2. Chapter 7 “Ionic and Metallic Bonding”

3. Valence Electrons are…? l The electrons responsible for the chemical properties of atoms, and are those in the outer energy level. l Valence electrons - The s and p electrons in the outer energy level –the highest occupied energy level l Core electrons – are those in the energy levels below.

4. Keeping Track of Electrons l Atoms in the same column... 1)Have the same outer electron configuration. 2)Have the same valence electrons. l The number of valence electrons are easily determined. It is the group number for a representative element l Group 2A: Be, Mg, Ca, etc. – have 2 valence electrons

5. Electron Dot diagrams are… l A way of showing & keeping track of valence electrons. l How to write them? l Write the symbol - it represents the nucleus and inner (core) electrons l Put one dot for each valence electron (8 maximum) l They don’t pair up until they have to (Hund’s rule) X

6. The Electron Dot diagram for Nitrogen l Nitrogen has 5 valence electrons to show. l First we write the symbol. N l Then add 1 electron at a time to each side. l Now they are forced to pair up. l We have now written the electron dot diagram for Nitrogen.

7. The Octet Rule l In Chapter 6, we learned that noble gases are unreactive in chemical reactions l In 1916, Gilbert Lewis used this fact to explain why atoms form certain kinds of ions and molecules l The Octet Rule: in forming compounds, atoms tend to achieve a noble gas configuration; 8 in the outer level is stable l Each noble gas (except He, which has 2) has 8 electrons in the outer level

8. Formation of Cations l Metals lose electrons to attain a noble gas configuration. l They make positive ions (cations) l If we look at the electron configuration, it makes sense to lose electrons: l Na 1s 2 2s 2 2p 6 3s 1 1 valence electron l Na 1+ 1s 2 2s 2 2p 6 This is a noble gas configuration with 8 electrons in the outer level.

9. Electron Dots For Cations l Metals will have few valence electrons (usually 3 or less); calcium has only 2 valence electrons Ca

10. Electron Dots For Cations l Metals will have few valence electrons l Metals will lose the valence electrons Ca

11. Electron Dots For Cations l Metals will have few valence electrons l Metals will lose the valence electrons l Forming positive ions Ca 2+ NO DOTS are now shown for the cation. This is named the “calcium ion”.

12. Electron Dots For Cations l Let’s do Scandium, #21 l The electron configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 l Thus, it can lose 2e - (making it 2+), or lose 3e - (making 3+) Sc = Sc 2+ Scandium (II) ion Scandium (III) ion Sc = Sc 3+

13. Electron Dots For Cations l Let’s do Silver, element #47 l Predicted configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 9 l Actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 1 4d 10 Ag = Ag 1+ (can’t lose any more, charges of 3+ or greater are uncommon)

14. Electron Dots For Cations l Silver did the best job it could, but it did not achieve a true Noble Gas configuration l Instead, it is called a “pseudo-noble gas configuration”

15. Electron Configurations: Anions l Nonmetals gain electrons to attain noble gas configuration. l They make negative ions (anions) l S = 1s 2 2s 2 2p 6 3s 2 3p 4 = 6 valence electrons l S 2- = 1s 2 2s 2 2p 6 3s 2 3p 6 = noble gas configuration. l Halide ions are ions from chlorine or other halogens that gain electrons

16. Electron Dots For Anions l Nonmetals will have many valence electrons (usually 5 or more) l They will gain electrons to fill outer shell. P 3- (This is called the “phosphide ion”, and should show dots)

17. Stable Electron Configurations l All atoms react to try and achieve a noble gas configuration. l Noble gases have 2 s and 6 p electrons. l 8 valence electrons = already stable! l This is the octet rule (8 in the outer level is particularly stable). Ar

18. Ionic Bonding l Anions and cations are held together by opposite charges (+ and -) l Ionic compounds are called salts. l Simplest ratio of elements in an ionic compound is called the formula unit. l The bond is formed through the transfer of electrons (lose and gain) l Electrons are transferred to achieve noble gas configuration.

19. Ionic Compounds 1)Also called SALTS 2)Made from: a CATION with an ANION (or literally from a metal combining with a nonmetal)

20. Ionic Bonding NaCl The metal (sodium) tends to lose its one electron from the outer level. The nonmetal (chlorine) needs to gain one more to fill its outer level, and will accept the one electron that sodium is going to lose.

21. Ionic Bonding Na + Cl - Note: Remember that NO DOTS are now shown for the cation!

22. Ionic Bonding l All the electrons must be accounted for, and each atom will have a noble gas configuration (which is stable). CaP Lets do an example by combining calcium and phosphorus:

23. Ionic Bonding CaP

24. Ionic Bonding Ca 2+ P

25. Ionic Bonding Ca 2+ P Ca

26. Ionic Bonding Ca 2+ P 3- Ca

27. Ionic Bonding Ca 2+ P 3- Ca P

28. Ionic Bonding Ca 2+ P 3- Ca 2+ P

29. Ionic Bonding Ca 2+ P 3- Ca 2+ P Ca

30. Ionic Bonding Ca 2+ P 3- Ca 2+ P Ca

31. Ionic Bonding Ca 2+ P 3- Ca 2+ P 3- Ca 2+

32. Ionic Bonding = Ca 3 P 2 Formula Unit This is a chemical formula, which shows the kinds and numbers of atoms in the smallest representative particle of the substance. For an ionic compound, the smallest representative particle is called a: Formula Unit

33. Properties of Ionic Compounds 1.Crystalline solids - a regular repeating arrangement of ions in the solid: Fig. 7.9, page 197 –Ions are strongly bonded together. –Structure is rigid. 2.High melting points l Coordination number- number of ions of opposite charge surrounding it

34. - Page 198 Coordination Numbers: Both the sodium and chlorine have 6 Both the cesium and chlorine have 8 Each titanium has 6, and each oxygen has 3 NaCl CsCl TiO 2

35. Do they Conduct? l Conducting electricity means allowing charges to move. l In a solid, the ions are locked in place. l Ionic solids are insulators. l When melted, the ions can move around. 3.Melted ionic compounds conduct. –NaCl: must get to about 800 ºC. –Dissolved in water, they also conduct (free to move in aqueous solutions)

36. - Page 198 The ions are free to move when they are molten (or in aqueous solution), and thus they are able to conduct the electric current.

37. Metallic Bonds are… l How metal atoms are held together in the solid. l Metals hold on to their valence electrons very weakly. l Think of them as positive ions (cations) floating in a sea of electrons: Fig. 7.12, p.201

38. Sea of Electrons ++++ ++++ ++++ l Electrons are free to move through the solid. l Metals conduct electricity.

39. Metals are Malleable l Hammered into shape (bend). l Also ductile - drawn into wires. l Both malleability and ductility explained in terms of the mobility of the valence electrons

40. - Page 201 1) Ductility2) Malleability Due to the mobility of the valence electrons, metals have: and Notice that the ionic crystal breaks due to ion repulsion!

41. Malleable ++++ ++++ ++++ Force

42. Malleable ++++ ++++ ++++ l Mobile electrons allow atoms to slide by, sort of like ball bearings in oil. Force

43. Ionic solids are brittle +-+- + - +- +-+- + - +- Force

44. Ionic solids are brittle + - + - + - +- +-+- + - +- l Strong Repulsion breaks a crystal apart, due to similar ions being next to each other. Force

45. Crystalline structure of metal l If made of one kind of atom, metals are among the simplest crystals; very compact & orderly l Note Fig. 7.14, p.202 for types: 1. Body-centered cubic: –every atom (except those on the surface) has 8 neighbors –Na, K, Fe, Cr, W

46. Crystalline structure of metal 2. Face-centered cubic: –every atom has 12 neighbors –Cu, Ag, Au, Al, Pb 3. Hexagonal close-packed –every atom also has 12 neighbors –different pattern due to hexagonal –Mg, Zn, Cd

47. Alloys l We use lots of metals every day, but few are pure metals l Alloys are mixtures of 2 or more elements, at least 1 is a metal l made by melting a mixture of the ingredients, then cooling l Brass: an alloy of Cu and Zn l Bronze: Cu and Sn

48. Why use alloys? l Properties are often superior to the pure element l Sterling silver (92.5% Ag, 7.5% Cu) is harder and more durable than pure Ag, but still soft enough to make jewelry and tableware l Steels are very important alloys –corrosion resistant, ductility, hardness, toughness, cost

49. More about Alloys… l Table 7.3, p.203 – lists a few alloys l Types? a) substitutional alloy- the atoms in the components are about the same size l b) interstitial alloy- the atomic sizes quite different; smaller atoms fit into the spaces between larger l “Amalgam”- dental use, contains Hg

50. Atoms and Ions - Naming l Atoms are electrically neutral. –Because there is the same number of protons (+) and electrons (-). l Ions are atoms, or groups of atoms, with a charge (positive or negative) –They have different numbers of protons and electrons. l Only electrons can move, and ions are made by gaining or losing electrons.

51. An Anion is… l A negative ion. l Has gained electrons. l Nonmetals can gain electrons. l 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)

52. A Cation is… l A positive ion. 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

53. Predicting Ionic Charges Group 1A: Lose 1 electron to form 1+ ions H1+H1+H1+H1+ Li 1+ Na 1+ K1+K1+K1+K1+ Rb 1+

54. Predicting Ionic Charges Group 2A: Loses 2 electrons to form 2+ ions Be 2+ Mg 2+ Ca 2+ Sr 2+ Ba 2+

55. Predicting Ionic Charges Group 3A: Loses 3 Loses 3 electrons to form 3+ ions B 3+ Al 3+ Ga 3+

56. Predicting Ionic Charges Group 4A: Do they lose 4 electrons or gain 4 electrons? Do they lose 4 electrons or gain 4 electrons? Neither! Group 4A elements rarely form ions (they tend to share)

57. Predicting Ionic Charges Group 5A: Gains 3 Gains 3 electrons to form 3- ions N 3- P 3- As 3- Nitride Phosphide Arsenide

58. Predicting Ionic Charges Group 6A: Gains 2 Gains 2 electrons to form 2- ions O 2- S 2- Se 2- Oxide Sulfide Selenide

59. Predicting Ionic Charges Group 7A: Gains 1 electron to form Gains 1 electron to form 1- ions F 1- Cl 1- Br 1- Fluoride Chloride Bromide I 1- Iodide

60. Predicting Ionic Charges Group 8A: Stable noble gases do not form ions! Stable noble gases do not form ions!

61. Predicting Ionic Charges Group B elements: Many transition elements Many transition elements have more than one possible oxidation state. have more than one possible oxidation state. Iron (II) = Fe 2+ Iron (III) = Fe 3+ Note the use of Roman numerals to show charges

62. Naming cations l Two methods can clarify when more than one charge is possible: 1)Stock system – uses roman numerals in parenthesis to indicate the numerical value 2)Classical method – uses root word with suffixes (-ous, -ic) Does not give true value

63. Naming cations l We will use the Stock system. l Cation - if the charge is always the same (like in the Group A metals) just write the name of the metal. l Transition metals can have more than one type of charge. –Indicate their charge as a roman numeral in parenthesis after the name of the metal (Table 9.2, p.255)

64. Predicting Ionic Charges Some of the post-transition elements also Some of the post-transition elements also have more than one possible oxidation state. have more than one possible oxidation state. Tin (II) = Sn 2+ Lead (II) = Pb 2+ Tin (IV) = Sn 4+ Lead (IV) = Pb 4+

65. Predicting Ionic Charges Group B elements: Some transition elements Some transition elements have only one possible oxidation state, such as these three: have only one possible oxidation state, such as these three: Zinc = Zn 2+ Silver = Ag 1+ Cadmium = Cd 2+

66. Exceptions: l Some of the transition metals have only one ionic charge: –Do not need to use roman numerals for these: –Silver is always 1+ (Ag 1+ ) –Cadmium and Zinc are always 2+ (Cd 2+ and Zn 2+ )

67. Practice by naming these: l Na 1+ l Ca 2+ l Al 3+ l Fe 3+ l Fe 2+ l Pb 2+ l Li 1+

68. Write symbols for these: l Potassium ion l Magnesium ion l Copper (II) ion l Chromium (VI) ion l Barium ion l Mercury (II) ion

69. Naming Anions l Anions are always the same charge l Change the monatomic element ending to – ide l F 1- a Fluorine atom will become a Fluoride ion.

70. Practice by naming these: l Cl 1- l N 3- l Br 1- l O 2- l Ga 3+

71. Write symbols for these: l Sulfide ion l Iodide ion l Phosphide ion l Strontium ion

72. Polyatomic ions are… l Groups of atoms that stay together and have an overall charge, and one name. l Usually end in –ate or -ite l Acetate: C 2 H 3 O 2 1- l Nitrate: NO 3 1- l Nitrite: NO 2 1- l Permanganate: MnO 4 1- l Hydroxide: OH 1- and Cyanide: CN 1- ?

73. l Sulfate: SO 4 2- l Sulfite: SO 3 2- l Carbonate: CO 3 2- l Chromate: CrO 4 2- l Dichromate: Cr 2 O 7 2- l Phosphate: PO 4 3- l Phosphite: PO 3 3- l Ammonium: NH 4 1+ Know Table 9.3 on page 257 If the polyatomic ion begins with H, then combine the word hydrogen with the other polyatomic ion present: H 1+ + CO 3 2- → HCO 3 1- hydrogen + carbonate → hydrogen carbonate ion (One of the few positive polyatomic ions)

74. Writing Ionic Compound Formulas Example: Barium nitrate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Ba 2+ NO 3 - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss- cross method to balance subscripts. Not balan ced! ( ) 2 Now balanc ed. = Ba(NO 3 ) 2

75. Writing Ionic Compound Formulas Example: Ammonium sulfate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! NH 4 + SO 4 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss- cross method to balance the subscripts. Not balan ced! ( ) 2 Now balanc ed. = (NH 4 ) 2 SO 4

76. Writing Ionic Compound Formulas Example: Iron (III) chloride (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Fe 3+ Cl - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss- cross method to balance the subscripts. Not balan ced! 3 Now balanced. = FeCl 3

77. Writing Ionic Compound Formulas Example: Aluminum sulfide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ S 2- 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss- cross method to balance the subscripts. Not balan ced! 23 Now balanced. = Al 2 S 3

78. Writing Ionic Compound Formulas Example: Magnesium carbonate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Mg 2+ CO 3 2- 2. Check to see if charges are balanced. They are balanced ! = MgCO 3

79. Writing Ionic Compound Formulas Example: Zinc hydroxide (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Zn 2+ OH - 2. Check to see if charges are balanced. 3. Balance charges, if necessary, using subscripts. Use parentheses if you need more than one of a polyatomic ion. Use the criss- cross method to balance the subscripts. Not balan ced! ( ) 2 Now balanced. = Zn(OH) 2

80. Writing Ionic Compound Formulas Example: Aluminum phosphate (note the 2 word name) 1. Write the formulas for the cation and anion, including CHARGES! Al 3+ PO 4 3- 2. Check to see if charges are balanced. They ARE balanced! = AlPO 4

81. Naming Ionic Compounds l 1. Name the cation first, then anion l 2. Monatomic cation = name of the element Ca 2+ = calcium ion l 3. Monatomic anion = root + -ide Cl  = chloride CaCl 2 = calcium chloride

82. Naming Ionic Compounds l some metals can form more than one charge (usually the transition metals) l use a Roman numeral in their name: PbCl 2 – use the anion to find the charge on the cation (chloride is always 1-) Pb 2+ is the lead (II) cation PbCl 2 = lead (II) chloride (Metals with multiple oxidation states)

83. Things to look for: 1) If cations have ( ), the number in parenthesis is their charge. 2) If anions end in -ide they are probably off the periodic table (Monoatomic) 3) If anion ends in -ate or –ite, then it is polyatomic

84. Practice by writing the formula or name as required… l Iron (II) Phosphate l Stannous Fluoride l Potassium Sulfide l Ammonium Chromate l MgSO 4 l FeCl 3