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AP Chem Review Packet

Test Format (new as of'06-'07)

The format of the AP chemistry test is pretty standard from year to year, having changed substantially for the fIrst time in at least a decade. It is broken into two parts (multiple choice and free response) with the free response section being broken up into two further parts.

• Multiple choice - 45% of final grade - 75 questions total 90 minutes no calculators allowed

• Free response 55% of fInal grade - 6 questions total - 90 minutes total calculators allowed on Part A only • Free response A 3 required questions - 55 minutes - calculators allowed ­

formula sheet provided o one equilibrium problem (20% of free response) o two other problems (each worth 20% offree response)

• Free response B-3 required questions - 40 minutes - no calculators allowed ­formula sheet provided

o One reaction question with three required reactions and a follow up question for each (10% of free response)

o Two other essay questions (each worth 15% offree response)

Test taking hints/strategies: Section I, Multiple Choice (90 minutes total)

1. Examine each question for a maximum of thirty seconds (on the average, some will take less time allowing more time for others).

a. Quickly determine the subject of the question. b. By the end ofthe thirty seconds ej:ner:

i. Mark the correct answer. ii. Mark a "Y" next to the questions that you know how to work but

need more time. iii. Mark a "N" next to the questions that you don't have any idea how

to work. L Points are deducted for incorrect answers. Don't guess

unless you can eliminate at least two choices!

c. Force yourself to move through twenty questions each ten minutes and the full seventy-five questions in forty minutes.

2. Now make a second pass concentrating on the "Y" questions only. Do not spend any time on the "N" questions. If you don't know the correct answer see if some key piece of knowledge will allow you eliminate two or three of the choices. Complete this pass in forty minutes.

3. Now make your third pass. Focus only on the ''Nil questions. Attempt to eliminate at least two choices. If you can, then make an intelligent guess. If not, leave it blank. You have only ten minutes, so make it count!

4. Before time expires, count the number that you have answered. You should answer at least sixty (60) questions.

Section II, Free Response (95 minutes total) Question 1. 1. Read all of Question 1 before doing any work. Items later in the problem may provide keys to earlier sections

2. Question 1 is always equilibrium. Determine which type (Gaseous equilibrium, acid/base, buffer, or precipitation). Look for key words and clues. a. AcidfBase: Look for the words acid or base, Ka or Kb, [H+], [OH-], or [H30+]. Any of these indicate an acidlbase problem.

b. Buffer: Look for the word buffer. Also, check for a weak acid and its conjugate base or for weak acid and its salt.

c. Precipitation: Look for Ksp or the word solubility or slightly soluble compound (solid).

d. Gas Equilibrium: Look for (g) on most of the reactants and products. 3. After determining the type ofreaction, write a reaction if one is not provided. Use the general forms given below:

Acid HA + H2O -----> H30+ or HA -----> H+ + A­Base A- + H2O -----> HA + OH-Precipitation MA(s) -----> M+(aq) 4- A-(aq)

4. Write an equilibrium constant expression. Leave out solids and liquids.

5. Solve the problem. THINK! Put in all of the given quantities in the equilibrium constant expression and solve for the unknown allowing the units to direct the problem.

Question 2 and 3: ClassifY the type ofproblem and answer them.

Part B: (40 minutes) Question 4 Reactions 1. Write the reactant in symbol form for all eight reactions showing each reactant in net ionic form as follows:

Strong Acids, bases, and soluble salts written as ions.

Weak acids, bases, and insoluble salts written as molecules.

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2. ClassifY the reactions as:

(A) Acid/Base - Look for H or OH or salts which could act as a weak acid or base.

(P) Precipitation Look for insoluble salts which could form as products according to the solubility rules.

(R) Redox If it is not acidlbase or precipitation, it is probably oxidation-reduction. Check for elements which could change oxidation states. Pay particular attention to the common oxidizing agents (N03-, Mn04-, Cr2072-, H202) and reducing agents (CI-, Br­, I- and elemental metals).

(0) Other - Anything else which doesn't fit above (usually either organic or complex­ion).

Questions 5 and 6

One of the essays will always be related to laboratory, while the topic of the other required essay will vary from year to year.

1. Be as specific as possible in your answer. Look for clues in the question as to what is really important.

2. Answer the question. State exactly what you are asked, not what you would like to answer.

3. Do not simply restate the question.

4. Remember that you will be getting partial credit. Answer any part about which you have any knowledge.

Helpful internet websites:

• ~ww.scicncegcek.netlAPchrn!i$l:r'yLAJ>U!1crs/directo~tml (Practice Problems)

• help). • (tons of useful info

• (review info and podcasts)

Topical Review:

Topic 1: Sig Figs, Conversions, Naming, and Stoichiometry (Ch. 1 3)

Goals: 1. Perform calculations using the correct number of sig figs. 2. Convert between units using dimensional analysis. 3. Name covalent compounds. 4. Name Ionic compounds. 5. Write the formulas for ionic and covalent formulas given the name. 6. Balance a chemical equation. 7. Use the mass ofreactants to determine the mass ofproducts formed. 8. IdentifY the limiting reagent when given masses of reactants. 9. Calculate the empirical formula ofa compound. 10. Use the molar mass and empirical formula to calculate the molecular formula.

Practice Problems:

1. A 2.000 gram sample containing graphite (carbon) and an inert substance was burned in oxygen and produced a mixture ofcarbon dioxide and carbon monoxide in the mole ratio 2.00:1.00. The volume of oxygen used was 747.0 milliliters at 1,092K and 12.00 atmospheres pressure. Calculate the percentage by weight ofgraphite in the original mixture.

2. Three volatile compounds X, Y, and Z each contain element Q. of element Q in each compound was determined. Some ofthe data below.

Percent by weight Molecular ofElementQ

64.8% Y 73.0% 104. Z 59.3% 64.0

(a) The vapor density of compound X at 27°C and 750. m:rh Hg was determined to be 3.53 grams per litre. Calculate the molecular weight of compound X.

(b) Determine the mass of element Q contained in 1.00 mole of each of the three compounds.

(c) Calculate the most probable value of the atomic weight of element Q. (d) Compound Z contains carbon, hydrogen, and element Q. When 1.00 gram of

compound Z is oxidized and all of the carbon and are converted to oxides, 1.37 2ralIlS of CO2 and 0.281 gram of water are Determine the most

molecular formula ofcompound Z.

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3. A sample of dolomitic limestone containing only CaC03 and MgC03 was analyzed. (a) When a 0.2800 gram sample of this limestone was decomposed by heating, 75.0

milliliters of COz at 750 mm Hg and 20°C were evolved. How many grams of CO2

were produced. (b) Write equations for the decomposition ofboth carbonates described above.

(c) It was also determined that the initial sample contained 0.0448 gram of calcium. What percent of the limestone by mass was CaC03?

(d) How many grams of the magnesium-containing product were present in the sainple in (a) after it had been heated?

Topic 2: Gases (Ch. 5) Goals:

L Use gas laws to predict and calculate the effects ofchanges in pressure, volume, and temperature on other variables ofthe gas.

2. Convert between degrees Celsius and Kelvin 3. Calculate the partial pres~ures of a mixture of gases when given their masses or

moles. 4. Determine the molar mass ofa gas using the ideal gas law. 5. Solve for variables ofa gas using the ideal gas law. 6. Explain rates ofeffusion and diffusion ofgas particles. 7. Use rates of effusion to determine the molar mass ofa gas. 8. Explain why real gases deviate from ideal behavior. 9. Use the Van der Waal's equation to calculate the pressure of a real gas.

Practice Problems: LA 6.19 gram sample of pas is placed in an evacuated 2.00 liter flask and is completely vaporized at 252°C.

(a) Calculate the pressure on the flask ifno chemical reaction were to occur. (b) Actually at 252°C the pas is partially dissociated according to the following

equation: PCls(g) ..... PCI3(g) + C12(g)

The observed pressure is found to be 1.00 atmosphere. In view of this observation, calculate the partial pressure ofPCIs and PC13 in the flask at 252°C.

2. Propane, C3Hs, is a hydrocarbon that is commonly used as fuel for cooking. (a) Write a balanced equation for the complete combustion of propane gas, which yields

COz(g) and HzO(l). (b) Calculate the volume of air at 30°C and 1.00 atmosphere that is needed to burn

completely 10.0 grams ofpropane. Assume that air is 21.0 percent O2 by volume.

(c) The heat of combustion of propane is -2,220.1 kJ/mol. Calculate the heat of formation,81JOr, of propane given that MJOr of HzO(l) = -285.3 kJ/mol and Ml~ of COz(g) = -393.5 kJ/mo!.

(d) Assuming that all of the heat evolved in burning 30.0 grams ofpropane is transferred to 8.00 kilograms of water (specific heat = 4.18 J/gK), calculate the increase in temperature ofwater.

3. A mixture of Hz(g), Oz(g), and 2 millilitres of HzO(1) is present in a 0.500 litre rigid container at 25°C. The number of moles of Hz and the number of moles of O2 are equal. The total pressure is 1,146 millimetres mercury. (The equilibrium vapor pressure of pure water at 25°C is 24 millimetres mercury.) The mixture is sparked, and Hz and Oz react until one reactant is completely consumed. (a) Identify the reactant remaining and calculate the number of moles of the reactant

remaining. (b) Calculate the total pressure in the container at the conclusion of the reaction if the

fmal temperature is 90°C. (The equilibrium vapor pressure of water at 90°C is 526 millimetres mercury.)

(c) Calculate the number ofmoles of water present .§ vapor in the container at 90°C.

Topic #3: Thermodynamics (Ch. 6 and 16)

Goals: I. Calculate the heat of a reaction using calorimetry data (q=mc.6T) 2. Use Hess's Law to calculate enthalpy ofa reaction. 3. Calculate work done by a reaction. 4. Identify a reaction as endothermic or exothermic. 5. Calculate the enthalpy ofa reaction given the enthalpies offormation ofthe

reactants and products. 6. Predict changes in entropy given the chemical equation 7. Predict if a reaction is spontaneous using entropy and gibbs free energy. 8. Describe the effect of temperature on reaction spontaneity 9. Calculate the free energy ofa reaction. 10. Describe the relationship between free energy and equilibrium. II. Explain the dependence offree energy on pressure.

Practice Problems: 1.

2 NO(g) + Oz - 2 N02(g) A rate expression for the reaction above is:

2dt =k[NOJ [02J

M-I,0 So 8G/ kcallmole cal/{mole)(K) kcallmole

NO(g) 21.60 50.34 20.72 Oz(g) o 49.00 o NOz(g) 8.09 57.47 12.39

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(a) For the reaction above, find the rate constant at 25°C if the initial rate, as defmed by the equation above, is 28 moles per liter-second when the concentration of nitric oxide is 0.20 mole per liter and the concentration of oxygen is 0.10 mole per liter.

(b) Calculate the equilibrium constant for the reaction at 25°C.

2. CO(g) + 2 Hz(g) -- CH30H(l) Mr -128.1 kJ

liH;0 (kJ mOrl)

tlG/ (kJ morl)

S" (J mOrl K l)

CO(g) -110.5 -137.3 +197.9 CH30H(l) -238.6 -166.2 +126.8

The data in the table above were determined at 25°C.

(a) Calculate tlG" for the reaction above at 25"C.

(b) Calculate ~ for the reaction above at 25°C.

(c) Calculate !.IS for the reaction above at 25"C.

(d) In the table above, there are no data for Hz. What are the values of liH;", tlG/, and of the absolute entropy, So, for H2 at 25°C?

3. C1z(g) + 3 Fz(g) -- 2 ClF3(g)

ClF3 can be prepared by the reaction represented by the equation above. For ClF3 the standard enthalpy offormation, liH;0, is -163.2 kilojouleslmole and the standard free energy offormation, tlG/, is -123.0 kilojoules/mole. (a) Calculate the value of the equilibrium constant for the reaction at 298K. (b) Calculate the standard entropy change, !!.S", for the reaction at 298K. (c) If CIF3 were produced as a liquid rather than as a gas, how would the sign and the

magnitude of!!.S for the reaction be affected? Explain. (d) At 298K the absolute entropies of Ch(g) and ClF3(g) are 222.96 joules per mole­

Kelvin and 281.50 joules per mole-Kelvin, respectively. (i) Account for the larger entropy of ClF3(g) relative to that ofC1z(g). (li) Calculate the value of the absolute entropy of Fz(g) at 298K.

4. Lead iodide is a dense, golden yellow, slightly soluble solid. At 25"C, lead iodide dissolves in water forming a system represented by the following equation.

PbIz(s) - Pbz.. + 2 I" Mf = +46.5 kilojoules

(a) How does the entropy of the system PbIz(s) + HzO(l) change as PbIz(s) dissolves in water at 25°C? Explain

(b) Ifthe temperature ofthe system were lowered from 25"C to 15°C, what would be the effect on the value ofKsp? Explain.

(c) If additional solid PbIzwere added to the system at equilibrium, what would be the effect on the concentration ofl" in the solution? Explain.

(d) At equilibrium, tlG O. What is the initial effect on the value of tlG of adding a small amount ofPb(NOJ)z to the system at equilibrium? Explain.

Topic #4: Atomic Structure and the Periodic Table (Ch. 7)

1. Identify metallic ions by the color flame they produce when burnt. 2. Explain how the Bohr model accounts for the different colors produced by

different metallic ions in flame tests. 3. Calculate the change in energy level of electrons based upon the color oflight

released. 4. Explain the location ofprotons, neutrons, and electrons in the quantum

mechanical model of the atom. 5. Write the electron configuration for any element. 6. IdentifY the quantum numbers for any given electron of an element. 7. Describe the shape and relative energy level of s, p, d, and f orbitals. 8. Use the periodic table to predict electronegativity, ionization energy, electron

affmity, atomic radius, and ion radius. 9. IdentifY and give characteristics ofthe following families on the periodic table:

alkali metals, alkaline earth metals, transition metals, metalloids, halogens, noble gases, lanthanides and actinides.

Practice Problems: 1. Suppose that a stable element with atomic number 119, symbol Q, has been

discovered. (a) Write the ground-state electron configuration for Q, showing only the valence-shell

electrons.

(b) Would Qbe a metal or a nonmetal? Explain in terms of electron configuration. (c) On the basis of periodic trends, would Q have the largest atomic radius in its group

or would it have the smallest? Explain in terms of electronic structure.

(d) What would be the most likely charge of the Q ion in stable ionic compounds? (e) Write a balanced equation that would represent the reaction of Q with water.

(f) Assume that Q reacts to form a carbonate compound.

(i) Write the formula for the compound formed between Q and the carbonate ion, C03

2-.

(ii) Predict whether or not the compound would be soluble in water. Explain your reasoning.

2. Use the details of modern atomic theory to explain each of the following experimental observations. (a) Within a family such as the alkali metals, the ionic radius increases as the atomic

number increases. (b) The radius of the chlorine atom is smaller than the radius of the chloride ion, CI".

(Radii: Cl atom:; O.99A; CI" ion 1.81 A) (c) The flIst ionization energy of aluminum is lower than the first ionization energy of

magnesium. (First ionization energies: 1zMg:; 7.6 ev; 13AI:; 6.0 ev)

)

(d) For magnesium, the difference between the second and third ionization energies is much larger than the difference between the first and second ionization energies. (Ionization energies for Mg: 1" = 7.6 ev; 2nd = 14 ev; 3'd = 80 ev)

3. Use principles of atomic structure and/or chemical bonding to answer each of the following. (a) The radius of the Ca atom is 0.197 nanometer; the radius of the Ca2

'" ion is 0.099 nanometer. Account for this difference.

(b) The lattice energy of CaO(s) is -3,460 kilojoules per mole; the lattice energy for K 20(s) is -2,240 kilojoules per mole. Account for this difference.

Ionization Energy (kJ/mol)

First 1 Second 419 13,050

ICa IK

590 11,140

(c) Explain the difference between Ca and K in regard to (i) their first ionization energies, (ii) their second ionization energies.

(d) The ftrst ionization energy of Mg is 738 kilojoules per mole and that of Al is 578 kilojoules per mole. Account for this difference.

Topic #5: Bonding (Ch. 8 and 9)

Goals: 1. IdentifY a compound as ionic or covalent. 2. Explain the difference between ionic, covalent, and metallic bonding. 3. Draw Lewis Structures given the chemical formula. 4. IdentifY a bond as covalent, polar covalent, or ionic based upon electronegativity. 5. IdentifY the shape ofa molecule based upon its Lewis structure. 6. IdentifY the electron geometry around a central atom in a molecule. 7. IdentifY the bond angles in a molecule. 8. Use the concept of resonance to describe bond length. 9. Describe how pi-bonds are formed. 10. Use bond energies to calculate enthalpy offormation. 11. Explain covalent bonding using the hybridization model

Practice Problems:

1. Suppose that a molecule has the formula AB3. Sketch and name two different shapes that this molecule may have. For each ofthe two shapes, give an example ofa known molecule that has that shape. For one of the molecules you have named, interpret the shape in the context ofa modem bonding theory.

\ ' 2. Butane, chloroethane, acetone, and I-propanol all have approximately the same molecular weights. Data on their boiling points and solubilities in water are listed in the table below.

Boiling Solubility Compound Formula Pt.(OC) in water

Butane CH3CH2CH2CH3 0 insoluble

Chloroethane CH3CH2Cl 12 insoluble 0Acetone 56 completelyII

miscibleCH3CCCH3 completely

I-Propanol CH3CH2CH20H 97 miscible

On the basis of dipole moments (molecular polarities) and/or hydrogen bonding, explain in a qualitative way the differences in the (a) boiling points ofbutane and chloroethane. (b) water solubilities ofchloroethane and acetone.

water solubilities of butane and I-propanol. boiling points ofacetone and I-propanol.

3. CF4 XeF4 C1F3

(a) Draw a Lewis electron-dot structure for each of the molecules above and identifY the shape of each.

(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each of these molecules.

4. Use simple structure and bonding models to account for each of the following. (a) The bond length between the two carbon atoms is shorter in C2~ than in C2Ho. (b) The H-N-H bond angle is 107.5°, in NH3•

(c) The bond lengths in S03 are all identical and are shorter than a sulfur-oxygen single bond.

(d) The 13' ion is linear.

Topic #6: Liquids and Solids (Ch. 10)

Goals:

1. IdentifY the intermolecular forces present in a molecule given its formula. 2. Use intermolecular forces to explain differences in melting and boiling points,

vapor pressure, and solubility. 3. Given the chemical formula ofa substance, identifY the type of solid that it forms.

4. Describe the structure ofa crystal. 5. Describe the crystalline structures of: network atomic solids, molecular solids,

and ionic solids.

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6. Explain how vapor pressure is related to changes of state. 7. Draw, explain, and interpret a phase diagram. 8. Explain how a phase diagram can be used to determine the relative densities of

the solid and liquid state of a substance.

Practice Problems: I. The state of aggregation of solids can be described as belonging to the fol1owing four types: (I) ionic (3) covalent network (2) metallic (4) molecular For each of these types of solids, indicate the kinds ofparticles that occupy the lattice points and identify forces among these particles. How could each type of solid be identified in the laboratory?

2.

.e "

J.­ /

z\.0 f! i' •

Ttmpcntllro (C)

The phase diagram for a pure substancc is shown above. Use this diagram and your knowledge about changes of phase to answer the following questions. (a) What does point V represent? What characteristics are specific to the system only at

point n. (b) What does each point on the curve between V and W represent? (c) Describe the changes that the system undergoes as the temperature slowly increases

from X to Yto Z at 1.0 atmosphere. (d) In a solid-liquid mixture ofthis substance, will the solid float or sink? Explain.

Topic #7: Solutions (Chapter 11)

Goals: 1. Calculate the strength of solutions using: molarity, molality, %mass, and

normality. 2. Calculate the heat of solution.

3. Relate the heat of solution to solubility ~d intermolecular forces. 4."ldentify the three factors affecting solubility. 5. Explain how these factors affect solubility. 6. Calculate the vapor pressure ofa solution. 7. Explain why non-ideal solutions do not behave according to Raoult's law. 8. Perform calculations using freezing point depression and boiling point elevation. 9. Calculate molar mass using freezing point depression and boiling point elevation

data. 10. Calculate osmotic pressure. II. Calculate molar mass from osmotic pressure. 12. Explain the phenomena of freezing point depression, boiling point elevation, and

osmotic pressure.

Practice Problems: 1. The freezing point and electrical conductivities of three aqueous solutions are given below.

Solution Freezing Electrical (0.010 molal) Point Conductivity

sucrose -0.0186°C almost zero formic acid -0.0213°C low

sodium formate -0.0361°C high Explain the relationship between the freezing point and electrical conductivity for each of the solutions above. Account for the differences in the freezing points among the three solutions.

2. (a) A solution containing 3.23 grams of an unknown compound dissolved in 100.0 grams of water freezes at -0.97°C. The solution does not conduct electricity. Calculate the molecular weight of the compound. (The molal freezing point depression constant for water is 1.86°C kg mok')

(b) Elemental analysis of this unknown compound yields the following percentages by weight H=9.74%; C=38.70%; 0=51.56%. Determine the molecular formula for the compound.

(c) Complete combustion of a 1.05 gram sample of the compound with the stoichiometric amount of oxygen gas produces a mixture of H,O(g) and CO,(g). What is the pressure of this gas mixture when it is contained in a 3.00 liter flask at 127°C?

3. The formula and the molecular weight of an unknown hydrocarbon compound are to be determined by elemental analysis and the freezing-point depression method. (a) The hydrocarbon is found to contain 93.46 percent carbon and 6.54 percent

hydrogen. Calculate the empirical formula of the unknown hydrocarbon.

(b) A solution is prepared by dissolving 2.53 grams of p-dichlorobenzene (molecular weight 147.0) in 25.86 grams of naphthalene (molecular weight 128.2). Calculate the molality of the p-dichlorobenzene solution.

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(c) The freezing point of pure naphthalene is determined to be 80.2°C. The solution prepared in (b) is found to have an initial freezing point of 75.7°C. Calculate the molal freezing-point depression constant of naphthalene.

(d) A solution of 2A2 grams of the unknown hydrocarbon dissolved in 26.7 grams of naphthalene is found to freeze initially at 76.2°C. Calculate the apparent molecular weight of the unknown hydrocarbon on the basis of the freezing-point depression experiment above.

(e) What is the molecular formula of the unknown hydrocarbon?

4. Concentrated sulfuric acid (l8A-molar H2S04) has a density of 1.84 grams per milliliter. After dilution with water to 5.2Q-molar, the solution has a density of 1.38 grams per milliliter and can be used as an electrolyte in lead storage batteries for automobiles. (a) Calculate the volume of concentrated acid required to prepare 1.00 liter of 5.20­

molar H2S04•

(b) Determine the mass percent ofH2S04 in the original concentrated solution. (c) Calculate the volume of 5.20-molar H2S04 that can be completely neutralized with

10.5 grams of sodium bicarbonate, NaHC03. (d) What is the molality ofthe 5.2Q-molar H2S04?

Topic #8: Kinetics (Chapter 12)

Goals: 1. Calculate the rate of a reaction. 2. Use experimental data to write rate laws for a chemical reaction.

3. Use the differential rate law to determine the integrated rate law. 4. Differentiate between integrated rate law and differential rate law. 5. Use a rate law and reaction rate to calculate the rate constant. 6. Determine the halflife of a reaction. 7. Determine how much reactant is left after a given amount of time using the

appropriate half-life equatioIL 8. Use graphs to determine the order ofthe reaction. 9. Use rate laws to determine if a reaction mechanism is plausible. 10. Use a reaction mechanism to write a rate law. 11. Explain what a catalyst does.

Practice Problems:

1. 2 NO(g) + O2 ->0 2 N02(g)

A rate expression for the reaction ab.Qve is: = k[NOJZIOz]

[NOW02{g)

:t'J02(g)

AHro kcallmole

So cal/(mole)(K)

AGr" kcal/mole

21.60 50.34 20.72 0 49.00 0 8.09 57A7 12.39

(a) For the reaction above, find the rate constant at 25°C if the initial rate, as defmed by the equation above, is 28 moles per liter-second when the concentration of nitric oxide is 0.20 mole per liter and the concentration of oxygen is 0.10 mole per liter.

(b) Calculate the equilibrium constant for the reaction at 25"C.

2. The decomposition ofcompound X is an elementary process that proceeds as follows: X(g)krkfA(g)B(g)H15 kilocalories

+ A 0=+

The forward reaction is slow at room temperature but becomes rapid when a catalyst is added.

(a) Draw a diagram of potential energy vs reaction coordinate for the uncatalyzed reaction. On this diagram label: (1) the axes (2) the energies of the reactants and the products (3) the ofthe activated complex (4) all energy differences

(b) On the same diagram indicate the change or changes that result from the addition of the catalyst. Explain the role ofthe catalyst in changing the rate of the reaction.

(c) If the temperature is increased, will the ratio kf/k,. increase, remain the same, or decrease? JustifY your answer with a one or two sentence explanation. [kr and k, are the specific rate constants for the forward and the reverse reactions, respectively.]

3. For a hypothetical chemical reaction that has the stoichiometry 2 X + Y Z, the->0

following initial rate data were obtained. All measurements were made at the same temperatul't:l.

~initial Rate of Formation ofZ, I Initial [X]"'

molL·1.sec·l) (moIL'I)

7.Ox 10"" 0.20 0.1 0 lAx 10'3 OAO 0.20 2.8xl0·3 OAO OAO 4.2xlO'3 0.60 0.60

(a) Give the rate law for this reaction from the data above. (b) Calculate the specific rate constant for this reaction and specifY its units. (c) How long must the reaction proceed to produce a concentration of Z equal to 0.20

molar, if the initial reaction concentrations are [X]o 0.80 molar, IY]o 0.60 molar and [Z]o = 0 molar?

)

(d) Select from the mechanisms below the one most consistent with the observed data, and explain your choice. In these mechanisms M and N are reaction intennediates. (1) X+Y--..M (slow)

X + M --.. Z (fast) (2) X + X ..... M (fast)

Y+M--..Z (slow) (3) Y --.. M (slow)

M + X --.. N (fast) N + X --.. Z (fast)

3. 2 NO(g) +2 H2(g) -+ N2(g) + 2 H20(g)

Experiments were conducted to study the rate of the reaction represented by the equation above. Initial concentrations and rates ofreaction are 'ven in the table below.

Initial Concentration

Initial Rate of Fonnation ofN2

4 0.0020 0.0060

(mol/I:rnin)

1.8 xl0-4

3.6xlo-" 0.30 X 10-4 1.2 X 10-4

(a) (i) Determine the order for each of the reactants, NO and H2, from the data given and show your reasoning. (ii) Write the overall rate law for the reaction.

(b) Calculate the value of the rate constant, k, for the reaction. Include units. (c) For experiment 2, calculate the concentration of NO remaining when exactly one­

half of the original amount of H2 had been consumed. (d) The following sequence of elementary steps is a proposed mechanism for the

reaction. 1. NO+NO ++ N20 2

II. N20 2 + H2 -+ H20 + N20 III. N20 + H2 -+ N2 + H20

Based on the data presented, which of the above is the rate-determining step? Show that the mechanism is consistent with (i) the observed rate law for the reaction, and (ii) the overall stoichiometry of the reaction.

Topic #9 Equilibrium and Solubility (Ch. #13)

Goals: 1. Define equilibrium 2. Describe the characteristics of a system at equilibrium 3. Write an equilibrium constant from a balance equilibrium equation 4. Calculate the equilibrium constant given initial concentrations ofreactants and

products. (Use ICE tables)

5. Calculate the equilibrium constant for gaseous systems using partial pressures. 6. Use the equilibrium constant ofa reaction to determine the extent ofthe reaction. 7. Use Le Chatelier's Principle to determine the shift of an equilibrium system when

stressed by temperature changes, pressure changes, and concentration changes. "" 8. Use Ksp to determine the solubility ofa solid in water. 9. Use solubility rules to determine the solubility of a solid in water.

Practice Problems: 1. ~CI(s) ..... NH3(g) + HCI(g)Ml = +42.1 kilocalories Suppose the substances in the reaction above are at equilibrium at 600K in volume V and at pressure P. State whether the partial pressure ofNH3(g) will have increased, decreased, or remained the same when equilibrium is reestablished after each of the following disturbances ofth.e original system. Some solid ~CI remains in the flask at all times. Justify each answer with a one-or-two sentence explanation. (a) A small quantity of~CI is added. (b) The temperature ofthe system is increased. (c) The volume of the system is increased. (d) A quantity of gaseous HCI is added. (e) A quantity of gaseous NH3 is added.

2. CO2(g) + H2(g) ++ H20(g) + CO(g)

When H2 (g) is mixed with C02W at 2,000 K, equilibrium is achieved according to the equation above. In one experiment, the following equilibrium concentrations were measured.

[H2] = 0.20 mollL [C02] = 0.30 mollL

[H20] = [CO] = 0.55 mollL

(a) What is the mole fraction of CO(g) in the equilibrium mixture? (b) Using the equilibrium concentrations given above, calculate the value of Kn the

equilibrium constant for the reaction. (c) Determine Kp in terms ofKc for this system. (d) When the system is cooled from 2,000 K to a lower temperature, 30.0 percent of the

CO(g) is converted back to CO2(g). Calculate the value of Kc at this lower temperature.

(e) In a different experiment, 0.50 mole of H2(g) is mixed with 0.50 mole of CO2(g) in a 3.0-liter reaction vessel at 2,000 K. Calculate the equilibrium concentration, in moles per liter, ofCO(g) at this temperature.

3. Ammonium hydrogen sulfide is a crystalline solid that decomposes as follows: N~HS(s) ,.... NH3(g) + H2S(g)

(a) Some solid ~HS is placed in an evacuated vessel at 25°C. After equilibrium is attained, the total pressure inside the vessel is found to be 0.659 atmosphere. Some solid N~HS remains in the vessel at equilibrium. For this decomposition, write the expression for Kp and calculate its numerical value at 25°C.

) ) )

(b) Some extra NH3 gas is injected into the vessel containing the sample described in part When equilibrium is reestablished at 25°C, the partial pressure ofNH3in the

twice the partial pressure of H2S. Calculate the numerical value of the partial pressure ofNH3and the partial pressure of H2S in the vessel after the NH3has been added and the equilibrium has been reestablished.

(c) In a different experiment, NH3 gas and H2S gas are introduced into an empty 1.00 liter vessel at 25°C. The initial partial pressure of each gas is 0.500 atmospheres. Calculate the number of moles of solid N~HS that is present when equilibrium is established.

4. At 25"C the solubility product constant, K.p, for strontium sulfate, srS04, is 7.6xlO·7•

The solubility product constant for strontium fluoride, SrF2, is 7.9xlO"lo.

(a) What is the molar solubility of SrS04 in pure water at 25°C? (b) What is the molar solubility of SrF2in pure water at 25"C? (c) An aqueous solution of Sr(N03h is added slowly to 1.0 litre ofa well-stirred solution

containIng 0.020 mole F' and 0.10 mole SOl' at 25°C. (You may assume that the added Sr(N03)2 solution does not materially aff'ectthe total volume of the system.)

1. Which salt precipitates first? 2. What is the concentration of strontium ion, Srl+, in the solution when the first

precipitate begins to form? (d) As more Sr(N03)2 is added to the mixture in (c) a second precipitate

At that stage, what percent of the anion ofthe first precipitate remains in

Topic #10: Acids and Bases (Ch. #14 and #15)

Goals: 1. Identify Arrhenius acids and Arrhenius bases in a chemical reaction. 2. Identify Bronsted-Lowry acids and bases in a chemical reaction. 3. Identify an acid, base, conjugate base, and conjugate acid in a dissociation

reaction using the Bronsted-Lowry definitions. 4. Write the equilibrium Ka, for the dissociation ofa weak acid or Kb

for the dissociation ofa weak base. 5. Calculate the pH, pOH, Ka, Kb, [H+], and [Off] given one ofthe other pieces of

information. 6. Calculate the pH of a strong acid or base concentration

Kb and the initial concentrations.

7. Calculate the pH of a weak acid or base

8. Calculate the percent dissociation of a weak acid or weak base. 9. Calculate Ka when given the percent dissociation. 10. Calculate the pH ofpolyprotic acids. 11. Describe the dissociation ofpolyprotic acids. 12. Determine the pH of a salt when mixed with water. 13. Define Lewis acid and Lewis base. 14. Identify a lewis acid or base in a chemical reaction. 15. Calculate the pH of a buffer solution given initial concentrations and Ka or Kb.

16. Calculate the pH of a buffer solution after having a strong acid or strong base added to the

17. Explain how buffering works. 18. Explain how an indicator works 19. Pick an appropriate indicator for a titration. 20. Draw a titration curve for the following combinations:

a. Strong acid with a strong base b. Strong base with a strong acid

~ c. Weak base with a strong acid d. Weak acid with a strong base

21. Draw the titration curve for a polyprotic acid with a strong base. 22. Calculate the pH of solutions during titrations. 23. Use a titration graph to calculate the Ka, pKa, and equivalence point.

Practice Problems: 1. (a) What is the pH ofa 2.0 molar solution of acetic acid. K.. acetic acid 1.8x10.5

(b) A buffer solution is prepared by adding 0.10 liter of2.0 molar acetic acid solution to 0.1 liter of a 1.0 molar sodium hydroxide solution. Compute the hydrogen ion concentration of the buffer solution.

(c) Suppose that 0.10 liter of 0.50 molar hydrochloric acid is added to 0.040 liter of the buffer prepared in (b). Compute the hydrogen ion concentration of the resulting solution.

2. A 5.00 gram sample of a dry mixture of potassium hydroxide, potassium carbonate, and potassium chloride is reacted with 0.100 liter of 2.00 molar HCl solution (a) A 249 milliliter sample of dry CO2 gas, measured at 22°C and 740 torr, is obtained

from this reaction. What is the percentage ofpotassium carbonate in the mixture? (b) The excess HCl is found by titration to be chemically equivalent to 86.6 milliliters of

1.50 molar NaOH. Calculate the percentages of potassium hydroxide and of potassium chloride in the original mixture.

3. The molecular weight of a monoprotic acid HX was to be determined. A sample of 15.126 grams of HX was dissolved in distilled water and the volume brought to exactly 250.00 millilitres in a volumetric flask. Several 50.00-millilitre portions of this solution were titrated against NaOH solution, requiring an average of38.21 millilitres ofNaOH.

The NaOH solution was standardized against oxalic acid dihydrate, H2C20 4'2H20 (molecular 126.066 gram mol· I

). The volume of NaOH solution required to neutralize grams ofoxalic acid dihydrate was 41.24 millilitres.

(a) Calculate the molarity of the NaOH solution. (b) Calculate the number ofmoles ofHX in a 50.00 millilitre portion used for titration. (c) Calculate the molecular weight ofHX.

(d) Discuss the effect of the calculated molecular weight of HX if the sample of oxalic acid dihydrate contained a nonacidic impurity.

') )

2

o I I II I o 5 10 15 20 25 30

4. millilitres ofNaOH

A 30.00 millilitre sample of a weak monoprotic acid was titrated with a standardized solution of NaOH. A pH meter was used to measure the pH after each increment of NaOH was added, and the curve above was constructed.

(a) Explain how this curve could be used to determine the molarity of the acid. (b) Explain how this curve could be used to determine the dissociation constant K.. ofthe

weak monoprotic acid. (c) If you were to repeat the titration using a indicator in the acid to signal the endpoint,

which of the following indicators should you select? Give the reason for your choice. Methyl red K. = lx1()S Cresol red K.. = lxl()8 Alizarin yellow K. = lx1()11

(d) Sketch the titration curve that would result if the weak monoprotic acid were replaced by a strong monoprotic acid, such as HCI of the same molarity. Identify differences between this titration curve and the curve shown above.

Topic #11: Electrochemistry (Ch. 17)

Goals: 1. Draw a galvanic cell 2. Calculate the cell potential ofa galvanic cell. 3. Use the half-reactions of a galvanic cell to draw and label the parts. 4. Use cell potential to calculate electrical work and free energy. 5. Calculate the cell potential when the concentrations are not 1 M. (Use the Nernst

Equation) 6. Describe an electrolytic cell. 7. Calculate the mass ofmetal plated using an electrolytic cell when given the curren

and time of the electroplating.

Practice Problems:

1. Br2 + 2 Fe2+(aq) --. 2 Bnaq) + 2 Fe3+(aq)

For the reaction above, the following data are available: 2 Br'(aq) --. Br2(1) + 2e- EO = -1.07 volts Fe2+(aq) - Fe3+(aq) + e- EO = -0.77 volts

S", callmoleK Br2(l) 58.6 Fel+(aq) -27.1 Br"(aq) 19.6 Fe3+(aq) -70.1

'(a) Determine ASo (b) Determine AGO (c) Determine Mlo

2. A steady current of 1.00 ampere is passed through an electrolytic cell containing a 1 molar solution ofAgN03 and having a silver anode and a platinum cathode until 1.54 grams of silver is deposited. (a) How long does the current flow to obtain this deposit? (b) What weight of chromium would be deposited in a second cell containing I-molar

chromium(lII) nitrate and having a chromium anode and a platinum cathode by the same current in the same time as was used in the silver cell?

(c) Ifboth electrodes were platinum in this second cell, what volume of ~ gas measured at standard temperature and pressure would be released at the anode while the chromium is being deposited at the cathode? The current and the time are the same as in (b)

3. An electrochemical cell consists of a tin electrode in an acidic solution of 1.00 molar Sn2+ connected by a salt bridge to a second compartment with a silver electrode in an acidic solution of 1.00 molar Ag+. (a) Write the equation for the half-cell reaction occurring at each electrode. Indicate

which half-reaction occurs at the anode. (b) Write the balanced chemical equation for the overall spontaneous cell reaction that

occurs when the circuit is complete. Calculate the standard voltage, E ", for this cell reaction.

(c) Calculate the equilibrium constant for this cell reaction at 298K. (d) A cell similar to the one described above is constructed with solutions that have

initial concentrations of 1.00 molar Sn2+and 0.0200 molar Ag+. Calculate the initial voltage, EO, ofthis cell.

Topic #12: Nuclear Chem (Ch. 18)

Goals: 1. Explain why the nucleus ofradioactive elements are unstable. 2. Describe the types ofradioactive decay in terms ofmass and protons. 3. Describe how penetrating the different types ofradioactive decay are.

) ) )

4. Write and balance nuclear equations. 5. Calculate halflife for nuclear reactions. 6. Explain the processes of fission and fusion. 7. Explain how the mass of the nucleus is less than the sum of the protons and

neutrons.

Practice Problems: See break homework

Topic #13: Organic Chem (Ch. 21)

Goals: 1. Name organic compounds. 2. Define isomer 3. Identify the functional groups oforganic compounds.

Practice Problems:

1. Consider the hydrocarbon pentane, CSHI2 (molar mass 72.15 g).

(a) Write the balanced equation for the combustion of pentane to yield carbon dioxide and water.

(b) What volume of dry carbon dioxide, measured at 2S·C and 785 mm Hg, will result from the complete combustion of 2.S0 g of pentane?

(c) The complete combustion of 5.00 of pentane releases 243 k:J of heat. On the basis of this information, calculate the of MI for the complete combustion of one mole of pentane.

(d) Under identical conditions, a sample of an unknown gas effuses into a vacuum at twice the rate that a sample of pentane gas effuses. Calculate the molar mass of the unknown gas.

(e) The structural formula of one isomer of pentane is shown below. Draw the structural formulas for the other two isomers of pentane. Be sure to include all atoms of hydrogen and carbon in your structures.

~~~~~ H-« -«-G -G-~-H

HHHHH

2.

Using the information in the table above, answer the following questions about organic compounds.

(a) For propanone,

(i) draw the complete structural formula (showing all atoms and bonds);

(ii) predict the approximate carbon-to-carbon-to-carbon bond angle.

(b) For each pair of compounds below, explain why they do not have the same value for their standard heat of vaporization, Mrvap. (You must include specific information about lmfu compounds in each pair.)

(i) Propane and propanone

(ii) Propanone and I-propanol

(c) Draw the complete structural formula for an isomer of the molecule you drew in, /part (a) (i).

(d) Given the structural formula for propyne below, H

+ H-e-esC-H

I H

(i) indicate the hybridization of the carbon atom indicated by the arrow in the structure above;

(ii) indicate the total number of sigma (0) bands and the total number of pi (n) bonds in the molecule

I Compound Compound Mrvap

Name Formula (kJmor1 )

I Propane CH3CH2CH3 19.0 ~~~-

32.0I Propanone CH3COCH3 ~~~-

I I-propanol CH3CH2CH20H 47.3