Alkali Metals

2. Increasing EA and IE, Decreasing Size +1 -3 -2 -1 +3 ns2np6 ns1 +2 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 ns2(n-1)d1 ns2(n-1)d10 Increasing EA and IE, Decreasing Size Halogens Inert Gases Alkali Metals Transistion Metals Alkaline Earth Metals

4. Atomic Radii

5. Fundamentals of General, Organic, and Biological Chemistry 5th Edition Chapter Four Ionic Compounds James E. Mayhugh Oklahoma City University 2007 Prentice Hall, Inc.

6. Outline • 4.1 Ions • 4.2 Periodic Properties and Ion Formation • 4.3 Ionic Bonds • 4.4 Some Properties of Ionic Compounds • 4.5 Ions and the Octet Rule • 4.6 Electron-Dot Symbols • 4.7 Ions of Some Common Elements • 4.8 Naming Ions • 4.9 Polyatomic Ions • 4.10 Formulas of Ionic Compounds • 4.11 Naming Ionic Compounds • 4.12 H+ and OH- Ions: An Introduction to Acids and Bases

7. 4.1 Ions • Atoms are electrically neutral because they contain equal numbers of protons and electrons. • By gaining or losing one or more electrons, however, an atom can be converted into a charged particle called an ion. • The symbol for an ion is written by adding the electrical charge as a superscript to the symbol for the element.

8. The lossof one or more electrons from a neutral atom gives a positivelycharged ion called a cation. • Sodium and other alkali metal atoms have a single electron in their valence shell. • By losing this electron, an alkali metal is converted to a cation with a full valence shell.

9. The gainof one or more electrons by a neutral atom gives a negativelycharged ion called an anion. • Chlorine and other halogens have 7 valence electrons. • Halogens easily gain an additional electron to fill out their valence subshell to form anions.

10. 4.2 Periodic Properties and Ion Formation • Ionization energy: The energy required to remove one electron from a single atom in the gaseous state. • Low ionization energies (IE) favor cation formation. • Electron affinity: The energy released on adding an electron to a single atom in the gaseous state. • High electron affinities (EA) favor anion formation. • Going from left to right on the periodic table, IE and EA values both increase.

11. Ionization energies (red) and electron affinities (blue) for the first 36 elements. • Alkali metals have the lowest ionization energies and lose an electron most easily. • Halogens have the highest electron affinities and gain an electron most easily.

12. The noble gases neither gain nor lose an electron easily. They have high IE values and low EA values. • Main group elements near the middleof the periodic table have intermediate values of both IE and EA. • These elements near group 4A lack a strong preference to either lose or gain electrons. • Later we will see that these elements tend not to form ionic bonds but form covalent bonds instead.

17. 4.3 Ionic Bonds When sodium metal reacts with chlorine gas crystals of sodium chloride or table salt are formed.

18. In a violent reaction, sodium atoms transfer electrons to chlorine atoms forming Na+ ions and Cl- ions. • The opposite electrical charges attract each other. Crystals of sodium chloride held together by ionic bonds result that are unlike either of the elements from which they are formed. • Ionic bond: The electrical attractions between ions of opposite charge in a crystal. • Ionic compound: A compound that contains ionic bonds.

19. Na+ and Cl- ions in a sodium chloride crystal. • Each Na+ion is surrounded by six Cl-ions, and each Cl-ion is surrounded by six Na+ions. • The crystal is held together by ionic bonds—the attraction between oppositely charged ions.

20. 4.4 Some Properties of Ionic Compounds • Ionic compounds are usually crystalline solids. • Ions in an ionic solid are held rigidly in place by attraction to their neighbors and cannot move about. • Once an ionic solid is dissolved in water or melted, the ions can move freely and conduct electricity. • High melting points and boiling points are observed for ionic compounds because the attractive forces are extremely strong. • Sodium chloride melts at 801ºC and boils at 1413ºC.

21. Ionic solids are not malleable and will shatter if struck sharply. • Ionic compounds dissolve in water if the attraction between water and the ions overcomes the attraction of the ions for one another. • Sodium chloride and some other familiar ionic compounds are very soluble and can be dissolved to make solutions of high concentration. • Other ionic compounds are not water-soluble, because water is unable to overcome the ionic forces in many crystals.

22. d- d+ H2O Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.

23. 4.5 Ions and the Octet Rule • Main group elements often combine in such a way that each winds up with an electron configuration like a noble gas. Usually 8 valence electrons or an electron octet. • Octet rule: Main group elements tend to undergo reactions that leave them with 8 valence electrons.

24. 4.6 Electron-Dot Symbol Electron-dot symbol: An atomic symbol with dots placed around it to indicate the number of valence electrons.

25. 4.7 Ions of Some Common Elements • Ionic charges of main group elements can be predicted using the group number and the octet rule. • For 1A, 2A, and 3A metals, charge = group number • For nonmetals in groups 5A, 6A, and 7A, anion charge = (group number) - 8.

27. Ions formed by elements in the first four periods.

28. Increasing EA and IE, Decreasing Size +1 -3 -2 -1 +3 ns2np6 ns1 +2 ns2np1 ns2np2 ns2np3 ns2np4 ns2np5 ns2 ns2(n-1)d1 ns2(n-1)d10 Increasing EA and IE, Decreasing Size Halogens Inert Gases Alkali Metals Transistion Metals Alkaline Earth Metals

29. 4.8 Naming Ions Main group metal cations in group 1A, 2A, and 3A are named by identifying the metal, followed by the word “ion” as shown below: • K+ Potassium ion • Mg2+ Magnesium ion • Al3+ Aluminum ion For transition metals which can form more than one type of cation, two systems are used.

30. In the old system, the ion with the smaller charge ends in -ousand the ion with the larger charge ends in -ic. • The new system uses Roman numerals to indicate the charge on the ion.

31. Anions are named by replacing the ending of the element name with -ide, followed by the word “ion.” • For example, the anion formed by fluorineis the fluorideion, and the anion formed by sulfuris the sulfideion.

32. 4.9 Polyatomic Ions • Ions that are composed of more than one atom are called polyatomic ions. • Most polyatomic ions contain oxygen and another element, and their chemical formulas show by subscripts how many of each type of atom are combined. • Sulfate ion, for example, is composed of one sulfur atom and four oxygen atoms, and has a charge of -2; the entire group of atoms acts as a single unit.

34. 4.10 Formulas of Ionic Compounds • A chemical formula shows the simplest ratio of anions and cations required for a total charge of zero. • A shortcut is to make the subscript of each ion equal to the charge on the other ion.

35. Once the numbers and kinds of ions in a compound are known, the formula is written using the following rules: • List the cation first and the anion second; for example, NaCl not ClNa. • Make sure to eliminate any common factors from the subscripts; for example, MgO not Mg2O2. • Do not write the charges of the ions; for example, KF not K+F- • Use parentheses around a polyatomic ion formula if it has a subscript; for example, Al2(SO4)3 not Al2SO43.

38. 4.11 Naming Ionic Compounds • Compounds are named by citing first the cation and then the anion, with a space between words. • For example, MgSO4 is called magnesium sulfate. • Some metals form more than one ion. We need to specify the charge on the cation in these compounds. The old and new ways to do this are shown in the examples below: • SnCl2 - Stannous chloride or Tin (II) chloride • SnCl4 - Stannic chloride or Tin (IV) chloride

42. 4.12 H+ and OH- Ions: An Introduction to Acids and Bases • The importance of the H+ cation and the OH- anion is that they are fundamental to the concepts of acidsand bases. • Acid: A substance that provides H+ ions in water; for example, HCl  H+ + Cl- • Base: A substance that provides OH- ions in water; for example, NaOH  Na+ + OH-

43. d- d+ H2O

46. Chapter Summary • Atoms are converted into cationsby the loss of one or more electrons and into anionsby the gain of one or more electrons. • Ionic compounds are composed of cations and anions held together by ionic bonds. Ionic compounds conduct electricity when dissolved in water, and they are generally crystalline solids with high melting points and high boiling points. • Main group elements tend to form ions in which they have gained or lost the appropriate number of electrons to reach a noble gas configuration.

47. Chapter Summary Cont. • Metals lose electrons more easily than nonmetals. As a result, metals usually form cations. • Nonmetals gain electrons more easily than metals. As a result, reactive nonmetals usually form anions. • The ionic charge can be predicted from the group number and the octet rule. For main group metals, the charge on the cation is equal to the group number. For nonmetals, the charge on the anion is equal to the group number - 8. • Ionic compounds contain appropriate numbers of anions and cations to maintain overall neutrality.

48. Chapter Summary Cont. • Cations have the same name as the metal, monatomic anions have the name ending -ide. • For metals that form more than one ion, a Roman numeral equal to the charge is used to name the ion. • The cation name is given first, with the charge of the metal ion indicated if necessary, and the anion name is given second. • An acidis a substance that yields H+ ions when dissolved in water, and a baseis a substance that yields OH- ions when dissolved in water.

49. End of Chapter 4