Molecular Compounds Chapter 8

1. Molecular CompoundsChapter 8 Honors Chemistry Greenville High School

2. Sect. 8.1 Compounds and Molecules • Compound: a substance that is made from the atoms of two or moreelements that are chemically bonded. • Notice: The type of bond is not important, can be ionic, covalent or metallic • Examples: • H2O, CO2, NaCl, C6H12O6 • Non-examples: • I2, O2, Na, Si

3. Compounds and Molecules • Molecule: a neutral group of a least two atoms held together by covalent bonds • Now the type of bond is important: • Only covalent bonds **Notice it only has to be two atoms** • It can have two or more atoms of the same element or two or atoms of different elements Examples: • H2O, CO2, F2, H2, C6H12O6 Non-Examples: • NaCl, MgO, Al2O3,

4. 3 Types of Chemical Bonds • Ionic Bonds – a metal cation transfers valence electrons to a nonmetal anion • Metallic Bonds – postive cations in a sea of mobile valence electrons • Covalent Bonds – the bonds we will study in this chapter Allthree types of chemical bonds are intramolecular forces : the forces between atoms within a compound

5. Covalent Bonds • Covalent Bonds – “Co-Workers” Nonmetal + Nonmetal • two atoms share valence electrons to form a stable octet • Examples: H2O, CO2, NO2, SF6 • Covalently bonded compounds are called molecules

6. Covalent Bonds • Molecular Formula: shows how many atoms of each element a molecule contains. • Examples: • Diatomic Elements - O2, H2, Cl2 • Molecules - CH4, NH3, H2O Oxygen molecule O2 Benzene C6H6

7. Molecular Formulas • The formula for water is written as H2O What do the subscripts tell us? • Molecular formulas do not tell any information about the….. structure! (the arrangement of the various atoms).

8. Covalent bonds • Why do nonmetals share electrons? • Remember Nonmetals • Hold on to their valence electrons • Cannot give away electrons to bond. • Still want to form a stable octet. • By sharing valence electrons both nonmetal atoms get to count the electrons toward a stable octet.

9. Showing Covalent bonding • Show the bonding of F2

11. Important Covalent Compounds • 7 Diatomic Elements *Memorize* O2 N2 F2 Cl2 Br2 I2 H2 These elements are NEVER found as individual atoms. Ex: The oxygen gas we breathe is O2

13. Sect. 8.2 Types of Chemical Bonds • Covalent Bonds • Nonmetals do not always equally share their electrons • Some nonmetals can have a stronger pull on the shared pair of electrons—like tug of war of e- • These 2 types of covalent bonds are called polar and nonpolar.

14. Sect. 8.2 cont.. • Covalent Bonds: Polar and Nonpolar • Polar: a covalent bond in which the bonded atoms have an unequal attraction for the shared pair of electrons • Nonpolar: a covalent bond in which the two bonding electrons are shared equally by the bonded atoms.

15. Sect. 8.2 cont… • Electronegativity: How bad an element wants an electron • Using electronegativity differences to predict polarity and the bond type • Electronegativity Difference: (in Packet p. 13) • 0.0 - 0.4 = Nonpolar Covalent • 0.4 – 1.7 = Polar Covalent • > 1.7 = Ionic

17. Partial negative: element is partially neg. • Partial positive: element is partially pos.

18. Types of Chemical Bonds • Examples:Determine the electronegativity difference, the bond type and indicate partial positive and partial negative charges. a.) H and I H= ___ I=___ , Δ = ____ Bond type=_______________ H - I b.) K and Br K=____ Br=_____, Δ = _____ Bond type=_______________ K - Br

19. Work on Packet pg. 1 and 2 Ex: Draw the electron dot diagram for the covalent bonds **Remember Hydrogen needs only 2 electrons to fill the outer shell. F2 CH4

20. Bonds • 2 valence electrons = 1 bond • Hydrogen can only form one single bond WHY??

21. Single Bond • Single bond: when atoms share 1 pair of electrons (2 electrons total) Draw lewis dot for H2O, then show bonds

22. ~Tips for writing lewis dot structures for molecules with more than 2 atoms: • Central atom: is the 1st element in the compound or molecule (except H) 1. **The central atom ALWAYS goes in the middle!!! *** 2. Rearrange dots so that every element has 8 valence electrons (H and He only need 2 val)

23. H O H Structural Formulas • structural formula: Showing bonds.

24. Double Bond **Two atoms can share more thanone pair of valence electrons. • Double bond: when atoms share 2 pairs of electrons (4 electrons total) Ex 1: Draw the lewis dot for CO2, then show structural formula

25. Double Bond cont… Ex 2: Draw the lewis dot for H2CO, then show structural formula.

26. Triple Bond ~ Triple bond: when atoms share 3 pairs of electrons (6 electrons total) Draw the lewis dot for HCN and show structural formula.

27. How to find the # of bonds in a lewis structure • Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e- needed (all have 8 or 2 e-) - # of val e- available = ____/2 to find the # of bonds

28. Find the total # of valence electrons. 2. Use the formula to find the number of bonds. # of val e- needed (all have 8 or 2 e-) - # of val e- available = ____/2 to find the # of bonds Ex: Find the number of bonds for each molecule or compound and write the lewis dot and structural formula: a.) CO b.) C2F4 c.) C2H6

29. Exceptions to Octet rule • For some molecules, it is impossible to satisfy the octet rule • Yet the stable molecules do exist • Two types of exceptions: • Atoms that cannot hold 8 valence electrons • Hydrogen, helium, beryllium, boron, aluminum • Atoms that can hold more than 8 valence electrons • Phosphorus, sulfur, iodine, xenon, krypton

30. Exceptions to the Octet Rule 1. Most covalent compounds of Beryllium: the number of valence electrons needed for Be is 4. • BeF2 2. Most covalent compounds of Group 13: Primarily Boron & Aluminum - the number of valence electronsneeded is 6 • AlF3 • BF3

31. I – I – I Exceptions to Octet rule 3.Sometime when Phosphorus, Sulfur, Iodine, Xenon & Krypton are the central element they can hold more than 8 electrons: • PCl5 • I3 • SF6

32. Review on charges on bonding: • Ionic Bonds: • Have a full positive or full negative charge. • Ionic bonds do NOT have partial charges. Why? • Polar Covalent Bonds: • Have partial positive or partial negative charges. Why? • Nonpolar Covalent Bonds: • Have NO partial positive or partial neg. charge. Why?

33. Intermolecular Forces (IMF) • Attractive forces betweenmolecules. • Much weaker than chemical bonds. • Intramolecular forces • are within a molecule. (bonds)

34. Types of IMF • London Dispersion Forces: • Occurs between nonpolar molecules (diatomics) • Caused by motion of electrons ( “e- sloshing” ), they create a temporary dipole (slight charge) • Weakest of all forces. View animation online.

35. + - Types of IMF • Dipole-Dipole Forces: • Occurs between polar molecules • Where one side is partial positive and one is partial negative. • Stronger than London Dispersion forces. View animation online.

36. Types of IMF • Hydrogen Bonding: • When Hydrogen bonds to Nitrogen, Oxygen or Fluorine (NOF) • Strongest of all intermolecular forces!

37. Types of intermolecular forces:

38. Examples of intermolecular forces:Classify as London, Dipole or Hbonding. • NCl (nonpolar) • CO (polar) • HF (polar)

39. Properties Molecular Compounds • Low melting points and boiling points. • The IMF between molecular compounds are weaker than ionic or metallic compounds • This means that only a small amount of energy is required break the bonds Strongest Bonds  Weakest Bonds

40. Heat and electrical conductors • Covalent bonds: poor electrical and thermal conductivity. • No mobile electrons to conduct current Review of bonds: Covalent: Ionic: Metallic:

41. Draw Lewis dot diagrams for polyatomic ions: p.6 in packet • SO42- 2. PO43-

42. Molecular Geometry Lewis structures fail to indicate three-dimensional shapes of molecules. The shape of a molecule controls some of its chemical and physical properties.

43. VSEPR Valence Shell Electron Pair Repulsion Theory - predicts the shapes of a number of molecules and polyatomic ions. • Electron pairs move to create the most stable arrangement. • -The repulsions between electron pairs causes molecular shapes to adjust so that the electron pairs stay as FAR APART as possible.

44. What are the ideal arrangements of electron pairs to minimize repulsions? • We need to identify the number of regions of high electron density, called the steric number,on the central atom. • Regions of high electron density include: • Single bonds • Double bonds • Triple bonds • Unshared (lone) pairs of electrons **Double and triple** bonds only count as ONE region of high electron density just like a single bond or a lone pair.

45. Examples: Draw the Lewis Dot Structure and fill in the following: 1. CH4 • Steric # ____ • # of lone pairs _____ 2. H2O • Steric # ____ • # of lone pairs _____ 3. CO2 • Steric # ____ • # of lone pairs _____

47. Examples: Use table to determine molecular shape and bond angle. 1. CH4 • Steric # 4 Molecular Shape: __________ • # of lone pairs 0 Bond angle: _________

48. 2. H2O • Steric # 4 Molecular Shape:_____________ • # of lone pairs 2 Bond angle:________________ 3. CO2 • Steric # 2 Molecular Shape:______________ • # of lone pairs 0 Bond angle: ______________

49. How does Molecular Geometry affect Polarity? • One polar bond on central atom • Molecule polar? • Molecule nonpolar? • 2. More than one polar bond on the central atom will cancel out polarities if they have equal electronegativities. • Molecule polar? • Molecule nonpolar?

50. How does Molecular Geometry affect Polarity cont.. • 3. One lone pair on the central atom- • Polar? Nonpolar? • Two or more lone pairs on the central atom • Polar? Nonpolar? Water (asymmetrical) Xenon tetrafluoride (symmetrical) Xenon difluoride (symmetrical)