1. Chapters 4 & 5 Chemical Formulas and Bonding

2. Compound Formation  Atoms form compounds to become more stable and achieve a full set of valence electrons.  Compound takes on a new set of properties from the original elements. Ex: water, sodium chloride, sucrose (sugar)

3. Ionic Bonding  Ionic bonds are the electrical attraction between positive and negative ions

4. Ionic Compounds  Consist of positive and negative charged ions Positively charged ions: cations Negatively charged ions: anions  Ionic compounds are electrically neutral Sodium ions + chloride ions  Lead (IV) ions + nitrate ions 

5. Properties of Ionic Compounds  Generally high melting points; solid at room temperature  Brittle  Dissociate in water Ions are pulled apart in water Causes the solution to become a good conductor of electricity  Represented by an empirical formula Shows the ratios of ions in the compound

6. Ion Formation  The Octet Rule Atoms tend to gain, lose, or share electrons in order to acquire a full set of valence electrons  Most atoms need 8 valence electrons

7. Lewis Dot Diagrams  Shows the valence electrons of an atom as dots around the element symbol. Shorthand way of representing the atom  Often used to represent changes to valence electrons during a chemical reaction

8. Lewis Dot Diagrams  Examples: Sodium chloride Aluminum fluoride

9. Writing Ionic Formulas  Examples: Lithium sulfate Potassium iodide Magnesium chlorate

10. Naming Ionic Compounds  Name of positive ion followed by name of negative ion  Examples: LiNO 3 FeCl 3 KMnO 4 CaCr 2 O 7 NaS 2 O 3

11. Homework  Worksheet 7.1

12. Covalent Bonding  Covalent bonds are formed by the sharing of a pair of electrons between 2 atoms A group of covalently bonded atoms is called a molecule  Covalent compounds are called molecular substances or compounds

13. Molecular (Covalent) Compounds  Shown by molecular formulas Shows the types of atoms and the actual number of each  Glucose = C 6 H 12 O 6  What would be the empirical formula?  The molecular formula for lactic acid is C 3 H 6 0 3  What would be its empirical formula?

14.  Sometimes, different molecules have the same molecular formulas Glucose is part of a family of sugars with the molecular formula C 6 H 12 O 6 Structural Formulas are used to specify how the atoms are actually bonded to each other.

15. Electronegativity in covalent bonds  Electrons in covalent bonds are not always shared equally  If one atom is significantly more electronegative than the other, the bond is called a polar bond  When atoms sharing electrons are similar in electronegativity, the bond is nonpolar

16. Bond Type by Electronegativity  Ionic bonds: large difference in electronegativity (>2.0) The elements are very far apart on the periodic table (metals and nonmetals)  Polar covalent bonds (.4-2.0) Elements are fairly close on the periodic table (very electronegative nonmetal with another nonmetal)  Nonpolar covalent bonds (<.4) Atoms are very close to each other

17. Naming molecular compounds  Prefixes are attached to the name of the atoms to indicate the number of atoms.  Less electronegative atom is listed first, more electronegative is second On the more electronegative atom, the end of the name is dropped and –ide is added.

18. Prefixes  Mono-  Di-  Tri-  Tetra-  Penta-  Hexa-  Hepta-  Octa-  Nona-  Deca-

19. Exceptions  1. Prefix mono- is normally left off the first word of a compound’s name  2. Diatomic molecules are called by their elemental name  3. Some molecules have common names: H 2 O = water NH 3 = ammonia

20. Practice  N 2 O 4  PCl 5  NF 3

21. Naming Hydrates  A hydrate is an ionic compound with water absorbed into its solid structure Anhydrous is the opposite (water- free)  Name the ionic compound followed by the proper prefix and the word hydrate.

22. Hydrates Practice  CuSO 4 * 5H 2 O  MgSO 4 * 7H 2 O

23. Acids  Acids are substances that dissolve in water to produce H + ions  Name acids according to anion: If anion ends in –ide  Hydro-root name of anion-ic acid If anion ends in –ate  Change the –ate to –ic acid If anion ends in –ite  Change the –ite to –ous acid

24. Acid practice  HCl  HBr  H 2 SO 4  HClO 3  HNO 2

25. Hydrocarbons  Covalently bonded organic compounds containing C and H  Name always ends in –ane  Prefix tells number of carbons  Number of hydrogens = carbonsX2 + 2

26. Hydrocarbon prefixes  Meth-  Eth-  Prop-  But-  Pent-  Hex-  Hept-  Oct-  Non-  Dec-

27. Hydrocarbon practice  Ethane  Nonane  C 8 H 18  C 5 H 12

28. Mixed Review  Name the following: H 2 SO 3 Mg(NO 2 ) 2 * 6H 2 O S 2 O 6  Write formulas for: Trinitrogen heptaoxide Heptane Potassium dichromate Hydrosulfuric acid