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Ionic Compounds

Ionic Compounds. Ions are atoms that have gained or lost electron(s) Atoms tend to make ions with characteristic oxidation states (charges) Metals are losers (+ ions) Nonmetals are gainers (- ions) Some atoms do not readily make ions (C, Si, many metalloids). Ionic Bonds.

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Ionic Compounds

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  1. Ionic Compounds • Ions are atoms that have gained or lost electron(s) • Atoms tend to make ions with characteristic oxidation states (charges) • Metals are losers (+ ions) • Nonmetals are gainers (- ions) • Some atoms do not readily make ions (C, Si, many metalloids)

  2. Ionic Bonds • An ionic bond is formed between two or more oppositely charged ions • Ionic compounds are made of a metal (+) and a nonmetal (-) • Ionic compounds are called salts • The overall charge on an ionic compound is zero • When a metal and nonmetal react, electrons are transferred

  3. Making Ionic Bonds - + Na Cl 17p+ 11p+ 11e- 17e- 10e- 18e-

  4. Making More Ionic Bonds +2 -2 Mg O 12 p+ 8p+ 12e- 8e- 10e- 10e-

  5. Formulas and Names • NaCl, sodium chloride • MgO, magnesium oxide • Binary salts • metal, then nonmetal • nonmetal ending changed to “ide” • no subscripts when ratio is 1:1

  6. Unequal charges 11 p+ +1 11e- Na -2 -1 10e- S +1 Na 11 p+ 11e- 16p+ 10e- 16e- Formula: Na2S 18e- Name: sodium sulfide Total + charge = 2, Total - charge = 2 Total charge overall = 1 + 1 + (-2) = 0

  7. Unequal charges Determine the formula of calcium bromide. -1 Br +2 +1 Ca -1 Br Formula: CaBr2

  8. Polyatomic ions Transition metal salts Salts of polyatomic ions

  9. Solubilities • Salts are soluble in water if ion-water interactions can supply enough energy to break apart the crystal lattice • Salts of lower-charged ions are more likely to be soluble (lower lattice energy) • All alkali metal and ammonium salts are soluble • All nitrates are soluble • All oxides are insoluble (alkali metal oxides react to form hydroxides)

  10. Ionic compound properties • Made of metal and nonmetal (except ammonium and organic base salts) • High MP (chemical bonds are broken in melting) • Crystal lattice • Brittle • Form ions in water solution (ionization) NaCl  Na+ + Cl - • Conduct electricity when melted

  11. Hydrates • Water can get trapped in crystal lattice of a crystallized salt Na2CO3.10H2O Sodium carbonate decahydrate CuSO4.5H2O Copper (II) sulfate pentahydrate Sodium acetate trihydrate NaC2H3O2.3H2O

  12. Hydrates • Some salts take water out of the air to become hydrates: hygroscopic • Example: Na2CO3 • Others take enough water to become solutions: deliquescent • Example: CaCl2

  13. Crystal Lattices and Energy • Regular repeating arrangement of ions is a crystal lattice • Energy holding lattice together is the lattice energy • Energy is released when lattice is formed (from gaseous ions) and absorbed when it is broken

  14. Crystal Lattices and Energy • Lattice energy is measured from the viewpoint of the system • When gaseous ions come together to form a crystal energy leaves the system • Since system energy is lower, lattice energy is always given as a negative value

  15. Crystal Lattices and Energy • Magnitude of lattice energy is directly proportional to charge density • Charge density is related to charge magnitude and ion size • Crystallization from gaseous ions is always negative; crystallization from solution can be negative or positive

  16. Metallic Bonds • Metals form molecular orbitals that cover the entire crystal • Electrons can move anywhere in the orbital, so metals conduct heat and electricity well • Metallic bonds are non-directional, so metals are malleable and ductile • Strength of metallic bonds depends on the number of mobile electrons in the bond per atom • Transition metals have mobile s and d electrons, so they are stronger and harder than alkali metals (only 1 s electron is mobile)

  17. Metal Alloys • Alloys are solid solutions of one or more metals • Substitutional alloy: made by metals with atoms of similar size • Interstitial alloy: made by metals with very different atomic sizes • Adding nonmetals (such as carbon to iron) makes directional bonds • Directional bonds make alloys harder, stronger and more brittle

  18. Covalent Bonds • Nonmetals of similar electronegativity cannot form ionic bonds • These atoms share electrons to complete their octet • Shared electrons “count” for both atoms • Each atom’s nucleus attracts the other atom’s electrons

  19. Forming Covalent Bonds Single bond, 2 electrons 8 e-! 2 e-! Cl H Shared!

  20. Multiple Bonds Pi (p) bond electron density above and below nuclei Double bond 8e-! Needs 1e-, makes 1 bond H C O 2e-! H Needs 2 e-, Makes 2 bonds Needs 4 e-, makes 4 bonds 8e-! Sigma (s) bond Electron density between nuclei

  21. Molecular Dot Structures • Count electrons – all valence electrons must appear in final structure • Follow octet rule • Remember how many bonds each type of atom makes (one for each extra electron needed)

  22. Polyatomic Ion Dot Structures • Same as molecular dot structures, except electrons must be added or subtracted to account for ion charge • Subtract electrons for + charge, add for – charge • Make all structures as symmetrical as possible

  23. Carbonate (CO3-2) Dot Structure Symmetry! -2 O O C O Count electrons! 6 + 6 + 6 + 4 + 2 = 24

  24. Molecular Substances • Made of molecules, which are loosely held together – van der Waals or London Dispersion forces • Tend to be liquids, gases or low melting solids • Melting molecular solids involves separating molecules from each other • Most are insulators

  25. Formulas and Names of Small Molecules • Many have common names (i.e. water, ammonia) • Systematic names use prefixes for each element • P2O5 – diphosphorus pentoxide • N2O – dinitrogen monoxide • “mono” is not used for the first element in a compound

  26. Formulas and Names of Small Molecules • CO2 – carbon dioxide • CO – carbon monoxide • SO3 – sulfur trioxide • CCl4 – carbon tetrachloride

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