1. Chapter 4 Compounds and Their Bonds 4.4 Polyatomic Ions 1 Copyright © 2009 by Pearson Education, Inc.

2. Polyatomic Ions A polyatomic ion is a group of atoms. has an overall ionic charge. Some examples of polyatomic ions are NH 4 + ammoniumOH − hydroxide NO 3 − nitrateNO 2 − nitrite CO 3 2− carbonatePO 4 3− phosphate HCO 3 − hydrogen carbonate (bicarbonate) 2

3. Some Compounds with Polyatomic Ions 3 Copyright © 2009 by Pearson Education, Inc.

4. Some Names of Polyatomic Ions The names of common polyatomic anions end in ate. NO 3 − nitratePO 4 3− phosphate with one oxygen less end in ite. NO 2 − nitritePO 3 3− phosphite with hydrogen attached use the prefix hydrogen (or bi). HCO 3 − hydrogen carbonate(bicarbonate) HSO 3 − hydrogen sulfite (bisulfite) 4

5. Names and Formulas of Common Polyatomic Ions 5

6. Naming Compounds with Polyatomic Ions The positive ion is named first, followed by the name of the polyatomic ion. NaNO 3 sodium nitrate K 2 SO 4 potassium sulfate Fe(HCO 3 ) 3 iron(III) bicarbonate or iron(III) hydrogen carbonate (NH 4 ) 3 PO 3 ammonium phosphite 6

7. Some Compounds with Polyatomic Ions 7

8. Learning Check Match each formula with the correct name. A. MgS1) magnesium sulfite MgSO 3 2) magnesium sulfate MgSO 4 3) magnesium sulfide B. Ca(ClO 3 ) 2 1) calcium chlorate CaCl 2 2) calcium chlorite Ca(ClO 2 ) 2 3) calcium chloride 8

9. Learning Check Name each of the following compounds: A.Mg(NO 3 ) 2 B.Cu(ClO 3 ) 2 C.PbO 2 D.Fe 2 (SO 4 ) 3 E.Ba 3 (PO 3 ) 2 9

10. Writing Formulas with Polyatomic Ions The formula of an ionic compound containing a polyatomic ion must have a charge balance that equals zero (0). Na + and NO 3 − -> NaNO 3 with two or more polyatomic ions has the polyatomic ions in parentheses. Mg 2+ and 2NO 3 − -> Mg(NO 3 ) 2 subscript 2 for charge balance 10

11. Learning Check Select the correct formula for each. A. aluminum nitrate 1) AlNO 3 2) Al(NO) 3 3) Al(NO 3 ) 3 B. copper(II) nitrate 1) CuNO 3 2) Cu(NO 3 ) 2 3) Cu 2 (NO 3 ) C. iron(III) hydroxide 1) FeOH2) Fe 3 OH3) Fe(OH) 3 D. tin(IV) hydroxide 1) Sn(OH) 4 2) Sn(OH) 2 3) Sn 4 (OH) 11

12. Learning Check Write the correct formula for each. A.potassium bromate B.calcium carbonate C.sodium phosphate D.iron(II) nitrite 12

13. Flowchart for Naming Ionic Compounds 13

14. Learning Check Name the following compounds: A. Ca 3 (PO 4 ) 2 B. FeBr 3 C. Al 2 S 3 D. Zn(NO 2 ) 2 E. NaHCO 3 14

15. Learning Check Write the formulas for the following: A. calcium nitrate B. iron(II) hydroxide C. aluminum carbonate D. copper(II) bromide E. lithium phosphate 15

16. Chapter 4 Compounds and Their Bonds 4.5 Covalent Compounds 16 Copyright © 2009 by Pearson Education, Inc.

17. Covalent Bonds Covalent bonds form when atoms share electrons to complete octets. between two nonmetal atoms. between nonmetal atoms from Groups 4A (14), 5A (15), 6A (16), and 7A (17). 17

18. Hydrogen Molecule A hydrogen molecule is stable with 2 electrons (helium). has a shared pair of electrons. 18

19. Forming Octets in Molecules In a fluorine, F 2,, molecule, each F atom shares 1 electron. attains an octet. 19

20. Carbon Forms 4 Covalent Bonds In a CH 4 (methane) molecule, 1 C atom shares electrons with 4 H atoms to attain an octet. each H atom shares 1 electron to become stable, like helium. 20

21. Multiple Bonds In a nitrogen molecule, N 2, each N atom shares 3 electrons. each N attains an octet. the bond is a multiple bond called a triple bond. the name is the same as the element. 21 Copyright © 2009 by Pearson Education, Inc.

22. Naming Covalent Compounds In the names of covalent compounds, prefixes are used to indicate the number of atoms (subscript) of each element. (mono is usually omitted) 22

23. Naming Covalent Compounds What is the name of SO 3 ? 1. The first nonmetal is S sulfur. 2. The second nonmetal is O, named oxide. 3. The subscript 3 of O is shown as the prefix tri. SO 3 -> sulfur trioxide The subscript 1 (for S) or mono is understood. 23

24. Naming Covalent Compounds Name P 4 S 3. 1. The first nonmetal, P, is phosphorus. 2. The second nonmetal, S, is sulfide. 3. The subscript 4 of P is shown as tetra. The subscript 3 of O is shown as tri. P 4 S 3 -> tetraphosphorus trisulfide 24

25. Formulas and Names of Some Covalent Compounds 25

26. Learning Check Select the correct name for each compound. A.SiCl 4 1) silicon chloride 2) tetrasilicon chloride 3) silicon tetrachloride B. P 2 O 5 1) phosphorus oxide 2) phosphorus pentoxide 3) diphosphorus pentoxide C.Cl 2 O 7 1) dichlorine heptoxide 2) dichlorine oxide 3) chlorine heptoxide 26

27. Learning Check Write the name of each covalent compound. CO 2 _____________________ PCl 3 _____________________ CCl 4 _____________________ N 2 O_____________________ 27

28. Example: Writing Formulas for Covalent Compounds Write the formula for carbon disulfide. STEP 1: Elements are C and S STEP 2: No prefix for carbon means 1 C Prefix di = 2 Formula: CS 2 28

29. Learning Check Write the correct formula for each of the following: A. phosphorus pentachloride B. dinitrogen trioxide C. sulfur hexafluoride 29

30. Learning Check Identify each compound as ionic or covalent, and give its correct name. A. SO 3 B. BaCl 2 C. (NH 4 ) 3 PO 3 D. Cu 2 CO 3 E. N 2 O 4 30

31. Study Tip: Ionic or Covalent A compound is ionic if the first element in the formula or the name is a metal or the polyatomic ion NH 4 +. K 2 O K is a metal; compound is ionic; potassium oxide covalent if the first element in the formula or the name is a nonmetal. N 2 O N is a nonmetal; compound is covalent; dinitrogen oxide 31

32. Learning Check Identify each compound as ionic or covalent and give its correct name. A. Ca 3 (PO 4 ) 2 B. FeBr 3 C. SCl 2 D. Cl 2 O 32

33. Learning Check Determine if each is ionic (I) or covalent (C ), and write the formula. A. calcium nitrate B. boron trifluoride C. aluminum carbonate D. dinitrogen tetroxide E. copper(I) phosphate 33

34. Chapter 4 Compounds and Their Bonds 4.6 Electronegativity and Bond Polarity 34 Copyright © 2009 by Pearson Education, Inc.

35. Electronegativity The electronegativity value indicates the attraction of an atom for shared electrons. increases from left to right going across a period on the periodic table. is high for the nonmetals, with fluorine as the highest. is low for the metals. 35

36. Some Electronegativity Values for Group A Elements 36 Low values High values ` Electronegativity increases ` Electronegativity decreases Copyright © 2009 by Pearson Education, Inc.

37. Nonpolar Covalent Bonds A nonpolar covalent bond occurs between nonmetals. is an equal or almost equal sharing of electrons. has almost no electronegativity difference (0.0 to 0.4). Examples: Electronegativity Atoms Difference Type of Bond N-N 3.0 - 3.0 = 0.0 Nonpolar covalent Cl-Br 3.0 - 2.8 = 0.2 Nonpolar covalent H-Si2.1 - 1.8 = 0.3 Nonpolar covalent 37

38. Polar Covalent Bonds A polar covalent bond occurs between nonmetal atoms. is an unequal sharing of electrons. has a moderate electronegativity difference (0.5 to 1.7). Examples: Electronegativity Atoms DifferenceType of Bond O-Cl 3.5 - 3.0 = 0.5Polar covalent Cl-C 3.0 - 2.5 = 0.5Polar covalent O-S 3.5 - 2.5 = 1.0Polar covalent 38

39. Comparing Nonpolar and Polar Covalent Bonds 39 Copyright © 2009 by Pearson Education, Inc.

40. Ionic Bonds An ionic bond occurs between metal and nonmetal ions. is a result of electron transfer. has a large electronegativity difference (1.8 or more). Examples: Electronegativity Atoms Difference Type of Bond Cl-K 3.0 – 0.8 = 2.2 Ionic N-Na 3.0 – 0.9 = 2.1 Ionic S-Cs2.5 – 0.7= 1.8 Ionic 40

41. Electronegativity and Bond Types 41

42. Predicting Bond Types 42

43. Learning Check Use the electronegativity difference to identify the type of bond [nonpolar covalent (NP), polar covalent (P), or ionic (I)] between the following: A. K-N B. N-O C. Cl-Cl D. H-Cl 43

44. Chapter 4 Compounds and Their Bonds 4.7 Shapes and Polarity of Molecules 44 Copyright © 2009 by Pearson Education, Inc. °

45. VSEPR In the valence-shell electron-pair repulsion theory (VSEPR), the electron groups around a central atom are arranged as far apart from each other as possible. have the least amount of repulsion of the negatively charged electrons. have a geometry around the central atom that determines molecular shape. 45

46. Four Electron Groups In a molecule of CH 4, there are 4 electron groups around C. repulsion is minimized by placing 4 electron groups at angles of 109°, which is a tetrahedral arrangement. the shape with four bonded atoms is tetrahedral. 46 Copyright © 2009 by Pearson Education, Inc.

47. Three Bonding Atoms and One Lone Pair In a molecule of NH 3, 3 electron groups bond to H atoms, and the fourth one is a lone (nonbonding) pair. repulsion is minimized with 4 electron groups in a tetrahedral arrangement. with 3 bonded atoms, the shape is pyramidal. 47

48. Two Bonding Atoms and Two Lone Pairs In a molecule of H 2 O, 2 electron groups are bonded to H atoms and 2 are lone pairs (4 electron groups). 4 electron groups minimize repulsion in a tetrahedral arrangement. the shape with 2 bonded atoms is bent. 48

49. Shapes with 4 Electron Groups Electron Pairs Bonded Atoms Lone Pairs Molecular Shape Example 440TetrahedralCH 4 431 PyramidalNH 3 422BentH2OH2O Copyright © 2009 by Pearson Education, Inc. 49

50. Learning Check State the number of electron groups, lone pairs, and use VSEPR theory to determine the shape of the following molecules or ions. 1) tetrahedral 2) pyramidal3) bent A. PF 3 B. H 2 S C. CCl 4 50

51. Study Tip To determine shape, 1. draw the electron-dot structure. 2. count the electron pairs around the central atom. 3. count the bonded atoms to determine shape. 4 electron pairs and 4 bonded atoms = tetrahedral 4 electrons pairs and 3 bonded atoms = pyramidal 4 electron pairs and 2 bonded atoms = bent 51

52. Polar Molecules A polar molecule contains polar bonds. has a separation of positive and negative charge called a dipole, indicated with  + and  -. has dipoles that do not cancel.  +  - H–Cl H — N — H dipole H dipoles do not cancel 52

53. Nonpolar Molecules A nonpolar molecule contains nonpolar bonds. Cl–Cl H–H or has a symmetrical arrangement of polar bonds. O=C=O Cl Cl–C–Cl Cl dipoles cancel 53

54. Determining Molecular Polarity STEP 1: Write the electron-dot formula. STEP 2: Determine the shape. STEP 3: Determine if dipoles cancel or not. Example: H 2 O.. H ─ O : H 2 O is polar │ H dipoles do not cancel 54

55. Learning Check Identify each of the following molecules as 1) polar or 2) nonpolar. Explain. A. PBr 3 B. HBr C. Br 2 D. SiBr 4 55

56. Chapter 4 Compounds and Their Bonds 4.8 Attractive Forces in Compounds 56 Copyright © 2009 by Pearson Education, Inc. °

57. Ionic Bonds In ionic compounds, ionic bonds are strong attractive forces. hold positive and negative ions together. 57 Copyright © 2009 by Pearson Education, Inc.

58. Dipole-Dipole Attractions 58 In covalent compounds, polar molecules exert attractive forces called dipole-dipole attractions. form strong dipole attractions called hydrogen bonds between hydrogen atoms bonded to F, O, or N, and other very electronegative atoms.

59. Dispersion Forces Dispersion forces are weak attractions between nonpolar molecules. caused by temporary dipoles that develop when electrons are not distributed equally. 59

60. Attractive Forces 60

61. Melting Points and Attractive Forces Ionic compounds require large amounts of energy to break apart ionic bonds. Thus, they have high melting points. Hydrogen bonds are the strongest type of dipole-dipole attractions. They require more energy to break than other dipole- dipole attractions. Dispersion forces are weak interactions and very little energy is needed to change state. 61

62. Melting Points of Some Substances 62

63. Learning Check Identify the main type of attractive forces for each: 1) ionic 2) dipole-dipole 3) hydrogen bonds 4) dispersion A. NCl 3 B. H 2 O C. Br-Br D. KCl E. NH 3 63