SCH3U: Chemical Bonding

1. SCH3U: Chemical Bonding Q: what did the bartender say when oxygen, hydrogen, sulphur, sodium and phosphorous walked into the room? A: OH SNaP!

2. Agenda • Questions from Last Class? • Ionic, Covalent and Metallic Bonding • Lewis Diagram • Single Atoms; hints for drawing Lewis diagrams • Compounds • Total Valence Electrons • Steps to draw Lewis Diagram for compounds • Practice Questions • Reminder (Homework; Quiz; Gizmo)

3. 1e- 3p+ 3p+ 17p+ 2e- 2e-1e- 7e- 8e- 2e- 17p+ 8e-8e-2e 4n0 18n0 4n0 18n0 Cl [Li]+ [ Cl ]– Li Ionic bonding (Li + Cl) • Ionic bonding (stealing/transfer of electrons) Li + Cl  [Li]+[Cl]– Chemical Formula Rutherford Diagram Lewis Diagram

4. 8p+ 8p+ 8e- 2e- 6e- 2e- 8n0 8n0 1e- 12p+ 12p+ 2e- 8e- 2e- 8e- 2e- 12n0 12n0 1e- O [Mg]2+ [ O ]2– Mg Ionic bonding (Mg + O) Chemical Formula Mg + O  [Mg]2+[O]2– Rutherford Diagram Lewis Diagram

5. 8p+ 8p+ 8e- 2e- 6e- 2e- 8n0 8n0 1e- 3p+ 3p+ 3p+ 3p+ 2e- 2e- 2e-1e- 2e-1e- 4n0 4n0 4n0 4n0 1e- O [Li]2+ [ O ]2– Li Li Ionic bonding: Li + O 2Li + O [Li]2+[O]2– or 4Li + O2 2[Li]2+[ O]2–

6. Ionic vs. Covalent

7. Types of Covalent Bonds • a. non Polar Covalent Bond: When EN difference is zero. Equal sharing of electrons. • b. Polar Covalent Bond: When EN difference is >0 and <1.7. Results in unequal sharing of electrons, one atom has a higher affinity for the electrons.

8. Polar Covalent Bonds

9. Metallic Bond (neither ionic or covalent) • Occurs between two metal atoms. • Valence electrons are released into a sea of electrons. • Atoms release their electrons in a shared pool leads to stronger metallic bonds. • Electrons are free to move and are not held together by a lattice formation (like that of ionic bonds • That is why pure metals such as copper, sodium and gold are soft elements

10. Metallic Bonds

11. Chemical Bonding Lewis Structures

12. Lewis Diagram • Single Atom: the chemical symbol and dots indicating the number of electrons in the valence energy level • Compound: representation of covalent bonding based on Lewis symbols; Shared electrons are shown as lines and lone pairs as dots

13. To Write Lewis Diagrams 1. Write the element symbol. Around this, draw dots one for each valence electron. 2. The dots should be spread over four sides. Dots are not paired until all sides have at least one dot. 3. It does not matter on which side dots are placed. For example, hydrogen can be drawn four ways: 4. The number of valence electrons is equal to the group number. For example, hydrogen is in group IA (group 1) and it has one valence electron. Neon is in 0 (group 8) and it has 8 valence electrons. The only exception is He which is in group 8 but has 2 valence electrons.

18. Example: AsBr3 STEP 1:place the single atom in the center and other atoms around it evenly spaced As Br Br Br

19. STEP 2: count the total # of valence e- for all atoms involved in the bonding Arsenic: 1 Arsenic with 5 valence electrons (1x5) = 5 AsBr3 5+21 =26 Bromine: 3 bromine with 7 valence electrons (3x7) = 21

20. STEP 3: place the electrons in pairs between the central atom and each non-central atom Br As Br Br

21. STEP 4: place the remaining electrons around the non-central atom until each has 8 electrons (H atoms have only 2e-) Ex: AsBr3 5 + (7x3) = 26 26 - 6 = 20 Br As Br Br

22. Br As Br Br Structural Formula STEP 5: if electrons remain they are placed in pairs around the central atom Ex: AsBr3 5 + (7x3) = 26 20 - 18 = 2 As Br Br Br

23. Exception: if the central atom is in group 14, 15, 16, 17 or 18, the octet rule must be satisfied by moving electron pairs from non-central atoms, creating multiple bonds. Ex: SO2 6 +(6x2) =18 O O S - O S = O Structural Formula

24. Resonance Structure When several structures with different electron distributions among the bonds are possible, all the structures contribute to the electronic structure of the molecule. These structures are called resonance structures EX) CO3 • When considering the resonance structure… • Formal charge closest to zero is favoured • Negative formal charges on the more negative elements are favoured

25. Summary • STEP 1:place the single atom in the center and other atoms around it evenly spaced • STEP 2: count the total # of valence e- for all atoms involved in the bonding • STEP 3: place the electrons in pairs between the central atom and each non-central atom • STEP 4: place the remaining electrons around the non-central atom until each has 8 electrons (H atoms have only 2e-) • STEP 5: if electrons remain they are placed in pairs around the central atom • Exception: if the central atom is in group 14, 15, 16, 17 or 18, the octet rule must be satisfied by moving electron pairs from non-central atoms, creating multiple bonds

26. Summary / Hints • When drawing Lewis Diagrams… • Symmetrical arrangements are favoured • Hydrogen is never a central atom • The least electronegative element is usually the central atom • When considering the resonance structure… • Formal charge closest to zero is favoured • Negative formal charges on the more negative elements are favoured

27. Practice Questions • Draw Lewis Structure/ Structural Formula for the following compounds • CCl4 • HCl • NH3 • H2O • OH • H2

28. Reminder • Do the practice problems and assigned homework • Gizmo activity is due this Thursday!