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  1. Outline • Review • Lewis Theory of Bonding • How to Draw Elemental Lewis Structures • Drawing Ionic and covalent Bonds Using Lewis Structures • Polyatomic Ions • Why are they important? • How to Draw them

  2. Lewis Theory of Bonding Chapter 4.1

  3. Learning Goal To be able to draw the Lewis structures of simple polyatomic ions including but not limited to NO3-, SO3, NH4+, and NO2-.

  4. Lewis Theory of Bonding (1916) • Atoms are stable if they have a noble gas electron configuration (Full valence shell) • Electrons are most stable when they are paired. Why? • Atoms form chemical bonds to achieve a stable octet of electrons • A stable octet may be achieved by an exchange of electrons between a metal and nonmetal atoms • A stable octet of electrons may be achieved by sharing of electrons between nonmetal atoms Nelson page224

  5. What is the Goal of a Chemical Reaction? • Atoms trade or share electrons to fill their valence electron shell (achieve a full octet) • Two ways to fill their octet: • Ionic Bonding • Covalent Bonding

  6. Ionic Bond • What is an Ionic Bond? • What do you think holds the two charged atoms together?

  7. Covalent Bonding • What is a covalent bond? • What holds the two atoms together?

  8. Elemental Lewis Dot Diagrams • How do you draw the Lewis dot diagram for chlorine? • If you need a refresher, read over page 225

  9. The Octet Rule • Elements 6-18 (carbon to argon) stable when they have 8 electrons in their outer shell. How can we explain this phenomenon? • How can we explain that hydrogen and helium are most stable with only 2 valence shell electrons?

  10. Lewis Structures for Ionic Bonds • Example: NaCl • Practice questions to complete as a group • Draw the Lewis Structures for the ionic compounds: • Lithium Bromide • Magnesium Chloride

  11. Lewis Structures for Covalent Bonds • Example HCl • Draw the Lewis Structures for the following covalent compounds: • Fluorine Gas (F2) • Methane Gas (CH4)

  12. One Last Practice question • As a group try to draw the structure for CO32- • Was it easy? • Did anyone get this? 2-

  13. Extending the Lewis Theory of Bonding • Original theory could not explain the structure of Polyatomic Ions (Such as CO32-) • NevilSidgwick expanded the Lewis Theory of bonding by proposing that • Each atom can contribute in bonding without contributing an electron • Octet is desirable, but not necessary for all atoms or polyatomic ions Nelson Page 227

  14. Why are Polyatomic Ions important? • Polyatomic ions are very common in everyday life • CO32- • Chalk 2-

  15. _ NaHCO3 - Sodium Bicarbonate NH4NO3 – Ammonium Nitrate

  16. Common Food Preservatives Sodium Erythorbate Sodium Nitrite

  17. Drawing Lewis structures - NASL method (use NO3-as example) 1.) Select a reasonable skeletal structure for the molecule or polyatomic ion to be drawn. In general, the least electronegative element or the element written first in the formula is the central element (never hydrogen), place the remaining atoms evenly spaced around the outside 2.) Calculate N (Needed) as the sum of electrons needed for all atoms by the octet rule. Remember the exceptions: H=2, Be=4, B=6;

  18. Drawing Lewis structures - NASL method (use NO3-as example) 3.) Calculate A (Available) as the sum of all valence electrons. For negatively charged ions, add the number of electrons equal to the charge of the anion to A. For positively charged ions, subtract the number of electrons equal to the charge of the cation to A; 4.) Calculate S (Shared) as the difference between N – A; 5.) Divide S by 2 to obtain the number of bonds to be extended from the central atom;

  19. Drawing Lewis structures - NASL method(use NO3-as example) 6.) Calculate L (Lone-pair electrons or simply “dots”) as the difference between A – S. Place the L dots into the skeleton as to fill the octet of every “A-group” element, except hydrogen. Remember that hydrogen has only one bond and NO dots; 7.) Check that the total number of used electrons is equal to A. • Charge on individual atoms = Valence of atom – Number of Bonds extending from atom – dots on atom.

  20. Summary of NASL method • Select central atom, evenly space other atoms around it • Calculate the number of electrons needed (N) • Calculate the number of electrons available (A). Add electrons for negative charges, subtract electrons for positive charges • Calculate shared (S) electrons (S=N-A) • Calculate number of bonds (#Bonds=S÷2) • Calculate lone pairs(L) (L=A-S), fill octets of outer atoms first, if any left over use on central atom • Check total number of used electrons = A

  21. Agenda – February 26 2013 • Reminders • Take up Gizmo • Complete Chapter 3 Quiz • Complete Self Evaluation • Continue with drawing polyatomic ions

  22. Practice Time!! • Write the Lewis structures for the following: a.) H2O b.) SO2 c.) CO2 d.) CO3–2 e.) CF4 f.) NH4+ g.) N2 h.) SOCl2 i.) CS2 j.) N2O k.) O2 l.)BeCl2

  23. CHECK ANSWERS • H-O-H O - S = O • O = C= O

  24. 3 More Practice Questions!!! 2- • SO42- • PO43- • SO3 3-

  25. Exceptions to the octet rule • When phosphorous and sulfur are bonded to multiple oxygen atoms (such as SO42-, PO43-, and SO3) they break the octet rule. • When phosphorous is bonded to multiple oxygen atoms, it needs 10 electrons to satisfy its octet • When sulfur is bonded to multiple oxygen atoms, it requires 12 electrons to satisfy its octet

  26. [ ] – O H C O Resonance structures • Lewis structures for certain atoms do not match experimental observations • For example, the bond lengths of CHO2– predicted by the Lewis structure are incorrect • The double CO bond should be shorter, and possess a greater bond energy (due to the higher concentration of e–s in a double bond) • Yet, experimentally, both bonds are the same WHY? • The reason is due to “resonance”

  27. O [ ] – O H C O O O Resonance structures • A resonance structure can be drawn for any molecule in which a double bond can be formed from two or more identical choices • Resonancestructurescanbedrawn2ways… [ ] – 1 H C [ ] – O 2 H C • Resonance implies that the bond flips back and forth. Really, it lies between extremes

  28. HOMEWORK • Read Chapter 4.1 and answer Q#10,12 on page 229 of the textbook. Also answer Q#4 on page 230 of the textbook.