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AP Chemistry 1st Semester Exam Review

Chemistry Basics

1. Rubidium has two naturally occurring isotopes, 85Rb (relative mass 84.9118 amu) and 87Rb (relative mass 86.9092 amu). If rubidium has an average atomic mass of 85.47 amu, what is the abundance of each isotope (in percent)?

2. An element has three isotopes. Given the abundances and relative masses, calculate the average atomic mass and determine (from the periodic table) which element it is.

Abundances Relative masses 0.005% 234.040947 amu0.720% 235.043924 amu99.275% 238.050784 amu

3. Calculate the mass of two oxygen molecules.

4. A block of aluminum occupies a volume of 15.0 mL and weighs 40.5 g. What is its density?

5. Mercury metal is poured into a graduated cylinder that holds exactly 22.5 mL. The mercury used to fill the cylinder weighs 306.0 g. From this information, calculate the density of mercury.

6. The simplest formula for an oxide of nitrogen that is 36.8 percent nitrogen by weight is…

7. What number of moles of O2 is needed to produce 14.2 grams of P4O10 from P? (Molar Mass P4O10 = 284)

Stoichiometry

1. A compound contains only carbon, hydrogen, nitrogen, and oxygen. Combustion of 0.157 g of the compound produced 0.213 g of CO2, and 0.0310 g of H2O. In another experiment it is found that 0.103g of the compound produces 0.0230g of NH3. What is the empirical formula of the compound?

2. A 15.67 g sample of a hydrate of magnesium carbonate was heated, without decomposing the carbonate, to drive off the water. The mass was reduced to 7.58 g. What is the formula of the hydrate?

3. Methane gas reacts with chlorine gas to form dichloromethane and hydrogen chloride, as represented by the equation below.

CH4(g) + 2 Cl2(g) CH2Cl2(g) + 2 HCl(g)

A 25.0 g sample of methane is placed in a reaction vessel containing 2.58 mol of Cl2(g).a. Identify the limiting reactant when the methane and chlorine gases are combined. Justify your

answer with a calculation.

b. If 90.0 g of CH2Cl2(g) is collected in the actual experiment conducted with the starting materials described above, what is the % yield.

Solution Stoichiometry

1. The following reactions

2K(s) + Br2(l) 2KBr(s) AgNO3(aq) + NaCl(aq) AgCl(s) + NaNO3(aq) HCl(aq) + KOH(aq) H2O(l) + KCl(aq)

are examples ofA) precipitation reactionsB) redox, precipitation, and acid-base, respectivelyC) precipitation (two) and acid-base reactions, respectivelyD) redox reactionsE) none of these

2. 6 I¯ + 2 MnO4¯ + 4 H2O(l) 3 I2(s) + 2 MnO2(s) + OH¯

Which of the following statements regarding the reaction represented by the equation above is correct?(A) Iodide ion is oxidized by hydroxide ion.(B) MnO4¯ is oxidized by iodide ion. (C) The oxidation number of manganese changes from +7 to +2.(D) The oxidation number of manganese remains the same.(E) The oxidation number of iodine changes from -1 to 0.

3. Which of the following does NOT behave as an electrolyte when it is dissolved in water?

(A) CH3OH (B) K2CO3 (C) NH4Br (D) HI (E) Sodium acetate, CH3COONa

2. When 100 mL of 1.0 M Na3PO4 is mixed with 100 mL of 1.0 M AgNO3 , a yellow precipitate forms and [Ag+] becomes negligibly small. Which of the following is a correct listing of the ions remaining in solution in order of increasing concentration?

(A) [PO4 3− ] < [NO3 − ] < [Na + ] (B) [PO4 3− ] < [Na + ]< [NO3 − ] (C) [NO3 − ] < [PO4 3− ]< [Na + ]

(D) [Na + ]< [NO3 − ] < [PO4 3− ] (E) [Na + ]< [PO4 3− ]< [NO3 − ]

3. In a qualitative analysis for the presence of Pb2+, Fe2+, and Cu2+ ions in a aqueous solution, which of the following will allow the separation of Pb2+ from the other ions at room temperature?

(A) Adding dilute Na2S(aq) solution (B) Adding dilute HCl(aq) solution (C) Adding dilute NaOH(aq) solution

(D) Adding dilute NH3(aq) solution (E) Adding dilute HNO3(aq) solution

4. The weight of H2SO4 (molecular weight 98.1) in 50.0 milliliters of a 6.00-molar solution is…

5. How many milliliters of 11.6-molar HCl must be diluted to obtain 1.0 liter of 3.0-molar HCl?

6. When 70. milliliters of 3.0-molar Na2CO3 is added to 30. milliliters of 1.0-molar NaHCO3 the resulting concentration of Na+ is…

7. The volume of distilled water that should be added to 10.0 mL of 6.00 M HCl (aq) in order to prepare a 0.500 M HCl(aq) solution is approximately…

8. A 20.0-milliliter sample of 0.200-molar K2CO3 solution is added to 30.0 milliliters of 0.400-molar Ba(NO3)2 solution. Barium carbonate precipitates. The concentration of barium ion, Ba2+, in solution after reaction is…

9, A student wishes to prepare 2.00 liters of 0.100-molar KIO3 (molecular weight 214). The proper procedure is

Reactions

1. Equal volumes of dilute equimolar solutions of sodium carbonate and hydrochloric acid are mixed.

2. Dilute sulfuric acid is added to a solution of barium acetate.

3. Solutions of potassium phosphate and zinc nitrate are mixed.

4. Dilute acetic acid solution is added to solid magnesium carbonate.

5. Solutions of ammonia and hydrofluoric acid are mixed.

6. Hydrogen sulfide gas is bubbled through a solution of potassium hydroxide.

7. A solution of sulfuric acid is added to a solution of barium hydroxide until the same number of moles of each compound has been added.

8. Potassium permanganate solution is added to concentrated hydrochloric acid.

9. Manganese (IV) oxide is added to warm, concentrated hydrobromic acid.

10. Chlorine gas is bubbled into cold dilute sodium hydroxide.

11. Solid iron (III) oxide is heated in excess carbon monoxide.

12. Aluminum metal is added to a solution of copper (II) chloride.

13. Hydrogen gas is passed over hot copper (II) oxide.

14. Small chunks of solid sodium are added to water.

15. Redox reactions 8-14, identify the oxidation and reduction half reactions. Indicate the initial and final oxidation state of the species that oxidizes.

Thermodynamics

11. Which of the following reactions has the largest positive value of ∆S per mole of Cl2 ?

(A) H2(g) + Cl2(g) 2 HCl(g) (B) Cl2(g) + 1/2 O2(g) Cl2O(g) (C) Mg(s) + Cl2(g) MgCl2(s)

(D) 2 NH4Cl(s) N2(g) + 4 H2(g) + Cl2(g) (E) Cl2(g) 2 Cl(g)

2. Which of the following must be true for a reaction that proceeds spontaneously from initial standard state conditions?

(A) ∆G° > 0 and Keq > 1 (B) ∆G° > 0 and Keq < 1 (C) ∆G° < 0 and Keq > 1

(D) ∆G° < 0 and Keq < 1 (E) ∆G° = 0 and Keq = 1

3. H2O(s) H2O(l)

When ice melts at its normal melting point, 273.16 K and 1 atmosphere, which of the following is true for the process shown above?

(A) ∆H < 0, ∆S > 0, ∆V > 0 (B) ∆H < 0, ∆S < 0, ∆V > 0 (C) ∆H > 0, ∆S < 0, ∆V < 0

(D) ∆H > 0, ∆S > 0, ∆V > 0 (E) ∆H > 0, ∆S > 0, ∆V < 0

4. For which of the following processes would ∆S have a negative value?

I. 2 Fe2O3(s) 4 Fe(s) + 3 O2(g)

II. Mg2+ + 2 OH− Mg(OH)2(s)

III. H2(g) + C2H4(g) 3 C2H6(g)

(A) I only (B) I and II only (C) I and III only (D) II only (E) I, II, and III

5. N2(g) + 3 H2(g) 2 NH3(g)

The reaction indicated above is thermodynamically spontaneous at 298 K, but becomes nonspontaneous at higher temperatures. Which of the following is true at 298 K?

(A) ∆G, ∆H, and ∆S are all positive. (B) ∆G, ∆H, and ∆S are all negative.

(C) ∆G and ∆H are negative, but ∆S is positive. (D) ∆G and ∆S are negative, but ∆H is positive.

(E) ∆G and ∆H are positive, but ∆S is negative.

6. Which of the following is a graph that describes the pathway of reaction that is endothermic and has high activation energy?

(A) (B)

(C)

(D) (E)

7.

H2(g) + 1/2 O2(g) H2O(l) ∆H° = x

2 Na(s) + 1/2 O2(g) Na2O(s) ∆H° = y

Na(s) + 1/2 O2(g) + 1/2 H2(g) NaOH(s) ∆H° = z

Based on the information above, what is the standard enthalpy change for the following reaction?

Na2O(s) + H2O(l) 2 NaOH(s)

(A) x + y + z (B) x + y − z (C) x + y − 2z (D) 2z − x −y (E) z − x −y

BondAverage Bond Energy

(kJ/mole)

I---I 150

Cl---Cl 240

I---Cl 210

8. I2(g) + 3 Cl2(g) 2 ICl3(g)

According to the data in the table above, what is the value of ∆ H° for the reaction represented above?

9. When a 25.7 g sample of NaI dissolves in 80.0 g of water in a calorimeter, the temperature rises from 20.5 oC to 24.4 oC. Calculate H for the process.

10. A 28.4 g sample of aluminum is heated to 39.4 oC, then is placed in a calorimeter containing 50.0 g of water. Temperature of water increases from 21.00 oC to 23.00 oC. What is the specific heat of aluminum?

11. CH4(g) + 2 O2(g) CO2(g) + 2 H2O(l); ∆Hrxn = −889.1 kJ Calculate the heat released when 2.00 L of CH4 is burned at STP.

12. WO3(s) + 3 H2(g) W(s) + 3 H2O(g)

Tungsten is obtained commercially by the reduction of WO3 with hydrogen according to the equation above. The following data related to this reaction are available:

WO3(s) H2O (g)

Hf (kilocalories/mole) -200.84 -57.8

Gf (kilocalories/mole) -182.47 -54.6

(a) What is the value of the equilibrium constant for the system represented above?

(b) Calculate S at 25C for the reaction indicated by the equation above.

(c) Find the temperature at which the reaction mixture is in equilibrium at 1 atmosphere.

13. Standard Heat of Absolute

Formation, Hf, Entropy, S,

Substance in kJ mol-1 in J mol-1 K-1

-----------------------------------------------------------------------

C(s) 0.00 5.69

CO2(g) -393.5 213.6

H2(g) 0.00 130.6

H2O(l) -285.85 69.91

O2(g) 0.00 205.0

C3H7COOH(l) ? 226.3

The enthalpy change for the combustion of butyric acid at 25C, Hcomb, is -2,183.5 kilojoules per mole. The combustion reaction is

C3H7COOH(l) + 5 O2(g) 4 CO2(g) + 4 H2O(l)

(a) From the above data, calculate the standard heat of formation, Hf, for butyric acid.

(b) Write a correctly balanced equation for the formation of butyric acid from its elements.

(c) Calculate the standard entropy change, Sf, for the formation of butyric acid at 25C. The entropy change, S, for the combustion reaction above is -117.1 J K-1 at 25C.

(d) Calculate the standard free energy of formation, Gf, for butyric acid at 25C.

Gases

1. As the temperature is raised from 20 °C to 40 °C, the average kinetic energy of neon atoms changes by a factor of…

(A) ½ (B) √313/293) (C) 313/293 (D) 2 (E) 4

2. When a sample of oxygen gas in a closed container of constant volume is heated until its absolute temperature is doubled, which of the following is also doubled?

(A) The density of the gas (B) The pressure of the gas (C) The average velocity of the gas molecules

(D) The number of molecules per cm3 (E) The potential energy of the molecules

3. Equal masses of three different ideal gases, X, Y, and Z, are mixed in a sealed rigid container. If the temperature of the system remains constant, which of the following statements about the partial pressure of gas X is correct?

(A) It is equal to 1/3 the total pressure

(B) It depends on the intermolecular forces of attraction between molecules of X, Y, and Z.

(C) It depends on the relative molecular masses of X, Y, and Z.

(D) It depends on the average distance traveled between molecular collisions.

(E) It can be calculated with knowledge only of the volume of the container.

4. When the actual gas volume is greater than the volume predicted by the ideal gas law, the explanation lies in the fact that the ideal gas law does NOT include a factor for molecular…

(A) volume (B) mass (C) velocity (D) attractions (E) shape

5. A compound is heated to produce a gas whose molecular weight is to be determined. The gas is collected by displacing water in a water-filled flask inverted in a trough of water. Which of the following is necessary to calculate the molecular weight of the gas, but does NOT need to be measured during the experiment?

(A) Mass of the compound used in the experiment (B) Temperature of the water in the trough

(C) Vapor pressure of the water (D) Barometric pressure

(E) Volume of water displaced from the flask

6. A gaseous mixture containing 7.0 moles of nitrogen, 2.5 moles of oxygen, and 0.50 mole of helium exerts a total pressure of 0.90 atmospheres. What is the partial pressure of the nitrogen?

7. 2 K + 2 H2O 2 K+ + 2 OH− + H2

When 0.400 moles of potassium reacts with excess water at standard temperature and pressure as shown in the equation above, the volume of hydrogen gas produced is…

8. The density of an unknown gas is 4.20 grams per liter at 3.00 atmospheres pressure and 127 °C. What is the molecular weight of this gas?

9. 3 Ag(s) + 4 HNO3 ⇄ 3 AgNO3 + NO(g) + 2 H2O

The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.10 moles of powdered silver is added to 10.0 milliliters of 6.0-molar nitric acid, the number of moles of NO gas that can be formed is…

Atomic Structure and Periodic Table Trends

Use these answers for questions 1-3

(A) Heisenberg uncertainty principle (B) Pauli exclusion principle (C) Hund's rule

(D) Shielding effect (E) Wave nature of matter

1. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic

2. Indicates that an atomic orbital can hold no more than two electrons

3. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron

4. The elements in which of the following have most nearly the same atomic radius?

(A) Be, B, C, N (B) Ne, Ar, Kr, Xe (C) Mg, Ca, Sr, Ba

(D) C, P, Se, I (E) Cr, Mn, Fe, Co

5. Which of the following represents the ground state electron configuration for the Mn3+ ion?

(A) 1s2 2s22p6 3s23p63d4 (B) 1s2 2s22p6 3s23p63d5 4s2 (C) 1s2 2s22p6 3s23p63d2 4s2

(D) 1s2 2s22p6 3s23p63d8 4s2 (E) 1s2 2s22p6 3s23p63d3 4s1

Use these answers for questions 6-9.

(A) 1s2 2s2 2p5 3s2 3p5 (B) 1s2 2s2 2p6 3s2 3p6 (C) 1s2 2s2 2p6 2d10 3s2 3p6

(D) 1s2 2s2 2p6 3s2 3p6 3d5 (E) 1s2 2s2 2p6 3s2 3p6 3d3 4s2

6. An impossible electronic configuration

7. The ground-state configuration for the atoms of a transition element

8. The ground-state configuration of a negative ion of a halogen

9. The ground-state configuration of a common ion of an alkaline earth element

10. What is the wavelength of a photon of red light (in nm) whose frequency is 4.55 1014 Hz?

11. What is the energy of a photon of blue light that has a wavelength of 453 nm?

Ionization Energies for element X (kJ mol−1)

First Second Third Fourth Five

580 1815 2740 11600 14800

12. The ionization energies for element X are listed in the table above. On the basis of the data, element X is most likely to be…

(A) Na (B) Mg (C) AI (D) Si (E) P

13. The effective nuclear charge experienced by the outermost electron of Na is different than the effective nuclear charge experienced by the outermost electron of Ne. This difference best accounts for which of the following?

(A) Na has a greater density at standard conditions than Ne.(B) Na has a lower first ionization energy than Ne.(C) Na has a higher melting point than Ne.(D) Na has a higher neutron-to-proton ratio than Ne.(E) Na has fewer naturally occurring isotopes than Ne.

Bonding

1. Molecules that have planar configurations include which of the following?

I. BCl3 II. CHCl3 III. NCl3

(A) I only (B) III only (C) I and II only (D) II and III only (E) I, II, and III

2. The electron-dot structure (Lewis structure) for which of the following molecules would have two unshared pairs of electrons on the central atom?

(A) H2S (B) NH3 (C) CH4 (D) HCN (E) CO2

3. Contains 1 sigma (σ) and 2 pi (π) bonds

(A) Li2 (B) B2 (C) N2 (D) O2 (E) F2

4. Has the largest bond-dissociation energy

(A) Li2 (B) B2 (C) N2 (D) O2 (E) F2

5. The SbCl5 molecule has trigonal bipyramid structure. Therefore, the hybridization of Sb orbitals should be...

(A) sp2 (B) sp3 (C) sp2d (D) sp3d (E) sp3d2

Use these answers for questions 6 - 8.

(A) hydrogen bonding (B) hybridization (C) ionic bonding

(D) resonance (E) van der Waals forces (London dispersion forces)

6. Is used to explain why iodine molecules are held together in the solid state

7. Is used to explain why the boiling point of HF is greater than the boiling point of HBr

8. Is used to explain the fact that the four bonds in methane, CH4, are equivalent

9. Draw the best Lewis Dot Structure for each of the following species. Also indicate the electronic arrangement and the molecular geometry for each.

Lewis Structure Electronic Arrangement Molecular GeometryBeF2

BCl3

CCl4

PBr5

SI6

BH2–

NI3

ClF4+

SF5–