Recall:

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.
 
 

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.

Examples:

Li ® Li⁺ + e^- ,

Mg ® Mg²⁺ + 2 e^- .

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

Examples:

Br + e^- ® Br^-

S + 2e^- ® S^2-

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^- ,

or

Na + Cl ® NaCl

NaCl is a compound (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:

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

However, crystals are really three-dimensional arrays of ions.

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.

What holds the ionic crystals together?

The octet rule turns out to be the main guide for determining how atoms come together to form compounds. G. N. Lewis invented a way to visualize how the octet rule 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.

H He Li Be B C N O F Ne Na Mg Al Si P S Cl Ar.
 

Covalent bonds (sharing electrons)

Elements that don't make ions easily can obtain the noble gas configuration by sharing electrons.

H· ·H gives H:H

Lewis dot formulas were invented to help us show how electrons are shared.

· = valence electron

- = shared pair = :

write H₂ as H:H or H-H.

(except for H, Li, Be, and possibly B).