1. IONS 7.1 Valence Electrons, The Octet Rule, and formation of Cations and Anions

2. Think about it… The most common example of an ionic compound is Sodium Chloride (table salt) What is it about sodium and chlorine atoms that cause them to combine? Why is the formula unit for sodium chloride NaCl and not Na 2 Cl? or NaCl 2 ?

3. Valence Electrons Elements within each group of the periodic table behave similarly because they have the same number of valence electrons Valence Electrons – the electrons in the highest occupied energy level of an element’s atoms. This can be determined by looking at the element’s electron configuration

4. For representative elements (not the d or f block), the number of valence electrons is the same as the group number Ex. Group 1A elements (H, Li, Na, K, etc.) have 1 valence electron. Carbon and Silicon (in Group 4A) have 4 valence electrons. Nitrogen and Phosphorous have 5 valence electrons. They are in group 5A. If you count noble gases as group 0, they are the exception. They have 8 valence electrons. Another way to figure out the number of valence electrons is to count the position of the element in it’s row on the periodic table. For example, Lithium is 1 st in in it’s row. It has one valence electron. Fluorine is 7 th in it’s row. It has 7 valence electrons.

5. Dot structures Valence electrons are usually the only electrons used in chemical reactions. Sometimes it’s helpful to see what electrons are doing in chemical reactions by drawing models. We show these electrons using electron dot structures (or also called Lewis Structures) Electron Dot Structures – diagrams that show valence electrons as dots surrounding the symbol of the element

6. Ex. Sulfur Electron Configuration 1s 2 2s 2 2p 6 3s 2 3p 4 Valence Electrons 6 Dot Structure is shown at right When writing a dot structure: Dots are placed around the element’s symbol with a maximum of 2 dots on each side The exact location of the dots is not critical, but we put on dot on each side first before pairing them up. It’s good to show them in pairs usually, but there are exceptions when you draw several together in a bond. We’ll look at those later.

7. Write the electron dot structures for the following. You can check your book for the answers. Fluorine Magnesium Phosphorous Krypton Oxygen Nitrogen Chlorine

8. The Octet and Duet Rules We know that Noble Gases like Neon and Argon are unreactive. This was explained by Gilbert Lewis who came up with the Octet Rule in 1916 as part of the Lewis Theory of Bonding Octet Rule – Atoms tend to achieve the electron configuration of a noble gas when forming compounds Octet – set of 8 electrons in the highest energy orbital Atoms with an octet are stable! You made electron dot structures on the last page for atoms that follow the octet rule Hydrogen, lithium, Beryllium and Helium are exceptions – their most stable electron configuration is known as a duet (2 electrons in their highest energy level) ●He● (this is the electron dot structure for helium, showing the how it follows the duet rule.)

9. Ion Formation: Cations Metals tend to lose/give away electrons to form cations. This leaves an octet or duet in the next-lowest energy level Example: Sodium Electron configuration: 1s 2 2s 2 2p 6 3s 1 Valence electrons: 1 (in the 3s orbital) Lewis Structure: Na ● Sodium tends to lose this electron, becoming Na + Na  Na + + e - New electron configuration: 1s 2 2s 2 2p 6 This is the same electron configuration as Neon New valence electrons: 8 (in the 2s and 2p orbitals) New Lewis Structure: (brackets, no dots, charge on outside) For transition metals, the charges of cations may vary. They sometimes form exceptions to the octet rule. We’ll learn about transition metal cations at a later time.

10. Formation of compounds: Cations Example: How does magnesium become more stable when forming a compound? Magnesium gives away 2 valence electrons to attain the same electron configuration as Neon (an octet) Mg  Mg 2+ + 2e - 1s 2 2s 2 2p 6 3s 2  1s 2 2s 2 2p 6

11. Ion Formation: Anions Nonmetals tend to take on electrons to form anions. This also creates an octet or duet in the next-lowest energy level. Example: Fluorine Electron configuration: 1s 2 2s 2 2p 7 Valence electrons: 7 (in the 2p orbital) Lewis Structure: (see right  ) Chlorine tends to gain one electron, becoming Cl - Cl + e -  Cl - New electron configuration: 1s 2 2p 6 This is the same electron configuration as Neon New valence electrons: 8 (in the 2s and 2p orbitals) New Lewis Structure: (brackets, full set of dots, charge on outside)

12. Formation of Compounds: Anions Example: How does chlorine become stable when forming a compound? Chlorine gains one electron to gain the same electron configuration as argon Cl + e -  Cl - 1s 2 2s 2 2p 6 3s 2 3p 5  1s 2 2s 2 2p 6 3s 2 3p 6 The ions produced when halogens gain electrons are called halide ions. All halide ions have charges of 1- All halide ions have a full octet in their outermost shell

13. Dot structures and bonding In the Lewis Theory of Bonding: A chemical bond involves the sharing or transfer of electrons to attain stable electron configurations (outer shells with 8 electrons) for all bonding atoms If the electrons are transferred, the bond is an ionic bond If the electrons are shared, the bond is a covalent bond. The electrons in the bonding atoms are arranged so that all atoms involved get stable electron configurations. Below is an example of an ionic bond. Covalent bonds will be shown later.