1. Ionic Compounds Chemistry – Chapter 8

2. Forming Chemical Bonds Chemical bond – force that holds 2 atoms together Amount of reactivity is directly related to valence electrons Formation of positive ions ▫Cation – positively charged ion ▫Group 1A elements lose 1 valence e-, forming an ion with a 1+ charge  Ex: By losing an e-, Na acquires the stable outer electron configuration of Ne ▫Group 2A elements lose 2 valence e-, forming an ion with a 2+ charge

3. ▫Transition metals commonly lose 2 valence e-, forming 2+ ions; however it’s also possible to lose an additional “d” electron, forming 3+ ion Formation of negative ions ▫Anion – negative ion  To designate anions, -ide is added to root of element  Ex: chloride, sulfide, etc. ▫Nonmetals form a stable outer electron configuration by gaining e- ▫Group 5A gain 3 e-, forming ions w/ 3- charge ▫Group 6A gain 2 e-, forming ions w/ 2- charge

4. Formation and Nature of Ionic Bonds Formation of an ionic bond ▫Ionic bond – electrostatic force that holds oppositely charged particles together  Oxide – ionic bond between metals and oxygen ▫# of e- lost must equal # of e- gained  Ex: Ca and F form the ionic compound CaF 2 ▫Criss-cross method

5. Properties of ionic compounds ▫Strong attraction of positive ions and negative ions in an ionic compound results in a crystal lattice ▫Solid ionic compounds are nonconductors of electricity b/c of the fixed positions of the ions ▫Liquid ionic compounds (or those dissolved in water) are conductors of electricity b/c ions are free to move  Electrolyte – an ionic compound whose aqueous solution conducts an electric current ▫Energy and the ionic bond  Endothermic – energy is absorbed during a chemical reaction  Exothermic – energy is released during a chemical rxn

6.  Formation of ionic compounds from positive and negative ions is always exothermic  Attraction of the positive ion for the negative ions close to it forms a more stable system that is lower in energy than the individual ions  Lattice energy – energy required to separate one mole of the ions of an ionic compound  The more negative the lattice energy, the stronger the force of attraction  Directly related to size of the ions bonded ▫Smaller ions have a more negative value b/c nucleus is closer to, and thus has more attraction for, the valence e-  Affected by the charge of the ion ▫Ions w/ larger positive or negative charges have a more negative lattice energy

7. Names and Formulas for Ionic Compounds Formulas for ionic compounds ▫Formula unit – simplest ratio of the ions represented in an ionic compound  Ex: KBr – 1:1 ratio; MgCl 2 – 1:2 ratio ▫Determining charge  Binary ionic compound – ionic compound formed by 2 ions (metal and nonmetal)  Monatomic ion – one-atom ion  Ex: Mg 2+ and Br -  Table 8-5 lists common ions of transition metals and groups 3A and 4A  Oxidation number – charge of a monatomic ion

8. ▫Compounds that contain polyatomic ions  Polyatomic ions – ions made up of more than one atom  A polyatomic ion acts as an individual ion  Exist as a unit, so NEVER change the subscripts of the atoms w/in the ion  Table 8-6 lists common polyatomic ions Naming ions and ionic compounds ▫Most polyatomic ions are oxyanions – polyatomic ion composed of an element, usually a nonmetal, bonded to one or more O atoms

9. ▫Rules for naming nonmetal-oxyanions (ex: N, S):  Ion w/ more O atoms is named using the root of the nonmetal plus suffix –ate  Ion w/ fewer O atoms is named using the root of the nonmetal plus suffix –ite  Ex: NO 3 - nitrate NO 2 - nitrite SO 4 2- sulfate SO 3 2- sulfite

10. ▫Rules for naming halogen-oxyanions:  Oxyanion w/ most O atoms is named using prefix per-, root of the nonmetal, and suffix –ate.  Oxyanion w/ one less O atom is named using root of the nonmetal and suffix –ate.  Oxyanion w/ 2 less O atoms is named using root of the nonmetal and suffix –ite.  Oxyanion w/ 3 less O atoms is named using prefix hypo-, root of the nonmetal, and suffix –ite.  Ex: ClO 4 - perchlorate ClO 3 - chlorate ClO 2 - chlorite ClO - hypochlorite

11. ▫Some things to remember:  Groups 1A and 2A metals have only one oxidation number  Transition metals and metals on the right side of the periodic table often have more than one oxidation number  Ex: IonsFormulaCompound Name Fe 2+ & O 2- FeOIron(II) oxide Fe 3+ & O 2- Fe 2 O 3 Iron(III) oxide

12. Metallic Bonds and Properties of Metals Metallic bonds – outer energy levels of the metal atoms overlap ▫Electron sea model – all the metal atoms in a metallic solid contribute their valence electrons to form a “sea” of electrons  Electrons in the outer energy levels of the bonding metallic atoms are not held by any specific atom and can move easily from one atom to the next; these are called delocalized electrons ▫Metallic bond – attraction of a metallic cation for delocalized electrons

13. ▫Properties of metals  Melting points vary greatly, but in general metals have high melting and boiling points  Ex: Hg is liquid at room temp.  Ex: W has a melting point of 3422 0 C, making it useful for light bulb filaments and spacecraft parts  Durable  Delocalized electrons in metal are free to move, keeping metallic bonds intact  Delocalized electrons move heat from one place to another more quickly than other materials

14. Metal alloys – mixture of elements that has metallic properties ▫Ex: steel is a mixture of iron and at least one other element ▫Substitutional alloy – has atoms of the original metallic solid replaced by other metal atoms of similar size  Ex: sterling silver ▫Interstital alloy – form when the small holes in a metallic crystal are filled with smaller atoms  Ex: carbon steel – holes in iron crystal are filled with carbon, making the solid harder and stronger