Name:______Date:______Period:____
REVIEW Unit 3Test (Chp 1-3): Matter, Measurement, & Stoichiometry
Multiple Choice (20 questions)(50% of score)
states of matter (solid, liquid, gas)
chemical vs. physical changes and properties
density
atomic structure (protons, neutrons, electrons, atomic #, mass #, isotopes, average atomic mass)
periodic table (metals, nonmetals, metalloids, groups, periods, etc.)
ionic vs. molecular compounds withwriting names formulas
polyatomic ions their charges
oxidation numbers of elements in compounds, like what is the ox. # of C in oxalate ion C2O42– ?
stoichiometry (limiting reactant, theoretical yield, and percent yield)
Free Response(32 points)(5 questions: 1 short, 3 medium, 1 long)(50% of score)
1)describe isotopes, calculate average atomic mass, convert atoms/grams of an isotope (6 pts)
2)lab question determining formula of a hydrate (5 pts) (study Lab handout & Chp 3 #53)
3)combustion analysis and empirical/molecular formula (7 pts)
4)write 3 reactions from verbal descriptions and balance them (some diatomics) (9 pts)
5)limiting reactantstoichiometric conversions (with mol to mol ratios) (5 pts)
Section I Multiple Choice
NOCALCULATOR
___1.A measured mass of an unreactive metal was dropped into a small graduated cylinderhalf filled with water. The following measurements were made.
Mass of metal = 19.611 grams
Volume of water before addition of metal = 12.4 milliliters
Volume of water after addition of metal = 14.9 milliliters
The density of the metal should be reported as
(A) 7.8444 grams per mL
(B) 7.844 grams per mL
(C) 7.84 grams per mL
(D) 7.8 grams per mL
___2.What is the oxidation number (charge) of chromium in the dichromate ion, Cr2O72– ?
(A) 2–
(B) 6+
(C) 7+
(D) 12+
___3.What number of moles of O2 is needed to produce 14.2 grams of P4O10 from P? (Molecular weight P4O10 = 284)
(A) 0.0500 mole
(B) 0.250 mole
(C) 0.0625 mole
(D) 0.500 mole
(E) 0.125 mole
10 HI + 2 KMnO4 + 3 H2SO4 5 I2 + 2 MnSO4 + K2SO4 + 8 H2O
___4.According to the balanced equation above, how many moles of HI would be necessary to produce 2.5 mol of I2, starting with 4.0 mol of KMnO4 and 3.0 mol of H2SO4?
(A) 20.
(B) 10.
(C) 8.0
(D) 5.0
(E) 2.5
Section II Free Response
CALCULATOR ALLOWED
CLEARLY SHOW THE METHODS USED AND STEPS INVOLVED IN YOUR ANSWERS.
It is to your advantage to do this, because you may earn partial credit if you do and little or no credit if you do not. Attention should be paid to significant digits.
- Analysis of naturally occurring ironby a mass spectrometer results in four peaks at 54, 56, 57, and 58. These four stable isotopes of iron are listed in the table below with Fe–56 as the reference isotope.
(a)Differentiate between the terms “element”, “atom”, and “isotope”. (2)
(b)Calculate the average atomic mass of iron. (1)
(c)How many grams of naturally occurring iron would contain 1.807 x 1024 atoms of Iron-57 ? (3)
2. Answer the following questions about BeC2O4(s) and its hydrate.
(a)Calculate the mass percent of carbon in the hydrated form of the solid that has the formula BeC2O4•3H2O. (1)
(b)When heated to 220.oC, BeC2O4•3H2O(s) dehydrates completely as represented below.
BeC2O4•3H2O(s) BeC2O4(s) + 3 H2O(g)
If 3.21 g of BeC2O4•3H2O(s) is heated to 220.oC, calculatethe mass of BeC2O4(s) formed. (1)
(c)Calculatethe moles of H2O(g) formed. (1)
(d) When a different student heated the hydrate, the following data were obtained.
Mass of empty container / 32.155 gInitial mass of sample and container / 36.382 g
Mass of sample and container after first heating / 34.011 g
Mass of sample and container after second heating / 33.962 g
The student then calculated the formula of the hydrateto be BeC2O4•2H2O.
Explain why the student’s calculations resulted in a formula with too few waters of hydrate.
Your explanation should make reference to the data. (2)
- Answer the following questions about a pure compound that contains only carbon, hydrogen, and oxygen.
(a) A 0.7549 g sample of the compound burns in O2(g) to produce 1.9061 g of CO2(g) and 0.3370 g of H2O(g).
(i)Calculate the individual masses of C, H, and O in the 0.7549 g sample. (3)
(ii)Determine the empirical formula for the compound. (2)
(iii)Determine the molecular formula for the compound if the molecular weight is 244 g∙mol–1. (1)
4.For each of the following three reactions, in part (i) write a balanced equation for the reaction. In part (i), coefficients should be in terms of lowest whole numbers. Assume that solutions are aqueous unless otherwise indicated. You may use the empty space at the bottom of the next page for scratch work, but only equations that are written in the answer boxes provided will be graded.
(a)Solid potassium chlorate is heated strongly forming potassium chloride and oxygen gas.
(i)Balanced equation: (3)
(b)Solid chromium is oxidized by potassium permangate to form potassium chromate and manganese(IV) oxide.
(i)Balanced equation: (3)
(c)Sulfur trioxide gasis bubbled through water producingsulfuric acid.
(i)Balanced equation: (3)
5. A 0.150 g sample of solid lead(II) nitrate is added to 0.905 g of sodium iodide in solution. Assume nochange in volume of the solution. The chemical reaction that takes place is represented by the following equation.
Pb(NO3)2(aq) + 2 NaI(aq) PbI2(s) + 2 NaNO3(aq)
(a)List an appropriate observation that provides evidence of a chemical reaction between the two compounds. (1)
(b)Calculate the number of moles of each reactant. (2)
(c)Identify the limiting reactant. Show calculations to support your identification. (2)
(d)What is the theoretical yield of PbI2(s)? (1)
(e)If 0.174 g of PbI2(s) was obtained by filtration and weighing, what is the percent yield of PbI2(s) for this reaction? (1)
Answer KEY
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3.2005#2
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5.2008B#3(a)(b)(c)
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