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  1. Summer Review Unit I: Chemical Foundations Unit II: Atoms, Molecules, Ions Unit III: Stoichiometry

  2. What you are already expected to know: • Sig Figs • Metric Conversions • Dimensional Analysis • Changing between units of temperature • Classification of matter (compound, element, homogeneous, heterogeneous) • Atomic Structure & Electron Configuration • Mole-molecules-gram calculations • Basic Bonding Theory & Assigning Oxidation #s • Naming compounds • Types of Rxns & Balancing Eqns • Basic Solution Chemistry

  3. What if you don’t know these things? • All of these concepts were part of your summer assignment packet • If you cannot complete some of the questions, then you need to come before/after school to receive additional assistance

  4. AP Chemistry Unit Layouts Semester 1 Semester 2 Unit 6: Kinetics (3 wks) Unit 7: Equilibrium (6 wks) Unit 8: Thermodynamics & Free Energy (4 wks) Unit 9: Electrochemistry (1 wks) AP REVIEW AP TEST: MONDAY, MAY 5th End of Year Project • Unit 1: Chemical Foundations (2 wks) • Unit 2: Stoichiometry, Reactions in Solutions (5 wks) • Unit 3: Atomic Structure & Periodicity (2 wks) • Unit 4: Bonding & Intermolecular Forces (5 wks) • Unit 5: Gases (2 wks)

  5. What will we be reviewing in class? UNIT 1: Chemical Foundations • Challenging dimensional analysis problems • Basic Bonding Theory • Naming Compounds • Limiting Reactants & Stoichiometry • Empirical & Molecular Formulas • Molarity Calculations • Separation Techniques

  6. The Fundamental SI Units

  7. Types of Error • Random Error (Indeterminate Error) - measurement has an equal probability of being high or low. • Systematic Error (Determinate Error) - Occurs in the same direction each time (high or low), often resulting from poor technique.

  8. Rules for Counting Significant Figures - Details • Exact numbershave an infinite number of significant figures. Can come from counting or definition. • 15 atoms • 1 inch = 2.54cm, exactly

  9. Rules for Significant Figures in Mathematical Operations • Multiplication and Division:# sig figs in the result equals the number in the least precise measurement used in the calculation. • 6.38  2.0 = • 12.76 13 (2 sig figs)

  10. Rules for Significant Figures in Mathematical Operations • Addition and Subtraction:# sig figs in the result equals the number of decimal places in the least precise measurement. • 6.8 + 11.934 = • 18.734  18.7 (3 sig figs)

  11. Universe Matter Energy Physical Change Homogeneous Heterogeneous PotentialEnergy KineticEnergy Solution Mixture Pure Substance Position Composition Chemical Change Element Compound Gravitational Electrostatic Electron Levels Nucleus Protons Neutrons Electrons

  12. SEPARATION OF MIXTURES • - mixtures can be separated into pure substances by physical means. • distillation • filtration • centrifuging • magnet • evaporation • chromatography

  13. Simple laboratory distillation apparatus.

  14. CENTRIFUGE

  15. Paper Chromatography Chromatography has two phases of matter: a stationary phase (the paper) and a mobile phase ( the liquid).

  16. Law of Conservation of Mass • Discovered by Antoine Lavoisier • Mass is neither created nor destroyed • Combustion involves oxygen, not phlogiston

  17. Other Fundamental Chemical Laws • A given compound always contains exactly the same proportion of elements by mass. • Carbon tetrachloride is always 1 atom carbon per 4 atoms chlorine. Law of Definite Proportion -- Joseph Proust

  18. Other Fundamental Chemical Laws • When two elements form a series of compounds, the ratios of the masses of the second element that combine with 1 gram of the first element can always be reduced to small whole numbers. • The ratio of the masses of oxygen in H2O and H2O2 will be a small whole number (“2”). Law of Multiple Proportions--John Dalton

  19. The Modern View of Atomic Structure • electrons • protons: found in the nucleus, they have a positive charge equal in magnitude to the electron’s negative charge. • neutrons: found in the nucleus, virtually same mass as a proton but no charge. The atom contains:

  20. The Mass and Charge of the Electron, Proton, and Neutron

  21. The Chemists’ Shorthand: Atomic Symbols 39 Mass number  K  Element Symbol 19 Atomic number 

  22. The Chemists’ Shorthand:Formulas • Chemical Formula: • Symbols = types of atoms • Subscripts = relative numbers of atoms CO2 • Structural Formula: • Individual bonds are shown by lines. O=C=O

  23. Ions • Cation: A positive ion Mg2+, NH4+ • Anion: A negative ion Cl, SO42 • Polyatomic: an ion containing a number of covalently bonded atoms acting as a single unit. • Ionic Bonding: Force of attraction between oppositely charged ions.

  24. Cations & Anions • Cations are positive ions. • Na ----> Na+ + e- • Anions are negative ions. • Cl2 + 2e- ----> 2Cl-

  25. Periodic Table • Elements classified by: • Properties -atomic number • Groups (vertical) 1A = alkali metals 2A = alkaline earth metals 7A = halogens 8A = noble gases • Periods (horizontal)

  26. The periodic table.

  27. Chemical Symbols • Symbols commonly missed. • A -- Al, Ar, As, Au, & Ag. • B -- Ba, Bi, B, Br, & Be. • C -- C, Ca, Cd, Cl, Cr, Co, Cs, & Cu. • M -- Mg, Mn, & Mo. • S -- S, Sb, Si, Sr, & Sn. • Latin -- Fe, Au, Ag, Sb, Pb, Na, K, Hg, & Cu. • German -- W

  28. Physical Properties of Metals Metals are: 1. efficient conductors of heat and electricity. 2. malleable (Can be hammered into thin sheets). 3. ductile (Can be pulled into wires). 4. lustrous (shiny). 5. tend to lose electrons and form cations. • Examples are: Na, Cu, Au, Ag, & Fe.

  29. Metalloids • substances with the properties of both metals and nonmetals. • also called semimetals • Lie along the zigzag line between metals and nonmetals • The six metalloids are: • B, Si, Ge, As, Sb, and Te.

  30. Physical Properties of Nonmetals Nonmetals are: 1. nonconductors of heat and electricity (insulators). 2. not malleable, but are brittle. 3. not ductile. 4. dull and without a luster. 5. tend to gain electrons to form anions. • Examples are: H, He, N, O, S, & P.

  31. The Chemists’ Shorthand: Atomic Symbols 39 1+ Mass number   Ion charge K  Element Symbol 19 Atomic number 

  32. Charges on Common Ions -4 -3 -2 -1 +1 +2 +3

  33. Common Names • sugar of lead • blue vitriol • quicklime • Epsom salts • milk of magnesia • gypsum • laughing gas lead(II) acetate copper(II) sulfate calcium oxide magnesium sulfate magnesium hydroxide calcium sulfate dinitrogen monoxide

  34. Naming Compounds 1. Cation first, then anion 2. Monatomic cation = name of the element • Ca2+ = calciumion 3. Monatomic anion = root + -ide • Cl = chloride • CaCl2 = calcium chloride Binary Ionic Compounds:

  35. Naming Compounds(continued) Binary Ionic Compounds (Type II): - metal forms more than one cation - use Roman numeral in name • PbCl2 • Pb2+is cation • PbCl2 = lead (II) chloride • plumbous chloride

  36. Naming Compounds(continued) - Compounds between two nonmetals -First element in the formula is named first. -Second element is named as if it were an anion. - Use prefixes - Never use mono- • P2O5 = diphosphoruspentoxide Binary compounds (Type III):

  37. NOMENCLATURE OF COMPOUNDS Binary -- 2 elements Ternary -- (3 elements) - Ionic (metal ion + polyatomic ion) Type I - Ionic (Type I metal + nonmetal) Group I, II, Al+3, Ag1+, Cd2+, & Zn2+ NaCl -- Sodium Chloride Ca3(PO4)2 -- calcium phosphate FeSO4 -- iron (II) sulfate -- ferrous sulfate Type II - Ionic (Type II metal + nonmetal) All other metals Fe2S3 -- iron (III) sulfide -- ferric sulfide Type III - covalent (2 nonmetals) CO2 -- carbon dioxide

  38. Chemical Nomenclature • Name each of the following: • CuCl • HgO • Fe2O3 • MnO2 • PbCl2 • CrCl3 copper(I) chloride cuprous chloride mercury(II) oxide mercuric oxide iron(III) oxide ferric oxide manganese(IV) oxide manganic oxide lead(II) chloride plumbous chloride chromium(III) chloride chromic chloride

  39. Chemical Nomenclature • Name each of the following: • P4O10 • N2O5 • Li2O2 • Ti(NO3)4 • SO3 • SF6 • O2F2 tetraphosphorusdecoxide dinitrogenpentoxide lithium peroxide titanium(IV) nitrate sulfur trioxide sulfur hexafluoride dioxygendifluoride

  40. Compounds: SO3 --Sulfur trioxide NO2 -- Nitrogen dioxide NO3 -- Nitrogen trioxide Polyatomic ions: SO32- -- Sulfite ion NO21- -- Nitrite ion NO31- -- Nitrate ion Common Nomenclature Mistakes

  41. A flow chart for naming acids. An acid is best considered as one or more H+ ions attached to an anion.

  42. Binary Acids • made up of two elements -- hydrogen and a nonmetal • named by using: • prefix hydro + root of nonmetal + ic + acid HCl-- hydrochloric acid H2Se -- hydroselenic acid

  43. Ternary Acids (oxyacids) • contain three elements -- hydrogen, nonmetal, and oxygen. • most oxygen per + root of nonmetal + ic + acid • less oxygen root of nonmetal + ic + acid • less oxygen root of nonmetal + ous + acid • least oxygen hypo + root of nonmetal + ous + acid

  44. Ternary Acids(continued) • HBrO4 perbromic acid • HBrO3 bromic acid • HBrO2 bromous acid • HBrOhypobromous acid • H3PO4 phosphoric acid • H3PO3 phosphorous acid • H3PO2 hypophosphorus acid

  45. Salt Nomenclature (Ionic compounds) • Binary salts (metal and nonmetal) • name of positive ion + root of nonmetal + ide NaCl -- sodium chloride K2S -- potassium sulfide

  46. Salt Nomenclature (continued) • Ternary salts ( metal and polyatomic ion) • name of positive ion + root of nonmetal + ate or ite • If the salt comes from anicacid, changeictoate. • H2CO3carbonicacid Na2CO3 sodiumcarbonate • H3PO4phosphoricacid K3PO4potassiumphosphate • If the salt comes from anousacid, changeoustoite. • H2SO3sulfurousacid Li2SO3 lithiumsulfite • HClOhypochlorousacid NaClO sodiumhypochlorite

  47. Avogadro’s number equals6.022  1023 units

  48. Calculating Moles & Number of Atoms • 1 mol Co = 58.93 g • (5.00 x 1020 atoms Co)(1mol/6.022 x 1023 atoms) = 8.30 x 10-4 mol Co (8.30 x 10-4 mol)(58.93g/1 mol) = 0.0489 g Co Moles are the doorway grams <---> moles <---> atoms