Transcript Chemical Bonding I: Lewis Theory

Chemical Bonding I:
Lewis Theory
How Valence Electrons Arrange
Themselves to Give
Chemical Bonds
Basic Ideas from the Previous
Chapter...
• Valence electrons are the electrons that
participate in chemical bonding.
• Configurations that result in inert gas cores or
half-filled or fully-filled subshells are
particularly stable.
Types of Chemical Bonds
Types of Atoms
Type of Bond
Characteristics of Bond
Metal & nonmetal
Ionic
Electrons transferred
Nonmetal & nonmetal
Covalent
Electrons shared
Metal & Metal
Metallic
Electrons pooled
Some pointers...
• Ionic bonds: Cations and anions form and are
held together electrostatically.
• Covalent bonds: Electrons are shared by atoms
(they travel back and forth between the atoms).
The most stable configurations have the electrons
between the nuclei (draw on board, Johnston!).
• Metallic bonds: Electrons move through the
entire crystal lattice of the metal.
Attraction & Repulsion
(Explain in detail!)
Electrostatic Repulsion
Electrostatic Attraction
Some examples of bonding types...
We can represent valence electrons as
dots...
Details—ionic bonds...
• Atoms either lose electrons to form cations
• or
• gain electrons to form anions.
• They try to attain an inert gas core, if possible!
For instance, consider potassium...
Octets...
• Except for H and He, most atoms try to get an
octet of electrons. (H makes “duets.”)
• This is called the “octet rule.”
• It holds true with just a few exceptions:
– Elements such as Be and B
– Elements with available empty d-orbitals can take
on more.
– We shall discuss all this soon when we get to
Lewis dot structures!
Anions & Cations Combine to Form a
Crystal Lattice
A quick aside...
• We can break up the process in easy to
visualize steps.
• The energies of the steps are additive.
• This is a sneak peak at the First Law of
Thermodynamics.
• We shall not say anything more here other
than energies add the same way as masses or
lengths!
The Born-Haber Cycle is the
vizualization!
Lattice Energy
• This is the main provider of energy to stabilize
the crystal!
• However, note that the electron affinity also
sneaks in her!
Trends in Lattice Energy (Size)
Ion Size
Examples
Metal Chloride
Lattice Energy
(kJ/mol)
LiCl
-834
NaCl
-788
KCl
-701
CsCl
-657
Trends in Lattice Energy (Charges)
Ion Charge
Examples
Compound
Lattice Energy
(kJ/mol)
NaF
-910
CaO
-3414
Just for pHun...
(The release of lattice energy)
• Let’s look at the reaction of sodium and
chlorine.
• All the steps occur together, of course.
• Be, here, you get to see lattice energy in all its
glory!
• Here is the
link...http://www.youtube.com/watch?v=Mx5JJWI2aaw
This is in accordance with Coulomb’s
Law...
Some Effects: Ions in Solution
A little demonstration...
Ionic Compounds by Themselves
Don’t Conduct Electricity
Ionic Compounds in Aqueous
Solution DO Conduct!
Ionic Compounds are Called “Strong
Electrolytes”
• Let’s look at a presentation by a kindly old
professor (as contrasted to your mean old
professor)...
• Here is the link:
http://www.youtube.com/watch?v=1XWnovm6JLs
A Short Critique of the Ionic Bond
Model...
• A great deal can be explained by assuming
simple electrostatic attractions/repulsions.
• These ideas allow one to develop methods to
calculate lattice energies of many types of
ionic compounds.
• The concept of electrolytes is introduced.
• In reality, however, there is no such thing as a
pure ionic bond! We shall discuss this shortly.
Covalent Bonding
• Here, we shall start with the Lewis model.
• We shall look first a just simple octets (and
duets) as done in organic chemistry.
• The model will have to be extended for some
compounds of elements in the 3rd row (and
further) in the periodic table.
• But, for now, we note that organic chemistry is
different than other endeavors→→→
Yes, organic chemistry is different...
Drawing Lewis Structures—Some Rules
• Here, we stick to just octets and duets.
• Extended structures will come later.
• Steps to follow:
– Count the valence electrons
– If you have ions, make allowances for this
– Distribute electrons so that H has duets and other
atoms have octets
– Exceptions to these will be discussed later
We construct a few on the board—I
shall construct as we go
Some terminology
Double & Triple Bonds
We shall construct some of these on the board...
O2
N2
CO
C2H2
Acetone
Benzene
Power of the Lewis Model
• Very good at showing what can and cannot
form!
• Predicts directionality.
• Works very well with organic compounds.
• Exceptions are interesting and will be
discussed later!
Electronegativity & Polar Bonds
• In all covalent bonds, electrons are shared.
• However, some elements attract electrons
more readily than others.
• This property is called electronegativity.
• Bonds formed this way are called polar bonds.
• Polar bonds have dipole moments.
HF: An example of a polar bond
(explain picture as we go)
How we can detect this property...
Trends in electronegativity
(back to the good old periodic table)
Note the values...
• 𝝌 = 4.0 for F
• 𝝌 = 3.5 for O
• These two are the most electronegative
• 𝝌 = 0.7 for Cs
• This is the least electronegative.
• Note that 𝝌 = 0 does not happen.
Bond Polarities
• Bond polarity depends on the
electronegativity differences between the
bonded atoms.
• We shall look at three cases on the next three
slides.
Identical Atoms: Always nonpolar
Ionic Bonds: The most polar
A very polar covalent bond
This table is useful...
Another way to look at these...
The dipole moment...
• When two atoms are separated by a distance,
r, equal and opposite charges of size q are
present on each atom.
• We usually represent this as the dipole
moment, μ.
• μ is in units of Debyes.
• Details are on the next slide...
Details about μ (“mu”)
A Clarification...
• Suppose we have two
ions with q = +e and -e.
• e=
1.60217653 x 10-19C.
• That is, these are ions
of charge +1 and -1.
• Let them be separated
by 100 pm.
• Then→→→→...
To get μ for a pure ionic bond...
Just multiply the bond
length in pm (divided by
100) times the value in
the previous slide!
This is easy to do (I hope)
since I have rounded
things!
We show this now→→→
We can thus define the % ionic
character of a bond!
Example from the book...
r = 130pm
∴ μionic = 4.803 x 130/100
= 6.24D
In the example,
μobs = 3.5D
So............
Some typical examples...
(Note the ∆𝝌 values!)
Some more examples...
Now the real pHun!
Writing Lewis Structures!!!
1) Write the correct skeletal structure for the
molecule.
2) Calculate the total number of electrons by
summing the valence electrons of each atom.
(Be sure to take ions into account!)
3) Distribute the electrons among the atoms giving
octets to all atoms other than H (duet for it).
4) If any atoms lack an octet, form double or triple
bonds as necessary.
On-the-board examples
(“Explain as we go” mode!)
CO2
NH3
NH4+
C2H6
C2H4
C2H2
N2O
NO2CN-
Resonance & Formal Charge
• Sometimes there are more than one possible
equivalent Lewis structures.
• In this part of the lecture we shall discuss the
phenomenon of resonance.
What is meant by “equivalent”
Two structures are equivalent if they can be
converted from one to the other by
a simple rotation of the entire molecule
or
a simple reflection of the entire molecule.
The relative positions of the electrons are left un
changed!
A very simple example: Ozone
The true structure is a “resonance
hybrid” (explain verbally)
More about this...
• Neither structure exists independently.
• The true structure is a linear combination of
the separate structures.
• This is an example electron delocalization.
Multiple Equivalent Structures are
Possible (e.g., nitrate)
Are nonequivalent Lewis structures
possible?
• Yes!
• It is possible to draw alternate structures that
obey the rules but are not equivalent.
• I shall draw CO2 structures as an example.
• This means that we have to choose a best
Lewis structure (or best set of resonance
structures).
• What is the key to this? →→→→
Formal Charge!
(Rules below)
Draw the structure.
Be sure to take into account anions and cations.
Start with the atom’s group #.
Subtract 1 for each electron completely “owned by the
atom.”
Subtract 1 for each pair of electrons shared by an atom.
This procedure gives the formal charges of each atom.
A simple first example: HF
Same for hydrogen. Here is the
summary...
Selecting the best Lewis structure
1) The sum of all formal charges on the atoms
must equal the total charge of the species.
2) The best structure always has the differences
in the formal charges minimized.
3) When formal charge cannot be zero, negative
formal charge should reside on the most
electronegative atom.
4) Now, we are ready for CO2!
Three CO2 structures!
•
•
•
•
We draw these on the board.
All the structures are “legal.”
But, one is obviously the best.
The minor structures, however, do contribute
to the calculated wave function.
• A bonus: We analyze also OCN-.
Exceptions to the Octet Rule
①Odd-Electron Species (“free radicals”)
②Incomplete octets (forced by formal charge)
③Expanded octets (when extra electrons can
be accommodated by d-orbitals)
Odd-electron species
• These do occur in nature but tend to be
unstable.
• We look at a few examples
– NO
– NO2 (compare this to N2O4)
– The t-butyl free radical
Incomplete Octets
• Be & B are strange!
• We explain BeCl2.
• BF3; why F cannot have a double bond and,
thus, B is forced into an incomplete octet.
• The very strange and curious case of BH3!
Expanded Octets
• Molecules such as AsF5, SF6, PCl5, and many
similar such exist.
• How do we accommodate these?
• We put electrons into vacant d-orbitals!
• We shall show how to do this with several
examples momentarily.
• In some cases, we shall also have to invoke
formal charges and resonance structures!
Some Verbal Examples
i.
ii.
iii.
iv.
v.
SF6
AsF5
SF4
XeF2
The sulfate anion (lots of resonance
structures and a BIG surprise!)
Bond Energies & Bond Lengths
• The chemical bond can be treated as a distinct
entity.
• It is a very powerful concept!
• We can assign properties to bonds.
1)
2)
3)
4)
Bond Energies
Bond Lengths
Bond Angles (next chapter)
Bond dipole moments (discussed in passim)
Bond Energies
• Chemical reactions (we have to restrict
ourselves to the gas phase here) can be
thought as occurring by the breaking of old
bonds and the forming of new bonds.
• Typical bond energies are shown in the next
slide.
• Note that these bond energies are averages
since they can—and do—vary slightly in
different molecules.
Humongous Table!
General Rules
• Bond breaking is endothermic; it takes energy
to break a bond!
• Bond forming is exothermic; you get energy
back if a bond is formed.
• Usually, if a net process is exothermic, the
reaction is favored.
• The next slide shows the equation...
The Equation
Two ways to do this...
• Brute force: Break ALL the bonds in the
reactant(s) and then form ALL the bonds in
the product(s).
• Finesse: Just look at the actual old bonds
broken and the actual new bonds formed.
• Use whichever way works best for you!
We shall now give YOU
the opportunity to do
some by yourself.
I shall then explain them!
Bond Lengths
• These are defined as the distance between the
nuclei of the bonded atoms.
• As with bond energies, these are averages
since there are slight variations according to
the molecular structure.
• The next few slides give some typical values.
• Nowadays, we use pm and the length unit.
• Before that, we used the Ångstrom
(1Å = 10-10m).
Another humongous table...
A quick note on trends...
• For a given atom pair,
single bonds are longer
than double bonds.
• And, of course, triple
bonds are shorter than
double bonds.
Bonding in Metals
• A common model is the
“electron sea model.”
• Sometimes, this is
called a “Fermi gas.”
• The electrons are
delocalized over the
enter metal chunk.
• Paired electrons can be
very far apart!