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UB CHE 101 - final review chem-2

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1 Final Exam Outline Chapter 1: - Pure or not pure?  Elements: 1 type of atom only, Ca, F2  Compounds: 2 or more different atoms in a FIXED ratio, NaCl, N2O4  Mixtures: No fixed ratio  Homogeneous - uniform  Heterogeneous – not uniform - Sig figs  Move left to right & start counting when you reach a number between 1 – 9. Zeros at the end of a number only count IF a DECIMAL point is present  Addition & Subtraction Rule = Column  Multiplication & Division Rule =Fewest  Round only when your next mathematical step is a different function than the one you just performed - Converting  Prefix definitions (Mega, milli, micro…)  Squaring and cubing conversion factors - Density = mass/volume  It is a conversion factor, and never start calculations with conversion factors2 Chapter 2: - Nuclear Symbols:  Mass # (A) = protons + neutrons  Atomic # (Z) = protons (& electrons in a NEUTRAL atom) - Isotopes: same protons, different neutrons  Use the abundance of each isotope for an element to calculate the atomic weight on the periodic table - Periodic Table design:  Families/Groups: Columns (alkali metals, Halogens…)  Periods: Rows  Types of elements:  Metals: Left of black line  Non-metals: Right of black line  Metalloids: Touching black line XAZ3 - Naming:  General Rule: what comes first stays the same, what comes second is altered  Diatomic gases: 7 exist  Molecular Compounds: nonmetals only  Use the general Rule (always end in –ide) & use di, tri, tetra prefixes  Ionic Compounds: metal & nonmetal (possibly polyatomic ions) HAS IONS!  Cations: name stays same, only use a roman numeral (RN) if capable of multiple charges (RN = charge)  Anions: altered - Monatomic = -ide - Polyatomic = -ate or –ite  Use CHARGES when forming (cris-cross when needed)  Acid Compounds: start with H  Binary Acids: general formula of hydro_____ic acid  Polyatomic Acids: depends on anion - -ate ending turns into –ic acid - -ite ending turns into –ous acid  Use CHARGES when forming (cris-cross when needed)4 Chapter 3: - Writing balanced equations  Form chemicals (IONICS & ACIDS use CHARGES)  Balance with coefficients (no fractions, reduce if possible) - Types of reactions: NOTE: use charges when forming products!  Combination: 2 natural state elements  1 compound  Decomposition: 1 compound  2 natural state elements  Combustion: C&H’s +O2(g)  CO2(g) + H2O(g)  Single Replacement: A + BC  AC + B (CHAPTER 4)  Use Activity Series to predict if rxn will occur  Free element (A) must be ABOVE the cation (B) in other reactant  Double Replacement: AB + CD  AD + CB (CHAPTER 4)  Reaction occurs if PRODUCE one of the following - Solid (aka. Precipitate, ppt) - Liquid (H2O, neutralization rxn, acid + base  water + salt) - Gas (H2CO3 or H2SO3 decompose into water & CO2 or SO2) - Weak electrolyte  PHASES: Don’t need to memorize entire solubility table, but should be familiar with the ions that are ALWAYS soluble  NH4+1  NO3-1  C2H3O2-1  Column 1 cations (Na+1, Li+1, K+1 …) - Mass of elements & compounds  Atomic weight = mass of 1 atom of an element  Formula mass(or weight) = mass of 1 molecule  Molar mass or Molecular weight = mass of 1 mole of an element or molecule (unit = g/mol)  Percent composition = 100compound of mass Totalcompoundin element of massx Unit=amu H2O: 2(1.0)+16.0=18.0amu Or 18.0 g/mol5 - Conversion factors  Ratio: chemical formula  Avogadro’s number: 6.022x1023  Molar Mass: from periodic table  Mol:Mol Ratio: Coefficients in balanced equation  Molarity = moles/L (Chapter 4)  Mol:Heat = need Hrxn (Chapter 5) NOTE: If given info about 2 different chemicals, write a balanced equation!! - Limiting Reactants  If given info about BOTH reactants, must find out which one is limiting  Convert BOTH reactants to the SAME product, compare, least amount came from limiting reactant  EXCESS = Initial  Used Amount  Calculate Used Amount by start with limiting reactant & convert to excess reactant - Percent yield = 100ltheoreticaactualx - Empirical vs. Molecular Formulas  Molar masses differ by some factor (x) Formula(x) Emperical of MassFormulaMolecular of Mass  Actual usually given Theoretical must be calculated (start w/reactant & convert to product)6 Chapter 4: - Molarity = moles/L - Dilution: 2211VMVM  (used for 1 chemical only!) - Titration: Given info about both reactants, but interested in solving for info about a reactant NOTE: If given info about 2 different chemicals, write a balanced equation!! - Electrolytes  Strong: 100% ions present  Ionic compounds, Strong Acids & Strong Bases  Weak: Partial ions present (more molecule than ion)  Weak acids & weak bases  Non: NO ions present  Molecular compounds - Total/Net ionic equations: describe how compounds actually exist in a beaker  Only break up a compound if AQUEOUS & STRONG ELECTROLYTE Chapter 5: - Thermodynamics: study of energy (E), unit = Joules; Potential & Kinetic energy  Types of energy (E):  Heat (q): Endothermic: q is positive; Exothermic: q is negative  Work (w): Done BY system is negative; Done ON system is positive - Enthalpy of reaction Hrxn = Products – Reactants  Exothermic, heat lost (product) H is negative  Endothermic, heat absorbed (reactant) H is positive  MOL: Heat conversion factor!! - Calorimetry: - Hess’s Law:  If reverse rxn, change sign of H,  If multiply (or divide) the rxn, multiply (or divide) the H - Standard Heat of Formation Hf : form 1 mole of product, from its elements in their natural state - Food Energy 2211smkgJwqE    TmCqqqsgssurroundinsystem7 Chapter 10: - Characteristics of gases  Fill container  Compressible  Homogeneous - Pressure Units: (atm, mmHg, torr, Pa) - STP: 0C & 1atm - Laws:  Boyle: P&V  Charles’s: V&T  Avogadro’s: V&n  IDEAL GAS LAW, make sure you have the correct units!  Combined gas law: Same gas at 2 different conditions - Mm & Density with ideal gas law - Partial Pressures - Kinetic Molecular Theory (KMT):  Pressure due to


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UB CHE 101 - final review chem-2

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