1. Compounds and Nomenclature

2. Bonding & Stability  Atoms want to be stable.  The Octet Rule states that a chemically stable atom contains 8 valence electrons.

3. Chemical Bonds  Atoms will gain, lose, or share electrons to obtain 8 valence electrons – this is called a chemical bond  There are 3 types of bonds: Metallic Ionic Covalent / Molecular

4. Ionic Bonds

5. Types of Compounds 1. Ionic Compounds: -involve the transfer of electron(s) between 2 oppositely charged ions (cation and anion) -metal and a nonmetal or a combination involving a complex ion -forms an ionic bond -exists as an ionic crystal lattice (not individual molecules) -known as a formula unit (eg. A formula unit of salt, not a molecule)

6. Crystal Lattice

7. Formula Unit  a chemical formula showing the simplest whole number ratio of cations to anions in an ionic compound.  Eg. NaCl – sodium chloride

8. Types of Compounds (continued) 2. Molecular compounds: -involve the sharing of electrons between nonmetals -forms a covalent bond - exists as individual molecules - Eg. Carbon dioxide CO 2 water H 2 O

9. Covalent Bond/Molecule

10. Ionic and Covalent Bonds

11. Properties of Ionic and Molecular Compounds 1.State at room temperature: -all ionic compounds are solids -molecular compounds may be a solid, liquid or a gas 2.Conductivity of solution: -ionic compounds conduct electricity (electrolytes) -molecular compounds do not conduct electricity (non-electrolytes)

12. 3.Solubility in water: -ionic compounds are soluble, to varying degrees (some better than others) and form colored or colorless solutions. -molecular compounds may or may not be soluble (colorless solutions if they form).

13. Nomenclature  Chemical nomenclature is the systematic naming of chemical compounds.  Science 1206 examines the naming of ionic compounds, molecular compounds and acids.  Compounds can be divided into two basic categories, those which are true binary compounds (contain only two types of elements), and those which contain more than two different types of elements.

14. Is there a metal or ammonium in the formula?

15. Ionic compounds Identify the type of ions: A. Monoatomic or simple ions B. Polyatomic or complex ions C.Multivalent ions D.Hydrates

16. Rules for Naming ionic compounds: A. Monoatomic or simple ions  Single atoms that have lost or gained one or more electrons  Form binary ionic compounds (2 simple ions)  Consist of cations and anions  Eg. Sodium + chlorine Na + Cl -

17.  Cations are written first, anions are second (name changes to “-ide” for the anion)  The total charge must be zero  Do not write charges in your final answer

18. Is there a metal or ammonium in the formula?

19. Rules for formulas: a. Write the symbols for the ions involved: eg. Silver and chlorine Ag+ and Cl-

20. Rules: b.Determine the lowest whole number ratio of ions which will provide an overall net charge of zero Ag 1+ Cl 1- becomes AgCl (silver chloride)

21. Example: potassium and oxygen potassium - K + oxygen -O 2- K 2 + O 1 2- becomes K 2 O potassium oxide

22. Practice – Writing Formulas Write formulas for the following compounds: Lithium bromide Potassium chloride Barium chloride Magnesium nitride Aluminum Fluoride Calcium Nitride

23. Is there a metal or ammonium in the formula?

24. Naming Ionic Compounds  Cation is named first  Anion is named second  Ending of anion is changed to “ide”  Practice: NaClBaCl 2 Al 2 O 3

25. B. Multivalent ions -certain transition metals can form more than one type of ion, each with a different charge. - eg. Cu 2+ - copper (II) Cu + - copper (I)

26. Multivalent ions (continued)  The transition metals have various electron configurations that will make them stable  Use a roman numeral after the cation to specify its charge (Stock naming system).  Eg. Iron (ii) oxideFeO Iron (iii) oxideFe 2 O 3

27. Stock vs. Classical FormulaStock NameClassical Name  Cu + copper(I) ionCuprous ion  Cu 2+ copper(II) ionCupric ion  Fe 2+ iron(II) ionFerrous ion  Fe 3+ iron(III) ionFerric ion

28. -the one written on top is the more common ion - eg. Fe 3+ - iron (III) Fe 2+ - iron (II)

29. Practice: CuSO 4 PbO uranium (vi) oxide uranium (iv) oxide

30. C. Polyatomic/Complex Ions  Polyatomic ion: atoms of 2 or more elements covalently bonded together with an overall charge eg. Nitrate NO 3 - AmmoniumNH 4 +  Complex ions are groups of atoms that are made stable by sharing electrons and which then become even more stable by gaining (usually) or losing electrons.

31.  The total positive charge in the formula must be equal to the total negative charge.  Rules: Name the cation, then name the anion Don’t change the ending of a polyatomic ion! Balance the charges If you need more than 1 complex ions, use brackets for that group

32. Practice NaNO 3 Al 2 (SO 4 ) 3 Mg(OH) 2 NaCH 3 COO

33. Is there a metal or ammonium in the formula?

34. D. Hydrates  Ionic compounds that contain water in their structure  eg. CuSO 4  H 2 O

35. Hydrate - Rules  Name the ionic part of the formula first  Name the water part second using a prefix system for the number of water molecules  Add prefix to “hydrate”  Prefixes: 1. mono 2. di 3. tri 4. tetra 5. penta 6. hexa 7. hepta 8. octa 9. nona 10. deca

36. Example: CuSO 4  5H 2 O Copper (II) sulfate pentahydrate

37. Review 1. Silver nitrate12. Copper(II) sulfate 2. Iron(II) phosphate13. Cobalt(II) iodide 3. Chromium(III) oxide14. Cesium phosphate 4. Nickel(II) fluoride15. Magnesium acetate 5. Copper(I) nitrate heptahydrate16. Potassium oxide 6. Lead(II) carbonate17. Strontium nitrate 7. Iron(II) fluoride18. Aluminum sulfate 8. Iron(III) hydroxide19. Calcium chlorate 9. Zinc phosphate dihydrate20. Rubidium cyanide 10. Potassium chlorate21. Tin(IV) oxide 11. Ammonium chromate22. Titanium(II) iodide

38. Ionic Formulas 1.BaCl 2 10. FeCl 3 2.Pb(NO 3 ) 2 11. Ca(CN) 2 3.TiI 3 12. Cu 2 S 4.K 2 CrO 4 14. Cd(ClO) 2 5.CoO15. SnO 2 6.Mg(ClO 4 ) 2 16. NaHCO 3 7.CuSO 4 17. Al(C 2 H 3 O 2 ) 3 8.Na 2 SO 3 18. Ni 3 (PO 4 ) 2

39. Molecular Compounds Chemical Bonding

40. Is there a metal or ammonium in the formula?

41. NO!!!!!!!!!

42. Molecular Compounds

43.  Binary molecular compounds form between 2 non-metals  Covalent bonds: shared electrons  Molecular formula: shows number and kind of atoms in a molecule

44. Naming  Use prefixes to specify number of atoms of each element in the molecule  Second element ends with “-ide”  No charges used in formula  The prefix “mono-” should not be used on the first element 1. mono 2. di 3. tri 4. tetra 5. penta 6. hexa 7. hepta 8. octa 9. nona 10. deca

45. Practice  Name the following: NO CO 2 N 4 O 9 N 6 O  Write formulas for the following: Boron trifluoride Sulfur hexafluoride Nitrogen monoxide Phosphorous pentachloride

46. Acids

47. ACIDS  HClHydrochloric 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

48. Acids have 2 criteria: 1.They must contain hydrogen (H + ) 2.They must be dissolved in water (aqueous); the formula will always contain the subscript aq.

49. Naming Acids Hydrogen is always the positive ion for an acid Ending Acid Name Example 1. -ide begins with hydro,HCl ends with -ic and acid 2. –ite ends with –ous and acid H 2 SO 3 3. –ate ends with –ic and acid H 2 SO 4