2. Chapter 7 Ionic and Metallic Bonding Na + Cl −

3. Section 7.1 - Ions l OBJECTIVES: –Determine the number of valence electrons in an atom of a representative element. –Explain how the octet rule applies to atoms of metallic and nonmetallic elements. –Describe how cations and anions form.

4. 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.

5. 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

6. 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

7. 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.

8. 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

9. Formation of Cations l Metals lose electrons to attain a noble gas configuration. –They make + ions (cations) l If we look at the electron configuration, it makes sense to lose electrons: –Na 1s 2 2s 2 2p 6 3s 1 1 valence e - –Na 1+ 1s 2 2s 2 2p 6 = a noble gas configuration (8 valence e - ), which is the most stable e - configuration

10. Electron (Lewis) Dot Diagrams For Cations l Most metals have few valence electrons; calcium has only 2 Ca

11. Electron (Lewis) Dot Diagrams For Cations l Metals lose valence electrons to become cations Ca

12. Electron Dot Diagrams 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”.

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+ (loses the 5s 1 electron to attain a pseudo-noble gas configuration)

14. Electron Dots For Cations l Silver did not achieve a true Noble Gas configuration l But pseudo-noble gas configuration (valence electrons consist of full d sublevel) is also quite stable

15. Electron Configurations: Anions l Nonmetals gain electrons to attain a stable, 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 halogens (group VII A) 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 tend to react until they achieve a noble gas configuration l Noble gases have 2 s and 6 p electrons in their valence shell l 8 valence electrons = most stable l This is the octet rule (8 in the outer level is particularly stable). Ar

18. Monoatomic/Polyatomic Ions l All ions we’ve seen so far are monoatomic = derived from one atom l There are also many polyatomic ions = groups of bonded atoms that have gained or lost electrons Ammonium (NH 4 + ) Nitrate (NO 3 − ) Phosphate (PO 4 3− ) Sulfate (SO 4 2− )

19. Polyatomic Ions l At right is a table of polyatomic ions that you need to know l Either memorize them, or always have a table ready to use on a test or quiz! Common Polyatomic Ions FormulaName C2H3O2−C2H3O2− Acetate H 2 PO 4 − Dihydrogen-phosphate HSO 4 − Hydrogen-sulfate HCO 3 − Hydrogen-carbonate NO 3 − Nitrate NO 2 − Nitrite OH − Hydroxide MnO 4 − Permanganate ClO − Hypochlorite ClO 2 − Chlorite ClO 3 − Chlorate SO 3 2− Sulfite SO 4 2− Sulfate CO 3 2− Carbonate CrO 4 2− Chromate PO 4 3− Phosphate NH 4 + Ammonium

20. 7.2 Ionic Bonds and Ionic Compounds l OBJECTIVES: – Explain the electrical charge of an ionic compound – Describe the formation of NaCl as a representative ionic compound – Describe three properties of ionic compounds.

21. Ionic Bonding l Anions and cations are held together by the attractive force of their opposite charges –That IS an ionic bond l Ionic compounds are also called salts. l Lowest, whole-number ratio of elements in an ionic compound is called the formula unit l A transfer of electrons (loss and gain) must occur before compound is formed l Electrons are almost always transferred to achieve noble gas configuration.

22. Ionic Compounds Made from a CATION and an ANION, in one of the following combinations: a)metal (+) with a nonmetal (−) b)metal (+) with a polyatomic anion c)ammonium (NH 4 + ) with a nonmetal or polyatomic anion

23. Ionic Bonding NaCl The metal (sodium) tends to lose its one electron from the outer level. The nonmetal (chlorine) accepts the electron that sodium loses.

24. 2 Na(s) + Cl 2 (g)  2 NaCl(s) Formation of NaCl

25. Ionic Bonding l All the electrons must be accounted for, and each atom will attain a stable, noble gas configuration CaP Lets do an example by combining calcium and phosphorus:

26. Ionic Bonding CaP

27. Ionic Bonding Ca 2+ P

28. Ionic Bonding Ca 2+ P Ca

29. Ionic Bonding Ca 2+ P 3- Ca

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

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

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

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

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

35. 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 sample of the substance For an ionic compound, the smallest representative sample is called a: Formula Unit

36. 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

37. Do they Conduct? 3.Solid ionic compounds do not conduct electricity, but melted ionic compounds do l Conducting electricity requires charged particles to move –In a solid, ions are locked in place l When melted, the ions can move around l Dissolved in water, they also conduct (ions free to move in aqueous solutions)

38. - 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.

39. 7.3 Bonding in Metals l OBJECTIVES: –Describe the arrangement of atoms and electrons in a metal –Explain the importance of alloys

40. Characterizing Metallic Bonds l Metals hold on to their valence electrons very weakly and the electrons are believed to move freely between metal atoms l Solid metals are thought of as positive ions (cations) held in place by a sea of surrounding electrons

41. Visualizing Metallic Bonding Metallic Bonding (Flash Media)

42. Characteristics of Metals l Metals are malleable –Can be hammered into shape l Also ductile - drawn into wires. l Both malleability and ductility explained in terms of the mobility of the valence electrons in metallic bonding

43. Malleable ++++ ++++ ++++ Force

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

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

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

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 – Brass: an alloy of Cu and Zn – Bronze: Cu and Sn

48. Why use alloys? l Properties are often superior to pure metal –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 (Fe + a bit of C + (in some cases) small amounts of other metals: Ni, Cr, Mn –corrosion resistant, ductility, hardness, toughness, cost

49. Selected Alloys NameCompositionProperties and Uses brassCu with 20–35% Zndecorative bronzeCu with 5–8% Snprimitive tools, coinage cast iron Fe, C up to 4.5% with Mn, P, Si, and S traces decorative work, weights dental alloyAg, Cu, Zn, Sntooth fillings duralim Al 93–95.5%, Cu 3.5–5.5%, Mn < 1%, Mg < 1% aircraft frames nichromeNi 60%, Fe 20%, Cr 20%filament wires pewterSn 80%, Pb 20%, traces of Sb modern pewter Pb replaced by Sb with some Cu and Bi solderSn 67%, Pb 33%plumbing stainless steelFe 78–86%, Cr 12–20%, Ni 2 corrosion resistant culinary, sterile vessels, decorative steel Fe, with C 0.5–1.5% + special purpose additions of Cr, Mn, Ni, Si, Cu, Co, Ti, W, V multifarious 18 k gold75% Au with Ag and Cujewelry Woods metalBi 50%, Pb 25%, Sn 12.5%, Cd 12.5%low m.p., fire sprinklers