1. Molecular Bonds (Putting Elements Together)

2. Molar Mass Each atom has an atomic mass Molar mass is the atomic mass of all the atoms in the molecule summed together For Example: H 2 O = 2 x Atomic Mass of H + 1 x Atomic Mass of O

3. Counting Atoms in a Molecule In the example, NH 3, the subscript 3 only applies to the hydrogen. –Therefore: there is 1 N and 3 H in ammonia In the example, 3Ca 3 (PO 4 ) 2, the number of atoms changes due to the Coefficient in front of the molecule The 3 is multiplied to the Ca, P and O The subscript 2, multiplies the P and O 3Ca 3 (PO 4 ) 2

4. This means that there are 3 x 3 Ca, 3 x 2 P and 3 x (4 x 2) O

5. Bonds...

6. No, not that kind – bonds between atoms to form molecules It all depends upon the atom’s valence (outer shell) electrons These are the e - in the last Energy Level (n = 1 through 7) Figure these out using the Periodic Chart and/or Lewis Dot Diagrams

7. The Roman Numerals Tell You How Many Valence Electrons for the Primary or Representative Elements; The Valence Electrons for the Transition Elements Vary I II III IV V VI VII VIII

9. Group I is monovalent; II is divalent; III is trivalent; IV is tetravalent; V is back to being trivalent (since three e - openings); VI is divalent; VII is monovalent and VIII has a complete octet, so these seldom react or bond

10. Bond Types (In General): Pure or Non-Polar Covalent Χ difference = 0 to 0.5 on the Pauling EN Scale The pair of e - shared are done so equally Two nonmetals bonded together Polar Covalent A shared pair of e -, but not equally χ difference = 0.5 to 1.6 Molecule has Partial + and – Charges Ionic Bonds χ difference = 1.7 or higher to the maximum of 4.0 Metal bonded with a nonmetal Metallic Bonds are similar to Ionic Bonds

11. Metallic Bonds Two or more metals mixed are called alloys Two major formats –Interstitial and Substitutional These bonds permit the roaming of e- which creates a sea of dissociated e - Called the Electron Sea Model

13. Ionic Bonds These are the bonds between a metal and a nonmetal The metal Ion is positively charged and called a cation The nonmetal Ion is negatively charged and called an anion The bonded molecule should be neutrally charged when finished

14. Knowing where the metals and nonmetals are on the table will make your life easier

15. Let’s take a moment to discuss polyatomic ions... This is a molecule that acts as a cation or anion For example: NH 4 + ammoniumN 3 - azide ClO 4 - perchlorateCN - cyanide HCO 3 - bicarbonateOH - hydroxide Cr z O 7 -2 chromateNO 3 - nitrate ClO 3 - chlorateC 2 H 3 O 2 - acetate Don’t PANIC – I gave a list to you!

16. In an Ionic Bond – one or more electrons are lost or gained by the atoms involved This allows the atoms to have a complete valence shell – following the octet rule

17. In an Ionic Compound – balance the molecule using the criss-cross rule Mg +2 + Cl -1 Mg Cl 2 The one is understood. This applies even if using a polyatomic ion

18. NH 4 + + O -2 (NH 4 ) 2 O The parentheses are used to keep the polyatomic together Pb +4 + CO 3 -2 Pb 2 (CO 3 ) 4 and this can be simplified by reducing the subscripts to Pb(CO 3 ) 2

19. Naming Ionic Compounds is really simple: 1. Name the cation (metal) using its proper name; if it is a polyatomic, do the same 2. Then, using the stem of the anion (nonmetal), simply add the suffix “ide” Zinc + Chlorine = Zinc Chloride Iron + Oxygen = Iron Oxide Lithium + Cyanide = Lithium Cyanide Ammonium + Fluorine = Ammonium Fluoride Cobalt + Phosphorous = Cobalt Phosphide

20. Transition Metals present an issue for balancing and naming molecules since they can have varying oxidation states For example: Manganese can be a +2 or +3 Iron can be a +2 or +3 Lead can be a +2, or even a +4 Copper is a +1 or +2 Gold is usually a +1 or +3 And Hydrogen is a +1 or a -1!

21. Transition Metals To determine the correct Roman Numeral to place after the metal: Roman Numeral = - (Charge # anion)(#anions) (# cations) This is needed because, for example, iron chloride can be either FeCl 2 or FeCl 3 ; or iron (II) chloride or iron (III) chloride

22. Therefore – Ionic Bonds are: Metal+ Nonmetal + ion - ion cation anion monatomic monatomic or (except NH 4 + ) polyatomic left of steps right of steps Reactions are Exothermic Form Crystal Lattice Structures

23. Covalent Compounds These can be monatomic or polyatomic compounds It is a bond between two nonmetals They share a pair of electrons They can be subgrouped into polar or nonpolar If a binary compound (2 atoms) – use the same naming rules as in Ionic Compounds

24. If it has more than two atoms – need to use the prefixes Number PrefixNumberPrefix 1 Mono 7Hepta 2 Di 8Octa 3 Tri 9Nona 4 Tetra 10Deca 5 Penta 11Undeca 6 Hexa 12Dodeca

25. Naming Covalent Compounds Process: 1.Prefix Indicating # + full name of first nonmetal 2.Prefix Indicating # + root name of second nonmetal + suffix “ide” 3.Watch for polyatomics and use their proper names

26. For Example: P 4 S 10 becomes Tetraphosphorous Decasulfide P 2 O 5 Becomes Diphosphorous Pentaoxide SF 6 becomes Sulfur Hexafluoride SiBr 4 becomes Silicon Tetrabromide

27. Covalent Bonds can be Polar or Nonpolar A nonpolar has no discernable negative or positively charged sides (EN difference is 0) A polar covalent bond means one side is negative and the other positive

28. Electronegativity Percent IonicBond Difference CharacterType 0.21 %Non-polar 0.44Covalent 0.5 -------------------------------------------------------------------------- 0.69 0.815 1.022Polar 1.230Covalent 1.439 -------------------------------------------------------------------------- 1.647Ionic if metal/nonmetal 1.855Polar Cov. if non/nonmetal 2.063 -------------------------------------------------------------------------- 2.270 2.476Pure Ionic 2.682 2.886 3.089 3.292

29. Some elements are able to form more than one oxyanion (polyatomic ions that contain oxygen), each containing a different number of oxygen atoms. For example, chlorine can combine with oxygen in four ways to form four different oxyanions: ClO 4 -, ClO 3 -, ClO 2 -, and ClO - (Note that in a family of oxyanions, the charge remains the same; only the number of oxygen atoms varies.) The most common of the chlorine oxyanions is chlorate, ClO 3 -. In fact, you will generally find that the most common of an element’s oxyanions has a name with the form (root)ate.

30. The anion with one more oxygen atom than the (root)ate anion is named by putting per- at the beginning of the root and -ate at the end. For example, ClO 4 - is perchlorate. The anion with one fewer oxygen atom than the (root)ate anion is named with -ite on the end of the root. ClO 2 - is chlorite. The anion with two less oxygen atoms than the (root)ate anion is named by putting hypo- at the beginning of the root and -ite at the end. ClO - is hypochlorite.

31. Oxyanion Example ClO - Hypochlorite ClO 2 - Chlorite ClO 3 - Chlorate ClO 4 - Perchlorate

32. Some compounds have common names as well as their scientific names – you should learn these and others! –NOnitrogen monoxidenitric oxide –H 2 Odihydrogen monoxidewater –NH 3 nitrogen trihydrideammonia –CH 4 carbon tetrahydridemethane –C 4 H 10 tetracarbon decahydride butane

33. Some atoms are Diatomic – KNOW THESE! H 2 N 2 O 2 F 2 Cl 2 Br 2 and I 2 and P is usually found as P 4 while Sulfur is found as S 8 Other elements will bond beyond the octet rule – like PCl 5, and the noble gas Xe bonds with F in XeF 6, XeF 2, XeF 4, and XeO 4 – and this is due to a thing called “hypervalence” or “expanded octet”

34. Molecular Geometry The 3-Dimensional Shapes of Molecules depend upon the valence e - ’s of the atoms involved Valence Bond Theory and VSEPR Model both use the same shapes –Basically – they focus on covalent bonds with the shared bonding pairs of electrons (BP) –The assumption is made that the molecule will adopt a geometry to minimize the repulsion between e - ’s

35. The General Shapes:

36. Basic Geometry Bond Angles Linear 180 o Trigonal Planar120 o Tetrahedral109.5 o Trigonal Bipyramidal90 o and 120 o Octahedral90 o

39. Molecular Orbital Theory MOT uses atomic orbitals (AO), e - λ’s and e - density regions to examine bonds

41. This is the end of Part I Next: –Van der Waals and London Dispersion Forces –Polarity –Intermolecular Forces –Lewis Dot Diagrams with Covalent Bonds –Determining Molecular Structure –Resonance Structures