bonding. ionic bonds: formed between ions. transfer of electrons occurring. covalent bonds:...
TRANSCRIPT
CHAPTER 8 AND 9
Bonding
TYPES OF BONDSIonic Bonds: Formed between ions. Transfer of electrons occurring.
Covalent Bonds: Molecules form between atoms that share electrons.
Metallic Bonds: Form multiple bonds using the free flowing electrons of a metal.
LEWIS SYMBOLS AND OCTET RULE
Atoms tend to gain and lose electrons until they have 8 valence electrons similar to noble gases. This is a stable electron configuration as shown experimentally by the high ionization energy and low electron affinity.
IONIC BONDING An atom with a low ionization energy reacts with an atom with high electron affinity.
The electron moves. Opposite charges hold the atoms together.
COULOMB’S LAW
Q is the charge. d is the distance between the
centers. If charges are opposite, E is
negativeexothermic
Same charge, positive E, requires energy to bring them together.
FORMING IONIC COMPOUNDS Lattice energy - the energy
associated with making a solid ionic compound from its gaseous ions.
M+(g) + X-(g) ® MX(s) This is the energy that “pays” for
making ionic compounds. Energy is a state function so we
can get from reactants to products in a round about way.
NA(S) + ½F2(G) ® NAF(S) First sublime Na Na(s) ® Na(g)
DH = 109 kJ/mol Ionize Na(g) Na(g) ® Na+(g) + e-
DH = 495 kJ/mol Break F-F Bond ½F2(g) ® F(g)
DH = 77 kJ/mol Add electron to F F(g) + e- ® F-(g)
DH = -328 kJ/mol Formation of solid Na+ + F- ® NaF (s)
DH = -575 kJ/mol Lattice energy Na(s) + ½F2(g) ® NaF(s)
DH = -928 kJ/mol
COVALENT BONDING Electrons are shared by atoms. Ionic and Covalent are extremes. In between are polar covalent bonds. The electrons are not shared evenly. One end is partially positive, the other
negative. Indicated using small delta .d
H - Fd+ d-
H - F
d+
d-
H - Fd+
d-
H -
F
d+
d-
H - F d+
d-
H - Fd+d-
H - F
d+d-
H - F d+
d-
+-
H - Fd+ d-
H - Fd+ d-
H - Fd+ d- H - F
d+ d-
H - Fd+ d-
H - Fd+ d-
H - Fd+ d-
H - Fd+ d-
- +
Electronegativity difference
Bond Type
Zero
Intermediate
Large
Covalent
Polar Covalent
Ionic
Covale
nt C
hara
cte
r d
ecre
ases
Ion
ic C
hara
cte
r incre
ases
COVALENT BONDS
The electrons in each atom are attracted to the nucleus of the other.
The electrons repel each other, The nuclei repel each other. They reach a distance with the lowest
possible energy. The distance between is the bond length.
0
En
erg
y
Internuclear Distance
Bond Length
0
En
erg
y
Internuclear Distance
Bond Energy
ELECTRONEGATIVITY The ability of an electron to attract
shared electrons to itself. Tends to increase left to right. Decreases as you go down a group. Noble gases aren’t discussed. Difference in electronegativity between
atoms tells us how polar.
DIPOLE MOMENTS A molecule with a center of negative
charge and a center of positive charge is dipolar (two poles),
We say that is has a dipole moment. Center of charge doesn’t have to be on
an atom. Will line up in the presence of an electric
field.
H - Fd+ d-
% I
on
ic C
hara
cte
r
Electronegativity difference
25%
50%
75%
LiBr
LiF
HCl
THE COVALENT BOND The forces that cause a group of
atoms to behave as a unit. Why? Due to the tendency of atoms to
achieve the lowest energy state. It takes 1652 kJ to dissociate a mole of
CH4 into its ions Since each hydrogen is hooked to the
carbon, we get the average energy = 413 kJ/mol
COVALENT BOND ENERGIES We made some simplifications in describing
the bond energy of CH4 Each C-H bond has a different energy.
CH4 ® CH3 + H DH = 435 kJ/mol
CH3 ® CH2 + H DH = 453 kJ/mol
CH2 ® CH + H DH = 425 kJ/mol
CH® C + H DH = 339 kJ/mol
Each bond is sensitive to its environment.
AVERAGES single bond one pair of electrons is shared. double bond two pair of electrons are shared. triple bond three pair of electrons are shared. More bonds, shorter bond length.
USING BOND ENERGIES We can find DH for a reaction. It takes energy to break bonds, and end
up with atoms (+). We get energy when we use atoms to form
bonds (-). If we add up the energy it took to break
the bonds, and subtract the energy we get from forming the bonds we get the DH.
Energy and Enthalpy are state functions.
FIND THE ENERGY FOR THIS
2 CH2 = CHCH3 +2NH3 O2+ ® 2 CH2 = CHC º N+ 6 H2O
C-H 413 kJ/mol
C=C 614kJ/mol
N-H 391 kJ/mol
O-H 467 kJ/mol
O=O 495 kJ/mol
CºN 891 kJ/mol
C-C 347 kJ/mol
LOCALIZED ELECTRON MODEL
Simple model, easily applied. A molecule is composed of atoms
that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms.
Three Parts1) Valence electrons using Lewis
structures2) Prediction of geometry using VSEPR3) Description of the types of orbitals
LEWIS STRUCTURE Shows how the valence electrons are
arranged. One dot for each valence electron. A stable compound has all its atoms with
a noble gas configuration. Hydrogen follows the duet rule. The rest follow the octet rule. Bonding pair is the one between the
symbols.
RULES1. Sum the valence electrons.2. Use a pair to form a bond between each pair
of atoms.3. Arrange the rest to fulfill the octet rule (except
for H and the duet).4. Place left over electrons on the central atom.5. If there are not enough electrons to give the
central atom an octet, try multiple bonds. H2O CO2
CN-
FORMAL CHARGE
The difference between the number of valence electrons on the free atom and that assigned in the molecule.
We count half the electrons in each bond as “belonging” to the atom.
Molecules try to achieve as low a formal charge as possible.
Negative formal charges should be on electronegative elements.
SO42-
RESONANCE Sometimes there is more than one
valid structure for an molecule or ion.
Use double arrows to indicate it is the “average” of the structures.
It doesn’t switch between them. Localized electron model is based
on pairs of electrons, doesn’t deal with odd numbers.
NO3-
NO2-
EXAMPLES
XeO3
NO4-3
SO2Cl2
EXCEPTIONS TO THE OCTET1. Odd numbers of electrons
NO2. Atoms with fewer than octet
Be and B often do not achieve octetBF3
3. Atoms with more than octet Third row and larger elements can exceed the
octet SF6
Use 3d orbitals?
EXCEPTIONS TO THE OCTET When we must exceed the octet, extra
electrons go on central atom.
ClF3
XeO3
ICl4-
BeCl2
VSEPR Lewis structures tell us how the atoms
are connected to each other. They don’t tell us anything about shape. The shape of a molecule can greatly
affect its properties. Valence Shell Electron Pair Repulsion
Theory allows us to predict geometry Molecules take a shape that puts
electron pairs as far away from each other as possible.
VSEPR Have to draw the Lewis structure to
determine electron domains. bonding nonbonding lone pair Lone pair takes more space. Multiple bonds count as one domain.
The number of domains determinesbond anglesunderlying structure
The number of atoms determines actual shape
VSEPRElectronDomains
BondAngles
UnderlyingShape
2 180° Linear
3 120° Trigonal Planar4 109.5° Tetrahedral
5 90° &120°
Trigonal Bipyramidal
6 90° Octagonal
ELECTRON DOMAIN GEOMETRY
ELECTRON DOMAIN VS. MOLECULAR SHAPEElectron Domain Just count the number of electron
domains and the geometry corresponds to that number.
Molecular Shape Electron domain geometry is many
times not the shape of the molecule. Here we are worried about just the atoms.
MOLECULAR GEOMETRY Multiple molecular geometries can result
from the same electron geometry.
NON-BONDING ELECTRONS Non-bonding pairs take up more space
than bonding pairs and therefore change the angles within the atom.
MULTIPLE BONDS Multiple bonds create a high electron
density on one side of the molecule which also changes angles.
DIFFERENT MOLECULAR GEOMETRIES Linear Trigonal Planar Bent Tetrahedral Trigonal pyramidal Trigonal bipyramidal Seesaw T-shaped Octahedral Square pyramidal Square planar
POLARITY Individual bonds can have bond dipoles
but the overall molecules polarity may be different.
POLARITY Individual bond dipoles (vectors) can be
added to give the entire molecular dipole.
HYBRIDIZATIONBeryllium: Has no single occupied orbitals therefore should not bond.
If it absorbs energy to excite an electron to 2p it has 2 electrons that can form a bond.
HYBRID ORBITALSThe s and p orbitals can combine to make a hybrid orbital called the sp hybridized orbital.
-sp orbitals have a linear position and is consistent with molecular geometries.-it has 2 degenerate orbitals.
HYBRID ORBITALS Other hybrid orbitals that can be made
include sp2 and sp3. This occurs when there is more than one p orbital overlapping with an s orbital.
SIGMA AND PI BONDS Sigma (σ) bonds overlap head to head
and electron density is in line with the internuclear axis.
Pi (π) bonds overlap side to side and electron density is above and below internuclear axis.
TYPES OF BONDS Single bonds are all sigma bonds. Multiple bonds have one sigma bond
and the rest are pi.
DELOCALIZED PI BONDING Pi electrons are
not localized over one N-O bond but delocalized over all 3 bonds.
This is related to resonance.
MOLECULAR ORBITAL (MO) THEORY In MO theory, we invoke the wave nature of
electrons. If waves interact constructively, the resulting
orbital is lower in energy: a bonding molecular orbital.
MO THEORY In hydrogen the
electrons go into the bonding orbital because it is lower in energy.
The bond order is one half the difference between the number of bonding and antibonding electrons.
For hydrogen, with two electrons in the bonding MO and none in the antibonding MO, the bond order is 2-0/2 = 1
MO THEORY
In the case of He2, the bond order would be
12
(2 − 2) = 0
• Therefore, He2 does not exist.
MO THEORY
For atoms with both s and p orbitals, there are two types of interactions:The s and the p orbitals
that face each other overlap in fashion.
The other two sets of p orbitals overlap in fashion.
MO THEORY
The resulting MO diagram looks like this (Fig. 9.41).
There are both s and p bonding molecular orbitals and s* and * antibonding molecular orbitals.
SECOND-ROW MO DIAGRAMS