Recall:

The Octet Rule

The Octet rule states that since the eight electron configuration _s²_p⁶ is particularly stable, atoms will tend to gain or lose electrons either by donation or by sharing in order to attain the octet configuration.

This rule provides the basis for understanding the structures of an enormous number of compounds.

The main exceptions to the octet rule occur in atoms with less than five electrons and in the transition metals. For the atoms of H, Li, and Be the nearest noble gas is helium.
 
 

Magnetism

There are three main classes of magnetism and they can all be explained by electrons in orbitals. The three main classes of magnetism are, diamagnetism, paramagnetism, and ferromagnetism.

We already know that the electron acts like a tiny magnet. The electron has a north pole and a south pole just like the magnets we are familiar with. We also know that when you try to push two magnets together the north poles repel each other, but the north pole of one magnet will attract the south pole of the other.

An electron alone in an orbital will act like a tiny magnet. If you apply an external magnetic field the electron will try to align itself with the magnetic field.

On the other hand, a pair of electrons in an orbital will not act like a tiny magnet because when the spins are paired the north pole of one electron is pointing the same way as the south pole of the other so that the net north/south effect of electron spin is cancelled.

Paired electrons are diamagnetic. Diamagnetic materials have a very very weak interaction with an external magnetic field. (In fact, they are very slightly repelled by a magnetic field.)

Fe and Ni are examples of ferromagnetic materials.

It was discovered empirically that there are groups of elements that have similar chemical properties.

If you line up the elements in order of atomic number you find that there is a repeating pattern of the properties of the elements. This pattern was not understood until quantum mechanics provided the electron configurations of the elements.

There is periodicity in several properties of the elements:

Size

Ionization energy

Chemical properties

Once we know the electron configurations of the elements we can understand the structure of the periodic table. Elements in the same column have the same outer electron configuration.
 

Bonding and Molecular Structure (Chapter 9)

Valence electrons and core electrons

Core electrons are all electrons in an atom that have the noble gas configuration plus any filled d or f subshells.

Electrons that are "outside" the core are called valence electrons.

Valence electrons are the electrons which are directly involved in chemical bonding. That is, the valence electrons determine the formulas of compounds that can be formed.

The core electrons do not contribute directly to bonding. They have an indirect contribution in that they determine the size of the atom and the size has an influence on the variety of bonds that can be formed.

Examples:

Li [He]2s¹

C [He]2s²2p²

O [He]2s²2p⁴

P [Ne]3s²3p³

Cl [Ne]3s²3p⁵

Ca [Ar]4s²

Cr [Ar]3d⁵4s¹

Br [Ar]3d¹⁰4s²4p⁵

La [Xe]4f ¹6s².

Reactivity and Bonding

We will now begin to use our knowledge of electronic structure (electron

configurations) to formulate principles which will allow us to predict how the elements will combine with each other to form compounds.

The main principles of reactivity and bonding are:

The electrons determine the chemical properties.

Closed-shell (filled shell) inner electrons, core electrons, have only a minor effect on bonding.

Electrons outside of the closed shell core, valence electrons, dominate the chemical properties and bonding.

The eight-electron (octet) valence shell configuration of the noble gases is particularly stable.

One way or another, atoms will try to reach the eight-electron valence electron configuration, either by gaining or losing electrons to form ions, or by sharing electrons to form covalent bonds.

Ions

One of the simplest ways atoms react is to form ions.

Elements near the "edges" of the periodic table can gain or lose electrons to form ions (the ions will then have the electron configuration of the nearest noble gas).

Metals (even the ones in the center of the periodic table) also lose electrons easily to form ions.

On the left side of the periodic table atoms can lose electrons to form positive ions.

[ ]ns¹ ® [ ]⁺ + e-

[ ]ns² ® [ ]²⁺ + 2 e- , etc.

Li ® Li⁺ + e- ,

Li = [He]2s¹, Li⁺ = (1s²)⁺

But it is NOT helium.

Mg ® Mg²⁺ + 2 e^- .

Mg = [Ne]3s², Mg²⁺ = (1s²2s²2p⁶)²⁺

(Al forms Al³⁺ & Sc forms Sc³⁺, etc.)

What noble gas electron configuration do these ions correspond to?

(Recall that ions have a number of electrons which is different from the number of protons.)

That is, the numbers of electrons and protons do not balance.

(The element involved is still determined by the number of protons - this is the atomic number - which does not change in forming an ion).

On the other side of the periodic table (right hand side) atoms gain electrons to form negative ions.

[ ]ns2np⁵ + e^-® [ ]ns²np^6(-)

[ ]ns²np⁴ + 2 e^- ® [ ]ns2np^{6 (2-)}, etc.

Examples:

Br + e^-® Br^-

Br = [Ar]3d¹⁰4s²4p⁵, Br^-= [Ar]3d¹⁰4s²4p^6(-)

(But it is NOT a krypton atom!)

Another example:

S + 2e- ® S^2-

S = [Ne]3s²3p⁴, S^2-= [Ne]3s²3p^{6 (2-)}

The S^2-ion has the same electron configuration as the argon ion.

(N^3-exists, but is less common.)

What is the electron configuration of N^3-?

A reminder:

Positive ions are called cations (they are attracted toward the "cathode.")

Negative ions are called anions (they are attracted toward the "anode.")
 

Transition Metal Ions

In making compounds out of elements that form ions easily the electrons lost by one element must equal the electrons gained by the other. (Compounds are neutral.)

This rule allows us to predict the composition or formulas of a great many ionic compounds.

Try it on a compound of Na and Cl:

Na likes to lose one electron to form Na⁺.

Cl likes to gain one electron to form Cl^- .

Na ® Na⁺ + e^-

Cl + e^- ® Cl^-

Na + Cl + e^- ® Na⁺ + Cl^- + e^- ,

NaCl is an ionic compound.

In ionic compounds the crystal is held together by electrostatic forces (the attraction between the + and - charges).

An ionic crystal is an array of positive and negative ions. In two dimensions this might look like this:

+ - + - + - + - + - + - + - + - +
- + - + - + - + - + - + - + - + -
+ - + - + - + - + - + - + - + - +
- + - + - + - + - + - + - + - + -
+ - + - + - + - + - + - + - + - +
- + - + - + - + - + - + - + - + -
+ - + - + - + - + - + - + - + - +
- + - + - + - + - + - + - + - + -
+ - + - + - + - + - + - + - + - +
- + - + - + - + - + - + - + - + -

There are no NaCl "molecules."

Another example:

A compound of Ca and F:

Ca likes to lose 2 electrons to form Ca²⁺.

F likes to gain an electron to form F^-.

One Ca provides enough electrons for two F atoms.

Ca + 2 F ® CaF₂

Ca + F₂ ® CaF₂

CaF₂ is ionic, i.e., there are no CaF₂ "molecules."

Another example:

A compound of Al and S:

Al likes to lose 3 electrons to form Al³⁺.

S likes to gain 2 electrons to form S^2-.

The electrons lost by the Al atoms must be gained by the S atoms. So it takes 2 Al atoms losing six electron which are donated to 3 S atoms.

The compound of Al and S is Al₂S₃, again the compound is ionic, there are no Al₂S₃ molecules.
 
 

Lewis Dot Symbols and Formulas

G. N. Lewis invented a way to visualize how the octet ruled worked by using what are now called "Lewis dot formulas".

In Lewis dot formulas the atom is represented by its symbol and the valence electrons are represented by dots placed around the symbol of the element.