CHEM 1411 - General Chemistry I

Course (Lecture and Lab) Syllabus

Summer 2007

 

 

Instructor: Dr. Shawn Amorde

e-mail address: samorde@austincc.edu

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

Office Hours: TThF, 7-8am

 

 

 

Chemical Bonding

 

Text Chapters (7.1-7.3, 7.5-7.12)

Recommended Problems (2, 4, 8, 14, 18, 20, 24, 30, 34, 42, 46, 52, 56, 74, 78)

 

            Chemical Bonding; Ionic Bonds

 

1.      Lewis Dot Formulas

2.      Formation of Ionic Compounds

 

Chemical Bonding; Covalent Bonds

 

3.      Formation of Covalent Bonds

4.      Lewis Formulas for Molecules

5.      Writing Lewis Formulas: The Octet Rule

6.      Formal Charges

7.      Limitations to the Octet Rule

8.      Resonance

9.      Polar and Non Polar Covalent Bonds

10.  Dipole Moments

11.  Range of Bonding Types

 

 

Chemical Bonding

 

We have just discussed the forces that hold elements together, now let us discuss the forces that hole molecules together.  There are two major classes of bonding, first ionic bonding results from electrostatic interactions among ions.  This type of bonding often results from the over net transfer of electrons from one atom to another.  Second, covalent bonding results from sharing electrons between atoms.

 

 

These bonding classes represent the extremes!!  All bonds have some of both bonding characteristics.  Just like our elements have different properties due to their outer most electrons, molecules have different properties bases on the type of bonding between the atoms.

There is a nice chart in the book on these properties, page 251.

 

 

 

 

 

 

 

 

 

Lewis Dot Formulas

 

We just learned the ground state electron configurations for elements, these electrons determine the chemical and physical properties of the elements and they determine the number and type of chemical bonds they form.  We write these electrons as Lewis dot formulas. 

Chemical Bonding involves the valence electrons, which are usually electrons in the outermost occupied shells.

We only write the outer most s and p orbital electrons!!

The transition metals have too many electrons to use this method!!

Let us start with writing Lewis dot formulas for a few elements.

 

 

 

 

 

 

 

 

Ionic Bonding

 

Remember that an ion is an atom with a charge!!

 

A positively charged ion is                                        A negatively charged atom is

 

 

 

 

 

 

A cation                                                                      An anion

 

 

 

Ionic bonding is the attraction of oppositely charged ions (cations and anions).

 

This type of bonding usually results in the formation of a solid like table salt.

Let us review electronegativity and ionization energy and electron affinity.  Let’s review the trends in the Periodic Table.

 

 

 

 

 

 

 

 

 

When the difference in electronegativity between two elements is large, the elements are likely to form an ionic bond.

Let us consider the Group 1A and 7A elements.  Let’s draw the electronic configurations for one of each.

 

 

 

 

 

 

 

 

 

If you form ions both atoms are now isoelectronic with a noble gas!!

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Remember the table?  The greatest difference in electronegativity is across the table, the greater this difference the more ionic the bond will be.

 

The energy associated with the attraction between opposite charges in an ionic solid, is called the crystal lattice energy for that solid.

 

 

 

 

 

 

 

 

 

 

 

What about Group 1A with 6A?

Let us write it out.

 

 

 

 

 

 

 

 

 

 

 

 

And Group 2A and 6A?

 

 

 

 

 

 

 

 

 

 

 

 

 

What about our transition metals?

Remember that the 4s orbital fills before the 3d orbital, and they are conversely the first electrons lost while forming ions and bonding.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Simple ions rarely have a charge greater than +3 or -3!! 

 

 

Covalent Bonding

 

We just discussed atoms bonding together with a large difference in electronegativity, but we also see atoms bonding together without a big difference in electronegativity.  This type of bonding is called covalent bonding.  The official definition of a covalent bond is when two atoms share one or more pair of electrons.

 

 Let us look at what it really means to “share” a pair of electrons.  We’ll start with H.  Let us draw out the electronic configuration of two hydrogen atoms.

 

 

 

 

 

 

 

 

 

 

 

 

The energy associated with “sharing” a pair of electrons is distance dependant!!

 

 

Let us draw this again including the s orbitals for both H atoms.

 

 

 

 

 

 

 

 

 

 

Remember we have competing forces at work here, attraction (+ & -) and repulsion (- & -, + & +)!!

 

Other nonmetal atoms also share electrons to form covalent bonds.  Most covalent bonds involve sharing 2, 4, or 6 electrons….i.e. 1, 2, or 3 pairs of electrons.

The number of pairs shared denotes the type of bond.  For 1 pair of shared electrons we call this a single bond (sigma bond), for 2 pairs we call this a double bond ( pi bond) and for 3 pairs we call this a triple bond (pi bond).  Sharing pairs of electrons is possible if the orbitals of the single atoms overlap and essentially form a new orbital.  This entire theory of bonding is called valence bond theory, or more commonly molecular orbital theory.

 

Let us draw this again.

 

 

 

 

 

 

 

 

 

Remember that there is a magic distance between the nuclei of bonded atoms and at that distance the atoms are “more stable”.  The statement “more stable” means the energy associated with the single atoms is more than the energy associated with the bonded atoms, there is a specific amount of energy associated with each bond!!  This is called the bond energy.

 

 

 

 

 

 

 

 

 

  Let’s not forget Lewis dot formulas!!  Let’s write a couple for covalently bonded molecules.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Let’s take a look at some more complicated molecules, including atoms with shared and unshared electrons.

 

How do we decide which electrons are shared between which atoms?  What is there are more than two atoms? 

 

Some rules;

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Formal Charge

 

The formal charge is a hypothetical charge on an atom in a molecule.  This is an extension of the concept of sharing electrons.  We must look at the overall total number of electrons being shared in a molecule and calculate the formal charge.  Here are some rules to follow;

 

FC = (group number)-[(number of bonds)+(number of unshared electrons)]

 

In our Lewis dot formula if an atom has the same number of bonds as its group number it has a FC of zero.

 

The sum of the FC in a molecule is always zero!!

In a polyatomic ion, the sum of the FC is the FC.

 

Let us write out some examples.

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Limitations to the Octet Rule

 

There are a few elements which do not fill the octet rule.

 

  1. Be usually only forms two bonds with its two valence electrons.
  2. The Group 3A elements tend to form only three bonds, especially B, and share six instead of eight electrons.
  3. Compounds or ions with an odd number of electrons.
  4. Compounds or ions who need to share more than eight electrons to hold all the available electrons.

 

These species are usually very reactive!!  For example, B forms BH3 and has an empty orbital available.  This molecule readily will accept a pair of electrons to share and form an ion, with the formal charge of +1.  This makes BH3 a Lewis acid. 

 

 

 

 

 

 

 

 

 

 

 

Resonance

 

A molecule in which two or more Lewis formulas with the same arrangement of atoms can be drawn is said to exhibit resonance.

 

This is usually seen with double or triple bonds!!  For example, CO2;

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Experimentally the CO bond length is between a single and a double bond!!  This means the electrons are actually shared between all of these bonds and are delocalized, giving a hybrid bond between for all three bonds!!

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Polar and Nonpolar covalent bonds

 

Remember covalent bonds are between atoms with similar electronegativities, but not always the same!!  This means the electrons being shared may be “pulled” toward ( or attracted to) the more electronegative atoms and not shared entirely equally.  This bond is called a polar covalent bond.  The term dipolar means “two poles” for us that refers to a positive and negative “pole”. 

Let’s look at an extreme case, HF

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

 

Dipole Moments

 

We can actually measure the polarity of a molecule, this number is called the dipole moment.  The dipole moment has this equation;

 

 

 

 

 

 

 

 

In general dipole moments are difficult to measure for each bond in polyatomic molecules and usually reflect the entire molecule.  We can write the dipole moments for poly atomic molecules as the sum of the individual dipole moments. 

For example, H2O