally shared).
The chlorine atom
in HCl adopts a partial
negative charge (d-)
and the hydrogen atom adopts a partial positive charge (d+).
Polar covalent bonds
can be represented by the following convention.
d+ d-
H--Cl
2.
Determine the central atom and draw a single bond to each of the atoms
surrounding it.
a. The central atom will most often be the atom with the lowest
electronegativity.
b. Molecules and polyatomic ions are symmetrical and compact.
c. In oxyacids, hydrogen is bonded to oxygen. {H2SO4
is an oxyacid}
d. When the noble gases Kr and Xe are present in a molecule,
they are central atoms.
3.
Complete the octet around each bonded atom.
Octet Rule Exceptions:
a. Hydrogen requires only 2 electrons. Hydrogen is always
a terminal atom and never a central atom.
b. Beryllium requires only 4 electrons.
c. Boron requires only 6 electrons.
d. Elements in Period 3 and beyond can exceed the octet
rule. In other words non-metallic atoms
with atomic numbers larger than 14 can hold more than 8 electrons.
4. Distribute any remaining electrons as lone pairs around the central atom (or atoms).
5.
If there are fewer than eight electrons on the central atom move one or
more electron pairs from a bonded atom(s)
to to form a multiple bond to the central atom(s).
Atoms that can form multiple bonds are C, N, O, and S.
NOTE: Lewis structures for ionic species should be enclosed in brackets [ ] and the charge indicated.
Rules for determining
formal charge
1. Half of the
electrons of a bond are assigned to each atom in the bond (counting each
dash as two electrons).
2. Both electrons
of a lone pair are assigned to the atom to which the lone pair belongs
3. Formal charge
= Valence electrons - assigned electrons.
4. Whenever
you can write several Lewis formulas for a molecule, choose the one having
the lowest magnitudes
of formal charge.
5. When two
proposed Lewis formulas for a molecule have the same magnitudes of formal
charges, choose the one
having the negative formal charge on the more electronegative atom.
: O - O=O |
|
O=O - O : |
One misconception which perhaps must be addressed is that electrons in bonds are not sticks or dots which are "stuck" in position but can be thought of as being "spread out" over several atoms. (Recall that quantum mechanics describes electrons in terms of three dimensional waves.) In particular, the data from the measurement of the bonds in ozone, O3, indicates that instead of one single bond and one double bond (as suggested by Lewis structure rules) the bonds are an intermediate between the two. The electrons in these bonds are said to be delocalized.
BOND AND ENERGY
Bond energy
(BE) the average enthalpy change for the breaking of an A-B bond in a molecule
in the gas phase.
Because it requires
energy to break bonds (endothermic process), bond energies are always positive
numbers.
When a bond is formed, the energy is equal to the negative of the bond energy (energy is released, exothermic).
Bond formation is exothermic. Bond breaking is endothermic.
In general, the enthalpy
of reaction (DH)
is (approximately) equal to the sum of the bond energies for bonds
broken minus the sum
of lthe bond energies formed.
DH = SBE(reactants) - SBE(products)
EXAMPLE
Using bond energies,
calculate the enthalpy of reaction for the formation of hydrogen chloride
gas from
hydrogen gas and chlorine
gas.
H2(g) + Cl2(g) ----> 2 HCl(g) DH = ?
Bond Energies: H-H 436 kJ/mol; Cl-Cl 243 kJ/mol; H-Cl 432 kJ/mol
DH = SBE(reactants) - SBE(products)
DH =[ (1 mol H-H bonds)(436 kJ/mol) + (1 mol Cl-Cl bonds)(243 kJ/mol) ] - [(2 mol H-Cl bonds)(432 kJ/mol)]
= -185 kJ
Valence-Shell Electron-Pair Repulsion Model: VSEPR
VSEPR is a simple
way of predicting the shapes of species that have main group elements as
central atoms.
According to this
theory, electron pairs in the valance shell repel each other and therefore
will be as far
apart from each other
as possible in three dimensional space.
The shape of a molecule
is determined by the positions of the atoms that are bonded to the central
atom.
The positions of the
bonded atoms depend upon the number of bonded atoms in the molecule and
upon
the steric number
(SN).
The steric number (also
referred to as total coordination number) is defined as the total number
of atoms and
sets of unshared electron
pairs around a central atom. To determine the steric number draw
the
Lewis structure and
count the number of bonded atoms and lone pairs on the central atom.
Each double
bond and each triple
bond count as 1 when determining steric number.
SN = number of bonded atoms + number of lone pairs on the central atom
Example 1
Calculate the steric
number for H2O, SiO2, and ClF3.
For H2O the steric number is 4. There are 2 bonded atoms and 2 lone pairs on the central atom.
For SiO2
the steric number is 2. There are 2 bonded atoms and no lone pairs
on the central atom.
Notice that a double
or triple bond count one each in determining the steric number.
VSEPR theory treats
multiple bonds as if they were a single pair of electrons between the bonded
atom and the central atom.
For ClF3 the steric number is 5. There are 3 bonded atoms and 2 lone pairs on the central atom.
VSEPR Formula
(Notation)
VSEPR theory predicts
shapes of molecules based upon the number of bonded atoms and lone pairs
on the central atom.
Molecules with a certain
number of bonded atoms and lone pairs are predicted to have a particular
geometry.
One convention used
to classify molecular shape predicted by VSEPR assigns a specific AXmEn
formula to each shape.
A represents
the central atom,
X represents
the bonded atom(s)
E represents
any lone pairs on the central atom.
The subscripts m
and n are the number of bonded atoms and the number of lone pairs
respectively.
Different molecules with identical VSEPR formulas will have identical molecular shapes.
Example 2
Use VSEPR notation
to describe H2O, SiO2, and ClF3.
H2O
AX2E2
SiO2 AX2
ClF3 AX3E2
Example 3
Predict the shapes
of H2O, SiO2, and ClF3.
For H2O the predicted shape is angular.
For SiO2 the predicted shape is linear.
For ClF3 the predicted shape is T-shaped.
Molecular Shapes Predicted by VSEPR Theory
Steric
Number |
Number of Bonding Pairs | Number of Lone Pairs | Molecular Geometry | VSEPR Notation | Bond Angles | Example |
2 | 2 | 0 | Linear | AX2 | 180o | BeCl2 |
3 | 3 | 0 | Trigonal Planar | AX3 | 120o | BH3 |
2 | 1 | Angular | AX2E1 | SO2 | ||
4 | 4 | 0 | Tetrahedral | AX4 | 109.5o | CH4 |
3 | 1 | Trigonal Pyramidal | AX3E1 | NH3 | ||
2 | 2 | Angular | AX2E2 | H2O | ||
1 | 3 | Linear | AX1E3 | HF | ||
5 | 5 | 0 | Trigonal Bipyramidal | AX5 | 90o,120o,180o | AsCl5 |
4 | 1 | Seesaw | AX4E1 | SF4 | ||
3 | 2 | T-shaped | AX3E2 | ICl3 | ||
2 | 3 | Linear | AX2E3 | ICl21- | ||
6 | 6 | 0 | Octahedral | AX6 | 90o, 180o | SF6 |
5 | 1 | Square Pyramidal | AX5E1 | BrF5 | ||
4 | 2 | Square Planar | AX4E2 | IF41- |
Sample molecules and ions for Lewis structures and VSEPR theory.
1 OSF4 2 COCl2 3 POCl3 4 SOCl2 5 OCN1- 6 SCN1- 7 XeO2F2 8 XeOF2 9 XeOF4 10 (CH3)2CO |
|
ACIDS
|