Understanding Chemical Bonding: Ionic and Covalent Bonds

1. Ch 8 Basic Concepts of Chemical Bonding

2. Atomic Stability Q. Why do atoms form compounds? • Search for outer shell stability by losing or gaining electrons! Nobel Gases  8 electrons in os = stable Ex: Na loses 1 e- to Cl Chemical Bond  attractive force that holds atoms together in a compound; when atoms gain, lose, or share e-, an attractive force pulls them together to form a compound

4. Oxidation Number  the number that tells you how many electrons an atom has gained, lost or shared to become stable

5. Ion • Def. Charged particle that has either gained or loses electrons • RESULT: Now has either more or fewer number of electrons than protons • Ex: sodium fluoride NaF active ingredient in toothpaste  Na loses 1e- to F; Na is +1; F is -1 • Ex: potassium iodide KI – ingredient in iodized salt  K loses 1 e -to I; K is +1; I is -1 • Cation = positive ion #P > #E • Anion = negative ion #P < #E

7. Ionic Bond • Def. Force of attraction between the opposite charges of the ions in an ionic compound (cations and anions) • Transfer of electrons btwn metals & nonmetals • Ex: magnesium chloride MgCl2 Mg loses 2 electrons to each Cl; Mg is +2; each Cl is -1 • RESULT neutral compound = sum of the charges of the ions equals 0 2(Mg) + -1(Cl) + -1(Cl) = 0

8. Ionic Bonding and Ionic Compounds • Ionic compound def. composed of + and – ions combined so that # of + and – charges are equal • most exist as crystalline solids= 3D structure of +/- ions attracted to each other • Formula unit def. simplest unit of atoms from which ionic compound can be established ex: NaCl • http://fikus.omska.cz/~bojkovsm/termodynamika/Obrazky/bindingstyper.swf

9. Formation of Ionic Compounds

10. Show formation of …KF, potassium fluorideNa20, sodium oxide

11. Characteristic of Ionic Bonding *Remember: Nature favors arrangements where potential energy is minimized. In an ionic crystal, ions minimize PE by combining in orderly arrangement crystal lattice

12. Energetics of Ionic Bond Formation What factors make formation ionic bonds so exothermic??? IE: It takes 496 kJ/mol (endo) to remove e- from Na Adding e- to Cl releases 349 kJ/mol (exo) 496-349 = 147 kJ/mol so why exothermic? A. Ionic compounds are stable bc attraction of ions of opposite charges attraction draws ions together releasing energy ions form crystal lattice

13. Lattice Energy def. energy required to completely separate 1 mole of a solid ionic compound into its gaseous ions negative values mean energy released when crystals are formed

14. Magnitude of Lattice Energies Potential Energy of 2 Interacting Charge Particles: E= k Q1Q2Q1& Q2 = charges on particles d d=distance between centers k= constant 8.99 x109 J m/C2 • LE increases as charge on ion increases • LE increases as radii decreases

15. Without consulting table, arrange following ionic compounds in order if increasing lattice energy: NaF, CsI, CaO • NaF Na+ F- • CsI  Cs+ I- • CaO Ca2+ O2- • Q1 x Q2 will be greatest for CaO • Cs larger than Na, I larger than F CsI < NaF < CaO Confirm: Check table 8.2 PRACTICE Which substance would you expect to have greatest lattice energy? MgF2, CaF2, ZrO2 A. ZrO2

16. Covalent Bond • Def. Attraction that forms between atoms when they share electrons • Atoms will be more stable by sharing e- rather than losing or gaining e-

17. Unequal Sharing • Electrons don’t always share equally between atoms in a covalent bond Strength of attraction of atom to electrons due to: • 1. size of atom • 2. size of the positive charge in the nucleus (a strong magnet will hold a metal better than a weak magnet) • 3. total # of electrons • 4. how far are the electrons from the nucleus being shared (a magnet has a stronger pull to a metal when it is next to it rather than a couple inches away)

18. Ex: HCl hydrochloric acid used to clean metal and found in your stomach to digest food • Cl – atoms have a stronger attraction for electrons than H atoms  electrons shared will spend most time near the chlorine atom • RESULT Cl atom has a partial negative charge (Greek delta)  H atom has a partial positive charge • VISUAL: Tug-of-war  stronger team pulls the rope towards them

19. Polar or Nonpolar • Polar molecule  molecule that has a slightly positive end and a slightly negative end although the overall molecule is neutral ex: water • Nonpolar molecule  molecule in which electrons are shared equal in bonds; doesn’t have oppositely charged ends; found in 2 identical atoms or molecules that are symmetric ex: CCl4

22. Ionic or Covalent??? • Bonding between atoms of different elements is rarely purely ionic or purely covalent. • Falls somewhere between 2 extremes depending on electronegativity measure of atom’s ability to attract e-

24. Electronegativity Values

25. less than 0.3 0.3-1.7 greater than 1.7

26. If you still need help understand polarity and electronegativity, watch this video: http://www.bing.com/videos/search?q=Examples+Polar+and+Nonpolar+Covalent+Bonds+With&Form=VQFRVP#view=detail&mid=295CAC087126B74D1565295CAC087126B74D1565

27. Covalent Bonding Molecular Compounds • molecule neutral group of atoms that are held together by covalent bonds • Red = O, White = H, Black = C • Molecular compound a chemical compound whose simplest units are molecules

28. chemical formula indicates what elements atoms and numbers of atoms in a chemical compound by using atomic symbols and numerical subscripts • molecular formula shows types and number of atoms combined in a single molecule of a molecular compound

29. Formation of a Covalent Bond • Nature favors chemical bonding most atoms have lower potential energy when bonded to other atoms than as independent atoms. • separated H atoms do not affect each other • PE decreases as atoms are drawn together by attractive forces • PE minimum when attractive forces are balanced by repulsion forces = ideal distance • PE increases when repulsion btwn like charges outweighs attraction between opposite charges

30. Octet Rule • Chemical compounds tend to form so that each atom has an octet of e- in outer energy level by gaining, losing or sharing e- • Draw Fluorine electron configurations:

31. Lewis Dot Diagrams Electron Dot Structure or Lewis Dot Diagram (Gilbert Lewis) • Def. A notation showing the valence electrons (electrons in outer energy level) surrounding the atomic symbol.

32. Lewis Structures • 1)Write the element symbol. • 2)Carbon is in the 4th group, so it has 4 valence electrons. • 3)Starting at the right, draw 4 electrons, or dots, counter-clockwise around the element symbol. On your sheet, try these elements on your own: • a)H b)P c)Ca d)Ar e)Cl f)Al

33. unshared pair- (lone pair) e- not involved in bonding and belong to one atom • Structural formula- shows bonds but not unshared pairs of e- in molecule

34. Single Covalent Bond • Made of 2 shared electrons • 1 comes from one atom in the bond and 1 comes from the other atom in the bond • Ex: water – O now is stable with 8 e in outer shell and H is stable with 2

35. Multiple Bonds • Bonds with multiple pairs of shared electrons • Ex: N2 Nitrogen has 5 e in os and needs 3 to be stable  shares 3 e (triple bond)

37. Lewis Structures for Molecules 1. Write the LS for each atom in the molecule. 2. Determine the total number of valence e- available. 3. Arrange atoms to form skeleton structure of molecule. When present C is central atom. Otherwise, least EN atom is central (Except H). 4. Complete the octets around all atoms bonded to the central atom. 5. Place any leftover e- around central atom even if results in more than 8 6. If there are not enough e- to give central atom octet, try multiple bonds. 5. Count e- to be sure that # of VE=number available. Check to see that atoms have octet.

38. Practice • Draw and build the Lewis Structure of NH3 • Draw and build the Lewis Structure of H2S • Draw and build the Lewis Structure of SiH4 • H-white, N-orange (should only have 3 holes) S-red, Si-black

39. Formal Charge • charge of the atom in a molecule; assuming atoms are all sharing e- equally regardless of electronegativity • ex: CO2 • RULE: Choose LS in which atoms formal charges are closest to 0 • RULE:Choose LS in which any negative charges reside on more EN atoms

41. Calculate formal charge on highlighted atom

42. Resonance Structures • def. bonding in molecules or ions that cannot be correctly represented by a single Lewis structure • resonance aka hybrids constantly alternating from one form to another O3 has a single structure that is avg of 2 structures • use double arrow to indicate resonance

43. Draw the 3 Resonance Structures for SO3

44. Draw the 3 Resonance Structures for NO3-

45. Draw the 3 Resonance Structures for HCO2-

46. Resonance in Benzene • Called "aromatic" initially because of its fragrance, aromaticity now refers to the stability of compounds that are considered aromatic, not only benzene.

47. Example of Aromatic Molecules

48. Exceptions to Octet Rules • #1. Odd Numbers of Electrons in molecules or polyatomic ions -Favors structure with lowest formal charge Ex: NO Draw both Lewis Structures Which is preferred?

49. #2 Molecules and polyatomic ions that have less than an octet of VE • Mostly occurs with Be and B • Draw Lewis Structure for BF3 • BF3 reacts energetically with molecules that have unshared pairs of e-

50. #3 Molecules and polyatomic ions that have more than an octet of VE • ex: PCl5 • ex: SF6 • Elements in period 2 only have 2s and 2p available for bonding= max 8 • Elements in period 3 and beyond have s, p, and d that can be used in bonding