Bonding & Chemical Nomenclature Chapter 8 & 9

1. Bonding & Chemical NomenclatureChapter 8 & 9

2. Some Key Terms • Chemical bond – a mutual electrical attraction b/w the nuclei and valence electrons of different atoms that binds the atoms together. • Ionic Bond – the electrostatic attraction of oppositely charged particles (cations & anions) • Ionic Compound – Molecules that contain ionic bonds where the positive and negative charges are equal. • Metallic bonding – chemical bond that results from the attraction b/w metal atoms and the surrounding sea of electrons

3. Key Terms cont. • Covalent bond – results from the sharing of electrons • Molecular compound – molecules that contain covalent bonds and whose simplest units are molecules. • Chemical Formula – a formula that indicates the relative #’s of atoms of each kind in a chemical compound. Uses atomic symbols and numerical subscripts. Ex. NaCl, Mg(OH)2 • Polyatomic Ion – A charged group of covalently bonded atoms. Contains two or more different elements

4. Key Terms cont. • Molecule – Neutral group of atoms held together by covalent bonds. • Oxidation number - # assigned to an atom in a molecular compound or ion that indicates its distribution of electrons (charge) • Binary Compound – A compound composed of 2 different elements • Octet rule – Chemical compounds tend to form so that each atom has an octet (8) electrons in the highest energy level except for a few of the smallest atoms.

5. What are ions? • Ions are atoms that have lost or gained electrons to become positively or negatively charged. • Metals tend to form positively charged ions called cations • Nonmetals tend to form negatively charged ions called anions. • To determine the charge of an ion we must look at the arrangement of VALENCE ELECTRONS.

6. Review of Lewis Dot Diagrams • Lewis dot diagrams show an elements valence electrons as dots around the element’s symbol. • To determine valence electrons for main group elements refer to the “A” group number. • REMEMBER – a valence electron is an electron in the outermost energy level

7. Formation of Ions • Label your periodic table with the following GENERAL charges above each group: • Group 1A = +1 • Group 2A = +2 • Group 3A = +3 • Group 4A = +/-4 • Group 5A = -3 • Group 6A = -2 • Group 7A = -1 • Group 8A = 0

8. Writing Ionic Formulas • CRISS-CROSS • Look up the charge for the first element, then look up the charge for the second element or group (polyatomic ion) and then criss-cross to get the subcripts • If the charges are equal and opposite they cancel and no criss-cross is needed. • Examples: • Sodium chloride • Potassium sulfide • Magnesium phosphate

9. Naming Binary Ionic Involves a metal and a nonmetal • Name the metal • Name the nonmetal and change the ending to –ide. Examples NaCl K2O BaF2

10. Naming Binary Ionic with transition metals Still Metal – Nonmetal. We use this method only if the metal can have a variable charge/oxidation number. The majority of these will be transition metals. Exceptions to memorize: Silver (Ag+1), cadmium (Cd+2), zinc (Zn+2) • Name the metal • Using a roman numeral write the oxidation number of the metal in parenthesis. To get the oxidation number uncriss-cross the charges. • Name the nonmetal and change the ending to –ide Examples CuBr2 Fe2O3

11. Naming Ionic Compounds with Polyatomic Ions • Name the metal • Name the polyatomic ion • Be sure to use a roman numeral if it is necessary (transition metals) Examples LiNO3 KMnO4 Na2CO3

12. Naming Binary Molecular (covalent) Compounds Involves a bond b/w 2 nonmetals • Name the 1st nonmetal. Use a prefix if there is more than one. • Name the 2nd nonmetal ALWAYS use a prefix. Change the ending to –ide. Examples NO P2O5 SO3 Prefixes 1 mono- 6 hexa- 2 di- 7 hepta- 3 tri- 8 octa- 4 tetra- 9 nona- 5 penta- 10 deca-

13. Writing Formulas for Molecular Compounds • Write the symbol for each of the elements involved • Look at the prefixes and decide what numbers are being represented. • Write those numbers as subscripts to the right of the elements symbol in the formula • Examples: • sulfur trioxide • Carbon monoxide • Dinitrogen monoxide

14. Naming Binary Acids • Acids generally start with Hydrogen • Binary Acids (H with one element) • Use the prefix hydro- • Then add the name of the second element and change the ending to –ic acid • HCl hydrochloric acid • HBrhydrobromic acid • HI hydroiodic acid

15. Naming Acids with Polyatomic Ions • Hydrogen with a polyatomic ion (group of atoms with a charge) • Identify the polyatomic ion • Change the ending accordingly: • -ate to –ic acid • -ite to -ous acid • (Do NOT use the prefix Hydro-) • H2SO4 sulfuric acid • HNO3 nitric acid • HNO2 nitrous acid

16. Writing Formulas for Acids • For acids you must look up the charges of the ions invovled. • Hydrogen is the first element – look up its charge • Then if it is a hydro- prefix acid look up the second element and charge and criss-cross • If it is an –ic acid or and –ous acid, you must look up the polyatomic ion and its charge then criss-cross • Example: • Hydrobromic acid • Sulfuric acid

17. Naming simple bases • Common bases for us will end in –OH or hydroxide ion • So name the first element then name the hydroxide • NaOH sodium hydroxide