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Test 3 Review

Test 3 Review. Electric charge. Mass #. Na. # of atoms. Atomic #. Ionic Compounds. Also referred to as a “salt” Formation involves a transfer of electrons Usually made up of a metal and a non-metal Are good conductors when they can be melted or dissolved

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Test 3 Review

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  1. Test 3 Review

  2. Electric charge Mass # Na # of atoms Atomic #

  3. Ionic Compounds • Also referred to as a “salt” • Formation involves a transfer of electrons • Usually made up of a metal and a non-metal • Are good conductors when they can be melted or dissolved • Typically have extremely high melting points

  4. Formation of an Ionic Bond e- jumps from Na to Cl Electron acceptor (Cl) meets electron donor (Na) Ions attract to form a neutral pair

  5. Structure • Smallest building blocks are ions, NOT MOLECULES • Large numbers of ions can attract to form clusters and eventually crystals Ion pair Ion cluster Crystal lattice

  6. Ions • Cations – positively charged ions • Na+ Ca2+ Al3+ • Anions – negatively charged ions • Cl- O2- • Polyatomic ions – ions made up of more than one type of atom • NO3- SO4-2 PO4-3

  7. Oxidation Number • The number of e- gained, lost or shared ub compound formations • Alkali metals +1 • Alkaline earth metals +2 • Oxygen group -2 • Halogens -1

  8. Write the formulas – ALWAYS put cation first • K+ and N3- • K3N • Ca2+ and N3- • Ca3N2 • Ba2+ and NO3- • Ba(NO3)2 • Criss-cross rule

  9. Naming Binary Ionic Compounds • Binary – made of 2 ions • Write cation first • Change anion ending to –ide • Na+ and Cl- • Sodium chloride • H+ and F- • Hydrogen fluoride • CaBr2 • Calcium bromide

  10. Naming Polyatomic Ionic Compounds • Name the cation • Polyatomic ion name is unchanged • NaNO3 • Sodium nitrate • Zinc carbonate • ZnCO3

  11. Molecular Compounds • Also called covalent compounds • A molecule is a neutral group of atoms that are held together by covalent bonds • The valence e- are shared by the atoms • Covalent bonding usually occurs between 2 non-metals • H2O, CO2, O2, NO

  12. Naming Molecular Compounds • Use prefixes 1 mono- 2 di- 3 tri- 4 tetra- 5 penta- 6 hexa- 7 hepta- 8 octa- 9 nona- 10 deca-

  13. Examples Tetraphosphorousdecoxide • P4O10 • N2O3 • As2O5 • OF2 Dinitrogen trioxide Diarsenicpentoxide Oxygen difluoride

  14. Diatomic Molecules • H2 • O2 • N2 • Cl2 • Br2 • I2 • F2 • 7 diatomic molecules • No noble gases • Halogens and N, O, H • They are all gases (not noble gases) except for Br and I • “Honclbrif”

  15. Try these. . . Sulfuric Acid • H2SO4 • HF • H3PO4 • H2SO3 • H2CO3 • HNO3 Hydrofluoric Acid Phosphoric Acid Sulfurous Acid Carbonic Acid Nitric Acid

  16. More Practice. . . CaBr2 • Calcium bromide • Chromium (III) acetate • Barium sulfate • Copper (I) sulfide • Sulfur hexafluoride Cr(C2H3O2)3 BaSO4 Cu2S SF6

  17. More Practice. . . Chromium (III) oxalate • Cr2(C2O4)3 • Hg(CN)2 • Cu(ClO4)2 • ZnC4H4O6 Mercury (II) cyanide Copper (II) perchlorate Zinc tartrate

  18. How to Calculate Molar Mass • The mass of a compound • In order to calculate molar mass (also called molecular weight) you add up the masses of each element in the compound • Be aware of subscript numbers that designate the amount of atoms per element • You get the masses from the periodic table • **be careful when rounding the mass

  19. Examples 58.5 g/mol • NaCl • Na = 23 g/mol • Cl = 35.5 g/mol • H2O • H = 1 g/mol (but there are 2) = 2 g/mol • O = 16 g/mol • HNO3 • H = 1 g/mol • N = 14 g/mol • O = 16 g/mol (but there are 3) = 48 g/mol • Ba(NO3)2 • Ba = 137.3 g/mol • N = 14 g/mol (but there are 2) = 28 g/mol • O = 16 g/mol (but there are 6) = 96 g/mol 18 g/mol 63 g/mol 261.3 g/mol

  20. Electron Sea Model • All metal atoms in a metallic solid contribute their valence e- to form a “sea” of e- • These e- move easily and freely because they are not tied to a specific atom • Delocalized electrons • Metallic cation is formed All empty space is evenly distributed v.e-

  21. Metallic Bonds • The attraction of a metallic cation for delocalized electrons • This accounts for a lot of theproperties of metals • Range of melting points • Malleability • Ductile • Durable • Hard to remove metallic cation because of the strong e- attraction • Mobile e- • Explains why they are good conductors

  22. Electronegativity and Bond Type • Find the difference in electronegativities of the two elements Non-polar Polar 0.5 1.7 • Pure • Covalent • share e- evenly • 2 non metals and/or metalloids • Polar • Covalent • Share e- but not evenly • One element holds e- more • Ionic • Metal and non-metal

  23. To create a Lewis Dot Diagram • Count total valence electrons available • Place electrons around atoms • Ensure each atom has an octet (8) • Or a pair for H (2)

  24. VSEPR Rules • Draw the Lewis Structure for the molecule • Count the total number of . . . • Bonded regions around the central atom • DOUBLE and TRIPLE bonds count as ONE REGION • Unshared e- pair • Count as ONE REGION

  25. Molecular Lewis Dot electron pairs around central atom Structure structuretotalsharedunshared H CH4 H-C-H 4 4 0 “tetrahedral” H NH3 H-N-H 4 3 1 “trigonal H pyramidal” H2O H-O-H 4 2 2 “bent”

  26. Total no. of electron pairs No. of shared pairs No. of unshared pairs Molecular shape Molecule

  27. Polarity • A molecule is polar if • There is a polar bond • It is ASSYMETRICAL (not symmetric) (+) H (-) O (+) H H C H H (+) (+) (+) H Polar (+) Non-Polar

  28. Typically. . . • Symmetric (non-polar) • Linear • Tetrahedral • Trigonal planar • If all elements around the center atom are the same • Asymmetric (polar) • Bent • Trigonal pyramidal

  29. Intermolecular Forces • Van der Waals forces (London Dispersion forces) • Weak forces between non-polar molecules • These forces determine volatility • Doesn’t take much nrg to break apart (liquid gas) • Most likely to be a gas • Like playing red rover and only holding pinkies together

  30. Intermolecular Forces • Dipole-Dipole • Attraction between polar molecules • Most likely to be a liquid • Play red rover and hold hands

  31. Intermolecular Forces • Hydrogen Bonding (H-Bonds) • Between hydrogen (H) and a highly electronegative element • F, O, N • Extreme case of dipole-dipole • Strongest of the intermolecular forces • Play red rover and link elbows • Needs A LOT of nrg to break bonds

  32. 1 mole of . . . • Carbon has a mass of 12 g • Oxygen has a mass of 16 g • H2O molecules has a mass of 18 g • How do these #’s relate to the atom or compound? • Atomic mass

  33. Avogadro’s Number • Amedeo Avogadro (1776-1856) • 1 mole = 6.0221415 x 1023 • Particles • Molecules • Atoms • Ions • Formula units • Etc, etc

  34. Percent Composition • Determine the mass percentage of each element in the compound.

  35. Empirical Formula • Gives the lowest whole # ratio of elements in a compound. • The empirical formula for C6H12O6 is • The empirical formula for C2H6 is • * most basic ratio of elements in the compound CH2O CH3

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