Monday! Period 1/2

1. Monday! Period 1/2 • Go over tests • Start Covalent Bonds! • Take notes • No HW!  • There will be a Quest/Test next Thursday 11 December

2. Monday! • Go over tests • Start covalent bonds! • No HW!  • There will be a Quest/Test next Thursday 11 December

3. Tuesday • Start Covalent Bonds! • Take notes • Naming a writing formulas of covalent bonds • It’s much easier than ionic! •  • HW: complete worksheet

4. Covalent Compounds! Chapter 6 Pgs. 188 - 221

5. Covalent Bond A bond formed when atoms share one or more pairs of electrons. Recall: it is a bond between nonmetals • Nonmetals hold onto their valence electrons. • Why?? • Because of their high electronegativity! They can’t transfer valence electrons to bond, they attract other electrons to themselves. • Still want noble gas configuration (Octet Rule), and they BOTH can’t gain! • Get it by sharing valence electrons with each other.

6. n = 2 - - - - Covalent Bond - - - - n = 1 - - - - - - - - + - - - - - - - - - - - - O [He]2s22p4 O [He]2s22p4 O2 Sharing of electrons to achieve a stable octet (8 electrons in valence shell).

7. Molecular Orbital • When shared electrons move in the space surrounding the nuclei of 2 atoms, the space these shared electrons move within is called a molecular orbital • Makes covalent bonds stronger than ionic • The orbital which is formed by the overlap of adjacent atomic orbitals • Ex: H2

8. Comparing Covalent and Ionic Bonds Video • http://www.youtube.com/watch?v=QqjcCvzWwww

9. Bond Length • The distance between two bonded atoms at their minimum potential energy; the average distance between the nuclei of two bonded atoms • Determined by potential energy

10. Potential Energy of Covalent Bonds • Recall: potential energy is energy of position (close vs far apart) • Unbonded atoms (far apart) tend to have high potential energy (except for noble gases) • Two atoms form a covalent bond at a distance where attractive and repulsive forces balance. At that distance, potential energy is at a minimum.

11. Analogy: a spring!

12. Bond Energy • The energy required to break a bond between two atoms It takes energy to break a bond, and energy is released when a bond is formed

13. Bond Length and Bond Energy • In general, shorter bonds are stronger than longer bonds (require more energy to break) • In general, multiple bonds are stronger and shorter than single bonds (require more energy to break)

14. Tuesday! • Properties of Covalent Compounds • Writing Formulas/Names Covalent Compounds! HW: Naming/Formula Worksheet

15. Properties of Covalent Compounds • Typically gases, liquids, or solids with low melting points • Most covalent compounds are insoluble in water • If they do dissolve in water, the solutions typically do not conduct electricity

16. Writing Formulas of Covalent Compounds • Covalent Compounds contain two types of nonmetals Key: FORGET CHARGES What to do: Use Greek prefixes to indicate how many atoms of each element, but don’t use “mono” on first element. 1 – mono 6 – hexa 2 – di 7 – hepta 3 – tri 8 – octa 4 – tetra 9 – nona 5 – penta 10 – deca

17. Binary Covalent Compounds Containing Two Nonmetals • Name the less electronegative non-metal first • With the Greek prefix indicating the number of atoms of that element present. • Don’t use mono on this element • Name the more electronegative non-metal • With the Greek prefix indicating the number of atoms of that element present • Ending should be replaced by the suffix –ide Greek prefixes you should know: Mono Di Tri Tetra PentaHexaHeptaOcta Nona Deca 1 2 3 4 5 6 7 8 9 10

18. H 2.1 Li 1.0 Be 1.5 B 2.0 C 2.5 N 3.0 O 3.5 F 4.0 Na 0.9 Mg 1.2 Al 1.5 Si 1.8 P 2.1 S 2.5 Cl 3.0 K 0.8 Ca 1.0 Sc 1.3 Ti 1.5 V 1.6 Cr 1.6 Mn 1.5 Fe 1.8 Co 1.8 Ni 1.8 Cu 1.9 Zn 1.7 Ga 1.6 Ge 1.8 As 2.0 Se 2.4 Br 2.8 Rb 0.8 Sr 1.0 Y 1.2 Zr 1.4 Nb 1.6 Mo 1.8 Tc 1.9 Ru 2.2 Rh 2.2 Pd 2.2 Ag 1.9 Cd 1.7 In 1.7 Sn 1.8 Sb 1.9 Te 2.1 I 2.5 * Cs 0.7 Ba 0.9 La 1.1 Hf 1.3 Ta 1.5 W 1.7 Re 1.9 Os 2.2 Ir 2.2 Pt 2.2 Au 2.4 Hg 1.9 Tl 1.8 Pb 1.8 Bi 1.9 Po 2.0 At 2.2 y Fr 0.7 Ra 0.9 Ac 1.1 * Lanthanides: 1.1 - 1.3 y Actinides: 1.3 - 1.5 Below 1.0 2.0 - 2.4 1.0 - 1.4 2.5 - 2.9 1.5 - 1.9 3.0 - 4.0 Electronegativities 1A 8A 1 1 3A 5A 7A 2A 4A 6A 2 2 3 3 2B 4B 6B 8B 1B 3B 5B 7B Period 4 4 5 5 6 6 7 Hill, Petrucci, General Chemistry An Integrated Approach 2nd Edition, page 373

19. Naming Covalent Compounds

20. Binary Covalent CompoundsContaining Two Nonmetals N2O dinitrogen monoxide N2O3dinitrogen trioxide N2O5dinitrogen pentoxide ICl iodine monochloride ICl3 iodine trichloride SO2 sulfur dioxide SO3 sulfur trioxide

21. Naming Covalent Compounds • ClO2 • P4S5 • SeF6 • O2 Chlorine Dioxide TetraphosphorusPentasulfide Selenium hexafluoride Oxygen (diatomic molecule)

22. Writing Formulas Nitrogen Trifluoride Phosphorus Triiodide DinitrogenTrioxide Ammonia (TrihydrogenMononitride) NF3 PI3 N2O3 NH3

23. Binary Covalent CompoundsContaining Two Nonmetals As2S3 • ________________ diarsenic trisulfide • ________________ sulfur dioxide • P2O5 ____________________ • ________________ carbon dioxide • N2O5 ____________________ • H2O ____________________ SO2 diphosphorus pentoxide CO2 dinitrogen pentoxide dihydrogen monoxide

24. Writing Formulas of Covalent Compounds EXAMPLES: carbon dioxide CO dinitrogen trioxide N2O5 carbon tetrachloride NI3 CO2 carbon monoxide N2O3 dinitrogen pentoxide CCl4 nitrogen triiodide

25. Please practice using the Formulas/Naming Covalent Compounds Worksheet

26. Wednesday!! • Go over Bonding Practice Worksheet • Electronegativity! • Polar vs. Nonpolar Molecules

27. Bond Polarity • The polarity of a bond depends on the electronegativities of the atoms involved • The greater the difference in electronegativity, the more polar the bond

28. Nonpolar Covalent Bond • A covalent bond in which the bonding electrons are equally attracted to both bonded atoms • Electronegativity difference: 0-0.5

29. Polar Covalent Bond • A covalent bond in which a shared pair of electrons is held more closely by one of the atoms • Electronegativity difference: 0.5-2.1 http://www.youtube.com/watch?v=rpVnUFL8Njk

30. Dipole • A molecule or part of a molecule that contains both positively and negatively charged regions • Can be shown in one of two ways

31. Thursday! • Covalent Compound Lewis Dot Structures!! • WOOOH! • HW: Finish Lewis Dot Structures Worksheet Don’t forget: Test Next Thursday

32. Ionic Bond • A compound in which electrons are completely transferred from one atom to another in the bond • Electronegativity difference: 2.1 and up

33. Structures of Covalent Bonds! • Next we will be drawing covalent bond Lewis Dot structures

34. Drawing Lewis Structures Recap: • Before drawing Lewis structures, become familiar with the following terms again: • Valence Electrons - Electrons that are found in the outermost shell of an atom and that determine the atom’s chemical properties • Lewis Structures - Structural formulas in which electrons are represented by dots; dot pairs or dashes between two atomic symbols represent pairs in covalent bonds • Unshared Pair - A nonbonding pair of electrons in the valence shell of an atom; also called a “lone pair”

35. Bonds • Single Bond (alkane) • A covalent bond in which two atoms share one pair of electrons • Double Bond (alkene) • A covalent bond in which two atoms share two pairs of electrons • Triple Bond (alkyne) • A covalent bond in which two atoms share three pairs of electrons

36. Drawing Lewis StructuresThe Grab Bag Method 1. Determine the total number of valence electrons. These go in the “grab bag.” 2. Determine the central atom. • Central atom is the atom that there is only one of, or the atom that will make the most bonds(or carbon if it is present). 3. Arrange the atoms around the central atom. 4. Form single bonds between each of the outer atoms and the central atom. Each single bond uses up two electrons. 5. Fill in the other electrons as valence electrons for the other atoms to complete their octets. 6. If you need more electrons than you have, form multiple bonds.

37. Drawing Lewis Structures Examples: CH3I CH4 SF6 PO43- BeH2 H2CO C2H6

38. Monday! • Go over Lewis Dot Structures • Resonance Structures • 1/2: • VSEPR Theory – Molecular Geometry Test = THURSDAY Start studying now!!

39. Resonance Structure • In chemistry, any one of two or more possible configurations of the same compound that have identical geometry but different arrangements of electrons

40. Resonance Structure • O3 • CO2

41. Tuesday! • Go over Resonance Structures • VSEPR Theory – Molecular Geometry • Molecular model kits/structures of covalent compounds! Test = THURSDAY!

42. Tuesday • Shapes of Molecules ~ VSEPR Theory • Molecular Model Kits! Test = THURSDAY

43. VSEPR Theory • A theory that predicts some molecular shapes based on the idea that pairs of valence electrons surrounding an atom repel each other

44. .. .. .. S O O O C O O S N F O O F F F F F F F P S F F F F F F The VSEPR Model The Shapes of Some Simple ABn Molecules SO2 Linear Bent Trigonal planar Trigonal pyramidal AB6 Trigonal bipyramidal Octahedral

45. VSEPR Theory

46. VSEPR Theory

47. VSEPR Theory

48. VSEPR Theory

49. VSEPR Theory

50. VSEPR Theory