Chapter 6: Chemical Bonds 6.1 – Ionic Bonding

Stable Electron Configurations  Atoms are stable when the highest energy level is filled with electrons Atom is not likely to react 8 valence electrons is the “magic” number  Atom is stable when there are 8 electrons in the outer orbit  Exception – Helium (He)  Why?  Helium has only 2 total electrons

Stable Electron Configurations Aluminum 3 Valence Electrons Unstable Neon 8 Valence Electrons Stable Helium 2 Valence Electrons Stable

Stable Electron Configurations  Remember how to find the number of valence electrons?

Stable Electron Configurations  Electron Dot Diagram – model of an atom that shows number of valence electrons Atomic symbol with dots outside the symbol that represent valence electrons Na. : :. Cl : : :.

Ionic Bonds  Ionic Bond – electrons are transferred from one atom to another Electrons are lost/gained Ions are formed  Ion – charged particle  Ex: Na +, Cl - Na. Cl : : :. + : : :. Na. + _ Protons: 11 Electrons: 11 Neutrons: 12 Protons: 17 Electrons: 17 Neutrons: 18 Protons: 11 Electrons: 10 Neutrons: 12 Protons: 17 Electrons: 18 Neutrons: 18

Ionic Bonds  2 types of ions: Cation – positively charged ion  Na + Anion – negatively charged ion  Cl -  Named by using part of the element’s name and adding the suffix -ide ChlorineChloride

Ionic Bonds  Chemical Bond – force that holds atoms or ions together as a unit

Ionic Bonds  Ionization Energy – amount of energy used to remove an electron Increases left to right and bottom to top of periodic table

Ionic Compounds  Ionic Compound – compound formed from an ionic bond Atoms are held together by the charge of the atoms after electrons are lost/gained  Chemical Formula – notation that shows what elements a compound contains and the ratio of atoms or ions of these elements in the compound Ex: NaCl (1:1 ratio) MgCl 2 (1:2 ratio)

Ionic Compounds  Some ionic compounds form crystal lattices Ions are in a fixed position called a lattice Crystals are formed when the particles of a solid form a lattice

Writing Ionic Formulas  Write the symbol for each atom  Identify the oxidation numbers  Do oxidation numbers cancel out? YES – write symbols of atoms NO – balance charges by using subscripts (usually use the criss cross method) Ex: calcium chloride CaCl +2 Ca +2 Cl -1 Ca +2 Cl -1 CaCl

Writing Ionic Formulas  Rewrite the symbols after the numbers have been switched Make sure you include the new subscripts Do NOT include the + or – symbols after you criss-cross If there is a 1 as a subscript, just write the symbol and do NOT write the 1 If the subscripts are the SAME number (i.e. Ca 2 O 2 ), simplify the formula by removing the numbers (CaO)

Writing Ionic Formulas  NOTE: If the name of an ion ends in “- ite” or “-ate”, this is a polyatomic ions (an ion that has more than one atom. There is a list of polyatomic ions on the back of your reference table. Use this list for the symbols and oxidation numbers of those ions. Keep the polyatomic ions in parentheses while writing the formula.

Writing Ionic Formulas Example 1: potassium sulfate K (SO 4 ) K +1 (SO 4 ) -2 K 2 (SO 4 ) 1 K 2 (SO 4 )

Writing Ionic Formulas  NOTE: Most transition metals will have roman numerals written with the name of that element. This represents the oxidation number of the element.  Example:iron (II)Fe +2 iron (III)Fe +3

Chapter 6: Chemical Bonds 6.2 – Covalent Bonding

Covalent Bonds  Covalent Bond – chemical bond in which atoms share electrons Atoms are held together by the attraction of the protons in the nucleus and the shared electrons orbiting the nucleus  Molecule – neutral group of atoms that are joined together by one or more covalent bonds  Diatomic Molecules – 2 atoms of the same element H 2, F 2, N 2, Cl 2, Br 2, I 2

Covalent Bonds Diatomic Elements

Covalent Bonds  Single Bond – 1 bond between diatomic molecules  Double Bond – 2 bonds between diatomic molecules  Triple Bond – 3 bonds between diatomic molecules N–N N=N N≡NN≡N

Sharing of Electrons  Polar Covalent Bond – electrons not shared equally Creates partial charge  Atom with stronger attraction has partial – charge  Atom with weaker attraction has partial + charge Ex: water, hydrogen fluoride H2OH2O HF

Sharing of Electrons  Non-polar Covalent Bond – electrons are shared equally No charge present Ex: Carbon dioxide CO 2

Attraction Between Molecules  Attractions between polar molecules are stronger than attractions between non- polar molecules Why?  Polar molecules have a charge  Non-polar molecules have no charge