Study Guide for Exame 2
CHAPTER 3: Stoichiometry * Stoichiometry – study of quantitative aspects of formulas and relations * The mole – SI unit for the amount of a substance. * The amount of matter that contains the same number of atoms as 12.0g of carbon -> 6.022 x 10^23 (Avogadro’s number) * Avogadro’s number – 6.022 x 10^23 * How to determine how many atoms of each element is in a compound: * (moles or grams)(6.022x10^23)(Number of atoms/1molecule) * Molar mass - Molar mass is the weight of one mole (or 6.022 x 1023 molecules) of any chemical compounds. * Mass % of an element in a compound: * ((Number of atoms of element)(atomic weight))/(Formula weight) * Empirical formula – Gives the lowest whole number ratio of atoms of each element in a compound (Grams)/(atomic weight) --- divide by lowest number on all * Molecular formula – gives actual whole number ratio of atoms of each element in each compound. (Molecular formula weight)/(Empirical formula weight) x compound * Formulas from analysis: * Structured formula – a formula that shows the atoms of a compound, their relative positions, and the bonds between them. * Isomers – compounds with the same molecular formula, but different properties and different arrangements of atoms * Writing chemical equations (symbols) : * + adding 2 or more chemicals together * -> Yields (Products) * (arrow forward and backward) reaction in both directions (denotes equilibrium) * (arrow up) gas evolved * (arrow down) solid precipitate forms * (s), (l), (g) solid, liquid, gas * (aq) aqueous solution * (Triangle) heating the solution * Limiting reagent – the reagent that is completely consumed in a reaction * Reagent that limits the amount of product possible * Theoretical yield – (Moles of limiting reagent) / (molecular weight of product) mol or g possible * Actual yield – mol or g of product attained * Percent yield – (actual yield)/(theoretical yield) x100 * Solutions – a homogenous mixture of 2 or more ingredients * Solvent – substances present in largest quantity * Solute – other substances present to a lesser extent * Concentration – the quantity of solute present in a given quantity of solvent or solution * Molarity - concentration measured by the number of moles of solute per liter of solution. (mol)/(L) * Dilution - To make thinner or less concentrated by adding a liquid such as water * Reactions in aqueous solutions – * Titration – method for determining concentration of an unknown solution using a known solution and monitoring the changes * Equivalence point – in acid-base titration, when mol H= mol OH
CHAPTER 4: Classes of Chemical Reactions * Solubility -the quantity of a particular substance that can dissolve in a particular solvent (yielding a saturated solution). * What happens when an ionic compound dissolves in water? * When an ionic compound is dissolved in water, it breaks into its constituent radical ions. say for example if you dissolve NaCl in water, then it decomposes into Na+ & Cl- * What happens when a covalent compound dissolves in water? * Covalent compounds consisting of long chains of carbon atoms do not dissolve in water. They do not have anything to attract either the Oxygen atoms or the Hydrogen atoms. * Electrolytes – solutes that exist as ions in solutions * Form ions -> electric current can pass through solution * No ions -> electric current will not pass * Precipitation reactions – a reaction where two soluble ionic compounds react to form an insoluble product * Predicting if a precipitate forms: * 1. Note the ions in reactants * 2. Consider all cation-anion combinations * 3. Check solubility rules * Solubility rules * All common compounds of Group 1A elements and NH4 are soluble * All common nitrates, acetates, and perchlorates are soluble. * Molecular equation – AB + CD -> AD + BC * Total ionic equation – When you add individual ions on both sides * Net ionic equation – when you write individual ions after canceling out spectators * Acid-base reactions: * Reactions involve water as reactant or product * Sometimes called neutralization reactions * Acids – substance that produces H ion when dissolved in water * Examples of strong acids: * HCl, HBr, HI, HNO3, H2SO4, HClO3, HClO4 * Bases – substance that produces OH ion when dissolved in water * Examples of strong bases: * LiOH, NaOH, KOH, RbOH, CsOH, Ca(OH)2, Sr(OH)2, Ba(OH)2 * Oxidation-reduction reactions - Redox reactions, or oxidation-reduction reactions, primarily involve the transfer of electrons between two chemical species * Oxidation: * It is a book-keeping system to monitor the gain and loss of electrons in a reaction * Oxidation number is the change an atom would have if electrons were completely transferred * How to assign oxidation numbers: * Pure element Oxidation number = 0 * Monatomic ion Oxidation number = charge on the atom * Molecule oxidation number = sum to 0 (of all atoms) * Polyatomic ion – sum of oxidation numbers for all atoms = charge on ion * Group 1A – ox # = +1 * Group 2A - +2 * H - +1 when combined with nonmetal (-1 combined with metal) * F – -1 * O - -2 (Except when bound to F or part of peroxide) * 7A - -1 * Be able to say what is oxidized and what is reduced in redox reaction * Balancing oxidation-reduction reactions * Assign ox# to all atoms * Identify reactants that are oxidized and reduced * Determine electrons lost and gained in oxidation and reduction * Total electrons gained = total elctrons lost * Complete the balancing by inspection * Combination reaction – two or more reactants combine, one of which is an element, combine to form a product * X+Y -> Z * Decomposition reaction – compound breaks down into 2 or more products, one of which is an element * 2KClO3 -> 2KCL + 3 O2 (thermal) * 2H2O -> 2H2 + O2 (electrochemical) * Displacement – 2 substances exchange ions or atoms * AB +CD -> AD +BC * X + YZ -> XZ + Y * Combustion – process of combining a substance with O2 often with release of heat or light * 2CO +O2 -> 2CO2 * 2C4H10 + 13O2 -> 8CO2 + 10H20
CHAPTER 5: Gases * Properties of gases * Indefinite shape and volume * Differ from solids and liquids by: * Expand to fill any container * Compressible * Will form homogeneous mixtures with any other gas * Different gases behave very similarly * Atmospheric pressure – The force that acts on a given area * Gases exert pressure on any surface they contact * Pressure units – atmospheres, mm Hg, torr, Pascals (conversions will be given) * Direct proportion – Both values increase * Indirect proportion - As one values increases, the other decreases * Boyle’s Law – Relates Pressure and volume * The volume of a fixed quantity of gas at a constant temperature is inversely proportional to temperature * Charle’s Law – Relates temperature and volume * The volume of a fixed amount of gas at a constant pressure is directly proportional to temperature * Avogadro’s Law – Relates n (moles) to volume * Based on observations made by Gay-Lussac * Gay-Lussac’s Law of Combining Volumes – at a given pressure and temperature, the volume of gases that react are in ratios of small whole numbers. * Avogadro proposed: equal volume of gas at the same pressure and temperature contain equal numbers of atoms or molecules of the gas * 1 mol of gas at STP occupy 22.41L (STP 1 atm and 273K) * Ideal gas equation – PV=nRT (R=.0821 L atm/ mol K) * Calculation of the density of a gas using the ideal gas equation – density = grams/volume * Calculation of molar mass using the ideal gas equation – Molar mass = (g)(R)(T)/(P)(V) * Combined gas law equation – (P1V1)/(T1) = (P2V2)/(T2) * Dalton’s law of partial pressures – (Ptotal)=(na+nb+nc)(RT/V) = ntotal(RT/V) * Partial pressures and mole fractions – Molar fraction = Partial pressure / total pressure