1. Chemical Bonds Ch. 3 Chemistry II Milbank High School

2. Sec. 3.1 Stable Electronic Configurations Objectives –To determine what electron configuration appears to be the most stable. –Define the octet rule

3. Background There are about 115 known elements There are about 20 million known chemical compounds Chemical Bonds: force that holds the atoms of elements together

4. Sec. 3.1 Stable Electron Configurations Which group is the most stable? Which group is the most reactive? What does this have to do with electrons? See Pg. 68

5. So what group has a stable electron configuration? Noble gases! They have a stable octet of electrons in the highest main energy level Octet Rule: atoms seek an arrangement that will surround them with eight electrons in the outer shell

6. Atoms and ions Atoms are electrically neutral. Same number of protons and electrons. Ions are atoms, or groups of atoms, with a charge (positive or negative) Different numbers of protons and electrons. Only electrons can move. Gain or lose electrons.

7. Anion A negative ion. Has gained electrons. Nonmetals can gain electrons. Charge is written as a superscript on the right. F 1- Has gained one electron (-ide is new ending= fluoride) O 2- Gained two electrons (oxide)

8. Cations l Positive ions. l Formed by losing electrons. l More protons than electrons. l Metals can lose electrons K 1+ Has lost one electron (no name change for positive ions) Ca 2+ Has lost two electrons

9. Ionic Compounds Ionic compounds- from joining metal cations and nonmetal anions- they are electrically neutral Usually solid crystals Melt at high temperatures

10. Sec. 3.2 Lewis Structures Objectives –Determine and then solve Lewis Structures

11. Lewis Structures A representation of an element in which the chemical symbol stands for the core of the atom and dots are placed around the symbol for its valence electrons (outer shell)

12. Examples What are the Lewis structures (or electron dot structures) of the following elements? –Nitrogen –Argon –Potassium

13. Sec. 3.3 Intro Ionic Bonds and Ionic Compounds Ionic Compounds: consists of oppositely charged ions held together by electrostatic attractions Ionic Bonds: attractive forces between positive and negative ions

14. Sec. 3.3 and Sec 3.4 The Sodium-Chlorine Reaction and Other Ionic Bonds Objectives –Investigate the sodium-chlorine reaction –Make some general considerations about ionic bonds

15. Background Sodium is HIGHLY reactive Chlorine not very reactive, other than it is irritating to the respiratory tract

16. Reaction… Cl 2 + 2Na 2 Cl - + 2Na + See Lewis Structure Name: NaCl More reactive substances become less reactive when in a compound –Release energy in the process Sodium atom becomes less reactive by losing an electron Chlorine atom becomes less reactive by gaining an electron

17. Ionic Bonds: Some General Considerations The Potassium Chlorine Reaction The Magnesium Oxygen Reaction Normally, metallic elements in Groups 1A, 2A, and 3A react with nonmetallic elements in Groups 5A, 6A, and 7A to form stable crystalline ionic structures

18. Sec. 3.5 Names of Simple Ions and Ionic Compounds Determine how ionic compounds are named

19. Naming Monoatomic positive ions (cations)—add “ion” –Na + is called a sodium ion Monoatomic negative ions (anions)—add “ide” to the usual ending –Cl - is called chloride ion –S 2- is called sulfide ion

20. Naming See Table 3.2 Charge on a Group 1A element is usually 1+ Charge on a Group 2A element is usually 2+ You can calculate the charge by subtracting 8 from the group number. –Oxygen is 6-8 = 2- No simple way to determine charge on B subgroup elements –Roman numerals usually indicate charge –Iron(II) ion means Fe 2+

21. Formulas for Ionic Compounds Electrically neutral ions –Potassium (K + ) combines with Bromide ions (Br - ) Name? KBr Ratio 1:1

22. Formulas for Ionic Compounds Non-neutral ions –One Calcium ion (Ca 2+ ) combines with two chloride ions (Cl - ) Ratio 1:2 Name? CaCl 2

23. Crossover Method The charge number for one ion becomes the subscript for the other What is the formula for aluminum oxide? –Charges Al 3+ and O 2- –Name Al 2 O 3

24. Names for Binary Ionic Compounds What is the name of the compound Na 2 S? Find the ions (Table 3.2) –Sodium? Sodium ion -- Na + –Sulfur? Sulfide ion – S 2+ –Name: sodium sulfide

25. Sec. 3.6 Covalent Bonds: Shared Electron Pairs Determine what a covalent bond is Show how to write covalent bonds using Lewis Structures

26. Covalent Bonds Covalent Bonds: bond formed by a shared pair of electrons between atoms Molecule: group of atoms that are chemically bonded together –H 2 represents a molecule of hydrogen

27. Examples H Cl –Bonding Pair (two shared electrons) –Lone Pairs (nonbonding)

28. Sec. 3.7 Multiple Covalent Bonds Determine the difference between single, double, and triple bonds

29. Covalent Bonds Single Bond: single pair of shared electrons Double Bond: a covalent linkage in which the two atoms share two pairs of electrons

30. Examples CO 2 Sometimes represented by dashes O=C=O In general, many nonmetals often form a number of covalent bonds equal to the eight minus the group number. Oxygen.. 8 - 6 = 2 covalent bonds

31. Sec. 3.8 Naming Covalent Compounds Determine the correct naming structure for covalent compounds

32. Names for covalent compounds Makes use of prefixes to name compounds Example: –N2O4–N2O4 dinitrogen tetraoxide

33. Prefixes Mono 1 Di2 Tri3 Tetra4 Penta5 Hexa6 Hepta7 Octa8 Nona9 Deca10

34. Sec. 3.9 Unequal Sharing: Polar Covalent Bonds Show another way in which atoms combine Determine the difference between nonpolar and polar covalent bonding

35. Hydrogen/Chlorine Reaction Both need an electron to achieve a noble gas configuration They share a pair and form a covalent bond Chlorine atom have a greater attraction for a shared pair of electrons than do hydrogen atoms –More electronegative than hydrogen –They hold the shared electrons more tightly

36. Polar vs Nonpolar Nonpolar Covalent Bond: covalent bond in which electrons are shared equally Polar Covalent Bond: electrons are drawn more closely to the more electronegative atom, creating a separation of charge

37. Representation Use δ+ and δ- to indicate that one end is partially positive and one end is partially negative Example H—Cl

38. Sec 3.10 Electronegativity Define electronegativity Show electronegatitive tendencies using the Periodic Table

39. Electronegativity The measure of an atom’s tendency to attract electrons in a covalent bond to itself The greater the electronegativity of an atom in a molecule, the more strongly the atom attracts the electrons in a covalent bond

40. Periodic Table Tendencies Within a period, elements generally become more electronegative from left to right Within a group, electronegativity decreases from top to bottom

41. Sec. 3.11 Rules for Writing Lewis Structures Discuss the rules for writing Lewis Structures Write Lewis Structures Go through the individual steps used to write a Lewis formula

42. Lewis Structures Skeletal structure—model of a molecule that tells us the order in which the atoms are attached to one another

43. Rules for Writing Lewis Structures 1. Hydrogen atoms form only one bond; they are shown at the end of a sequence of atoms 2. Polyatomic molecules and ions often consist of a central atom surrounded by more electronegative atoms –Hydrogen exception—always on outside, even when bonded to a more electronegative element

44. Steps after you have the skeletal formula 1. Calculate the total number of valence electrons 2. Write the skeletal structure 3. Place electrons about outer atoms so that each has an octet 4. Subtract the number of electrons assigned so far from the total calculated in Step 1. Assign the rest to the central atom 5. Move lone pairs to form a multiple bond on central atom.

45. Examples Example 3.12 Example 3.13

46. Sec. 3.12 The VESPR Theory Predict and construct shapes of many molecules by using the VESPR theory

47. VESPR Theory Valence shell electron theory— describes the geometrical shape of a molecule or polyatomic ion based on the mutual repulsions among electon groups surrounding the central atom(s) in the structure Minimizes repulsion between the like-charged particles

48. Steps in determining shapes 1. Draw a Lewis Structure –Use dots to indicate lone pairs 2. To determine the shape, use the steric number (number of atoms bonded to the central atom + number of lone pairs on the central atom) 3. Draw shape based upon steric number 4. Place electron pairs as far apart as possible. –See Table 3.4

49. Examples Use table 3.4 Examples 3.14 –BH 3 –SCl 2

50. Sec. 3.13 Polar and Nonpolar Molecules Define dipole Show the correct way to represent dipoles

51. Dipoles A molecule is a dipole if it has a positive and negative end Represented by an arrow with a plus end H—Cl

52. Sec. 3.14 Polyatomic Ions Define polyatomic ions Determine formulas for polyatomic ions

53. Polyatomic Ions Charged particles containing two or more covalently bonded atoms Table 3.5 – some common polyatomic ions Example 3.15

54. Sec. 3.15 Exceptions to the Octet Rule Determine what a free radical is Show expanded valence shells

55. Three exceptions to the octet rule 1. Molecules in which the total number of valence electrons is an odd number 2. Molecules in which the total number of valence electrons is too low to allow all atoms to have a filled valence shell 3. Molecules involving elements from the third and higher periods can have an expanded valence shell

56. Molecules with odd numbers of valence electrons Free Radicals—a highly reactive atom or molecular fragment characterized by having one or more unpaired electrons See examples on pg. 91

57. Expanded Valence Shells Refer to situations in which the central atom in a Lewis Structure is able to accommodate more than the usual octet of electrons in its valence shell Usually encountered when the central atom is a nonmetal of the third period or beyond