Bonding

1. Bonding Two or more atoms join together to form a stable group. There are several types of forces (BONDS) which hold the atoms together in these groups. IONIC (ionic compounds) COVALENT (molecular compounds) METALLIC (pure elements or homogeneous mixtures) WEAK INTERMOLECULAR (ie.: the forces that hold water molecules together)

2. Ionic Compounds High melting point Conduct electricity when molten or dissolved in water Molecular Compounds Low melting point Does NOT conduct electricity when molten or dissolved in water. Differences in properties

3. An ion is an atom which has “lost” or “gained” one or more electrons, thus obtaining a “net” positive or negative charge. What is an Ion?

4. Cations - ions formed when e- are lost from an atom/group of atoms Positive Ions (+) Usually formed from metal elements Anions - ions formed when e- are gained by an atom/group of atoms Negative Ions (-) Usually formed from nonmetal elements Cations vs. Anions Covalent and ionic bonding

5. Ions in Water A Matter of Ions

6. Ions vs. Moleculeslose or gain e- sharing of e- (between atoms)

7. Valence Electrons • Valence Electrons are the electrons in the Outermost Energy Shell. • The number of valence electrons matches the group number for the representative elements (s & p). • Transition (d) and Inner-transition (f) elements typically have “2” valence electrons. • However, some of their “outer ‘d’” electrons may participate in bonding.

8. Electron Dot Structures for Representative Elements Electron dot structures for selected representatives and noble-gas elements

10. Octet Rule • Atoms will attempt to gain or lose electrons to attain the same electron configuration as a Noble Gas. • Noble Gases have 8 valence electrons (except Helium) • What actually happens is dependent on the elements involved in the reaction! • Metal & Nonmetal = Ionic Bond (eg.: NaCl) • Nonmetal & Nonmetal = Covalent Bond (eg.: H2O)

11. Predicting Ionic Charges Periodic table in which the metallic elements that exhibit a fixed ionic charge are highlighted.

12. Monatomic Ions vs. Polyatomic Ions • Monatomic IONS are comprised of only ONE element. • Ex. Na+, Cl-, Al3+, O2-, etc. • Polyatomic IONS are comprised of two or more elements/atoms. • NH4+, NO3-, SO42-, CO32-, etc.

13. Polyatomic Ions!

14. Ionic Compounds Ionic Compounds are typically Crystalline in form • (a) fluorite and • (b) ruby.

15. Ionic compoundsoccur when cations and anions bond together. These elements join together because of electrostatic attractions between charged IONS. This “joining” occurs in 3D, thus, crystalline in form. Remember, cations are usually metallic elements (or “ammonium”) and anions are usually nonmetallic elements (or MnO4-, CrO4-2, Cr2O7-2, MoO4-2).

16. Which part of the ionic compound is responsible for the typical physical properties? Copper (II) oxide is black, whereas copper (I) oxide is reddish brown. Iron (II) chloride is green, whereas iron (III) chloride is bright yellow.

17. Sodium chloride (a,b) Two-dimensional cross section and a three-dimensional view of sodium chloride. (c) sodium chloride crystals

18. Why NaCl? • Cross section of the structure of the ionic solid NaCl.

19. Reactions of Ionic Compounds (an important example) • Tooth Enamel Demineralization Ca10(PO4)6(OH)2 + 8H+<==> 10Ca2+ + 6PO3- + 2H2O

20. Predicting formulas using the Periodic Table and valence electrons Metal atoms LOSE valence electrons to form cations Nonmetal atoms GAIN valence electrons to form anions Na (1 v.e-) --> Na+ (looks like neon) Cl (7 v.e-) --> Cl- (looks like argon)

21. Naming Ionic Compounds Nomenclature of ionic compounds.

22. Formation of ionic compounds Loss of an electron from a sodium atom leaves it with one more proton than electrons, so it has a net electrical charge of +1.

23. Tests for Ionic Compounds • Flame Tests • Ex: Ba, Na, K • Precipitation Tests • Ex. Pb(NO3)2 + KI