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South Pasadena AP Chemistry [Keep for Reference] 1 Matter and Measurement BLUFFER’S 1. Matter Normally exists in 3 physical states: LiquidFixed volume, Fluid; Takes on the shape of lower part of container; well-defined surface Solid Rigid Shape; very little volume change as temperature and pressure change Gas Volume expands to fill the container; volume varies according to temperature and pressure Kinetic Molecular Theory The idea that matter consists of molecules or atoms that are in constant, random motion. Kinetic Energy = Energy of motion; higher temperature = more motion Macroscopic – seen with the eyes. Microscopic – seen with a microscope Particulate or Submicroscopic – Structures at the atomic level (what we think about) Mixtures Heterogeneous Mixture – A mixture where the properties of the mixture vary throughout. (Like an Oatmeal cookie, the different components are visible) Homogeneous Mixture – Also called a solution, where the components mix at a molecular level; different properties of the mixtures are unnoticeable. Purification – The separation of a mixture into its components. (techniques: distillation, filtration, & chromatography) 2. Elements GUIDE 3. Physical Properties Properties of a lone sample (ex. mass, volume, boiling temp, melting temp, conductivity, etc.) Density is the physical property that relates the mass of an object to its volume Mass Density = /Volume Extensive Property – Properties, like mass and volume, that depend on the amount of substance Intensive Properties – Properties like color and density; independent of the amount of substance Temperature -- how hot a substance is; physical properties (like density) vary with temp Celsius 0C for freezing point of water and 100C for melting point of water. Kelvin – same scale as Celsius; 0C = -273 K; 0 K = no motion; Celsius o + 273 = Kelvin 4. Chemical Properties How substance interacts with other substances. Ex. forms gas with acid; burns in air, etc. 5. Physical and Chemical Change Physical Change – where the identity of all the substances remain unchanged (melting, boiling, grinding, pounding into sheets, etc.) Chemical Change (Reaction) – atoms rearrange to convert one substance into another Chemical Equation – A representation of the chemical reaction taking place Example: P4 + 6Cl2 4PCl3 6. Measurements/Calculations Accuracy – how close to a “true value”; quantified by percent error. A substance that cannot be decomposed further by chemical means Precision – how close measurements are to each other. Measured by significant figures or ± notation. [I assume you know metric system.] Names given by symbols: Example: Helium = He, Gold = Au, Aluminum = Al Dimensional Analysis – use of a conversion factor to change units (ex: metric conversions, mass & volume, time units, etc.) South Pasadena • AP Chemistry Name ________________________________ Period ___ Date ___/___/___ 2 • Atoms and Elements STUDY LIST The Development of the Atomic Theory: Molar Mass Calculations: I can: State the four “signpost scientists”, their experiments, what they added to the atomic theory, and the name of their model. Define the three theories that Dalton explained in terms of atoms: o Law of Conservation of Matter o Law of Definite/Constant Proportions o Law of Multiple Proportions Give examples and solve calculation problems related to each of the three theories. Sketch a cathode ray tube as demonstrated in class and state how J.J. Thomson’s experiments led to the idea that atoms have positive and negative parts, the negative parts are all the same, and the negative parts (called electrons) have a certain charge/mass ratio. Define cathode rays. State the factors that determine how much a moving charged particle will be deflected by an electric or magnetic field. Explain Millikan’s oil drop experiment & how it added to the atomic theory. Sketch the set-up used by Ernest Rutherford (the gold-foil experiment), show what he observed, and explain how these observations led to the idea that most of the mass of the atom is concentrated into a tiny, amazingly massive, positively-charged nucleus. Calculate the isotopic mass of an atom given the resting mass of protons and neutrons. Explain that a mole of any element is actually made up of various isotopes in a constant percentage abundance. Calculate the average atomic mass of an element using the percent abundance and mass of each isotope. Calculate the percent abundance of isotopes given the average atomic mass and isotopic masses of an element. Parts of the Atom: State the three particles that make up an atom, their symbol, their charge, their mass, and their location. State the number of protons, neutrons, and electrons in any atom or ion. Explain that isotopes are two atoms with the same atomic number (number of protons) but different mass numbers (number of nucleons—protons + neturons). Represent the nucleus with isotopic notation, 220 such as: 86 Rn Recognize when two nuclei are isotopes of each other. The Families of the Periodic Table: List the common families of the periodic table and recognize to which family any element belongs. Recognize metals, non-metals, and metalloids (semi-metals) on the periodic table. State and define the terms conductivity, malleability, ductility, and sectility. State some element facts such as which elements are too radioactive to exist, which is the largest non-radioactive element, which element has the greatest density, and which element has the highest melting point. Explain how Dmitri Mendeleev put together the periodic table and why we give him credit for the table even though others were working along the same lines. List the three elements that Mendeleev predicted and where they are located on the periodic table. A Little Nuclear Chemistry: State that Henri Becquerel discovered radioactivity and Marie Curie studied it. List the three “Becquerel rays” (alpha, beta, and gamma) and state why alpha particles were the perfect tool for Ernest Rutherford to study the structure of atoms. State that the alpha particle is the same as a helium nucleus, a beta particle is a high-speed electron, and a gamma ray is a high-energy form of light. South Pasadena AP Chemistry Name __________________________________ Period ___ Date ___/___/___ 3 Chemical Formulas STUDY I can: Formulas Look at a formula and state how many elements and atoms are in that compound. Calculate the molecular mass or molar mass of any compound. State that the mass of a molecule is measured in amu’s and the mass of a mole is measured in grams. Give examples of empirical formulas, molecular formulas, and structural formulas. Identify a formula as empirical, molecular, or structural. Ionic Compounds I can: List the names and formulas of 60 ions. State whether a compound is an ionic compound or a nonmetal compound. Write the formula of an ionic compound given the two ions or its name. Know when to use parentheses. Name an ionic compound given the formula. Determine the charge on an ion from information in an ionic formula. Nonmetal Compounds aka Molecular Compound Write the formula of a binary nonmetal compound (molecular compound) given its name. Name a binary nonmetal compound (molecular compound) given its formula. LIST Percent Composition Calculate the percent composition (by mass) for any compound. Calculate the empirical formula from percent composition data. Determine the molecular formula of a compound given its empirical formula and molar mass. Hydrates Give examples of hydrates and anhydrous compounds. Calculate the formula of a hydrate from dehydration data. The Mole State the significance of the mole. State the three mole facts for any substance (molar volume, molar mass, Avogadro’s number) 1 mole = 22.4 Liters @ STP (gases only) 1 mole = 6.02 x 1023 particles (particles = molecules or atoms) 1 mole = gram molecular mass of chemical Use dimensional analysis to convert between moles, mass, volume, and number of particles for a chemical. Use density as a conversion factor in mole problems. Use gas density to calculate molar mass. South Pasadena AP Chemistry Name __________________________________ Period ___ Date ___/___/___ 4 Chemical Equations and Stoichiometry STUDY I can: Chemical Equations Limiting Reactant Problems Give examples of products and reactants in a chemical equation. State that Antoine Lavoisier introduced the law of conservation of matter. Combustion State that combustion is another name for burning. Write an equation for a combustion reaction given only the fuel that is burned. Correctly label substances in an equation as solid (s) , gas (g), liquid (l), or aqueous (aq) Balancing Equations Balance equations by adding coefficients. Recognize when an equation is balanced. State that the formulas of reactants and products should not be changed in order to balance equations. Stoichiometry Problems Use the stoichiometric factor ( of the problem) to convert from moles of one substance to moles of a different substance. (i.e. In the equation: N2 + 3H2 2NH3, 3 mol H2 2 mol NH3) Convert between the quantities of mass, volume, molecules and moles using dimensional analysis (i.e. use 1 mol = 22.4 L, 1 mol = 6.02 x 1023 molecules, and 1 mol = gram molecular mass) Show the units of molar mass as grams/mol or g·mol-1. LIST Recognize that a problem with two “given values” is a limiting reactant problem. Determine the limiting reactant and excess reactant in a problem. Solve problems involving Limiting Reactants Calculate how much excess chemical is left over after a reaction. Percent Yield Problems Use stoichiometry to calculate the theoretical yield (mass of a product) in a problem. State that actual yields are usually given in a problem. Use the theoretical yield and actual yield to calculate the percent yield. Chemical Analysis Problems Calculate the mass of each element in a given compound given data such as the masses of CO2 and H2O formed in a combustion reaction. Use mass and mole information to calculate the empirical formula of an unknown substance. Use percent composition to equalize mass and mole information derived from different samples. Continuous Variation Data Use “continuous variation” laboratory data to determine the correct mole ratio of an equation. (Micro Mole Rockets Lab) South Pasadena AP Chemistry 5 Reactions in Aqueous Solution Properties of Aqueous Solutions Define solute, solvent, and solution. Give examples. Define electrolytes. Give operational and theoretical definitions of electrolytes. Know that soluble ionic compounds and strong acids are strong electrolytes. Ionic compounds of low solubility [e.g. Mg(OH)2] and weak acids/bases are weak electrolytes. Know that molecular compounds (except acids) are non-electrolytes. Know that alcohols (e.g. CH3OH )are not ionic hydroxides. Bases are usually metallic hydroxides. Know the solubility rules. State whether an ionic compound is soluble in water. Precipitation Reactions Know that ppt reactions are double replacement reactions that produce an insoluble product. Given two ionic compounds in solution, correctly determine the products. (Know your ions). Determine which products are precipitates. Use (aq) and (s) symbols correctly. Correctly write the ions in a soluble ionic compound. [e.g. CaCl2(aq) Ca2+ + 2Cl] Identify spectator ions. Write molecular, detailed ionic, and net ionic equations for a ppt reaction. Acids and Bases Give operational (cabbage juice) and theoretical (ions) definitions of acids and bases. Know that acids increase the H+ ion concentration in an aqueous solution. (Theoretical definition) Memorize the 8 strong acids. Know that acids are molecular compounds that form ions when in aqueous solution. Be able to name acids according to their anion. S T U D Y L I S T [ide hydro__ic acid; ate __ic acid; ite __ous acid; S: add “ur”; P: add “or”] Know that bases increase the OH ion concentration in an aqueous solution. (Theoretical definition) Memorize the soluble hydroxides (except NH4OH) that are the strong bases. Understand that ammonia(aq), H2O NH3 + NH4+ + OH forms a weak basic solution. Know that metal oxides form bases [CaO + H2O Ca(OH)2] while nonmetal oxides form acids [CO2 + H2O H2CO3] Know that acids react with bases to form H2O and a salt. (Neutralization) Write equations for acid-base reactions including NH3 (example on page 199) as the base. Know that strong acids and strong bases are written as ions in the ionic equations. Gas Forming Reactions Recognize the six products that turn into gases. Memorize the gases formed. Organizing Reactions in Aqueous Solution Double Replacement reactions (text calls them exchange reactions) (FredWilma/Barney-Betty reactions) also have the old fashioned name: metathesis reactions. Know the three examples of double replacement reactions and the “driving force” for each. Precipitate reactions form an insoluble product. Acid-Base reactions form water (a very weak electrolyte therefore, a very stable product). Gas-forming reactions form a gas. Know that a driving force is something that keeps the new combinations of ions from reforming the old combinations of ions. Oxidation-Reduction is a fourth type of reaction driven by the transfer of electrons. Oxidation-Reduction Reactions Fe2O3 + 3 CO 2 Fe(s) + 3 CO2 Know that an important type of reaction gets its name from atoms that combine with oxygen. During the refining of iron, carbon monoxide combines with oxygen (from the iron ore), CO CO2 and is oxidized. Large masses of iron ore (Fe2O3) are reduced to a smaller amount of iron metal. Understand that since CO helps the iron ore to be reduced, CO is called the reducing agent. Since Fe2O3 causes the C to be oxidized, iron ore is called the oxidizing agent. What ever is oxidized acts as the reducing agent. What ever is reduced acts as the oxidizing agent. Know that oxidation-reduction (redox) is driven by the transfer of electrons. Mnemonics to help: GROL (Gain=Reduce / Oxidize=Lose); LeO the lion says GeR (Losing e’s = Oxidation / Gaining e’s = Reduction); OIL RIG (Oxidation is Losing e’s / Reduction is Gaining e’s) A redox reaction can be divided into two half-reactions. The oxidation halfreaction has electrons as a product. The reduction half-reaction has electrons as a reactant. Be able to assign oxidation numbers to any atom in any substance. Learn the rules on page 207. Recognize redox reactions because oxidation numbers change. (# = oxidation / # = reduction), electrons are gained or lost, or oxygen atoms are gained or lost. Know several common oxidizing agents and reducing agents and what they turn into. Measuring Concentrations of Cmpds. in Solution Know the definition of molarity, M, as one way to communicate concentration of solute. Know that the symbol [X] means the concentration of X in moles/Liter. Be able to determine the concentration of ions in an ionic compound. For example, in 0.25 M AlCl3 [AlCl3] = 0.25 M [Al3+] = 0.25 M [Cl] = 0.75 M Use the molarity formula to calculate moles, mass, volume, or molarity of a solution. Know that Volume x Molarity = moles of solute. Dilution problems use ViMi = VfMf. Describe how to make a solution correctly. Know what a volumetric flask is. Stoichiometry of Reactions in Aqueous Solution Use molarity as another conversion factor to solve stoichiometry problems. Know that titration is a technique called quantitative chemical analysis because you are measuring. It is also called volumetric analysis (because you are measuring volumes). [Note: qualitative analysis involves no measurements such as using solubility rules to determine the identity of an unknown ionic compound.] Understand the terms indicator, equivalence point, standardization, and primary standard. [Note: you saw a titration being done in the Measurement video early in the summer. Chloride ion from the Chesapeake Bay was being titrated against silver nitrate to determine the salinity (saltiness) of the water. Yellow K2CrO4 was used as an indicator because it formed the reddish-brown ppt, Ag2CrO4 (which looked pink) when all the chloride ion was used up.] Know common indicators such as phenolphthalein for titrations with strong bases. Understand that a titration can be done with an acid-base reaction or a redox reaction. In each case, some sort of indicator must be used to tell when equivalent amounts of reactants have been mixed. South Pasadena • AP Chemistry [Keep for Reference] 6 • Energy and Chemical Reactions Driving Forces I can… state that product-favored (spontaneous) reactions tend toward maximum entropy, S, and minimum enthalpy, H. state the sign of H based on observation of warming or cooling of the surroundings. correctly apply the terms exothermic and endothermic to situations where the surroundings are warming or cooling. draw a PE curve (uphill or downhill) based on information about warming or cooling of the surroundings. Measuring Heat state the units of heat capacity, specific heat, and molar heat capacity as well as the significance of each. convert between the heat units of calories and Joules. (4.184 J = 1 calorie) use calorimetry (q=mCT) to calculate heat changes during temperature changes. calculate the heat transferred when two objects, at different temperatures, come into contact. Energy = Heat and Work state the difference between work and heat energy. state the difference between system and surroundings. recognize the system and the surroundings in a chemical or physical system. calculate the change in internal energy based on changes in heat absorbed by the system and work done by the system. state that H is a more general (and useful) measure of energy than E and that H = q when a reaction occurs at constant pressure. S T U D Y L I S T Chemical Work = Expanding Gases relate physical work (w=F·d) and chemical work (w=P·V). calculate PV work done by an expanding gas. state that no work is done in a constant volume situation such as a bomb calorimeter. Calculating H -- Hess’s Law state the definition of a state function. list examples of properties that are and are not state functions. write the equation for the heat of formation of a substance. state that the heat of formation of an element under standard conditions has a value of zero. use Hess’s Law to calculate the energy of a chemical or physical change. Calculating Heat During Phase Changes – Heats of Fusion and Vaporization use heats of vaporization or heats of fusion to calculate heat changes during phase changes. write an equation showing the heat of fusion or heat of vaporization. South Pasadena • AP Chemistry [Keep for Reference] 12 • Gases and Their Properties £ STUDY LIST Know the pressure of the atmosphere at sea level measured in atm, kPa, mmHg, torr, psi Convert one pressure unit into another Understand how to measure pressure using a U-tube manometer, open-end manometer, and a barometer Convert gauge pressure into total pressure £ £ £ £ £ £ Sketch a P vs. V graph Manipulate P V data so a straight- line graph is obtained Form a mathematical law from straight- line graph State Boyle’s Law Recognize situations of Boyle’s Law Do Boyle’s Law problems £ £ £ £ £ £ Sketch a V vs. T graph Graphically determine a value for absolute zero State Charles’s Law Explain why temperatures must be in K Do Charles’s Law problems Recognize situations of Charles’s Law £ £ £ £ £ Know the General Gas Law (P,V&T) Know Avogadro’s Law (V&n) Know the Ideal Gas Law Given the molar volume of a gas (22.414 L at STP) determine values of R, the ideal gas constant Do Ideal Gas Law problems £ £ £ £ Derive the gas density equation from the Ideal Gas Law £ Do gas density problems £ Calculate molar mass from P, V, and T data £ Do Gas Laws and Stoichiometry problems £ Know Dalton’s Law of Partial Pressures £ Do Partial Pressure problems £ Apply this to gases collected over water £ £ £ £ £ Know the principal features of the Kinetic Molecular Theory of gases Be able to explain why each of the gas laws works in terms of the Kinetic Molecular Theory Understand the significance of the Maxwell-Boltzmann distribution curve s on pages 566-567 Derive Graham’s Law of Effusion from rms or KE of two gases Do Graham’s Law problems £ Compare van der Waal’s equations for Real gases with the Ideal Gas Law £ Know the correction factors that appear in the Real Gas Law South Pasadena • AP Chemistry [Keep for Reference] 15 • Chemical Kinetics: Rates of Reaction • How to talk about Rate rate = ∆[chemical]/∆time rate of disappearance of reactant or rate of appearance of product use coefficients to change one rate to another watch your signs (∆[React.] = -∆[Prod.]) instantaneous rate is slope of [R] vs. time graph. Initial rate is often used. • How to Speed Up a Reaction [Use Collision Theory, Kinetic Molecular Theory] increase the concentration of reactants - increase molarity of solutions - increase partial pressure of gases [collision model: more collisions] more surface area between unlike phases [collision model: more collisions] increase the temperature [collision model: more & harder collisions] add a catalyst - homogeneous catalyst (used & reformed) - heterogeneous catalyst (surface catalyst) [collision model: alternate mechanism that requires lower energy collision or ensures that correct particles collide] • Two Important Diagrams PE energy profile of a reaction +50 c +35 PE a 0 d b –35 –50 e reaction coordinate ? H of the reaction relates reactant and product PE’s / exo- or endothermic/ downhill, -∆H, or uphill, +∆H activation energy (Ea) = energy barrier • activated complex (at the peak) • whether a reaction is fast or slow depends on the activation energy in the PE profile • PE profile does not change with change in temperature of the reactants? • adding a catalyst lowers the Ea A BLUFFER’S GUIDE The KE distribution of a substance threshold energy KE - temperature is the average KE -increasing temperature spreads out curve to the right, increases average KE threshold energy KE - adding a catalyst moves the threshold energy to the left. threshold energy KE How do these two picture relate to each other (turn the KE on its side... the particles use their KE to provide the needed PE to react) • Reaction mechanisms - step-by-step...two particles at a time - example overall: 4 HBr + O2 → 2 Br2 + 2 H2O mechanism: HBr + O2 → HOOBr HOOBr + HBr → 2 HOBr HOBr + HBr → Br2 + H2O HOBr + HBr → Br2 + H2O [note: HOOBr and HOBr are not in the overall reaction because they are neither reactants nor products, they are “reactive intermediates”] - overall reaction is sum of steps - slowest step is rate-determining step Rate Laws - what they mean - how to determine them - how they relate to the rate determining step - how they help you choose a mechanism General Form: Equation: A + B → C Rate = k [A]x[B]y k is the “specific rate constant” Use experimental data to determine x, y, and k. The Rate Law CANNOT be determined from the overall reaction. It MUST be determined experimentally because the rate law reflects only the “rate determining step.” Rate law can be determined from the initial rates. See Example 15.3 and Exercise 15.3 Rate Law matches the Mechanism Examples for: 2A + 3B → C (fill in from lecture) Rate Determining Step Rate Law in the mechanism Rate = k [A][B] Rate = k [A]2 2 Rate = k [A] [B] Rate = k order of rxn - first and second order reactions - what these look like graphically - how you can graphically tell the order of a reaction order straight-line plot Slope 0 [R]t vs. t -k 1 ln[R]t vs. t -k 2 1/[R]t vs. t k - how this relates to the rate law half-life - relationship to radioactivity (a first order reaction) - the equation [ A]o ln = kt [ A]t - the special case of half-life ln(2) = 0.693 = kt½ chain reactions (fill in from lecture & video) - initiation steps - propagation steps A + B → X (slow) - termination steps A + A → X (slow) examples: - H2 + Cl2 → 2 HCl - polymerization reactions (addition) - ozone depletion A + A D X (fast) B + X → Y (slow) Each step is usually bimolecular. A third order overall reaction often comes from a fast equilibrium before a slow step. This could be a mechanism that depends on a catalyst only. The concentrations would not matter. ozone layer - specifics on why CFC’s are dangerous to the ozone layer and are economically desirable here on the surface Determining Ea from calculations using the Arrhenius Equation South Pasadena • AP Chemistry [Keep for Reference] 16 • Chemical Equilibria BLUFFER’S GUIDE 1. aA +bB + . . . rR +sS + . . . r s [R] [S] ⋅ ⋅ ⋅ Kc = [A] a [B] b ⋅ ⋅ ⋅ and for gases: (PR ) r (PS ) s Kp = (PA ) a (PB ) b 2. K > 1 products favored K < 1 reactants favored 3. Excluded: solids; pure liquids; water (in aqueous solutions) because their [ ]’s do not change. 4. Convert from Kc to Kp Kp = Kc(RT)∆n where ∆n = moles of gaseous product – moles of gaseous reactant. 5. Typical question: Given Kc and the starting concentrations of reactants, find concentrations of products at equilibrium. 10. If out of equilibrium: Calculate the reaction quotient (Q) similar to the way an equilibrium constant would be found. If: Q < K forward reaction occurs to reach equilibrium Q > K reverse reaction occurs to reach equilibrium 11. Problem solving: • Set up problems using the “magic box” (or ICE box) C = “change” or ∆. Example: A B+C A B C initial 5.0 M 0M 0M ∆ equilibrium “∆” row only follows the stoichiometry of the equation. • Example: Kc for acetic acid = 1.8 x 10-5. What is the equilibrium concentration of [H+] in a 0.100 M solution of the acid? 6. Equilibrium constant for a reverse reaction = 1 K the value of the forward reaction. Learn when to make an approximation (needed for multiple choice questions!) 5% rule usually works when value of K is 103 smaller than value of known concentrations. Example: A B+C K = 3.0 x 10-6 if [A] = 5.0M initially; find [C] at equilibrium. • 7. Equilibrium constant for a doubled reaction = K2 . If greater than 5% use the quadratic equation: (not usual on the AP exam) ax2 + bx + c = 0 x= − b ± b 2 − 4ac 2a 8. When using Hess’s Law: Koverall = K1 x K 2 • 9. Le Châtelier’s Principle: effect of changes in concentration, pressure, & temperature. Equilibrium always “shifts” away from what you add. “Stress” means too much or too little: chemical, heat, or room. Based on a handout by William Bond, Snohomish HS Another easy to solve situation is the perfect squares situation. Example: H2 + I2 2HI K = 3.5 x 102 Calculate [HI] when [H2 ] = [I2 ] = 0.10 M South Pasadena • AP Chemistry [Keep for Reference] 17 • Acid-Base Equilibria BLUFFER’S GUIDE 1. H2 O H+ + OH− Kw = [H+][OH−] = 10−14 pH = -log[H+ ] pH+pOH = 14 [H+] =10−pH Convert between pH, pOH, [H +], & [OH−] 2. Acid Ionization Constant (K a): HA + H2O H3O+ + AKa = [A- ][H 3O+]/[HA] H3O+ + F- Example: HF + H2O Ka = [F- ][H3O+]/[HF] 3. Typical question: Given Ka and the starting concentrations of acid, find concentrations (or pH) of [H+] at equilibrium. Example: Ka for acetic acid = 1.8 x 10-5. Find the pH of 0.100M acetic acid. 4. Polyprotic Acids: H3PO4, H2SO4, H2C 2O4, etc. The 1st dissociation is strong for H2 SO4 . When using Hess’s Law with a polyprotic acid: Koverall = Ka1 x Ka2 Calculating pH, use Ka1 5. Bronsted-Lowry Definitions. Acids = H+ donors; Bases = H+ acceptors Conjugate acid-base pairs. 6. Base Ionization Constant (K b): B + H2O BH+ + OHKb = [BH+ ][OH- ]/[B] Example: F- + H2O HF + OHKb = [HF][OH- ]/[F- ] 7. Salt solns can have pH’s ≠ 7 (hydrolysis) ions from weak acids → basic solutions C2 H3O2 − + H2O HC2 H3 O2 + OH− ions from weak bases → acidic solutions NH4 + + H2 O NH4 OH + H+ 8. Ka x Kb = Kw = 10-14 only applies for conjugate acids & bases! Example: Ka HC2 H3 O2 = 1.8 x 10-5 Kb C2 H3O2- = 10-14 / 1.8 x 10-5 9. Percent ionization = [H+]eqilibrium /[HA]initial x 100 10. Acid Strength-know the 6 strong acids: HCl, HBr, HI, HNO3, HClO 4, and H2SO4 (removal of the first H+ only) (a) binary acids - acid strength increases with increasing size and electronegativity of the “other element”. ( NOTE: Size predominates over electronegativity in determining acid strength.) Examples: H2 Te > H2O & HF > NH3 (b) Oxoacids - Acid strength increases with increasing: (1) electronegativity (2) number of bonded oxygen atoms (3) oxidation state of the “central atom”. Example: HClO 4 or [O 3Cl(OH)] is very acidic NaOH is very basic Acid strength also increases with decreasing radii of the “central atom”. Example: HOCl (bond between Cl and OH is covalent--making HOCl acidic) HOI (bond between I and OH is ionic-making HOI basic) 11. Lewis Acids and Bases: (This applies to coordinate covalent bonds.) Lewis Acid--electron pair acceptor Lewis Base--electron pair donor “Have Pair…Will Share” – Lewis Base In complex ion formation, metal ions are Lewis acids, and ligands are Lewis bases. Example: Cu2+ + 4NH3 Cu(NH3) 42+ Cu2+ acts as an acid; NH3 acts as a base. 12. Strong Bases: amide ion, NH2 − hydride ion, H−, methoxide ion, CH3 O− Based on a handout by William Bond, Snohomish HS South Pasadena AP Chemistry [Keep for Reference] 18 Acid-Base Reactions BLUFFER’S 1. Buffers: A solution that resists a change in pH when small amounts of acid or base are added. GUIDE Similarly, for bases with conjugate acids: B + H2O HB+ + OH- [OH ][HB ] [B] [B] [OH-] = Kb [HB ] [B] pOH = pKb - log [HB ] Kb = Buffers are a mixture of a weak acid & its conjugate base or a weak base & its conjugate acid. Examples: HC2H3O2 & C2H3O2or NH3 & NH4+ Your blood is a buffer. The equilibrium is the acid equilibrium: H3O+ + Ain which both the acid and its conjugate base are available to counteract the stress of adding acid or base (Le Châtelier’s). The equilibrium shifts, but is almost completely counteracted by the proton donor or acceptor. HA + H2O Similarly, for bases with conjugate acids: B + H2O HB+ + OH- The best buffer contains large, equal amounts of the proton donor and the proton acceptor. [HA] = [A-] (Note: they cancel out of Ka) Buffers can also be formed by changing a weak acid into its conjugate base by neutralizing some of the acid.. HA + OH- H2O + AThe same can be done with a weak base: B + H+ HB+ So, a weak acid and some strong base can form a buffer. A weak base and some strong acid can also form a buffer. 2. Titration: A carefully measured neutralization. Acid + Base H2O + Salt Since volumes are measured, this is a “volumetric analysis.” Consider: [H ][A - ] [HA] + The best buffer, Ka = [H ]; pH = pKa. The pH of a buffer can be adjusted by changing the ratio of acid and base. HA H+ + A- Ka = [HA] [H+] = Ka [A - ] [HA] pH = pKa - log [A - ] ============================== HA + OH- H2O + AThe salt is the conjugate base of the weak acid or the conjugate acid of the weak base. A “halftitration” (neutralizing half the weak acid or weak base) forms a buffer. In that case the pH = pKa of the acid. (pOH = pKb) 3. Equivalence Point: The point in a titration when stoichiometric amounts of acid and base have reacted. Note that the salt solution that is formed may have a pH >, <, or = 7. (Remember hydrolysis.) Indicators for a titration are selected based on the pH at the equivalence point. acid STRONG STRONG weak weak base STRONG weak STRONG weak pH at eq. pt. pH = 7 pH < 7 pH > 7 it depends 4. Titration Curves: This graph shows how the pH changes as a titration occurs. (A) Strong acid/Strong Base HCl + NaOH H2O + NaCl NOTE: Graph should have “pH” as the vertical axis and “added base” as the horizontal axis. The graph should be in an “S” shape. The middle of the “S” is the equivalence point (pH = 7). The top part of the “S” levels off at the pH of the base solution. (B) Weak acid/Strong Base HA + OHA- + H2O NOTE: The middle of the lower part of the “S” indicates the point of maximum buffering where [HA]/[A-] = 1. The middle of the “S” is the equivalence point (above pH = 7) and [HA] = 0. The top part of the “S” levels off at the pH of the base solution. (C) Weak base/Strong acid B + H3O+ BH+ + H2O NOTE: Graph should have “pH” as the vertical axis and “added acid” as the horizontal axis. The graph should be in a “backwards S” shape. The middle of the upper part of the “backwards S” indicates the point of maximum buffering where [B]/[HB+] = 1. The middle of the “backwards S” is the equivalence point (below pH = 7) and [B] = 0. The bottom part of the “backwards S” levels off at the pH of the acid solution. (D) Weak diprotic acid/Strong base H2A + OHHA- + H2O HA- + OHA2- + H2O NOTE: Graph should have “pH” as the vertical axis and “added base” as the horizontal axis. The graph should be in a “double S” shape. The middle of the lower part of the “first S” indicates the point of maximum buffering of the first buffering zone where [H2A]/[HA-] = 1. The middle of the “first S” is the first equivalence point where [H2A] = 0. The top part of the “first S” (i.e. the lower part of the “second S”) indicates the point of maximum buffering of the second buffering zone where [HA-]/[A2-] = 1. The middle of the “second S” is the second equivalence point where [HA-] = 0. The top part of the “second S” levels off at the pH of the base solution. From a handout by William Bond, Snohomish HS South Pasadena AP Chemistry [Keep for Reference] 18 Acid-Base Reactions STUDY LIST I can: describe how a pH buffer behaves when small amounts of acid or base are added. explain why a buffer works (buffering capacity) based on the presence of the weak acid (H+ donor) and conjugate base (H+ acceptor). I can show mathematically that diluting the buffer does not change the pH of the buffer; but it reduces its buffering capacity. calculate the pH of the best buffer you can make from a given acid and its conjugate base given Ka’s of weak acids (or Kb’s of weak bases) choose the acid / conjugate base needed to get a buffer of specified pH. (Given Ka’s of acids.) choose pairs of substances that will make a buffer: --weak acid & its conjugate base or --weak base & some strong acid --weak acid & some strong base calculate the pH of a buffer using the ICE box or the Henderson-Hasselbach equation. solve titration equivalence point problems using VH+ MH+ = VOH- MOHexplain that at the endpoint of a weak acid titration the solution only contains the conjugate base of the acid. I can calculate the concentration of the conjugate base and the pH at the endpoint of a titration. explain that weak acids and strong acids require the same amount of base to be neutralized because the weak acids will dissociate during neutralization. determine the equivalence point (end point) of the titration by looking at a titration curve. determine the pKa of the weak acid being titrated by looking at a titration curve. do the eight calculations that will allow me to sketch the pH curve for a weak acid or weak base. --weak base & its conjugate acid pH of the weak acid solution initially amount of based needed for titration concentration of conjugate base at endpoint pH of the solution at the endpoint pH halfway to the equivalence point (e.p.) pH a little before halfway to the e.p. pH a little after halfway to the e.p. pH after all of the acid has been neutralized translate all of my knowledge and skills from a weak acid titration to a weak base titration. South Pasadena • AP Chemistry [Keep for Reference] 19 • Precipitation Reactions 1. Solubility Rules Review/memorize these rules. They can be split into four groups: ALWAYS SOLUBLE: alkali metal ions (Na+, K+, Li+, Rb+, Cs+ ), NH4 +, NO3-, C2 H3 O2 -, ClO 3-, ClO 4 USUALLY SOLUBLE: chlorides, bromides, iodides (Cl- , Br-, I-) except “AP/H” (Ag+, Pb2+, Hg2 2+) sulfates (SO4 2-) except “CBS/PBS” (Ca2+, Ba2+, Sr2+, Pb2+) fluorides (F-) except “CBS/PM” (Ca2+, Ba2+, Sr2+, Pb2+, Mg2+) USUALLY INSOLUBLE: oxides/hydroxides (O 2-, OH-) except “CBS” ((Ca2+, Ba2+, Sr2+) NEVER SOLUBLE: CO32-, PO4 3-, S2-, SO 32-, CrO 42-, C2 O4 2except alkali metals & NH4 + 2. Solubility Product (Ksp) This type of equilibrium involves solids of low solubility. A saturated solution is a solution at equilibrium. The constant has no denominator. Example: Co(OH)2 (s) Co2+ + 2OH2+ 2 Ksp = [Co ][OH-] = 2.5 x 10-16 What is the pH of a saturated solution? Let x = the amount (moles) of solid that will just saturate 1 L of solution. Co(OH)2 (s) x -x 0 Co2+ + 2OH0 0 +x +2x x 2x (x) (2x) 2 = 4x3 = 2.5 x 10-16 x = 3.97 x 10-6 [OH-] = 2x = 7.94 x 10-6 pOH = 5.1 pH = 14- pOH = 8.9 BLUFFER’S GUIDE 3. Solubility vs. Ksp “Molar solubility” is the concentration of the saturated solution in moles/Liter. (Solubility is sometimes reported in g/100 mL of water.) As in the example, for a 1:2 compound, Ksp = 4x3 (where x = solubility) 1:1 Ksp = x2 1:2 Ksp = 4x3 1:3 Ksp = 27x4 2:3 Ksp = 108x5 4. Will a Precipitate Form? Ion Product (Q sp ) = “reaction quotient”. Qsp < Ksp more solid will dissolve Qsp = Ksp solution is saturated Qsp > Ksp ppt will form until Qsp = Ksp Note: Be sure to calculate concentration of DILUTED ions. Example: 50. mL of 2.0 x 10-4 M Co(NO3 )2 is mixed with 200 mL of 1.0 x 10-3 M NaOH. Will a precipitate form? [Note:K sp given in other example problem.] 50 = 4.0 x 10-5 M 250 200 [OH-] = 1.0 x 10-3 M x = 8.0 x 10-4 M 250 Qsp = (4 x 10-5 ) (8 x 10-4 )2 = 2.56 x 10-11 Qsp > Ksp ; a precipitate will form! [Co2+] = 2.0 x 10-4 M x 5. Solubility can be influenced by pH. If the anion came from a weak acid, the salt will be more soluble in a solution of strong acid. Example: CaCO3 (s) Ca2+ + CO3 2- In a strong acid, H+ combines with CO32to re- form the weak acid, H2 CO3 (which may decompose into CO2 & H2 O). More CaCO3 (s) will dissolve to reach equilibrium. South Pasadena AP Chemistry [Keep for Reference] 20 Entropy and Free Energy BLUFFER’S 1. There are two driving forces for reactions. Reactions tend toward: minimum Enthalpy, H (heat energy) H , H<0, downhill maximum Entropy, S (randomness) S +, S>0, uphill 2. Recognize whether S >0 or < 0. Entropy increases, S +, S > 0: from solid to liquid to gas fewer moles (g) to more moles (g) simpler molecules to more complex molecules smaller molecules to longer molecules ionic solids with strong attractions to ionic solids with weaker attractions separate solute & solvent to solutions gas dissolved in water to escaped gas 3. Product or Reactant favored reactions depend on H, S, and absolute Temp Product-Favored… H S + + + + at higher temperatures at lower temperatures at all temperatures never (reactant-favored at all temps) GUIDE 4. Many books use the term “spontaneous” for “product-favored.” A spontaneous reaction does not necessarily mean a fast reaction. The SPEED of a reaction is Kinetics (Ch 15)… we are discussing whether a reaction CAN OCCUR which is Thermodynamics (Ch 6 and Ch 20). 5. Gibbs Free Energy, G, puts the effects of H, S, and Temperature together. G = H - TS G<0, G , product-favored reaction G>0, G +, reactant-favored reaction G=0, reaction is at equilibrium Important: Note that H is usually in kJ/mol S is usually in J/mol·K 6. Convert between K, G, and E using equations given on the AP Exam. South Pasadena • AP Chemistry [Keep for Reference] 21 • Electrochemistry BLUFFER’S GUIDE 1. Electrochemistry is all oxidationreduction chemistry. Leo Ger OIL RIG Oxidation: loss of e− ; ox # increases Reduction: gain of e−; ox # decreases example: Fe2+ + 2e− → Fe(s) (reduction) 2. In a reaction, the oxidizing agent gets reduced; the reducing agent gets oxidized. 3. Balancing redox reactions: oxidation number method § assign ox #’s to every atom § determine changes in ox # § balance changes § balance all atoms except H & O § balance O’s (add H2 O’s) § balance H’s (add H+’s) § adjust for basic solution if needed half-reaction method. § determine oxidation & reduction § write two separate half-reactions § balance all atoms except H & O § balance O’s (add H2 O’s) § balance H’s (add H+’s) § add e− ‘s to more positive side § balance e-‘s between half-reactions § combine half- reactions § adjust for basic solution if needed 4. Electricity can either cause a reaction (electrolysis, electrolytic cell) or can be produced by the reaction (Galvanic cell, electrochemical cell, Voltaic cell). 5. Electrolysis / Electroplating coulomb (C) = an amount of charge amp = current = charge per second 1 amp · 1 second = 1 Coulomb 1 C / amp·s Faraday constant, F: 1 mole e- = 96,500 C 6. Electrolysis calculations begin with amp·s Example: How many moles of copper metal can be plated using a 10 amp circuit for 30 s? 10amp x 30s x 1C x 1 mol e- x 1 mol Ag = 1 amp·s 96500C 1 mol e-3 = 3.1 x 10 mole Ag 7. Spontaneous redox reactions (unlike electrolysis/electroplating) can simply occur (as in the ornament lab) or can be separated so the oxidation and reduction occur in different containers (half- cells). In this way, the electrons must move through an outside wire (this is an electrochemical cell—a battery). 8. Every atom has a different “potential” to accept electrons… “reduction potential” Ag+(aq) + e¯ → Ag(s) E° = +0.80 v 2+ Cd (aq) + 2e¯ → Cd(s) E° = −0.40 v These are measured by comparing every chemical to the same “standard half-cell.” The reduction with the more positive E° value will occur as written; the other reaction will reverse (oxidation). Ex: 2Ag+ + Cd 2Ag + Cd2+ The difference in the E° values is the voltage of a cell made using these two reactions. Ex: +0.80 v – (-0.40 v) = 1.20 volts NOTE that you do not multiply the Cd voltage by 2. Comparing every cell to the same standard cell accounts for this. 9. Any change that drives the reaction forward will increase the cell’s voltage. 10. In all electrochemical cells: Oxidation occurs at the Anode Reduction occurs at the Cathode South Pasadena • AP Chemistry Name __________________________________ Period ___ Date ___/___/___ 21 • Electron Transfer Reactions General Terms I can… o Determine the oxidation number of any element. o State that oxidation number is the charge an STUDY LIST o Explain that during the electrolysis of an ionic solution, either the + ion can be reduced or water can be reduced. In the same way, either the – ion can be oxidized or water can be oxidized. atom would have if all of the shared electrons were assigned to the more electronegative atom. o Use a reduction potential chart to determine is gaining or losing electrons (LeO GeR). o State that electrical current is measured in o Identify for any element in a reaction whether it o Explain that when oxidation occurs, reduction must also occur (RedOx). o Correctly apply the terms oxidizing agent and reducing agent to a redox reaction. o State that there are two big topics in electrochemistry, (1) Electrolysis—in which electricity (moving electrons) causes chemical change, and (2) Electrochemical Cells—in which chemical changes cause a flow of electrons (electricity). Electrolysis which of two substances is more likely to be reduced or oxidized. Coulombs and 1 Coulomb = 1 amp·1 sec. o State that 1 Faraday (F) = 1 mole of electrons = 96,500 Coulombs. o Use the Faraday, amps and seconds to quantify electrolysis problems. Electrochemical Cells (Voltaic Cells & Galvanic Cells) I can… o State that oxidization always occurs at the anode and reduction always occurs at the cathode. o Draw a simple electrochemical cell: I can… o State that during electrolysis, electricity applied to a solution causes ions to migrate to the electrodes. o State that an electrode is the part of the conductor that touches the solution. o State that reduction always occurs at the cathode (red cat). o State that oxidation always occurs at the anode (an ox). o Write equations for the reactions that occur at the electrodes when water undergoes electrolysis (memorize how to derive these). (−) cathode: 2 H2 O(l) + 2 e− → H2 (g) + 2 OH− (+) anode: 2 H2 O(l) → O2 (g) + 4 H+ + 4 e− o Use the reduction potential chart to determine which chemical is the anode (smaller E°) and which chemical is the cathode (larger E°). o State that standard conditions are 25°C, solutions are 1 M, and gases are 1 atm. o Calculate the voltage of a standard cell as the difference in the two E° values. (not like Hess) o State that the anode is the (−) electrode because − the chemicals are being oxidized (losing e ’s). o State that for non-standard cells, changes that drive the reaction forward increase the voltage. (The Nernst equation allows you to calculate this voltage for a non-standard cell.)