Chemical Bonds Section 1 Introduction to Chemical Bonding

1. Chemical BondsSection 1 Introduction to Chemical Bonding

2. Objectives • Define chemical bond. • Explain why most atoms form chemical bonds. • Describe ionic and covalent bonding. • Explain why most chemical bonding is neither purely ionic or covalent. • Classify bonding type according to electronegativity differences.

3. 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

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

6. 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. 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-

9. 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)

10. 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

11. 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

14. 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-

16. Electronegativity Values

17. less than 0.3 0.3-1.7 greater than 1.7

18. 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

19. HOMEWORK Section Review pg 177 #1-5

20. Objectives • Define molecules and molecular formula. • Explain relationships among potential energy, bond length, and bond energy. • Learn the basic steps used in writing Lewis structures. • Explain how resonance structures are used to represent molecules.

21. Section 2 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

22. 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

23. 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

24. Characteristic of the Covalent Bond • Bond length average distance between 2 bonded atoms H-H 75 pm • Form CB= H atoms release energy= amt of energy equals drop in PE • Bond energy energy required to break a chemical bond and form neutral isolated atoms (kJ/mol)

26. 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:

27. Objectives • Learn the basic steps used in writing Lewis structures. • Explain how resonance structures are used to represent molecules.

28. 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.

29. 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

30. 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

31. 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

32. 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)

34. Lewis Structures for Molecules Draw the LS for CH3I 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. Add unshared pairs of e- to each nonmetal (except H) such that each is surrounded by 8. 5. Count e- to be sure that # of VE=number available. Check to see that atoms have octet.

35. 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

36. 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

37. Draw the 3 Resonance Structures for SO3

38. Homework • Electron Dot Diagrams and Lewis Structures Worksheet

39. Objectives • Compare and contrast a chemical formula of molecular and ionic compounds • Compare and contrast properties of ionic and molecular compounds • Write Lewis structures for polyatomic ions

40. Section 3 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

41. Formation of Ionic Compounds

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

43. 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

44. Lattice Energy def. energy released when one mole of an ionic crystalline compound is formed from gaseous ions negative values mean energy released when crystals are formed

47. Polyatomic Ions def. a charged group of covalently bonded atoms • combine w/ ions of opposite charge to form ionic compounds • https://www.youtube.com/watch?v=A9qzV1j3NuE

48. Metallic Bonding • Occurs bc e- move freely among a metal’s + charged ions e- form a cloud around the metal ion • RESULT: ductility & malleable • Ex: metal hammered into sheets doesn‘t break bc ions are in layers that slide past each other w/out losing their attraction to the electron cloud • Ex: good conductor of electricity bc outer-level electrons are held weakly • http://cd1.edb.hkedcity.net/cd/science/chemistry/resource/animations/metallic_bond/metallic.html • https://www.ck12.org/physical-science/Metallic-Bonding-in-Physical-Science/enrichment/Metallic-Bond-Animation/?referrer=concept_details

49. Metallic Bonding • Ductility- ability of a substance to be drawn, pulled, or extruded through a small opening to produce a wire • Malleability- ability of a substance to be hammered into thin sheets

50. Homework • Chemical Speed Dating Profile • Chapter Review