MC_Chem1A_Sp08_LecExamFinal

Multiple Choice
Identify the choice that best completes the statement or answers the question.

____	1.	You can identify a metal by carefully determining its density. A 8.44 g piece of an unknown metal is 1.25 cm long, 2.50 cm wide, and 1.00 cm thick. What is a possible identity of the element?
a.	nickel, 8.90 g/cm3	
b.	aluminum, 2.70 g/cm3	
c.	titanium, 4.54 g/cm3	
d.	zirconium, 6.51 g/cm3	
e.	chromium, 7.20 g/cm3	

____	2.	What is the correct answer to the following expression: (72.61 - 68.59) ´ 18.76? Carry out the subtraction operation first.
a.	7 ´ 101	
b.	75	
c.	75.4	
d.	74.542	
e.	75.5415	

____	3.	Express 0.38300 in exponential notation.
a.	3.8300 ´ 10-1	
b.	3.83 ´ 10-1	
c.	3.830 ´ 10-2	
d.	3.83 ´ 101	
e.	3.8300 ´ 101	

____	4.	The reaction of 10.0 g H2(g) with 10.0 g O2(g) yields 8.43 g H2O(g). What is the percent yield of this reaction?
a.	9.43%	
b.	27.3%	
c.	42.2%	
d.	66.8%	
e.	74.9%	

____	5.	Bromine has two naturally occurring isotopes. The average mass of bromine is 79.904 u. If 50.54% of bromine is found as bromine-79 (78.9183 u), what is the mass of the other isotope?
a.	79.82 u	
b.	79.97 u	
c.	80.91 u	
d.	81.93 u	
e.	82.91 u	

____	6.	All of the following statements are true EXCEPT
a.	for any neutral element, the number of protons and electrons are equal.	
b.	electrons and protons have equal mass, but opposite charges.	
c.	the mass number is the sum of the number of protons and neutrons.	
d.	the atomic number equals the number of protons.	
e.	isotopes of an element have identical atomic numbers.	

____	7.	The density of lithium is 0.546 g/cm3. What volume is occupied by 1.96 ´ 1023 atoms of lithium?
a.	0.0859 cm3	
b.	0.596 cm3	
c.	4.14 cm3	
d.	5.63 cm3	
e.	39.0 cm3	

____	8.	Which of the following formulas is not correct?
a.	Al2(SO4)3	
b.	NaClO3	
c.	Ba2O3	
d.	Mg(NO3)2	
e.	KH2PO4	

____	9.	Which of the following compounds will be soluble in water: KNO3, Ca3(PO4)2, CuCl2, and Fe2S3?
a.	KNO3 and Ca3(PO4)2	
b.	KNO3 and CuCl2	
c.	Ca3(PO4)2 and Fe2S3	
d.	CuCl2 and Fe2S3	
e.	KNO3, Ca3(PO4)2, and Fe2S3	

____	10.	Which of the following species are strong acids: HNO3, HBr, H3PO4,CH3CO2H, and HSO4-?
a.	HNO3 and HBr	
b.	HNO3 and H3PO4	
c.	HBr and HSO4-	
d.	H3PO4, CH3CO2H, and HSO4-	
e.	HNO3, H3PO4, and HSO4-	

____	11.	What is the oxidation number of each atom in potassium nitrate, KNO3?
a.	K = +1, N = -3, O = -2	
b.	K = +1, N = +5, O = -2	
c.	K = +1, N = -3, O = +2	
d.	K = -1, N = +3, O = -2	
e.	K = 0, N = 0, O = 0	

____	12.	The pH of an aqueous KOH solution is 12.38. What is the hydrogen ion concentration of this solution?
a.	8.1 ´ 10-14 M	
b.	4.2 ´ 10-13 M	
c.	8.1 ´ 10-6 M	
d.	0.21 M	
e.	1.1 M	

____	13.	If 5.00 mL of 0.314 M KOH is diluted to exactly 125 mL with water, what is the concentration of the resulting solution?
a.	5.02 ´ 10-4 M	
b.	1.26 ´ 10-2 M	
c.	0.127 M	
d.	0.281 M	
e.	7.85 M	

____	14.	Hydrazine, N2H4, is a liquid used as a rocket fuel. It reacts with oxygen to yield nitrogen gas and water.

N2H4(l) + O2(g) ® N2(g) + 2 H2O(l)

The reaction of 3.80 g N2H4 evolves 73.7 kJ of heat. Calculate the enthalpy change per mole of hydrazine combusted.
a.	-8.74 kJ/mol	
b.	-19.4 kJ/mol	
c.	-2.80 ´ 102 kJ/mol	
d.	-622 kJ/mol	
e.	-8.98 ´ 103 kJ/mol	

____	15.	If the energy of 1.00 mole of photons is 441 kJ, what is the wavelength of the light?
a.	271 nm	
b.	451 nm	
c.	680 nm	
d.	732 nm	
e.	891 nm	

____	16.	What type of orbital is designated n = 4, l = 3, ml = -3?
a.	4s	
b.	4p	
c.	4d	
d.	4f	
e.	none	

____	17.	Which of the following orbitals might have ml = -2?
a.	s	
b.	s and p	
c.	p and d	
d.	d and f	
e.	p, d, and f	

____	18.	How many electrons can be described by the quantum numbers n = 7, l = 1?
a.	0	
b.	2	
c.	6	
d.	7	
e.	14	

____	19.	What is a possible set of quantum numbers for the unpaired electron in the orbital box diagram below?

a.	n = 3, l = 1, ml = -1, ms = +1/2	
b.	n = 3, l = 2, ml = -1, ms = -1/2	
c.	n = 3, l = 2, ml = -2, ms = +1/2	
d.	n = 4, l = 0, ml = 0, ms = +1/2	
e.	n = 4, l = 1, ml = -1, ms = +1/2	

____	20.	Which one of the following statements is INCORRECT?
a.	Ionization energy is always a positive value.	
b.	Ionization energy is the energy required to remove an electron from a gaseous atom.	
c.	For any element, the second ionization energy is larger than the first ionization energy.	
d.	Ionization energy decreases across a periodic of the periodic table.	
e.	Ionization energy decreases down a group of the periodic table	

____	21.	Which of the following are resonance structures for the formate ion, HCO2-?

a.	1 and 2	
b.	2 and 3	
c.	3 and 4	
d.	1, 3, and 4	
e.	2, 3, and 4	

____	22.	To prepare 1.000 kg of 5.5% by mass KCl(aq), one may dissolve
a.	994.5 g KCl in 5.5 g of water.	
b.	5.5 g KCl in 994.5 g of water.	
c.	5.5 g KCl in 1000.0 g of water.	
d.	55 g KCl in 945 g of water.	
e.	55.0 g KCl in 1000.0 g of water.	

____	23.	What is the mole fraction of 1.98 m Fe(NO3)3(aq)? The molar mass of Fe(NO3)3 is 241.9 g/mol and the molar mass of water is 18.02 g/mol.
a.	0.0345	
b.	0.0641	
c.	0.324	
d.	0.479	
e.	0.863	

____	24.	The vapor pressure of pure water at 45 °C is 71.9 mm Hg. What is the vapor pressure of a mixture of 21.0 g sucrose (C12H22O11, molar mass 342.3 g/mol) and 79.0 g water?
a.	0.992 mm Hg	
b.	15.1 mm Hg	
c.	19.1 mm Hg	
d.	55.8 mm Hg	
e.	70.9 mm Hg	

____	25.	Which of the following aqueous solutions should have the lowest freezing point?
a.	pure H2O	
b.	1 m NaBr	
c.	1 m Ca(NO3)2	
d.	1 m NH3	
e.	1 m C6H12O6	

____	26.	The freezing point depression constant for water is -1.86 °C/m. At what temperature will a solution containing 7.99 g BaCl2 and 52.0 g H2O begin to freeze? Assume that no ion-pairing occurs between Ba2+ and Cl-.
a.	-4.12 °C	
b.	-2.75 °C	
c.	-1.37 °C	
d.	-0.857 °C	
e.	-0.248 °C	

____	27.	The osmotic pressure of blood is 7.65 atm at 37 °C. What mass of glucose (C6H12O6, molar mass = 180.2 g/mol) is needed to prepare 5.00 L of solution for intravenous injection? The osmotic pressure of the glucose solution must equal the osmotic pressure of blood. (R = 0.08206 L·atm/mol·K)
a.	1.50 g	
b.	54.2 g	
c.	126 g	
d.	271 g	
e.	2270 g	

____	28.	What mass of He contains the same number of atoms as 5.0 g Kr?
a.	0.24 g	
b.	0.80 g	
c.	1.2 g	
d.	5.0 g	
e.	1.0 ´ 102 g	

____	29.	All of the following are named correctly EXCEPT
a.	LiClO4; lithium perchlorate.	
b.	CaHPO4; calcium hydrogen phosphide.	
c.	NaCN; sodium cyanide.	
d.	Mg(OH)2; magnesium hydroxide.	
e.	CaSO3; calcium sulfite.	

____	30.	What is the mass percent of each element in dichloromethane, CH2Cl2?
a.	10.06% C, 60.24% H, 29.70% Cl	
b.	20.00% C, 20.00% H, 60.00% Cl	
c.	24.10% C, 3.11% H, 72.79% Cl	
d.	33.87% C, 0.22% H, 65.91% Cl	
e.	14.14% C, 2.37% H, 83.48% Cl	

____	31.	The compound P4S3 is used in matches. It reacts with oxygen to produce P4O10 and SO2. The unbalanced chemical equation is shown below.

P4S3(s) + O2(g) ® P4O10(s) + SO2(g)

What mass of O2 will react with 0.641 grams of P4S3?
a.	0.0466 g	
b.	0.0932 g	
c.	0.186 g	
d.	0.372 g	
e.	0.746 g	

____	32.	What is the net ionic equation for the reaction of aqueous lead(II) nitrate with aqueous sodium bromide?
a.	Pb(NO3)2(aq) + 2 NaBr(aq) ® PbBr2(aq) + 2 NaNO3(s)	
b.	Na+(aq) + NO3-(aq) ® NaNO3(s)	
c.	Pb2+(aq) + 2 Br-(aq) ® PbBr2(s)	
d.	Pb2+(aq) + 2 Na+(aq) ® PbNa2(s)	
e.	Pb(NO3)2(aq) + 2 NaBr(aq) ® PbBr2(s) + 2 NaNO3(aq)	

____	33.	If 50.0 g H2O at 13.6 °C is combined with 85.0 g H2O at 93.7 °C, what is the final temperature of the mixture? The specific heat capacity of water is 4.184 J/g·K.
a.	26.1 °C	
b.	40.0 °C	
c.	56.1 °C	
d.	64.0 °C	
e.	80.1 °C	

____	34.	Which of the following sets of quantum numbers refers to a 4p orbital?
a.	n = 1, l = 1, ml = -1	
b.	n = 1, l = 2, ml = -1	
c.	n = 4, l = 1, ml = 0	
d.	n = 4, l = 2, ml = -1	
e.	n = 4, l = 3, ml = +2	

____	35.	What 1+ ion has the following ground state electron configuration?

a.	K+	
b.	Rb+	
c.	In+	
d.	Tl+	
e.	Y+	

____	36.	Use VSEPR theory to predict the molecular geometry of CS2.
a.	linear	
b.	bent	
c.	trigonal-planar	
d.	tetrahedral	
e.	octahedral	

____	37.	What volume of CO2 at 15 °C and 1.50 atm contains the same number of molecules as 0.410 L of O2 at 35 °C and 3.00 atm?
a.	0.192 L	
b.	0.219 L	
c.	0.767 L	
d.	0.820 L	
e.	1.14 L	

____	38.	If a gas effuses 1.618 times faster than Kr, what is its molar mass?
a.	18.02 g/mol	
b.	28.01 g/mol	
c.	32.01 g/mol	
d.	51.79 g/mol	
e.	65.88 g/mol	

____	39.	Ammonia, NH3, is used as a refrigerant. At its boiling point of -33 °C, the enthalpy of vaporization of ammonia is 23.3 kJ/mol. How much heat is required to vaporize 355 g of ammonia at -33 °C?
a.	1.12 kJ	
b.	152 kJ	
c.	251 kJ	
d.	486 kJ	
e.	8.27 ´ 103 kJ	

____	40.	Which of the following phase transitions is endothermic?
a.	gas to solid	
b.	liquid to gas	
c.	liquid to solid	
d.	gas to liquid	
e.	none of the above	

MC_Chem1A_Sp08_LecExamFinal
Answer Section

MULTIPLE CHOICE

	1.	ANS:	B	PTS:	1	OBJ:	1.4 Physical Properties

	2.	ANS:	C	PTS:	1	OBJ:	1.8 Mathematics of Chemistry

	3.	ANS:	A	PTS:	1	OBJ:	1.8 Mathematics of Chemistry

	4.	ANS:	E	PTS:	1	
OBJ:	4.4 Reactions in which One Reactant is Present in Limited Supply	

	5.	ANS:	C	PTS:	1	OBJ:	2.4 Atomic Weight	

	6.	ANS:	B	PTS:	1	OBJ:	2.3 Isotopes

	7.	ANS:	C	PTS:	1	OBJ:	2.5 Atoms and the Mole

	8.	ANS:	C	PTS:	1	OBJ:	3.3 Ionic Compounds: Formulas | Names | and Properties

	9.	ANS:	B	PTS:	1	OBJ:	5.1 Properties of Compounds in Aqueous Solution

	10.	ANS:	A	PTS:	1	OBJ:	5.3 Acids and Bases	

	11.	ANS:	B	PTS:	1	OBJ:	5.7 Oxidation-Reduction Reactions

	12.	ANS:	B	PTS:	1	OBJ:	5.9 pH | A Concentration Scale for Acids and Bases

	13.	ANS:	B	PTS:	1	OBJ:	5.8 Measuring Concentrations of Compounds in Solution

	14.	ANS:	D	PTS:	1	OBJ:	6.5 Enthalpy Changes for Chemical Reactions

	15.	ANS:	A	PTS:	1	OBJ:	7.2 Planck | Einstein | Energy | and Photons

	16.	ANS:	D	PTS:	1	OBJ:	7.5 Quantum Mechanical View of the Atom

	17.	ANS:	D	PTS:	1	OBJ:	7.5 Quantum Mechanical View of the Atom

	18.	ANS:	C	PTS:	1	OBJ:	8.2 The Pauli Exclusion Principle

	19.	ANS:	E	PTS:	1	OBJ:	8.4 Atomic Electron Configurations

	20.	ANS:	D	PTS:	1	OBJ:	8.6 Atomic Properties and Periodic Trends

	21.	ANS:	C	PTS:	1	OBJ:	9.5 Resonance		

	22.	ANS:	D	PTS:	1	OBJ:	14.1 Units of Concentration

	23.	ANS:	A	PTS:	1	OBJ:	14.1 Units of Concentration

	24.	ANS:	E	PTS:	1	OBJ:	14.4 Colligative Properties

	25.	ANS:	C	PTS:	1	OBJ:	14.4 Colligative Properties

	26.	ANS:	A	PTS:	1	OBJ:	14.4 Colligative Properties

	27.	ANS:	D	PTS:	1	OBJ:	14.4 Colligative Properties

	28.	ANS:	A	PTS:	1	OBJ:	2.5 Atoms and the Mole

	29.	ANS:	B	PTS:	1	OBJ:	3.3 Ionic Compounds: Formulas | Names | and Properties

	30.	ANS:	E	PTS:	1	OBJ:	3.6 Describing Compound Formulas

	31.	ANS:	E	PTS:	1	
OBJ:	4.3 Mass Relationships in Chemical Reactions: Stoichiometry	

	32.	ANS:	C	PTS:	1	OBJ:	5.2 Precipitation Reactions

	33.	ANS:	D	PTS:	1	OBJ:	6.2 Specific Heat Capacity and Heat Transfer

	34.	ANS:	C	PTS:	1	OBJ:	7.5 Quantum Mechanical View of the Atom

	35.	ANS:	C	PTS:	1	OBJ:	8.5 Electron Configurations of Ions

	36.	ANS:	A	PTS:	1	OBJ:	9.7 Molecular Shapes	

	37.	ANS:	C	PTS:	1	OBJ:	12.2 Gas Laws: The Experimental Basis

	38.	ANS:	C	PTS:	1	OBJ:	12.7 Effusion and Diffusion

	39.	ANS:	D	PTS:	1	OBJ:	13.5 Properties of Liquids

	40.	ANS:	B	PTS:	1	OBJ:	13.5 Properties of Liquids