Download Advanced Placement Chemistry 1st Semester Review

Advanced Placement Chemistry 1st Semester Review
1.
Answer the following questions for each of the listed elements.
a. Charge of each possible ion that forms in aqueous solution.
b. Formula for any possible polyatomic ion that forms, if any at all.
c. Formula of common oxides that form. Identify which oxides yield acidic solution in water and which yield
basic solution in water.
Element
Ion(s)
Ion Formula(s)
Oxide formula
Oxide Rxn w/
Acid
Oxide Rxn w/
Base
(A) F
(B) S
(C) Mg
(D) Ar
(E) Mn
2.
3.
List common uses for each of the following substances. Also, list common properties (oxidizer, reducer,
amphoteric properties, acidic, basic, neutral) for each.
(A) Hydrofluoric acid
_____________________________________________________________
(B) Carbon dioxide
_____________________________________________________________
(C) Aluminum hydroxide
_____________________________________________________________
(D) Ammonia
_____________________________________________________________
(E) Hydrogen peroxide
_____________________________________________________________
Define the following terms. Give examples of each type of solid.
(A) A network solid
______________________________________________
(B) A molecular solid with zero dipole moment ______________________________________________
4.
(C) A molecular solid with hydrogen bonding
______________________________________________
(D) An ionic solid
______________________________________________
(E) A metallic solid
______________________________________________
Assume that you have an "unknown" consisting of an aqueous solution of a salt that contains one of the ions
listed above.
a. Provide the color of the ion in solution.
b. Identify any characteristic odor when mixed with a base then with an acid.
c. Identify any ppt that forms from the reaction with the SO42- ion and with the OH1- ion.
Formula
Name
Color of ion (aq) Odor
PPT (color) w/ PPT (color)
SO42w/ OH12
(A) CO3 ¯
(B) Cr2O72¯
(C) NH4+
(D) Ba2+
(E) Al3+
©T. M. Chipi, chemistry instructor
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5.
The liquefied hydrogen halides have the normal boiling points given below. Explain the relatively high boiling
point of HF as compared to the other halogens.
Hydrogen Halide Normal Boiling Point, °C
HF
+19
HCl
- 85
HBr
- 67
HI
- 35
6.
Provide examples below representing a pair of isotopes.
Atomic Number Mass Number Element Name
I.
(A)
II.
I.
(B)
II.
I.
(C)
II.
I.
(D)
II.
7.
What is the effect observed on each of the following when a gas in a closed container is heated at constant
volume until its absolute temperature is 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 cm
3
(E) The potential energy of the molecules
______________________________________________
______________________________________________
______________________________________________
8.
Which element is represented by the electronic configuration shown, 1s2 2s2 2p6 3s2 3p3? __________________
What is the formula of the compound that forms when this element reacts with Group 1A, 2A, & 3A elements?
9.
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? (R = 0.0821 liter-atm / mole-K)
10. The critical temperature of a substance is the _____________________________________________________
11. Write the nuclear equation when Po-214 decays; the emission consists consecutively of an alpha particle, then
two beta particles, and finally another alpha particle.
__________________________________________________________________________________________
12. What is the color of each of the following ions in solution?
(A) Fe(H2O)42+
____________________
2+
(B) Cu(NH3)4
____________________
(C) Cr2O72____________________
+
(D) K
____________________
(E) Br¯
____________________
(F) NO3¯
____________________
(G) I¯
____________________
2+
H) Co
____________________
©T. M. Chipi, chemistry instructor
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13. Write the net ionic equation for the reaction between calcium carbonate and hydrochloric acid.
_________________________________________________________________________________________
14. A(s) + B(g) <==> C(s) + D(g); ΔH = (-)
How can the substances in the equation above at equilibrium, at pressure P and temperature T, be shifted to
favor the products?
_________________________________________________________________________________________
15. What is needed in order to convert from molarity to molality? ____________________________________
16. What process of radioactive decay changes C-14 to N-14? _______________________________________
17. Equal masses of (3) different ideal gases, X, Y, and Z, are mixed in a sealed rigid container at constant
temperature. What is true about the partial pressure of gas X? How can it be calculated?
_________________________________________________________________________________________
18. Draw the geometry of the SO3 molecule.
19. Give the type of covalent bond found in each of the following molecules and rank them from the shortest to the
longest bond length, and from the strongest to the weakest bond.
Element
Covalent Bond (S,D,T)
Bond Length (1-5)
Bond Strength (1-5)
(A) N2
(B) O2
(C) Cl2
(D) Br2
(E) I2
20. Metallic copper is heated strongly with concentrated sulfuric acid. Write the reaction.
_________________________________________________________________________________________
21. What is the observed trend for atomic radius across the periods for the transition metals?
_________________________________________________________________________________________
22. Write the reaction producing tetraphosphorus decoxide from the burning of white phosphorus. How many
moles of O2 is needed to produce 17.0 grams of P4O10? _____________________________________________
23. What is the general formula of an alkene? If 0.3.00 gram of any alkene is burned in excess oxygen, what
number of moles of H2O is formed? ____________________________________________________________
CH4(g) + 2 O2(g)  CO2(g) + 2 H2O(l); ΔΗ° = - 889.1 kJ
ΔHf° H2O(l) = - 285.8 kJ / mole
ΔHf° CO2(g) = - 393.3 kJ / mole
24. What is the standard heat of formation of methane, ΔHf° CH4(g), as calculated from the data above?
25. Two flexible containers for gases are at the same temperature and pressure. One holds 0.50 gram of hydrogen
and the other holds 8.0 grams of oxygen. What can be said regarding each of the following for each gas?
(A) The volume of the containers.
________________________________
(B) The number of molecules in each container.
________________________________
(C) The density of each sample of gas.
________________________________
(D) The average kinetic energy of the molecules of each gas.
________________________________
(E) The average speed of the gas molecules for each gas.
________________________________
©T. M. Chipi, chemistry instructor
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26. Where does Pi bonding occur in a species? ______________________________________________________
3 Ag(s) + 4 HNO3 <==> 3 AgNO3 + NO(g) + 2 H2O
27. The reaction of silver metal and dilute nitric acid proceeds according to the equation above. If 0.30 mole of
powdered silver is added to 30.mL of 3.0-molar nitric acid, how many moles of NO gas can be formed?
28. What is always true about the phase diagram of any one-component system with respect to
a. the slope of the curve representing equilibrium between phases
______________________________
b.
the temperature at the triple point vs. that at the normal freezing point ______________________________
c.
the pressure at the triple point compared to normal pressure?
______________________________
29. At 20. °C, the vapor pressure of toluene is 22 millimeters of mercury and that of benzene is 75 millimeters of
mercury. An ideal solution, equimolar in toluene and benzene, is prepared. At 20. °C, what is the mole fraction
of benzene in the vapor in equilibrium with this solution?
30. Ice is added to some hot water in a Thermos-bottle, which is then sealed. What has happened to the total energy
and the total entropy when the system reaches equilibrium?
31. Write the ground state electron configuration for the Mn3+ ion. _______________________________________
32. When 70.mL of 3.0-molar Na2CO3 is added to 30.mL of 1.0-molar NaHCO3 the resulting [Na+] is
33. State whether each of the following has a dipole moment or not (a zero dipole moment).
(A) HCN
(B) NH3
(C) SO2
(D) NO2
(E) PF5
_________________
_________________
_________________
_________________
_________________
34. A solution of potassium dichromate is added to an acidified solution of iron (II) sulfate. Write the redox
reaction.
_________________________________________________________________________________________
35. What is the proper procedure used in the titration of an acid with a base and vise-versa?
_________________________________________________________________________________________
36. Write the net ionic equation for the reaction that occurs during the titration of nitrous acid with sodium
hydroxide ________________________________________________________________________________
37. Determine which of the following species functions as an oxidizing agent and which as a reducing agent.
(A) Cr2O72¯ __________________________
(B) MnO4¯ __________________________
(C) NO3¯
__________________________
(D) S
__________________________
(E) I¯
__________________________
38. Determine which of the following elements are paramagnetic and which are diamagnetic.
a. Ca
__________________________
b. V
__________________________
c. Co
__________________________
d. Zn
__________________________
e. As
__________________________
©T. M. Chipi, chemistry instructor
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39. What is the proper procedure in the preparing of a molar solution of a soluble solid in water?
_________________________________________________________________________________________
40. 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
41. What is the mole fraction of ethanol, C2H5OH, in an aqueous solution in which the [C2H5OH] = 4.6 molal?
42. Use a set of four quantum numbers to describe one of the outermost electrons in a barium atom in the ground
state. ___________________________________________
43. What is needed to calculate the molecular weight of a gas collected by displacing water in a water-filled flask
inverted in a trough of water, but does NOT need to be measured during the experiment?
_________________________________________________________________________________________
44. A 27.0-gram sample of an unknown hydrocarbon was burned in excess oxygen to form 88.0 grams of carbon
dioxide and 27.0 grams of water. What is a possible molecular formula of the hydrocarbon?
45. How many moles of NaF must be dissolved in 1.00 liter of a saturated solution of PbF2 at 25 °C to reduce the
[Pb2+] to 1 x 10¯6 molar? (Ksp of PbF2 at 25 °C = 4.0 x 10¯8)
46. Consider the equilibrium: HgO(s) + 4 I¯ + H2O <==> HgI42¯ + 2 OH¯; ΔH < 0. What is the effect of each of the
following changes on the concentration of HgI42¯?
(A) Increasing the concentration of OH¯
______________________________________________
(B) Adding 6 M HNO3
______________________________________________
(C) Increasing the mass of HgO present
______________________________________________
(D) Increasing the temperature
______________________________________________
(E) Adding a catalyst
______________________________________________
47. In a titration experiment based on the equation 5 Fe2+ + MnO4¯ + 8 H+ <==> 5 Fe3+ + Mn2+ + 4 H2O, 25.0mL of
an acidified Fe2+ solution requires 14.0 milliliters of standard 0.050-molar MnO4¯ solution to reach the
equivalence point. The concentration of Fe2+ in the original solution is
48. Draw structures for each of the following, including resonance structures necessary to describe the bonding
satisfactorily.
(A) H2S
(B) SO2
(C) CO2
(D) OF2
(E) PF3
49. Organize the following aqueous solutions from the highest boiling point to the lowest boiling point.
(A) 0.10 M potassium sulfate
(B) 0.10 M hydrochloric acid
(C) 0.10 M ammonium nitrate
(D) 0.10 M magnesium sulfate
(E) 0.20 M sucrose
________________
________________
________________
________________
________________
50. Calculate the volume of hydrogen gas produced at standard temperature and pressure when 2.00 grams of
aluminum metal is added to an excess of hydrochloric acid.
©T. M. Chipi, chemistry instructor
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Review of AP Chemistry topics: 2nd year course in chemistry, first semester
Equilibrium:
Equilibrium tables
Kc, Kp, ksp, keq
Reaction Quotient (Qc, Qp, Qip)
Le Chatelier’s principle
Solutions:
Molarity, molality, % by mass, % by volume, mass/volume %
Electrolytes (strong and weak) and non-electrolytes
Mole fraction
Colligative properties (vapor pressure, osmotic pressure, freezing point depression and
boiling point elevation)
van’t Hoff factor
Ideal and non-ideal solutions
Solubility, solubility of gases vs. solubility of solids, solubility curves
Henry’s Law (solubility vs. pressure); S1P1 = S2P2
Thermochemistry:
Enthalpy
Heat of formation
Entropy (definition): solid < liquid < gas
First Law of Thermo
Review of Chemistry 1 topics: 1st year course in chemistry
Atomic structure
Ionization energy
3- -1
+1
3+
Isoelectronic series [N -F , Ne, Na -Al ]
Electron affinity
Electron configuration (s, p, d, f)
Electronegativity
Orbital filling
Ionic and atomic radii
Paramagnetism
Mole concept
Chemicals and their use
Dimensional analysis
Bonding: ionic & covalent, σ & π
Empirical formula
Dipole moment
Colors of ppt (solids)
Hydrogen bonding (H-F, H-N, H-O)
Bond length/strength
Types of crystalline solids
Net-ionic equations
Molecular geometry
Neutralization based on
% composition
stoichiometric coefficients
Kinetic Molecular Theory (gases)
Conservation of mass
Density-Molar Mass-ideal gas equation
Solubility
Names & formulas
Chemical reactions (redox, etc.)
Isotopes
Colors of ions in solution
Charge of ion in complex ion
Naming/formula complex ion
Periodicity
Normal boiling point
Temperature = Average KE of particles
Density from PV = nRT
Quantum numbers
Resonance structures
Stoichiometry and general problem solving Organic compounds
Phase diagrams: critical point, triple point, critical temperature, phase change
Behavior of gases: temperature, pressure, volume, amount or concentration
Nuclear Chemistry (alpha, beta, positron, e- capture)
Lab procedures (titration, making molar soln, collection of gas over water)
©T. M. Chipi, chemistry instructor
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