American Chemical Society Cumulative Exam (Chapters 11-16)

Chapter 11 • Phase diagram and identify melting, boiling, and sublimating processes • Identify compound, element, mixtures • Relating vapor pressure and b.p., • Identifying intermolecular forces, and relating to b.p., m.p., • Hydrogen bonding Notes: 1) States of Matter a. gas= compressible fluids b. liquids= relatively incompressible fluids c. solids= nearly incompressible and rigid d. solids & liquids have higher density than gases so their molar volumes are smaller i. solid and liquid states have similar densities ii. molecules in solids and liquids are in close contact with each other but in gases they are far apart e. phase transition= a change of substance from one state to another i. heat is gained/ lost 2) Characteristics of Matter a. vapor pressure= partial pressure of vapor over the liquid which is measured at equilibrium i. rate of condensation= rate of vaporization ii. dependent on temperature 1. as temp increases vapor pressure increases 2. volatile= liquids and solids with relatively high vapor pressures at normal temperatures b. c. d. e. f. g. iii. the VP of a liquid at particular temperature is the partial pressure of the vapor over the liquid measured at equilibrium intermolecular forces= weaker, less directional, and long range intramolecular forces= strong, directional, and short range boiling point= temp at which the VP of a liquid equals the pressure on the liquid i. as temp increases the vapor pressure of water increases until it reaches atmospheric pressure (boiling) ii. normal boiling point= boiling point at 1 atm 1. lower boiling point= higher vapor pressure freezing point= temperature at which a pure liquid changes to a crystalline solid melting point= temperature at which a solid becomes a liquid heat of phase transition i. deltaH (fusion)= heat needed to melt 1 mol of pure substance ii. deltaH (vaporization)= heat needed to boil 1 mol of pure substance 3) IMF a. dipole-dipole forces= forces through which polar molecules attract one another i. polar molecules tend to align themselves from (+) to (-) end b. london forces= weak attractive forces resulting from instantaneous dipoles due to the distortion of electron clouds around a molecule i. forces increase with molecular weight ii. all covalent molecules exhibit some london forces c. when IMF in a liquid are strong VP is low d. normal BP is lowest for liquids with the weakest IMF e. hydrogen bonding= a H atom covalently bonded to a very electronegative atom and a lone pair of electrons on a very electronegative atom i. only N, O, and F are electronegative enough ii. molecules with hydrogen bonding have extremely high boiling points compared to molecules with similar forces f. solid units are held together by IMFs and chemical bonds 4) Solids a. 4 types of solids: i. molecular (VW forces)= held together by IMFs 1. ex: dry ice ii. metallic (metallic bonding)= (+) cores of atoms held together by a surrounding sea of electrons 1. ex: iron and copper iii. ionic (ionic bonding)= cations and anions held together by electrical attraction 1. ex: cesium chloride, sodium chloride iv. covalent solid (covalent bond)= atoms held together in large networks or chains by covalent bonds b. melting point and structure of solids: i. ionic and covalent solids= MP are high because melting involves breaking chemical bonds ii. ionic solids= MP increases with strength of ionic bond which is stronger when the charge is high and the ions are small iii. metals= often have high melting points 1. middle of transition metals have the highest melting points c. crystalline solids have well-defined, ordered structures in 3 dimensions d. amorphous solid has a disordered structure e. crystal lattices have geometric arrangement of lattice points in a crystal f. unit cell= smallest box like unit for constructing a crystal by stacking the units in 3D g. 7 basic shapes possible for unit cells are used to classify crystals Chapter 12 • Concept of dissolution • Calculating solution concentration. • Dilution calculation. • Energetics Notes: 1) “likes dissolve likes” principle= two substances with similar intermolecular forces and magnitude are most likely soluble 2) solution= homogeneous mixture of ions/ molecules a. solute= the substance dissolved b. solvent= the dissolving medium 3) solubility= the amount of a substance that dissolves in a solvent a. saturated solution= one holding as much solute as is allowed at a stated temperature 4) Solvents a. polar solvents dissolve polar (or ionic) solutes i. immiscible fluids= two fluids that do not mix ii. miscible fluids= fluids that dissolve in each other b. nonpolar solvents dissolve nonpolar solutes i. miscible fluids= fluids that dissolve in each other ii. immiscible fluids= two fluids that do not mix 5) Solutions a. molecular solutions i. polar molecules interact well with polar solvents ii. dipole-dipole interactions of water are explained as electrostatic attraction iii. nonpolar solutes interact with nonpolar solvents (london forces) 1. heptane and octane are both nonpolar components of gasoline and are completely miscible liquids 2. gasoline and water are nearly immiscible b. ionic solutions i. polar solvents interact well with ionic solutes ii. electrostatic attractions of water for cations and anions c. for dissolved gases in liquids: solubility decreases as temperature increases d. for most solids dissolved in liquids: solubility increases as temperature increases e. heat of solution= heat can be evolved/ absorbed when an ionic compounds dissolves in water i. when NaOH dissolves in water it gets very warm (exothermic) ii. when ammonium nitrate dissolves in water is gets very cold (endothermic) 6) Henry’s law= solubility of a gas in a liquid is proportional to partial pressure of the gas in contact with the liquid 7) Properties of Solutions a. colligative properties of solutions depend on solute concentrations i. vapor pressure reduction ii. freezing point depression iii. boiling point elevation iv. osmosis b. ways of expressing concentration i. ratio of the amount of solute to the amount of solvent or solution 1. quantity of solute, solvent, or solution can be expressed in volumes or in molar or mass amounts c. molarity= moles of solute per liter of solution d. molality= moles of solute per kilograms of solvent e. radult’s law= vapor pressure of a solution containing a solute is proportional to the mole fraction of the solvent i. each volatile solute contributes to the VP of the solution ii. volatile compounds can be separated using fractional distillation iii. BP elevation= since VP is decreased in the presence of a nonvolatile solute, a greater temperature must be reached to achieve boiling 1. the change in BP is proportional to the molal concentration iv. FP depression is also proportional to the molal concentration of the solute Chapter 13 • Kinetics • Calculating rxn rate • Determining rate law from initial rates • Calculations for 0, 1st and 2nd order rxns • Writing rate eqn for an elementary rxn • Deriving rate law from a rxn mechanism (elementary rxn slow vs. fast, vs. equilibrium, rate-determining step, overall rxn) • Arrhenius equation (k=Ae^(-Ea/RT)) Notes: 1) reaction rate= the increase of molar concentration of a product per unit of time a. can also be expressed as the decrease in molar concentration of a reactant per unit of time 2) factors that affect reaction rate: a. concentration of reactants= reaction rate increases when concentration increases b. temperature= reactions speed up when temp is increased c. catalyst= increases reaction rate 3) rate law= an equation that relates reaction rate to concentration of reactions ( and catalyst) raised to various powers 4) reaction order for each species= the exponent of concentration in the rate law a. overall order= sum of the orders 5) determining rate law: a. if doubling the concentration doubles the rate then it is a first order b. if doubling the concentration quadruples the rate then it is a second order c. if doubling the concentration 8-folds the rate then it is a third order d. zero order means no effect on reaction rate 6) half life= time required for reactant concentration to decrease to one-half of its initial value 7) reaction kinetics= measure the rate, determine the rate law, propose a mechanism 8) reaction mechanism= the “pathway” of a reaction 9) elementary reactions= the individual steps in the overall reaction “pathway” a. rate determining steps= slowest elementary reaction b. unimolecular= one reactant c. bimolecular= two reactants d. termolecular= three reactants i. the coefficient of elementary reactions becomes the power in rate laws 10) collision theory a. homogeneous catalysts= use of a catalyst in the same phase as the reacting species b. heterogeneous catalysts= a solid catalyst in contact with a liquid or gaseous solution c. concentration of a catalyst= a substance that increases reaction rate without being consumed in the overall reaction i. catalysts speed up reaction rate by decreasing activation energy ii. catalysts provide an alternative pathway with a faster rate Chapter 14 • Equilibrium • Writing eq.-constant expression (stoichiometry) • Calculating eq. Constant (K) • Using rxn quotient (Q) to predict • Calculating eq. Conc & Solving eq. Problems (exact & approximate) • Applying “change-rxn-condition” principle Notes: 1) chemical equilibrium= state in which the concentration of all reactant and all products remains constant a. forward rate= reverse rate 2) Concentration a. large kc means a large concentration of products b. small kc means a large concentration of unreacted reactants 3) reaction quotient Qc a. if Qc > Kc the reaction will shift left towards the reactants b. if Qc < Kc the reaction will shift right towards the products c. if Qc = Kc the reaction is at equilibrium 4) change reaction condition (Le Chatelier’s Principle) = when equilibrium is disturbed by change in T, P, or C it will shift in a way to counteract this change a. effect of pressure change: i. reducing volume favors the side with the least amount of moles and vice versa ii. if the reactants and products have the same number of moles then there is no effect b. effect of concentration change: i. remove products and reaction shifts to the right ii. add reactant and the reaction shifts to the right c. effect of temperature change: i. for exothermic reactions Kc is smaller at higher temperatures ii. for endothermic reactions, Kc is larger at higher temperatures 5) Changes in equation effect on Kc a. if you reverse a reaction, invert the value of Kc b. if you multiply an equation by 2,3,.., raise Kc to that same power c. if you divide an equation by 2,3,..., take the root of Kc d. when you add individual equations together, take the product of equilibrium constants to obtain overall Kc Chapter 15-16 • Acid-Base • Electrochemistry and Redox • Ionic eqn. for neutralization rxn • Calculating pH from H+ concentration, & vice versa • Calculating concentration and pH in weak acid or base solution using Ka or Kb • Calculating concentrations of species in a salt solution, and common-ion effect • Calculating pH of a buffer (H-H eqn) • Calculating pH at equivalence point in titration of a weak acid with strong base (titration curve) Notes: 1) acid-base concept a. How to identify i. acids produce hydronium ion in water ii. bases produce hydroxide ions in water b. strong acids completely ionize to give H+ and an anion i. ex: perchloric acid and hydrochloric acid c. strong bases ionize completely to give OH- an a cation i. ex: sodium hydroxide d. weak acids are not completely ionized and exist in reversible reactions i. ex: acetic acid and hydrofluoric acid e. weak bases are not completely ionized and exist in reversible reactions i. ex: ammonium hydroxide 2) Bronsted -Lowry definition of Acid and Base= donation/ acceptance of a proton a. acids donate protons b. bases accept protons 3) Lewis definition a. acid= electron-pair acceptor b. base= electron-pair donor 4) acid strength: a. bond polarity increases as acid strength increases b. bond strength decreases as acid strength increases c. more polarized= more easily removed proton= greater acid strength d. larger atom X= weaker the bond= greater the acid strength 5) polyprotic acids= acid strength of a polyprotic acid and its anions decreases with increasing negative charge a. common-ion effect 6) neutralization reactions a. titration (volumetric analysis)= a procedure to determine the amount of acid/ base quantitatively by adding acid/ base into base/acid until the acid-base reaction is complete i. equivalence point= the point when a stoichiometric amount of base has been added b. buffer= a solution capable of resisting changes in pH when limited amounts of acid or base are added to it i. if a strong base is added, it is neutralized by the weak acid in the buffer