Chapter 15 Ionic Bonding 15.1 Objectives Use the periodic table to infer the number of valence electrons in an atom and draw its electron dot (Lewis dot) structure. Describe formation of cations from metals and anions from nonmetals California Standards 1d. Students know how to use the periodic table to determine the number of electrons available for bonding. 2e. Students know how to draw Lewis dot structures.

ELECTRONS AVAILABLE FOR BONDING Valence Electrons: ELECTRONS AVAILABLE FOR BONDING (the red ones)

Valence Electrons Valence electrons are electrons in the outmost shell (energy level). They are the electrons available for bonding. The number of valence electrons largely determines the chemical properties of that element. For Groups 1A-7A, the number 1-7 is the number of valence electrons for that atom. Group 0 is an exception – you can think of it as group 8A because all the noble gases (except He) have 8 valence electrons.

Group 1 (alkali metals) have 1 valence electron

Group 2 (alkaline earth metals) have 2 valence electrons

Group 13 elements have 3 valence electrons

Group 14 elements have 4 valence electrons

Group 15 elements have 5 valence electrons

Group 16 elements have 6 valence electrons

Group 17 (halogens) have 7 valence electrons

Group 18 (Noble gases) have 8 valence electrons, except helium, which has only 2

Transition metals (“d” block) have 1 or 2 valence e- . Why?

Lanthanides and actinides (“f” block) have 1 or 2 valence electrons

Valence Electrons Valence electrons are usually the only e- used to bond to other atoms. Therefore you usually only show the valence e- in electron dot structures. Electron dot structures are diagrams that show valence e- as dots.

Generic Dot Notation An atom’s valence electrons can be represented by electron dot (AKA Lewis dot) notations. 1 valence e- X 2 valence e- 3 valence e- 4 valence e- 5 valence e- 6 valence e- 7 valence e- 8 valence e- ?

Li Be B C N O F Ne Dot Notations – Period 2 Lewis dot notations for the valence electrons of the elements of Period 2. lithium Li beryllium Be boron B carbon C nitrogen N oxygen O fluorine F neon Ne

Electron Dot Structures Note how two dots per side are drawn x four sides = 8 dots maximum. Note how each side gets one before any side gets two. See how the number of dots is the same for each element within a group (column).

Octet Rule The Octet Rule was created by Gilbert Lewis in 1916. That’s why these diagrams are sometimes called Lewis dot structures. In forming compounds, atoms tend to achieve the e- configuration of a noble gas, 8 valence e-. An octet is a set of 8. Each noble gas (except He) has 8 valence electrons in their highest principle energy level, and the general configuration is ns2np6 (like 2s22p6 or 3s23p6)

Metallic vs. Nonmetallic Elements Atoms of the metallic elements (including column 1A and 2A) tend to lose their outer shell valence e- so they can have a complete octet at the next energy level down. Atoms of nonmetallic elements tend to gain e- (steal e-) or share e- with another nonmetallic element to achieve their complete octet. There are exceptions but the octet rule usually applies to most atoms in compounds.

Cations and Anions Before Na 1s2 2s2 2p6 3s1 After Na+ 1s2 2s2 2p6 If an atom loses a valence e- = cation If an atom gains/steals a valence e- = anion Metals create cations because they start with 1 to 3 e- and usually get all of the valence e- stolen so they can get down to a full lower level octet. Example: Sodium loses 1 e- Before Na 1s2 2s2 2p6 3s1 After Na+ 1s2 2s2 2p6 (note the 8 e- in the n=2 shell) Like Ne 1s2 2s2 2p6 (Neon has 8 e- in the n=2 shell) The change is written as follows: Na· Na+ + e-

Cations Cations of group 1A alkali metals +1 Cations of group 2A alkaline earth metals +2 ·Mg· Mg2+ + 2e- For transition metals, the charges on the cations may vary, refer to roman numerals. Example: Fe has two: iron(II) or Fe2+ iron(III) or Fe3+ Some atoms formed by transition metals do not have noble-gas electron configurations and are therefore exceptions to the octet rule.

Exceptions Example: Ag Silver 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1 4d10 (oddball) Silver would have to lose 11 electrons to get down to noble gas Krypton’s configuration. To gain enough e- to get to Xenon’s configuration, it would have to gain 7 electrons. Neither one is likely. But if Ag loses its one 5s1 electron, then it has an outer shell with 18 e- (the 4 shell), which is a full shell, and relatively favorable. Therefore Ag always forms the Ag+ cation.

Exceptions Other elements that behave similarly are Copper(I) – Cu+ , Gold(I) – Au+ ,Cadmium Cd2+ Mercury(II) – Hg2+ All have pseudo-noble gas configurations. Note: The column Cu and Ag and Au is in is labeled column 1B in your book on page 392. Cd is column 2B. Can you now see why?

Anions Anions are atoms or groups with a negative charge (extra electrons). Atoms of nonmetallic elements have relatively full valence shells and are looking to steal e- to make their shells full. Cl 1s2 2s2 2p6 3s2 3p5 neutral atom Cl- 1s2 2s2 2p6 3s2 3p6 anion Ar 1s2 2s2 2p6 3s2 3p6 now Cl- has Ar config.

Anions The ions that are produced when atoms of chlorine and other halogens gain electrons are called halide ions. Here are some common anions:

Section 15.2 Ionic Bonding California Standards Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds. Students know salt crystals, such as NaCl, are repeating patterns of positive and negative ions held together by electrostatic attraction.

Ionic vs. Covalent Bonds Bonds: Forces that hold groups of atoms together and make them function as a unit. Ionic bonds – transfer of electrons Covalent bonds – sharing of electrons (this will be Ch. 16)

Ionic Bonding Na: 1s22s22p63s1 now Na+ 1s22s22p6 Cl: 1s22s22p63s23p5 now Cl- 1s22s22p63s23p6

Aluminum has three valence e- to steal, and the Bromine atoms would each like to steal one e-. So the Aluminum atom gives up three electrons and the Bromine atoms each receive one.

Examples of Ionic Compounds Mg2+ + 2Cl1- → MgCl2 Magnesium chloride: Magnesium loses two electrons and each chlorine gains one electron 2Al3+ + 3S2- → Al2S3 Aluminum sulfide: Each aluminum loses three electrons (six total) and each sulfur gains two electrons (six total)

Two K atoms lose 1 e- each => One O atom gains 2 e-

3 Mg atoms lose x 2 e- each => 2 N atoms gain 3 e- each

Metal Monatomic Cations Ion name Lithium Li+ Sodium Na+ Potassium K+ Magnesium Mg2+ Calcium Ca2+ Barium Ba2+ Aluminum Al3+

Nonmetal Monatomic Anions Ion Name Fluorine F- Fluoride Chlorine Cl- Chloride Bromine Br- Bromide Iodine I- Iodide Oxygen O2- Oxide Sulfur S2- Sulfide Nitrogen N3- Nitride Phosphorus P3- Phosphide Recall that anions end in –ide.

Sodium Chloride crystal lattice Ionic compounds form solid crystals at ordinary temperatures. Ionic compounds organize in a characteristic crystal lattice of alternating positive and negative ions. All salts are ionic compounds and form crystals.

NaCl – Which is a cation? Anion? Why do they get bigger/smaller?

Question: What is a “formula unit”?

Coordination Number The coordination number of an ion is the number of ions of opposite charge that surround the ion in a crystal. Face-centered cubic – coordination number of 6 (each Na surrounded by 6 Cl atoms) Simple cubic – coordination number of 8

Properties of Ionic Compounds Structure: Crystalline solids Melting point: Generally high Boiling Point: Electrical Conductivity: Excellent conductors, molten and aqueous (not as a solid) Solubility in water: Generally soluble

15.3 Metallic Bonding Students know atoms combine to form molecules by sharing electrons to form covalent or metallic bonds or by exchanging electrons to form ionic bonds.

Metals – sea of electrons Metals are made up of closely packed cations which are surrounded by mobile valence electrons. The mobile valence electrons are often referred to as a “sea of electrons”. Valence electrons do not “belong” to any one cation.

Metallic Bonding Metallic bonding is the chemical bonding that results from the attraction between positively charged metal cations and the surrounding sea of electrons. The sea of drifting electrons insulates the metal cations from one another. The metal cations easily slide past one another like ball bearings immersed in oil. Vacant p and d orbitals in the metal's outer energy levels overlap, and allow outer electrons to move freely throughout the metal.

Properties of Metals Metals are good conductors of heat and electricity Metals are malleable, they can be hammered or forced into shapes Metals are ductile – they can be drawn into wires Metals have high tensile strength Metals have luster

Ductile, Malleable Metals The cations in the ionic crystal aren’t insulated like the metal cations, so upon pressure, it shatters.

Metallic Bonding Strong forces of attraction are responsible for the high melting point of most metals.

Crystalline Structure in Metals Body centered cubic - Every atom has 8 neighbors Na, K, Fe, Cr, W Face centered cubic Every atom has 12 neighbors Cu, Ag, Au, Al, Pb Hexagonal close packed Packed in hexagons Mg, Zn, Cd

Packing in Metals Model: Packing uniform, hard spheres to best use available space. This is called Hexagonal close packed. Each atom has 12 nearest neighbors. Does this look like how oranges are stacked at the grocery store? It should.

Alloys Most metals you use every day are not actually a pure single metal, but an alloy, or blend of metals. Alloys are often prepared by melting a mixture of ingredients together and then cooling the mixture. In semiconductors, we used to sputter a mixture of aluminum, copper and silicon from a target and then alloy the metals together in a 400 C furnace.

Aluminum sputtering target and close up view of aluminum alloy grain structure. If you are interested in metals, here’s a great source to review their properties: http://www.metallographic.com/Technical/Metallography-Intro.html

Alloys Sterling silver Ag 92.5%, Cu 7.5% Stainless steel Fe 80.6%, Cr 18.0%, C 0.4%, Ni 1.0% Cast iron Fe 96%, C 4% Brass Cu and Zn (% varies) Bronze Cu and Sn (% varies)

Metal Alloys Substitutional Alloy: some metal atoms replaced by others of similar size. Note that brass is an alloy of copper and zinc.

Metal Alloys Interstitial Alloy: Interstices (holes) in closest packed metal structure are occupied by small atoms. Note that steel is an alloy of iron, plus other elements such as carbon, molybdenum, chromium or nickel.