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Regents Review

Regents Review

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Regents Review

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  1. Regents Review Unit 5 – Bonding Unit 6 – Naming and Moles Unit 7 – Reactions and Stoichiometry Unit 8 – Heat and Phases

  2. Unit 5Bonding Chapters 7-8

  3. Octet Rule • Atoms tend to lose or gain electrons to achieve a full valence shell (8) • Exceptions H (0,2) and He (2)

  4. Electron Dot Structures • Diagrams that show valence electrons, usually as dots • AKA Lewis Electron Dot Diagrams • Rules • Start on any side • First two get paired together • Next three are separated • Fill in as needed O

  5. Ions • Atoms that have gained or lost electrons, and now have a charge • Must show charge Na+ F- O-2

  6. Compounds • Two Main Types of Compounds • Ionic • Molecular (Covalent) • Based on type of bonding involved

  7. Bonding • Bond • Shared or exchanged electrons that hold two atoms together • Three Main Types • Ionic • Covalent • Metallic

  8. Bonding • Ionic Bond • Electrons are transferred from one atom to another (one gives, one takes) • Metal and nonmetal, NaCl • Large electronegativity difference • Polyatomic ion, Mg(NO3)2 • More than 2 elements

  9. Ionic Compounds • Ionic compounds are electrically neutral, even though they are composed of charged ions • Total positive charge equals total negative charge

  10. Determining Formulas • Must be electrically neutral • Total positive charge must equal total negative charge • Use oxidation numbers from Periodic Table • Group 1  +1 Group 2  +2 • Group 13  +3 Group 15  -3 • Group 16  -2 Group 17  -1

  11. Determining Formulas • Determine number of each ion to balance out charge • Use as subscript for element symbol • Ex: CaCl2, Na3PO4, Mg(NO3)2 • Write Positive Ion First • Formula must be smallest whole-number ratio

  12. Example • Calcium and Fluorine Ca+2 F- F- Ca1F2 CaF2

  13. Polyatomic Ions • Group of atoms that collectively have gained or lost electrons (Table E) • Sodium and Nitrate Na+ (NO3)- Na1(NO3)1 NaNO3

  14. Short-cut (criss-cross method) • Magnesium and Carbonate Mg+2 CO3-2 Mg2(CO3)2 Must Simplify MgCO3

  15. Dot Structures • Shows valence electrons • Must show charge for Ions • NaCl Na+ Cl-

  16. Covalent Bonds • Electrons are shared between two atoms • Each atom will try to achieve a full valence shell • 2 nonmetals • Two types of covalent bonds • Non-Polar Covalent – Shared equally • Polar Covalent – Shared unequally

  17. Covalent Bonding Single Bond Triple Bond • H2 • O2 • N2 Double Bond O O H H N N

  18. Covalent Bonding • H2O • CO2 Single Bond O H Single Bond H O O C

  19. More Examples • HCl • NH3 • CH4 H Cl H H C H H N H H H

  20. Determining Bond Type • Bond type is based on electronegativity difference (ΔEN) between two bonding atoms • Nonpolar Covalent Bond • 2 of the same nonmetals • Polar Covalent Bond • 2 different nonmetals • Ionic Bond • Metal and a nonmetal

  21. Determining Bond Polarity • The larger the difference in electronegativity, the more polar the bond. • Which is more polar? Most Polar Biggest Cl H F I H H H Br ΔEN 1.8 1.0 0.8 0.5

  22. Properties • Ionic Compounds • Most ionic compounds are hard, crystalline solids at room temperature • High melting points • Mostly soluble in water • Can conduct an electric current when melted or dissolved in water(aq).

  23. Properties • Covalent Compounds • Most molecular compounds tend to have relatively lower melting and boiling points than ionic compounds.

  24. Polyatomics • Compounds with polyatomic ions contain BOTH ionic and covalent bonds • Example: NaNO3 - N O O Na+ O

  25. Allotropes • Two or more different molecular forms of the same element in the same physical state • Different properties because they have different molecular structures • O2 vs O3 • Diamond, Graphite, Fullerenes (pictured on next slide)

  26. Allotropes

  27. Metallic Bonding • Bonding within metallic samples is due to highly mobile valence electrons • Free flowing valence electrons • “Sea of Electrons”

  28. Network Solids • All atoms in a network solid are covalently bonded together • very high melting and boiling points • Examples • Diamonds ( C ) • Graphite ( C ) • Silicon Dioxide (SiO2) • Silicon Carbide (SiC)

  29. Bond Energy • When two atoms form a bond, energy is released • Example: Cl + Cl  Cl2 + energy • Energy needs to be added to break a bond • Example: Cl2 + energy  Cl + Cl

  30. Molecular Polarity • Polar Molecule • one end of a molecule is slightly negative(δ-) and the other end is slightly positive(δ+). • Asymmetrical charge distribution • Nonpolar Molecule • Can not be separated into different ends • Symmetrical charge distribution

  31. Polar Molecules δ- • H2O • HCl • NH3 N H H H O H H Cl δ- δ+ δ+ H δ- δ+ N δ+ δ+ H δ+

  32. Nonpolar Examples δ+ • CH4 • CO2 H δ- δ+ δ+ H C H δ+ H δ+ O=C=O δ- δ-

  33. Polarity • Ionic Compounds are Ionic • Nonpolar Covalent Bonds always indicate Nonpolar Molecules • Polar Covalent Bonds • Determine Symmetry

  34. Polarity • Nonpolar Molecules • CH4, CO2, H2, N2, O2, … • Polar Molecules • H2O, HCl, HBr, NH3, …

  35. “Like Dissolves Like” • Polar and Ionic substances will dissolve in other Polar Substances • Nonpolar substance will dissolve in other nonpolar substances

  36. Intermolecular Forces • Intermolecular Forces of Attraction • attraction between two molecules or ions that hold them together (not a bond) • Determines melting and boiling points of compounds • Stronger intermolecular forces, higher melting and boiling points

  37. Intermolecular Forces • Van der Waals • Dispersion • Dipole-Dipole • Molecule-Ion • Hydrogen Bonding Weakest Strongest

  38. Hydrogen Bonding • Hydrogen bonded to N, O, or F, is attracted to the N, O, or F of another molecule. • Not actual bond, just attraction Hydrogen “Bond” H H F F

  39. Unit 6Naming and Moles Chapter 9-10

  40. Naming Ions • Positive Ions, cations, simply retain their name. • Na+Sodium Ion • Mg2+Magnesium Ion • Negative Ions, anions, change ending of element to –ide • Cl- Chloride Ion • Br- Bromide Ion

  41. Polyatomic Ions • Selected polyatomic ions are on Table E in the Reference Tables. • Polyatomic ions keep their names • Except in Acids

  42. Naming Systems • Ionic System • Metals and Nonmetal, more than 2 elements • Stock System (Roman Numerals) • Use when the metal element has more than one positive oxidation number • Roman Numeral is the charge of the metal ion • Binary Covalent System (Prefixes) • 2 nonmetals (including metalloids) • Second element ends in –ide

  43. Naming Ionic Compounds • Name positive ion first, then negative ion. • NaCl  Sodium chloride • Mg(OH)2Magnesium hydroxide

  44. -1 -1 Cl Cl -1 -1 Cl Cl Stock System Example • SnCl4 • Tin(IV) Chloride +4 Sn +4 -4

  45. Roman Numerals

  46. Binary Covalent Example • N2Cl3 • DinitrogenTrichloride • CO2 • Carbon Dioxide • PCl5 • Phosphorus Pentachloride

  47. Prefixes

  48. Avogadro’s Number • 6.02 x 1023 • Number of representative particles in a mole • 1 mol He = 6.02 x 1023 atoms • 1 mol H2 = 6.02 x 1023 molecules

  49. Gram Formula Mass • Mass of the formula in g/mol • Simply add the atomic masses of each element in the formula together • H2O = 1 + 1 + 16 = 18 g/mol • Also known as gram atomic mass, gram molecular mass, molar mass

  50. Mole - Mass Conversion • Example: 96 g of Oxygen gas = ? mol