Ionic Bonding + 

Valence Electrons An atom can have many electrons. How many of them take part in chemical bonding? The number can vary, but in all cases the same principle is in effect: The electrons in the highest occupied energy level of an element’s atoms are the ones involved in bonding. These electrons are called valence electrons.

Determining Valence Electrons The number of valence electrons of an atom is usually obvious from examining the electron configuration of the element. The valence electrons determine the chemical properties of the element.

The Group Number and Valence For representative elements (s and p blocks) the number of valence electrons is equal to the last digit of the group number: Group 1 = 1 Group 2 = 2 Group 13 =3 Group 14 = 4 Group 15 = 5 Group 16 = 6 Group 17 = 7 Group 18 has no valence electrons, because they are not generally available for making chemical bonds

Electron Dot Structures Valence electrons are often represented as dots surrounding the element’s symbol. For example, going across period 2, All of the elements of a given group (except He) have the same number of electron dots in their structures.

The Octet Rule In 1916, Gilbert Lewis (UC Berkeley) came up with a simple principle to explain why atoms form many kinds of ions and molecules – the octet rule. The octet rule – atoms in compounds tend to have the electron configuration of a noble gas.

How the Octet Rule Works Remember, each noble gas (except He) has the electron configuration ns2 np6 in its highest energy level. Metallic elements obey the octet rule by losing electrons. The loss of valence electrons produces a positive ion – a cation. Non-metallic elements obey the octet rule by gaining electrons to become negative ions – anions Non-metals can also share electrons with another non-metals to satisfy the octet rule, making covalent bonds.

Some Examples of the Octet Rule in Action Al and Br2  AlBr3 Na and Cl2  NaCl In both cases, the metal donates electrons to the non-metal. The metal forms a positive ion (cation) to form an octet The non-metal forms a negative ion (anion) to form an octet

Ionic Bonding Worksheet Complete the following table by providing the electron configurations for the outermost energy level, the number of valence electrons, and the electron dot diagrams for each of the elements given.

Electron Configuration Electron Dot Diagrams Element Electron Configuration Valence e- Electron Dot Lithium Nitrogen Silicon Bromine

Electron Configuration Write the formula and the complete electron configuration for each of the following: Formula Electron Configuration Sodium ion Fluoride ion Potassium ion Strontium ion Sulfide ion

Atom Ion Na, Na ion Cl, Cl ion P, P ion Ca, Ca ion Write the electron dot structures for each of the following atom-ion pairs. Atom Ion Na, Na ion Cl, Cl ion P, P ion Ca, Ca ion

Representing Ionic Compounds with Lewis Diagrams Lewis structures are useful tools to keep track of the valence electrons in elements and compounds. In sodium fluoride, we can represent it in a number of ways: - + Brackets are often used; The octet from the lower energy level of the metal is often left out [Na]+[ ]-

Write the electron dot formula for each ion in the ionic compound: Chemical Formula Na, F Mg, Cl Ca, S

K, O Al, Br Rb, N Al, O

Properties of Ionic Compounds Solids at room temperature Very high melting points (>> 300◦C) Do not conduct electricity in solid state Many are highly soluble in water Solutions of ionic compounds are electrolytes – they conduct electricity

Ionic compounds do NOT form molecules! NaCl – formula unit water – H2O – molecular formula Chemical structure of ionic compounds are known as formula units – the simplest whole number ratio of the ions.

The Octet Rule in Action sodium, (2-8-1) or 1s2 2s2 2p6 3s1 when it loses its electron, it has an octet in its highest (2nd) energy level. 3s  loss of valence electron 2p          2s    1s    +

Electron Configurations of Anions When non-metal atoms gain valence electrons, they become negative ions - anions. For fluorine (and all group 17 elements) form halide ions: 1s2 2s2 2p5 + e-  1s2 2s2 2p6 an octet in n=2! 2-7 + e-  2-8 2p       2s  + e-   1s   -

Models of Ionic Compounds