Unit 4 Bonding


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Unit 4 Bonding

  1. 1. UNIT 4 - Bonding (Ch. 6 & 7) Let’s get molecular!!!!
  2. 2. Why do elements form compounds? <ul><li>What’s the attraction? </li></ul><ul><li>Molecules - made up of 2 or more atoms to form ONE unit or compound </li></ul><ul><li>Molecular compounds are usually gases or liquids </li></ul><ul><ul><li>Consist of 2 or more nonmetals </li></ul></ul><ul><ul><ul><li>(Where are nonmetals located on P.T?) </li></ul></ul></ul>
  3. 3. Examples <ul><li>Water….. H 2 O </li></ul><ul><ul><li>2 molecules of hydrogen </li></ul></ul><ul><ul><li>1 molecule of oxygen </li></ul></ul><ul><li>Sucrose (sugar)…. C 6 H 12 O 6 </li></ul><ul><ul><li>6 molecules of Carbon </li></ul></ul><ul><ul><li>12 molecules of Hydrogen </li></ul></ul><ul><ul><li>6 molecules of Oxygen </li></ul></ul>
  4. 4. Properties of Covalent Bonding (molecular bonding) <ul><li>Share electrons </li></ul><ul><li>Low melting point </li></ul><ul><li>Do not conduct electricity when dissolved </li></ul><ul><li>Bond with 2 or more NONMETALS ONLY </li></ul>
  5. 5. Diatomic Elements (molecules) <ul><li>KNOW THESE!!!!! </li></ul><ul><ul><li>H 2 , N 2 , O 2 , F 2 , Cl 2 , Br 2 , I 2 </li></ul></ul><ul><li>These 7 elements exist in nature as DIATOMIC!!! </li></ul><ul><ul><li>Want a stable octet soooooo badly they will bond with themselves! </li></ul></ul>
  6. 6. Single vs. Double vs. Triple bonds <ul><li>Draw Lewis Dot structure for: </li></ul><ul><ul><li>H N O Cl </li></ul></ul><ul><li>How many electrons will each element need to “share” to be “HAPPY?” </li></ul>
  7. 7. <ul><li>Single bond – one pair of electrons shared </li></ul><ul><li>Double bond – two pairs of e- are shared </li></ul><ul><li>Triple bond – three pairs of e- are shared </li></ul><ul><li>Which bond would be the strongest? Weakest? </li></ul>Single vs. Double vs. Triple bonds
  8. 8. Single vs. Double vs. Triple Bonds <ul><li>Which diatomic molecules will form a single bond? Double? Triple? </li></ul>
  9. 9. Writing/Naming Covalent <ul><li>Know your Prefixes </li></ul><ul><ul><li>1 – mono </li></ul></ul><ul><ul><li>2 – di </li></ul></ul><ul><ul><li>3 – tri </li></ul></ul><ul><ul><li>4 – tetr(a) </li></ul></ul><ul><ul><li>5 – pent(a) </li></ul></ul><ul><li>(Table 9.4) </li></ul><ul><ul><li>6 – hex(a) </li></ul></ul><ul><ul><li>7 – hept(a) </li></ul></ul><ul><ul><li>8 – oct(a) </li></ul></ul><ul><ul><li>9 – non(a) </li></ul></ul><ul><ul><li>10 – dec(a) </li></ul></ul>
  10. 10. Rule 1: Ex. N 2 O <ul><li>If the 1 st nonmetal has MORE than ONE molecule, use the prefix. </li></ul><ul><li>Always use prefix when naming 2 nd nonmetal </li></ul><ul><li>Change ending to “ide” </li></ul><ul><li>(oxygen….oxide) </li></ul><ul><li>diNitrogen monoxide </li></ul>
  11. 11. Rule 2: Ex. PCl 3 <ul><li>If the 1 st nonmetal has only ONE molecule, do NOT use prefix. </li></ul><ul><li>Always use prefix when naming 2 nd nonmetal </li></ul><ul><li>Change ending to “ide” </li></ul><ul><ul><li>(Chlorine…..Chloride) </li></ul></ul><ul><li>Phosphorous trichloride </li></ul>
  12. 12. Your Turn: <ul><li>Name the following: </li></ul><ul><ul><li>SO 3 </li></ul></ul><ul><ul><li>N 2 O 4 </li></ul></ul><ul><ul><li>P 2 O 5 </li></ul></ul><ul><ul><li>CO </li></ul></ul><ul><ul><li>NH 3 </li></ul></ul>
  13. 13. More Practice <ul><li>Write the formula for the following covalent compounds: </li></ul><ul><ul><li>Carbon disulfide </li></ul></ul><ul><ul><li>Disulfur dichloride </li></ul></ul><ul><ul><li>Oxygen trifluoride </li></ul></ul><ul><ul><li>Tetraphosphorous decoxide </li></ul></ul>
  14. 14. Welcome to Ionic Bonding You better learn your Polyatomic Ions!!
  15. 15. IONIC BONDING <ul><li>Properties: </li></ul><ul><ul><li>Bond btwn Metal and Nonmetal </li></ul></ul><ul><ul><ul><li>Cations/Anions???? </li></ul></ul></ul><ul><ul><li>Transfer electrons </li></ul></ul><ul><ul><li>Strong bond b/c of opposite charges </li></ul></ul><ul><ul><li>Forms Crystal Lattice structure </li></ul></ul><ul><ul><li>High melting pt. </li></ul></ul><ul><ul><li>Conduct electricity when dissolved in water </li></ul></ul><ul><ul><ul><li>Breaks bonds apart so only “ions” remain </li></ul></ul></ul>
  16. 16. Review Ion formation <ul><li>Cations: </li></ul><ul><li>Na </li></ul><ul><ul><li>Na +1 </li></ul></ul><ul><li>Mg </li></ul><ul><ul><li>Mg +2 </li></ul></ul><ul><li>Naming: </li></ul><ul><li>Name of metal </li></ul><ul><ul><li>Na +1 …sodium </li></ul></ul><ul><ul><li>Mg +2 …magnesium </li></ul></ul><ul><li>Anions: </li></ul><ul><li>Cl </li></ul><ul><ul><li>Cl -1 </li></ul></ul><ul><li>O </li></ul><ul><ul><li>O -2 </li></ul></ul><ul><li>Naming: </li></ul><ul><li>Change ending to “-ide” </li></ul><ul><ul><li>Cl -1 …chloride </li></ul></ul><ul><ul><li>O -2 …oxide </li></ul></ul>
  17. 17. Writing Binary Ionic Formulas <ul><li>Total “charge” of formula must equal ZERO </li></ul><ul><li>Always write Metal FIRST! </li></ul><ul><ul><li>Ex. Sodium Chloride </li></ul></ul><ul><ul><li>Ex. Calcium Chloride </li></ul></ul><ul><ul><li>Ex. Potassium Sulfide </li></ul></ul>
  18. 18. Naming Binary Ionic <ul><li>DO NOT use PREFIXES!!! (covalent only) </li></ul><ul><li>Write name of cation FIRST, then change ending of anion to “-ide” </li></ul><ul><ul><li>Ex. LiBr </li></ul></ul><ul><ul><li>Ex. MgCl 2 </li></ul></ul><ul><ul><li>Ex. Sr 3 N 2 </li></ul></ul>
  19. 19. Naming using Roman Numerals <ul><li>Use R.N for ALL metals EXCEPT groups IA, IIA, Al, Zn, Ag </li></ul><ul><ul><li>Ex. Copper(II) = Cu 2+ </li></ul></ul><ul><ul><li>Ex. Lead(IV) = Pb 4+ </li></ul></ul><ul><ul><li>Ex. Chromium(III) = ??? </li></ul></ul>
  20. 20. Naming using Roman Numerals <ul><li>First Ask, do I need to use a roman numeral? </li></ul><ul><li>What does the R.N tell me? </li></ul><ul><ul><li>Ex. CuCl 2 </li></ul></ul><ul><ul><li>Ex. PbS </li></ul></ul>