3/14 Bell work

1. 3/14 Bell work • Pull out Types of Reactions/Predicting products practice. • Find the percent composition of chlorine in the following compounds: Lithium chloride Carbon tetrachloride Magnesium chlorate Aluminum chlorite

2. Agenda • Bell work • Review of types of reactions/predicting products practice • Pass back HW • Ionic and covalent bonding • Lewis dot • Lewis structures • HW: Lewis structures practice, study for test

3. Chemical Bonding Atoms bond through their electrons; therefore, we started by learning about how an atom’s electrons are organized.

4. Electron Shell Model Remember, this is a conceptual model and NOT what the atom actually looks like (not a physical model). Rather, it helps us to understand how the electrons within atoms behave.

5. The Bohr/Shell Model Diagram Each atom has its own configuration of electrons. Elements in the same group (vertical column) have similar configurations, which is why they have similar properties.

6. Bonding • Atoms obey the octet rule and fill their outermost electron shell with 8 valence electrons • Hydrogen is an exception to this rule and can only have 2 electrons in its outermost shell

7. Bonding • Ionic: transfer of electrons (steal/give away) to fulfill octet rule, form oppositely charged ions that are then attracted to one another • Covalent: share electrons to obtain octet, are held together by mutual attraction to electrons • Metallic: electron sea model, electrons belong to no atom in particular, move around like in a sea

8. Types of Bonds IONIC COVALENT e- are transferred from metal to nonmetal e- are shared between two nonmetals Bond Formation Type of Structure true molecules crystal lattice Physical State liquid or gas solid Melting Point low high Solubility in Water yes usually not yes (solution or liquid) Electrical Conductivity no Other Properties brittle odorous Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

9. H H O H O H H H Cl Cl Cl Cl O Ionic vs. Covalent • Ionic compounds form repeating units. • Covalent compounds form distinct molecules. • Consider adding to NaCl(s) vs. H2O(s): Na Na Na Na Cl Cl Na Na • NaCl: atoms of Cl and Na can add individually forming a compound with million of atoms. • H2O: O and H cannot add individually, instead molecules of H2O form the basic unit.

10. Ionic vs Covalent • This is why you must always reduce the formulas of ionic compounds • They do not form actual molecules, they are composed of formula units with the lowest ratio • You CANNOT reduce the formulas of covalent compounds because they form distinct molecules

11. - + Ionic Bonding transfer of electron Na Cl NaCl

12. Ionic Bonding • All the electrons must be accounted for! Ca +2 Ca P -3 +2 Ca P -3 +2

13. P3- P 3- Ca2+ Ca2+ Ca2+ Ionic Bonding Ca2+ Ca3P2 Ca2+ P3- Ca2+ Formula Unit P3-

14. F F 8 Valence electrons 8 Valence electrons Covalent bonding • Fluorine has seven valence electrons • A second F atom also has seven • By sharing electrons • Both end with full orbitals (stable octets)

15. x x x x o o x x F H x x x x x x x x x o o x F H o o C C x x x x F F H F H H H F x x x x x x x x x o o C C x x x o o x x x x x x o o x x o o x x x x x x x x o o x x x H F x x x x x x x H F x x x x x x x carbon tetrafluoride (CF4) methane (CH4)

16. x x x x x x x x O x x x O x O = C = O x x x x x x x x x x x x x x O O x x x x x x x x x x x I x x x o o x o x o x x o o x I x o x x o No I I I No x x x x o o C C x x x x x x x o o x x o x o x x x x x o o x x I x x x x x nitrogen triiodide (NI3) carbon dioxide (CO2)

17. Types of Bonds METALLIC e- are delocalized among metal atoms Bond Formation Type of Structure “electron sea” Physical State solid Melting Point very high Solubility in Water no yes (any form) Electrical Conductivity malleable, ductile, lustrous Other Properties Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem

18. - + - + + - + - - + - + + - + - - + - + Ionic solids are brittle Strong repulsion breaks crystal apart. Force

19. + + + + + + + + + + + + Metals are Malleable • Hammered into shape (bend). • Ductile - drawn into wires. • Electrons allow atoms to slide by.

20. Types of Bonds Metallic Bonding - “Electron Sea”

21. Metallic Bonds In metals, valence shells of atoms overlap, so v.e– are free to travel between atoms through material. Not so in metals. In insulators (like wood), the v.e– are attached to particular atoms.

22. ductile conduct heat and electricity malleable Properties of Metals All due to free-moving v.e–.

23. Questions to ponder • Compare and contrast ionic and covalent bonds. • Given two elements, how could you predict which type of bond they would form?

25. Valence Electrons Valence Electrons are the electrons in the outermost shell of an atom. These are the ones that participate in and affect/determine chemical bonding.

26. Valence electrons • How can we find the number of valence electrons an atom has?

27. Electron-dot Structure Electron-dot structure: A notation showing the valence electrons surrounding the atomic symbol.

28. Elements within the same group have the same electron-dot structure. Elements with 4+ shells: only look at outermost s and p electrons.

29. Practice • Draw the Lewis dot structure for the following atoms: H O N F Ar • How many paired electrons does each atom have? How many unpaired electrons?

30. Bonding • The number of unpaired electrons allows you to determine the number of covalent bonds an atom can form • Example: Carbon

31. The Covalent Bond The number of covalent bonds an atom can form equals its number of unpaired valence electrons.

33. Ionic Bonds • Ionic bond formation depends on how many electrons an atom needs to complete their octet. • You must balance the charges so each atom has a full octet or zero valence electrons • They are then attracted to one another because of their opposite charges

34. Lewis Structures 1) Count up all valence electrons 2) Connect all atoms with a single bond • “Multiple” atoms on the outside • “Single” atoms on the inside • C always in center; H always on outside 3) Complete octets on exterior atoms (not H) 4) Check • Valence electrons math from step 1 (not too many or too few) • All atoms (except H) have an octet, if not try multiple bonds • Any extra electrons? Put on central atoms

35. Multiple Bonds • Sometimes atoms share more than one pair of valence electrons. • A single bond is when atoms share one pair (2 total) of electrons • A double bond is when atoms share two pair (4 total) of electrons. • A triple bond is when atoms share three pair (6 total) of electrons.

36. Practice • F2 • H2O • NH3 • O2 • CO2 • N2

37. The Covalent Bond The covalent bond is represented using a straight line. F — F F F

38. Practice • F2 • H2O • NH3 • O2 • CO2 • N2

39. H O Water Each hydrogen has 1 valence electron Each hydrogen wants 1 more The oxygen has 6 valence electrons The oxygen wants 2 more They share to make each other happy

40. H O Water • Put the pieces together • The first hydrogen is happy • The oxygen still wants one more

41. H O H O H H Water • The second hydrogen attaches • Every atom has full energy levels • A pair of electrons is a single bond

42. The Covalent Bond The number of covalent bonds an atom can form equals its number of unpaired valence electrons.

43. Practice • F2 • H2O • NH3 • O2 • CO2 • N2

44. Examples • NH3 • N - has 5 valence electrons wants 8 • H - has 1 valence electrons wants 2 • NH3 has 5+3(1) = 8 • Connect all atoms with single bond N H

45. Examples • Draw in the bonds • Subtract # of electrons used from total • Put extras on central atom • Check octets • Check to make sure all electrons used H H N H

46. The Covalent Bond The number of covalent bonds an atom can form equals its number of unpaired valence electrons.

48. Practice • F2 • H2O • NH3 • O2 • CO2 • N2

49. x x O O x x O O x x x x x x x x Formation of Multiple Covalent Bonds By combining more than one unpaired electron at a time, a double bond is formed. Both oxygen atoms end up with eight valence electrons.

50. Practice • F2 • H2O • NH3 • O2 • CO2 • N2