1. Chapter 6.1 Introduction to Chemical Bonding  Molecule – smallest electrically neutral unit of a substance that still has the properties of the substance  Made up of two or more atoms

2. Molecular Compound  Compounds composed of molecules made up of different atoms

3. Molecular Compounds  Usually low melting and boiling points  Usually liquids or gases at room temperature  Usually made up of two or more nonmetals

4. Ions  Atoms or groups of atoms that have a positive or negative charge  Forms when an electron is lost or gained

5. Cation  Any atom or group of atoms that has a positive charge  Lose electrons  Usually Metals

6. Anion  Any atom or group of atoms that has a negative charge  Gains electrons  Usually Nonmetals

7. Ionic Compounds  Compound composed of anions and cations  Usually a metal cation and a nonmetal anion

8. CharacteristicMolecular Compound Ionic Compound Representative Unit MoleculeFormula Unit Type of Elements NonmetalsMetal and Nonmetal Physical StateSolid, Liquid or Gas Solid Melting PointLow below 300 degrees Celsius High above 300 degrees Celsius

9. Chapter 6.2 Representing Chemical Compounds  Chemical Formulas  Molecular Formulas  Formula Units  Laws of Proportions

10. Chemical Formulas  Show the kinds and numbers of atoms in the smallest representative unit of the substance

11. Monatomic  Represent their chemical formula by writing their symbol  Copper – Cu  Helium - He

12. Diatomic  If the molecules of an element have more than one atom, use a subscript to indicate how many  Hydrogen – H 2  Oxygen – O 2

13. The Seven Diatomic Elements You must Memorize  Hydrogen – H 2  Fluorine – F 2  Oxygen – O 2  Nitrogen – N 2  Chlorine – Cl 2  Bromine – Br 2  Iodine - I 2

15. Molecular Formulas  Shows the kinds and numbers of atoms present in a molecule of a compound  Does not show anything about the structure

16. Ammonia NH 3

17. Formula Units (Ionic Compounds)  Does not represent a molecule  No separate molecular units  Arranged in an orderly pattern

18. Formula Unit  Lowest whole number ratio of ions in a compound  NaCl 1:1  MgCl 2 1:2

19. The Law of Definite Proportions  In samples of any chemical compound, the masses of the elements are always in the same proportion.

20. The Law of Multiple Proportions  Whenever two elements form more than one compound, the different masses of one element that combine with the same mass of the other elements are in small whole number ratios.

21. Chapter 6.3 Ionic Charges  Monatomic Ions  Polyatomic Ions

22. Monatomic Ions  Ions consisting of only one atom

23. Monatomic Ions – Cations  Group 1A ions all have a 1+ charge  Group 2A ions all have a 2+ charge  Group 3A ions all have a 3+ charge  Only for metals

24. Monatomic Ions – Cations  ? the group number from 8 to get number of the charge

25. Monatomic Ions - Anions  Group 7A ions all have a 1 – charge  Group 6A ions all have a 2 – charge  Group 5A ions all have a 3 – charge  Only for nonmetals

26. Monatomic Ions - Anions  Subtract group number from 8 to get number of the charge

27. Naming ions with more than one oxidation numbers  Stock System  Use a roman numeral after the symbol to indicate charge  Iron –  Fe 2+ - Iron (II) ion  Fe 3+ - Iron (III) ion

28. Polyatomic Ions  Tightly bonded group of atoms that behave as a unit and carry a charge  Sulfate – SO 4 2-

29. Polyatomic Ions  Most end in “ite” or “ate”  Look at table E  Three exceptions  Ammonium  Cyanide  Hydroxide

30. Polyatomic Ions  “ite” indicates one less of oxygen  “ate” indicates one more of oxygen  Does not tell how many oxygens

31. Chapter 6.4 Ionic Compounds  Writing Formulas  Naming Binary Compounds  Ternary Ionic Compounds

32. Anions  Change ending to “ide”  Oxygen – Oxide  Sulfur - Sulfide

33. Writing Formulas for Binary Ionic Compounds  Binary Compound – composed of two elements  The positive charge of the cation must balance the negative charge of the anion.

34. Writing Formulas for Binary Ionic Compounds  The total net charge must equal 0.  Potassium Chloride  K + and Cl -  KCl  Rust - Iron(III) Oxide  Ions - ?  Formula ?

35. Crisscross method  Rust Iron(III) Oxide  Fe 3+ and O 2-

36. Naming Binary Ionic Compounds  Cation first, anion second  Make sure charges balance, if not, use subscripts to balance  If the cation has more than one oxidation number, use roman numerals to indicate charge (ONLY Positive)

37. Ternary Ionic Compounds  Contains atoms of three different elements  Calcium Carbonate  CaCO 3

38. Ternary Ionic Compounds  “ate” or “ite” ending indicates a polyatomic anion containing oxygen  Calcium Nitrate  Ca 2+, NO 3 -  Ca (NO 3 ) 2

39. Chapter 6.5 Molecular Compounds and Acids  Binary Molecular Compounds  Naming Common Acids

40. Binary Molecular Compounds  Composed of two nonmetallic elements  Prefixes are used to indicate the number of each atom present

41. Binary Molecular Compounds  Prefixes PrefixNumber Mono-1 Di-2 Tri-3 Tetra-4 Penta-5 Hexa-6 Hepta-7 Octa-8 Nona-9 Deca-10

42. Binary Molecular Compounds  All end with “ide”  Drop the vowel at the end of the prefix if the element begins with a vowel  CO  Carbon Monooxide  Carbon Monoxide

43. Binary Molecular Compounds  If the first element has just one atom, drop the mono-  CO  Monocarbon monoxide  Carbon monoxide

44. Binary Molecular Compounds  Tetraiodine nonoxide  ??  N 2 0  ??

45. Naming Common Acids  Compounds that produce hydrogen ions when dissolved in water  Anions connected to as many H + ions to make the molecule neutral

46. Naming Common Acids  HCl  Hydrochloric Acid  H 2 SO 4  Sulfuric Acid  HNO 3  Nitric Acid  HC 2 H 3 O 2  Acetic Acid  H 3 PO 4  Phosphoric Acid  H 2 CO 3  Carbonic Acid

47. Chapter 6.6 Summary of Writing and Naming PxQyPxQy P = Hydrogen