CHAPTER 6

1. CHAPTER6 • Ionic Bonds and • Some Main-Group Chemistry

2. Chemical Properties • Chemical properties and the reactivities of various elements can be predicted based upon their electronic configurations. • Only the outermost electrons or valence electrons are involved in chemical reactivity and properties.

3. Noble Gases • The noble gases are characterized by having completely filled outer shells. • This configuration is very stable and the gases have very low reactivity (virtually inert). • He 1s2 • Ne 1s22s2 2p6 • Ar 1s2 2s2 2p63s2 3p6 • Kr 1s2 2s2 2p6 3s2 3p64s2 4p6 • etc.

4. Representative Elements • The elements in A groups on periodic chart are known as representative elements. • This is because they are easiest to predict the properties of and best represent what we know about elemental structure and periodicity.

5. d - Transition Elements • The elements in B groups on periodic chart are known as transition metals or d-transition elements. All are metals with a valence shell of ns1orns2. There is little variation of properties since inner d-electrons are being added to the elements. They are the ‘transition’ from pure metals to the nonmetals.

6. f - Transition Elements • These are known as the inner transition metals or f-transition metals. All have an outer valence of ns2. There is very little variation in the properties of these elements since their electrons are being added 2 shells below the valence shell! Moseley’s triumph

7. radii increases radii increases Atomic Radii (A Group Elements Only) • Atomic radii describes the size of atoms. • This increases as you go from the right to the left. (opposite what you would think since more e– are added.) • This increases from top to bottom. • (as expected – more shells are being added)

8. Atomic Radii All radii are in angstroms, Å.

9. Atomic Radii • The decreasing radii across a period is due to the shielding or screening effect of the inner electrons [He] or [Ne], etc. • Consequently the outer electrons feel a stronger effective nuclear charge than expected. • Li [He] shields effective charge is +1 • Be [He] shields effective charge is +2 • F [He] shields effective charge is +7 • Na [Ne] shields effective charge is +1

10. Ionization Energy - Potential • Ionization energy is the minimum amount of energy required to remove the most loosely held electron from an isolated gaseous atom. • This is the measure of an element’s ability to form positive ions. • first ionization energy • Atom(g) + energy ® ion+(g) + e- • Mg(g) + 738kJ/mol ® Mg+ + e-

11. Ionization Energy - Potential • second ionization energy • This is the energy required to remove a 2nd electron from an gaseous ion. • ion+ + energy ® ion2+ + e- • Mg+ + 1451 kJ/mol ® Mg2+ + e- • Additional ionization energies are possible, these increase exponentially with charge.

12. IE increases IE decreases Ionization Energy - Potential • Ionization energy, IE, generally increases as you go across a period (smaller radii, e– more tightly held). • Important exceptions at Be and Mg, N, and P, because of Hund’s rule Ionization energy generally decreases as you go down a group (larger radii, e– more loosely held).

13. Ionization Energy The unexpected dips in these trends are due to Hund’s rule. In each case the loss of an electron leads to a more stable half filled orbital.

14. Electron Affinity • Electron affinity, EA, is amount of energy absorbed when an electron is added to an isolated gaseous atom to from a –1 ion. • EA has a negative value when energy released (more stable anion formed) and a positive value when energy is absorbed (unstable anion formed). • EA is a measure of an atom’s ability to form negative ions – the more negative the value the more readily these are formed. atom(g) + e–+ EA  ion–(g)

15. EA more negative EA more negative Electron Affinity EA becomes more negative as you go up the periodic table. • EA becomes more negative as you go from left to right. • Exceptions exist when you add e– to half-filled or filled orbitals and produce a less stable ion. • Mg(g) + e–+ 231 kJ/mol ® Mg-(g) EA = +231 kJ/mol • Br(g) + e–® Br–(g) + 323 kJ/mol EA = –323 kJ/mol

16. Electron Affinity The peaks seen in the plot are where electrons are added to elements with filled or half-filled atomic orbitals.

17. Ionic radii increases Ionic radii increases Ionic Radii • Ionic radii increases as you go from right to left across the periodic table. Ionic radii increases as you go from down the periodic table.

18. Ionic Radii • Cations (+ ions) are always smaller than their neutral atoms. • The larger the + charge, the smaller the ion. • Anions (- ions) are always larger than their neutral atoms. • The larger the – charge, the larger the ion. • For the same number of electrons, anions are larger than cations.

19. Electronegativity increases EN increases Electronegativity • Electronegativity, EN, is the measure of the tendency of an atom to attract electrons to itself in compounds (to gain electrons). • EN increases as you go from left to right (to F). EN increases as you go up the table.

20. Chemical Bonding • Chemical bonds are classified into two types: Ionic bonding results from electrostatic attractions among ions, which are formed by the transfer of one or more electrons from one atom to another. Ionic bonds are common between metals and nonmetals. Covalent bonding results from sharing one or more electron pairs between two atoms. Covalent bonds occur between nonmetals. • A compound may contain both ionic and covalent bonds.

21. The Octet Rule • Representative elements achieve noble gas configurations in most of their compounds. • Metals tend to lose electrons to drop back to a noble gas configuration. • Nonmetals tend to gain or share electrons to attain a noble gas configuration. • With the exception of He, all noble gases have 8 valence electrons. • This tendency to achieve a valence “noble gas” configuration is called the octet rule. • The octet rule was considered absolute until the late 1960s.

22. Comparison of Ionic and Covalent Compounds • IonicCovalent • Melting Pt HighLow • (> 700 °C)(< 300 °C) • Solubility • polar solventsSolubleInsoluble • nonpolar solventsInsolubleSoluble • Conductivity • solidLowLow • moltenHighLow • aqueousHighLow

23. DEN • Bonds between metals and non-metals tend to be ionic (DEN >1.7) • Bonds between non-metals tend to be covalent (DEN <1.7) • The magnitude of DEN establishes the direction and strength of the dipole moment in a compound or ion.

24. Net reaction Born-Haber Cycle for NaCl -348.6 kJ/mol 495.8 kJ/mol 122 kJ/mol -787 kJ/mol 107.3 kJ/mol -411 kJ/mol This is Hess’ Law coming in Chapt 8

25. Intermolecular Attractions and Phase Changes • Ion-Ion interactions • The force of attraction between two oppositely charged ions is determined by Coulomb’s law

26. Intermolecular Attractions and Phase Changes • The energy of attraction between two ions is given by: Generally it is the magnitudeof the charge and not the number of ions with that charge that are considered when evaluating this strength of attraction.

27. Intermolecular Attractions and Phase Changes • Coulomb’s Law and the Attraction Energy determines: • melting points of ionic compounds • boiling points of ionic compounds • the solubility of ionic compounds • the heat of solvolysis of ionic compounds

28. Intermolecular Attractions and Phase Changes • Example: Arrange the following ionic compounds in the expected order of increasing melting and boiling points. • NaF, CaO, CaF2 • Consider the relative charges on the ions. (the greater the charge, the greater the attraction) • Consider the relative size of the ions (this equals separation of charge).

29. Reactions of Selected Main Group Elements

30. Alkali Metals (Group IA) • Alkali metals are extremely reactive, thus they are not found free in nature. • The alkali metals have only one valence electron (ns1). • Group IA metals have low first ionization energies. • Ionization energies decrease with increasing size of the alkali metals. • All the alkali metals form stable 1+ ions. • All the IA salts are soluble (virtually no exceptions)

31. Reactions of Group IA Metals • The alkali metals are strong reducing agents that form 1+ cations by losing one electron per metal atom. • Consider the reaction of lithium with oxygen, a Group VIA nonmetal, to form lithium oxide, Li2O, an ionic compound.

32. Reactions of Group IA Metals • Sodium reacts with oxygen to form sodium peroxide, Na2O2. The charge on the peroxide ion, O2-2, is -2, and the oxidation number of each oxygen is -1. Potassium and the heavier alkali metals react with oxygen to form superoxides, MO2. The charge on the superoxide ion, O2-, is -1 ; the oxidation number of each oxygen is -1/2 .

33. Reactions of Group IA Metals • The alkali metals react readily with water to form hydroxides and H2. • The reaction of sodium with water is (molecular equation) The ionic equation for this reaction is

34. Reactions of Group IA Metals • The alkali metals react with H2 to form the metal hydride. • The reaction is shown below All alkali metals form a hydride. Typically the metal is reacted while in the molten state.

35. Reactions of Group IA Metals • All alkali metals react with N2 to form a nitride. Lithium nitride is the most stable of the alkali nitrides. The nitrides react explosively with water to yield ammonia and 3 equivalents of the metal hydroxide

36. Reactions of Group IA Metals • The alkali metals react with the halogens to form the corresponding salt. • The general reaction is shown below. The alkali metals react with the other VIA elements to yield the expected product.

37. Reactions of Group IA Metals • The alkali metals react with P to form phosphides (exactly like the nitrides) The alkali metals react with P to form These react explosively with water to yield phosphine and 3 equivalents of the metal hydroxide

38. Reactions of Group IA Metals Dissolving Metal Reactions • The alkali metals dissolve/react with liquid ammonia to make M+ and dissolved e A small quantity of sodamide, NaNH2 (or other metal amide) is also generated. Li and Na are most commonly used.

39. The Group IA Metals • Lithium • Greek lithos meaning stone • Discovered by Arfvedson (Sweden) in 1817 • Lightest of metals • Excellent heat transfer agent • Used in organic synthesis (dissolving metals, organolithium compounds) • Used in glass manufacture • Medicine – treatment for manic depression

40. The Group IA Metals • Sodium • soda or sodanum, headache remedy, Latin natrium • discovered by Davy • Most widely used of the alkali metals • Used in sodium street lamps • Organic synthesis, dissolving metal reactions and sodium hydride • NaK used for inert atmospheres and heat transfer • Manufacture of soap, water softening, industry

41. The Group IA Metals • Potassium • potash, ashes; Latin kalium • discovered by Davy • Common in plants • Used for finest soaps • Highly soluble salts • Used as a sodium replacement, Lite Salt • Used extensively in organometallic chemistry • Violently reactive metal

42. The Group IA Metals • Rubidium • Latin rubidius meaning dark or deepest red • possible use in "ion engines" for space vehicles • used as a "getter" in vacuum tubes • photocell component • used for making special glasses • RbAg4I5 has the highest room temperature conductivity of any known ionic crystal. possible use in thin film batteries

43. The Group IA Metals • Cesium • Latin, caesius, meaning sky or heavenly blue • Discovered by Kirchhoff and Bunsen • Most common in Manitoba, Canada, lepiodite • Strongest base • Used in organic synthesis, catalyst • Used in atomic clocks • Used in IR lamps • Possible rocket fuel for proposed “ion” propulsion systems • Liquid at room temperature

44. The Group IA Metals • Francium • named after France, discovered in 1939 at the Curie Institute • (picture of uraninite which contains Fr due to constant radioactive decay) • Heaviest IA element • Highly radioactive • No measurable amount ever produced • Predicted to have properties similar to Cs

45. Reactions of Group IA Metals • Diagonal Similarities exist between elements in successive groups near the top of the periodic table. • IA IIA IIIA IVA • Li Be B C • Na Mg Al Si • Li and Mg have similar charge densities and electronegativities. • Li compounds are similar to Mg compounds

46. Reactions of Group IA Metals • The Group IA metal oxides are basic. • They react with water to form strong soluble bases. • The molecular reaction of sodium oxide with water is The ionic equation for this reaction is

47. Group IIA Metals • Alkaline earth metals are silvery white, malleable, ductile, and somewhat harder than Group IA metals. • All have two valence electrons (ns2). • The ionization energies are greater for Group IIA than for Group IA metals.

48. Group IIA Metals • Most compounds of Group IIA metals are ionic • Be compounds show a great deal of covalent character similar to compounds of aluminum in Group IIIA. • The Group IIA metals show the +2 oxidation state in all their compounds. • Ease of formation of 2+ ions increases from Be to Ra.

49. Group IIA Metals • Group IIA metals are less reactive than those of Group IA.Still too reactive to occur free in nature. • The metals are prepared by various methods Ca & Mg are abundant in minerals Be, Sr & Ba are less common Ra is radioactive and is rare

50. Reactions of Group IIA Metals • Except for Be, all the Group IIA metals are oxidized in air to oxides. • The IIA oxides (except for BeO) are basic and react with water to give hydroxides. • Calcium oxide or lime, CaO, reacts with water to give calcium hydroxide or slaked lime, Ca(OH)2.