CHEM 1411 - General Chemistry I

Course (Lecture and Lab) Syllabus

Fall 2007

 

 

Instructor: Dr. Shawn Amorde

e-mail address: samorde@austincc.edu

web page: http://www.austincc.edu/samorde

Office Hours: TThF, 1-2pm TTh, F 12-1pm

 

 

 

Molecular Structure and Covalent Bonding Theories

 

Text Chapters (8.1-8.15)

Recommended Problems (2, 6, 8, 11, 14, 22, 24, 25, 26, 29, 34, 38, 44, 48, 52, 54, 55, 58, 64, 85, * You should write out the Lewis structure, central atom hybridization, electronic and molecular geometry, polarity of each bond, molecular polarity, and if you would expect the molecule to vary from the predicted molecular geometry and why for each molecule in problems 22, 24, 34, 38, 44, and 48)

 

            Covalent Bonding Theories

 

1.      Valence Shell Electron Pair Repulsion Theory

2.      Polar Molecules and their Molecular Geometry

3.      Valence Bond Theory

 

Molecular Shapes and Bonding

 

4.      Linear Electronic Geometry

5.      Trigonal Planar Electronic Geometry

6.      Tetrahedral E. G.

7.      Trigonal Bipyramidal E. G.

8.      Octahedral E. G.

9.      Compounds with Double Bonds

10.  Compounds with Triple Bonds

 

 

Covalent Bond Theories

 

We have just discussed ( in chapter 7) how to arrange the atoms in a molecule on paper in 2-D using Lewis dot formulas, essentially which atoms are connected to which atoms and with how many bonds.  Now let us determine the spatial arrangement of the atoms in 3-D including, the shape and energy of the bonding orbitals.  We will rely on two theories and one physical property to determine the shapes and relative energies of these bonding orbitals. 

 

 

 

 

 

Valence Shell Electron Pair Repulsion Theory

 

This is a general rule for the arrangement of electrons in covalently bonded molecules.  The rule states;

Every pair of electrons around the central atom in a molecule will be as far away from the other pairs of electrons as possible.

·         In simple terms every bond and every lone pair will be placed as far apart as possible around the central atom and still have the orbital overlap necessary for bonding.

·         The number of atoms bonded to the central atom and the number of lone pairs on the central atom dictate the 3-D orientation of the groups in space. 

·         There are a set of general shapes designated for each number of atoms and lone pairs on each central atom.

 

 

For example, 

 

 

 

 

 

 

 

 

 

Electronic Geometry for each Number of Atoms or Lone Pairs on each Central Atom

 

2                                            Linear, 180o                      

3                                            Trigonal Planar, 120o

4                                            Tetrahedral, 109.5o

5                                            Trigonal Bipyramidal, 90o, 120o, 180o

6                                            Octahedral, 90o, 180o

 

 

 

Valence Bond Theory

 

The valence bond theory is guideline to explaining how bonding works.  We have already discussed that bonding needs to have overlapping orbitals, but what if the orbitals are not in the same energy level?  The same shape?  Can we still bond atoms together?  Yes, we know this it true by all the compounds we have isolated. 

 

When atoms come together to form molecules the valence orbitals from both atoms hybridize, that is they kind of melt together and take on an average shape for all orbitals involved. 

 

This process is called hybridization and these orbitals are called hybrid orbitals.

We have a system for labeling these orbitals to designate the number and type of orbitals which have hybridized to form the bonding orbitals in molecules.  Again, there is a general guideline which associates the shape of the orbital to shape of the molecules.

 

For example,

 

 

 

 

 

 

 

 

 

 

 

 

 

Here is a handy chart for you to use when predicting shape and hybridization;

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Molecular Polarity

 

Recall that bonding between atoms with different electronegativities results in a polar bond, this polar bond can lead to a polar molecule and a dipole.  We need to examine the polar forces at work in each molecule to determine the overall dipole moments.  The molecular geometry is influenced by polarity and therefore we need to add this bit of information to our 3-D models of orbital hybridization.

 

For example,

 

 

 

 

 

 

 

 

 

 

 

 

Individual dipole moments can add and subtract within a molecule, we will use electronic vectors to show this relationship.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Now, let’s put all of these factors together and draw examples of each of these in molecules.

 

 

 

 

 

 

Molecular Structure and Covalent Bonding Theories Cont.

 

Text Chapters (7.1-7.12, 8.1-8.15, 4.4)

Recommended Problems ( )

 

            Lewis Dot Formulas

1.      Expanded Valence, beyond the octet rule

2.      Ions

3.      Formal Charge

4.      Oxidation Number

5.      Nomenclature

 

Beyond the Octet Rule

 

We have just discussed the valuable theory of the octet rule, which is extremely useful in predicting the bond connectivity and geometry of molecules.  However, the octet rule is a general guide line and we now need to expand our rule book to include some more guidelines for compounds that do not fit the octet rule.

 

Period 2 elements tend to follow the octet rule almost exclusively, but Period 3 elements can accept more than 8 electrons in their outer valence shells.  These elements usually share 8, 10, or even 12 electrons.  For example, let us look at some P compounds.

 

 

 

 

 

 

 

 

 

 

 

 

 

This is called an expanded valence.  There are two general theories on why the Period 3 elements and beyond expand their valence shells.  The first is that the presence of the d orbitals with in range of bonding is responsible for the “overflow” of electrons.  The second is that “size matters”.  Essentially the Period 2 elements are more tightly packed together than the Period 3 elements and beyond, therefore the geometry of fitting more electrons around a bigger molecule is a little easier for the Period 3+ elements.

Recent studies have shown that the later argument is probably closer to what is =really happening.

 

Let’s look at Sulfur

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Now what do we do?  We can draw two structures for this molecule.  Which one is more accurate?  Lewis dot formula rules do not help us here.  This is where the concept of Formal charge can help us.

 

Let’s draw both structures again.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Formal charge is a method of bookkeeping for electrons!!  Becareful here!!

Formal charge is only for covalently bonded atoms and the total charge must be zero!!

Formal charge is based on sharing pairs of electrons!! 

Formal charge is assuming perfect covalent bonding in which one electron from one atom is shared with one electron from another atom, evenly.

This is an exaggerated covalent character, in which the structure with the lowest charges wins.

 

 

We have already seen there are at least 2 equations for calculating formal charge on an atom.  Here is much easier way!!

 

 

 

 

 

 

 

 

 

 

 

 

 

In general if your atom is “sharing” more than just what is normal for the element, you have a positive charge, and vice versa for negative.

 

Here is a chart to compare what is “normal” for some atoms.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Let us look at the ammonia example again.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Ions

 

Let’s now draw some Lewis dot formulas for some common ions.  How about NO3-?

Remember, total number of available electrons includes ions charges!!

So, Let’s write it out.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Here is the second “helper” to the Lewis dot formulas,  we have no way to predict which bonds are double and single in this ion.

Real life experiments tell us the bond lengths are all the same here!!  This is due to resonance.  We can now write this ion as resonance structures with a double headed arrow, or as the dotted bonds.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Oxidation Number

 

We have just discussed the “exaggerated” covalent bonds and using Formal Charge as a bookkeeping method.  The same concept is needed for our ionic bonds, this is called oxidation number.

 

Oxidation number is the charge an atom would have if the more electronegative atom took both electrons completely!!

 

What about NO3?

 

 

 

 

 

 

 

 

 

 

 

 

And MgO

 

 

 

 

 

 

 

 

NaH

 

 

 

 

 

 

 

 

 

NaF

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Let’s revisit Molecular Orbital Theory

 

Now that we have seen atoms that can bond to more than four atoms we need to predict their shape.

 

Our hybridized orbitals now include the d orbitals.  Therefore, past four bonded atoms or lone pairs we now have sp3d and sp3d2 hybridization.

 

The 5 atoms sp3d takes the shape of trigonal bipyramidal.

All the same rules apply, only now we have two angles to look at and a choice for lone pair placement.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Remember the lone pairs repel more than bonded atoms, so the require more space.  We have a choice in where to place them…..

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

What about sp3d2?  Octahedral, SF6

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

What about lone pairs?

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Rules and Steps for Lewis Dot Formulas

 

·         In most compounds the elements in each atom achieve the isoelectronic configurations of the closest noble gas

 

  1. Figure out which element is the least electronegative, except H.
  2. Hydrogen is always terminal and can only form one bond.
  3. Place the least electronegative atom in the center, this is called the central atom.
  4. Even space the remaining atoms around the central atom, keeping in mind the following guidelines;

 

 

H always only forms one bond, so it should be terminal

Atoms tend to form (8 – group number) of bonds, so….

Halogens (Group 7A) usually form one bond

Oxygen (Group 6A) usually forms 2 bonds

Nitrogen (Group 5A) usually forms 3 bonds

Carbon (Group 4A) usually forms 4 bonds

Boron (Group 3A) usually forms 3 bonds (does not follow the octet rule)

Beryllium (Group 2A) usually forms two bonds (does not follow the octet rule)

Lithium (Group 1A) usually forms one bond (mostly ionic bonds though!!)

 

Non metals can form single, double, or triple bonds.

Oxygen forms single or double bonds.

 

  1. Add up the total number of electrons needed for each atom to achieve noble gas configuration, N
  2. Add up the total available electrons for each atom, A
  3. Subtract A from N, this is the total number of shared electrons, S
  4. Divide S by 2, this is the total number of electron pairs
  5. Start by placing the shared electrons as single bonds around the central atom
  6. Place any remaining single bonds to atoms not bonded to the central atom
  7. How many bonds do you have? Subtract this from the total (S).  If zero then go to step 14.  If not go to step 12.
  8.  Check the octet rule for each atom, starting with the central atom.
  9.  Fill in double or triple bonds where needed
  10.  Have all the pairs been used?  Check the number of available electrons.  Subtract the number of shared electrons from the available electrons (A-S), these “extra” electrons are used to fill in lone pairs.
  11. Fill in the “extra” electrons as pairs around the appropriate atoms. (i.e. O)
  12. Now Check your drawing does it make sense?

 

 

What if you can draw more than one double or triple bond and still satisfy the octet rule and fit the Lewis dot guidelines?

 

We use Formal Charge for neutral atoms only!!  This is not for ions!!

 

There are two general formulas to use to calculate FC on each atom.

 

  1. FC = (group number) – [(number of bonds) + (unshared electrons)]
  2. FC = group number – number of lone pairs – ˝ number of bonding pairs

 

 

Let’s fix the example from yesterday

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

These charges help us pick the “best” placement for the double bonds.  They can also mean the actual molecule shares these electrons equally among the bonds you are comparing, this is called resonance.

 

Here are some examples.

 

 

 

 

 

 

What about the book examples? 

 

There is an additional way to look at formal charge.  Formal charge helps you keep track of which electrons are being shared as pairs from where……in other words we have been considering bonds as one electron from each atom forms a bond, this is not always the case, for example your book describes the ammonium ion as an example of formal charge.

What does this really mean?