2015-2016 AP CHEMISTRY MIDTERM EXAM Review The midterm exam follows the format of the AP Chemistry exam. The actual AP Chemistry exam consists of 2 sections: Section 1 is 60 multiple choice in a time of 90 minutes (50% of grade) Section 2 is 3 long free-response and 4 short free-response in 90 minutes (50% of grade) Due to the shortened time, the midterm is modified for a total time of 2 hours: Section 1 is 46 multiple choice (50% of grade) Section 2 is 3 long free-response and 2 short free-response (50% of grade) Multiple Choice: You will be given the periodic table and formula sheet, but may NOT your calculator. Free Response: You may use your calculator for this part
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2015-2016 AP CHEMISTRY MIDTERM EXAM Review2015-2016 AP CHEMISTRY MIDTERM EXAM Review The midterm exam follows the format of the AP Chemistry exam. The actual AP Chemistry exam consists
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2015-2016
AP CHEMISTRY
MIDTERM EXAM Review
The midterm exam follows the format of the AP Chemistry exam.
The actual AP Chemistry exam consists of 2 sections:
Section 1 is 60 multiple choice in a time of 90 minutes (50% of grade)
Section 2 is 3 long free-response and 4 short free-response in 90 minutes (50% of grade)
Due to the shortened time, the midterm is modified for a total time of 2 hours:
Section 1 is 46 multiple choice (50% of grade)
Section 2 is 3 long free-response and 2 short free-response (50% of grade)
Multiple Choice: You will be given the periodic table and formula sheet, but may NOT your calculator.
Free Response: You may use your calculator for this part
SECTION I: Multiple Choice
Please mark the best choice for each question.
Note: For all questions involving solutions and/or chemical equations, assume that the system is in pure water at room tem-
perature unless otherwise noted.
1. A combination of sand, salt, and water is an
example of a ___________________.
a. Homogeneous mixture d. Pure substance b. Heterogeneous mixture e. Solid c. Compound
2. Which state or states of matter are significantly
compressible?
a. gases d. liquids and gases b. liquids e. solids and liquids c. solids
3. If matter is uniform throughout and cannot be
separated into other substances by physical
means, it is _____.
a. a compound d. a heterogeneous mixture b. either an element or a
compound e. an element
c. a homogenous mixture
4. Of the following, only _______________ is a
chemical reaction.
a. Melting of lead d. Crushing of Stone b. Dissolving sugar in water e. Placing a penny in water c. Tarnishing of Silver
5. Of the following, _____ is smallest in mass.
a. 25 kg d. 2.5 x 109 fg
b. 2.5 x 10-2
mg e. 2.5 x 1010
ng c. 2.5 x 10
15 pg
6. Precision refers to _______________.
a. How close a measured value is to other measured numbers b. How close a measured number is to the true value c. How close a measured number is to the calculated value d. How close a measured number is to zero e. How close a measured number is to the determined mean
7. A scientific __________ is a concise statement or
an equation that summarizes a broad variety of
observations. It has been extensively tested.
a. law d. trend b. hypothesis e. pattern c. theory
8. Which one of the following is NOT one of the
postulates of Dalton’s atomic theory?
a. Each element is composed of tiny, indivisible particles
called atoms b. All atoms of a given element are identical to each other
and different from those of other elements c. During a chemical reaction, atoms are changed into atoms
of different elements d. Compounds are formed when atoms of different elements
combine e. Atoms of an element are not changed into different types
of atoms by chemical reactions
9. The charge on an electron was determined in the
______.
a. cathode ray tube by JJ Thompson b. Rutherford gold foil experiment c. Millikan oil drop experiment d. Dalton atomic theory e. radioactive decay of Carbon by James Chadwick
10. Which combination of protons, neutrons, and
electrons is correct for the 63
29Cu isotope of
Copper?
a. 29 protons, 34 neutrons, and 29 electrons b. 29 protons, 29 neutrons, and 63 electrons c. 63 protons, 29 neutrons, and 63 electrons d. 34 protons, 29 neutrons, and 34 electrons e. 34 protons, 34 neutrons, and 29 electrons
11. In the Rutherford nuclear-atom model,
a. The heavy subatomic particles, protons and neutrons,
reside in the nucleus b. The three principle subatomic particles (protons, neutrons,
and electrons,) all have essentially the same mass. c. The light subatomic particles, protons and neutrons, reside
in the nucleus d. Mass is spread essentially uniformly throughout the atom e. The three principle subatomic particles(protons, neutrons,
and electrons) all have essentially the same mass and mass
is spread essentially uniformly throughout the atom.
12. In the periodic table, the elements are arranged in
_____________.
a. Alphabetical order b. Order of increasing atomic number c. Order of decreasing metallic properties d. Order of increasing neutron content e. Chemically families based on bit physical and chemical
properties
13. Which pair of elements would you expect to
exhibit the greatest similarities in the physical and
chemical properties?
a. Oxygen and Sulfur d. Hydrogen and Helium b. Carbon and Nitrogen e. Silicon and Phosphorous c. Potassium and Calcium
14. The elements in groups 1A, 6A, and 7A are called
_______________, respectively.
a. Alkaline earth metals, halogens, and chalcogens b. Alkali metals, chalcogens, and halogens c. Alkali metals, halogens, and noble gasses d. Alkaline earth metals, transition metals, and halogens
15. Which one of the following does not occur as a
diatomic molecule in elemental form?
a. oxygen d. hydrogen b. nitrogen e. bromine c. sulfur
16. When a metal and a nonmetal react, the ________
tends to lose electrons and the _______ tends to
gain electrons.
a. Metal, metal d. Nonmetal, metal b. Nonmetal, nonmetal e. None of the above, these
elements share electrons c. Metal, nonmetal
17. The correct formula for molybdenum (IV)
hypochlorite is ____________.
a. Mo(ClO3)4 d. Mo(ClO4)4 b. Mo(ClO)4 e. MoCl4 c. Mo(ClO2)4
18. When the following equation is balanced, the
coefficients are ____________.
NH3 + O2 → NO2 + H2O
a. 1, 1, 1, 1 d. 1, 3, 1, 2 b. 4, 7, 4, 6 e. 4, 3, 4, 3 c. 2, 3, 2, 3
19. When the following equation is balanced, the
coefficient of Al is __________.
2 3 2Al (s) + H O (l) Al(OH) (s) + H (g)
a. 1 d. 4 b. 2 e. 5 c. 3
20. Predict the product in the combination
reaction below.
Al (s) + N2 (g) → ____________
21. Pentacarbonyliron (Fe(CO)5) reacts with
phosphorous trifluoride and hydrogen to
release carbon monoxide. The reaction of 5.0
mol of Fe(CO)5, 8.0 mol of PF3, and 6.0 mol of
H2 will release ________ mol of CO.
Fe(CO)5+2PF3+H2→Fe(CO)2(PF3)2H2+3CO
a. 15 d. 6.0 b. 5.0 e. 12 c. 24
22. When H2SO4 is neutralized by KOH in
aqueous solution, the net ionic equation is:
a. SO42-
+ 2K+ → K2SO4 (aq)
b. SO42-
+ 2K+ → K2SO4 (s)
c. H+ + OH
- → H2O (l)
d. H2SO4 + 2OH
- → 2H2O (l) + SO4
2-
e. 2H+ + 2KOH → 2H2O (l) + 2K
+
23. When aqueous solutions of ___________ are
mixed, a precipitate forms.
a. NiBr2 and AgNO3 d. KOH and Ba(NO3)2
b. NaI and KBr e. Li2CO3 and CsI
c. K2SO4 and CrCl3
24. Which one of the following is a correct
expression for molarity?
a. mol solute / L solvent d. mol solute / kg solvent
b. mol solute / mL Solvent e. μmol solute / L solution
c. mmol solute / mL
solution
25. How many moles of Co2+
are present in
0.200L of a 0.400 M solution of CoI2
a. 2.00 d. 0.0800
b. 0.500 e. 0.0400
c. 0.160
26. There are _____ paired and _______ unpaired
electrons in the Lewis Symbol for a
phosphorous atom. a. 4, 2 d. 4, 3
b. 2, 4 e. 3, 2
c. 2, 3
27. The halogens, alkali metals, and alkaline
earth metals have ____ valence electrons,
respectively.
a. 2, 4, and 6 d. 7, 1, and 2
b. 1, 5, and 7 e. 2, 7, and 4
c. 8, 2, and 3
28. Elements from opposite sides of the periodic
table tend to form ___________.
a. Covalent compounds
b. Ionic compounds
c. Compounds that are gases at room temperature
d. Diatomic compounds
e. Metallic compounds
Questions 29-32
(A) Heisenberg uncertainty principle
(B) Pauli exclusion principle
(C) Hund's rule (principle of maximum multiplicity)
(D) Shielding effect
(E) Wave nature of matter
29. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic
30. Explains the experimental phenomenon of electron diffraction
31. Indicates that an atomic orbital can hold no more than two electrons
32. Predicts that it is impossible to determine simultaneously the exact position and the exact velocity of an electron
Questions 33-35 refer to the phase diagram below of a pure substance.
(A) Sublimation
(B) Condensation
(C) Solvation
(D) Fusion
(E) Freezing
33. If the temperature increases from 10° C to 60° C at a constant pressure of 0.4 atmosphere, which of the processes occurs?
34. If the temperature decreases from 110° C to 40° C at a constant pressure of 1.1 atmospheres, which of the processes occurs?
35. If the pressure increases from 0.5 to 1.5 atmospheres at a constant temperature of 50° C, which of the processes occurs?
Questions 36-38 refer to the following diatomic species.
(A) Li2
(B) B2
(C) N2
(D) O2
(E) F2
36. Has the largest bond-dissociation energy
37. Has a bond order of 2
38. Contains 1 sigma (σ) and 2 pi (π) bonds
39. In a molecule in which the central atom exhibits sp3d
2 hybrid orbitals, the electron pairs are directed toward the corners of
(A) a tetrahedron
(B) a square-based pyramid
(C) a trigonal bipyramid
(D) a square
(E) an octahedron
40. Which of the following sets of quantum numbers (n, l, ml, ms) best describes the valence electron of highest energy in a
ground-state gallium atom (atomic number 31) ?
(A) 4, 0, 0, 1/2
(B) 4, 0, 1, 1/2
(C) 4, 1, 1, 1/2
(D) 4, 1, 2, 1/2
(E) 4, 2, 0, 1/2
41. CH3CH2OH boils at 78 °C and CH3OCH3 boils at - 24 °C, although both compounds have the same composition. This differ-
ence in boiling points may be attributed to a difference in
(A) molecular mass
(B) density
(C) specific heat
(D) hydrogen bonding
(E) heat of combustion
42. Based on concepts of polarity and hydrogen bonding, which of the following sequences correctly lists the compounds below
in the order of their increasing solubility in water?
X = CH3-CH2-CH2-CH2-CH3
Y = CH3-CH2-CH2-CH2-OH
Z = HO-CH2-CH2-CH2-OH
(A) Z < Y < X
(B) Y < Z < X
(C) Y < X < Z
(D) X < Z < Y
(E) X < Y < Z
43. Concentrations of colored substances are commonly measured by means of a spectrophotometer. Which of the following
would ensure that correct values are obtained for the measured absorbance?
I. There must be enough sample in the tube to cover the entire light path.
II. The instrument must be periodically reset using a standard.
III. The solution must be saturated.
(A) I only
(B) II only
(C) I and II only
(D) II and III only
(E) I, II, and III
44. The data below were gathered in order to determine the density of an unknown solid. The density of the sample should be re-
ported as
Mass of an empty container = 3.0 grams
Mass of the container plus the solid sample = 25.0 grams
Volume of the solid sample = 11.0 cubic centimeters
(A) 0.5 g/cm3
(B) 0.50 g/cm3
(C) 2.0 g/cm3
(D) 2.00 g/cm3
(E) 2.27 g/cm3
45. Which of the following pairs of compounds are isomers?
46. It is suggested that SO2 (molar mass 64 grams), which contributes to acid rain, could be removed from a stream of waste gas-
es by bubbling the gases through 0.25-molar KOH, thereby producing K2SO3. What is the maximum mass of SO2 that could
be removed by 1,000. liters of the KOH solution?
(A) 4.0 kg
(B) 8.0 kg
(C) 16 kg
(D) 20. kg
(E) 40. kg
47. I2(g) + 3 Cl2(g) 2 ICl3(g)
According to the data in the table below, what is the value of ∆H° for the reaction represented above?
Bond Average Bond Energy
(kilojoules / mole)
I---I 150
Cl---Cl 240
I---Cl 210
(A) - 870 kJ
(B) - 390 kJ
(C) + 180 kJ
(D) + 450 kJ
(E) + 1,260 kJ
48. 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
49. Which of the following molecules has a dipole moment of zero?