Download Chem 171 Review - Exam 1

Chem 171
Review - Exam 1
The following is a list of concepts and terms with which you should be well-acquainted. This is not a substitute for your notes.
This list of review topics is not meant to be exhaustive, it is only meant to help you identify the major areas we have discussed.
You are responsible for all material covered in lecture and in the text book.
Chapter 1: Chemistry and Measurement
•
Chemistry and the scientific process
•
SI units
base units: mass, length, time, temperature, amount, current, luminous intensity
vs.
derived units: derived from the base units; ex: density (mass / vol), vol (length • width • height), velocity (length / time)
•
metric prefixes and multiples; use in unit conversions: mega-, kilo-, deci-, centi-, milli-, micro-, nano-, picoex: 1 pm = 1012 m; 1 mg = 103 g
•
conversion factors and dimensional analysis
you should be comfortable with intrasystem conversions (metric to metric, English to English), intersystem
conversions (metric to English and vice versa), and combined unit conversions
Remember: when you raise a unit to some power, you must raise the entire conversion factor to that power
ex: (1 cm)3 = (10−2 m)3 therefore, 1 cm3 = 10−6 m3
•
uncertainty in measurements – indicated by the number of recorded significant figures
significant figures: all certain digits plus the 1st uncertain digit
given a number, determine how many significant figures it has
understand when zeros are and are not significant
determination of the number of significant figures to record when making a measurement with an instrument
rules for significant figures in the answer to calculations based on measured values; rule for addition and
subtraction is not the same as the rule for multiplication and division
•
chemical vs. physical properties and changes
•
intensive vs. extensive properties
•
classification of matter: elements, compounds (fixed composition), and mixtures (variable composition)
•
elements - you should know the name (properly spelled) and symbol for elements 1 – 56, W, Pt, Au, Hg, Pb, and Bi
Chapter 2: Atoms, Molecules, and Ions
•
atoms and atomic structure
early experiments and scientists that established our modern understanding of the atom and its structure
protons, electrons, neutrons - know the characteristics of these particles including relative masses and electrical charge
atomic structure - where do the various subatomic particles reside?
relative size and density of the atom and the nucleus
atomic # vs. mass # - what are they, and what information do they provide?
isotopes
•
the periodic table
arrangement
periods vs. groups (families)
recognize main group elements, alkali metals, alkaline earth metals, transition metals, halogens, noble (or inert) gases
metal / nonmetal / metalloid (or semimetal) distinctions
•
atomic mass
atomic mass unit scale for mass of single atom vs. gram unit scale
calculation of the average atomic mass of an element based on that element's isotopes and percent abundances
•
molecular compounds vs. ionic compounds
molecules vs. ions
cations vs. anions; ion formation; monatomic vs. polyatomic ions; p, n, and e count for a monatomic ion
•
NOMENCLATURE: I cannot stress this enough - you should know this stuff cold! Refer to the nomenclature handout.
you should be able to provide chemical formulas or IUPAC names for binary molecular compounds, ionic compounds
and acids
•
writing and balancing chemical equations; what do the stoichiometric coefficients represent?
Chapter 3: Calculations with Chemical Formulas and Equations
•
the mole and molar mass
What does a mole represent?
Avogadro's number, NA = 6.022 x 1023 mol1
molar mass of elements or compounds: mass (in g) per mole of an element or compound; units g/mol.
molar mass of compound = ∑molar mass of its elements
interpretation of chemical formula in terms of mol of atoms, or ions per mol of compound; compound stoichiometry
•
mass percent composition
•
chemical formula determination from mass % data and combustion analysis data
empirical formula vs. molecular formula; how are these related?
•
chemical reaction stoichiometry:
interpretation of stoichiometric coefficients to determine the mole ratios between reactants and products
using mole ratios from chemical equations as conversion factors
•
reaction yield (actual yield) vs. theortetical yield vs. percentage yield - what are they, how do you calculate them?
•
limiting reactants and reaction tables
completion of reaction table; mol reactant & product before reaction, change during reaction, remaining after reaction
determination of limiting reactants within a reaction
calculation of theoretical yield based on limiting reactant
identification of the excess reactant and how much excess remains after reaction
Chapter 4: Chemical Reactions
•
the nature of solutions
speciation and stoichiometry associated with the dissolution of ionic vs. molecular compounds in water
strong vs. weak vs. nonelectrolytes: what ware the distinctions? what types of compounds fall in each category?
solution compostion: molarity (M) or molar concentration = mol solute/L sol'n
calculations of molarity of solutions, molarity of ions in a solution
molarity as a conversion factor relating mol of solute and volume of solution
•
calculations that involve molar concentration of solutions:
dilution calculations, MiVi = MfVf, or you can think of it in terms of MconcVconc = MdilVdil
solution stoichiometry in chemical reactions; ex. volume of sol’n required for reaction with a given amount of reactant, or
to form a given amount of product
DO NOT USE M1V1 = M2V2 FOR REACTION STOICHIOMETRY CALCULATIONS!! You must demonstrate to me
that you understand that a chemical reaction is occurring, and that you understand the mole ratios of reactants and products
in that reaction.
The following are some of the more important conversion factors that we have discussed:
•
density offers you a relationship between mass and volume: density = mass / volume
•
Avogadro's number offers you a relationship between the number of particles and mol
NA = 6.022 x 1023 mol–1 (things per mole where things can be atoms, molecules, ions, formula units, etc.)
•
molar mass (of elements or compounds) offers you a relationship between mass of substance and mole quantity
molar mass = g (of an element or a compound) / mole
Note: you determine the mass of an individual atom, molecule, or formula unit by combining Avogadro’s number and
molar mass as conversion factors. Follow the conversion string: atom (molecule, ion, formula unit)  mol  mass in g
•
mole ratios from a BALANCED CHEMICAL EQUATION gives you a relationship between amounts of reactants being
consumed and amounts of products being formed in a given chemical reaction. This allows you to convert from moles of
reactant A to moles of reactant B, or moles of reactant A to moles of product C, etc.
•
molar concentration, or molarity offers a relationship between the number of moles of solute dissolved per L of solution
Review Problems:
These review problems are only representative of the kinds of problems with which you should feel comfortable. This is not
meant to be representative of the exam questions - in other words, DO NOT EXPECT THAT THE EXAM WILL LOOK
LIKE THIS!
1.
The density of a substance is 8.4 g/cm3. Express this density in SI units.
2.
Calculate the volume (in cm3) occupied by 25.0 g of aluminum. For aluminum, d = 2.70 g/cm3.
3.
Give the symbol and name for:
a.
the alkali metal in the same period as chlorine
b.
a halogen in the same period as magnesium
c.
the heaviest alkaline earth metal
d.
a noble gas in the same period as carbon
4.
How many significant figures are in each of the following numbers?
a.
1.5003
b.
0.0070
c.
5.7
e.
5 x 103
f.
5.0005
g.
22.9898
d.
h.
2.00 x 107
0.0040
5.
Perform the indicated calculations and express the answer to the correct number of significant figures. Use scientific
notation where appropriate.
a.
17.2 x 12.55 =_____
b.
(1.4 x 1.11) / 42.33 = _____
c.
18.33 x 0.0122 = _____
d.
25.7 - 25.25 = _____
e.
13.51 + 0.00459 = _____
f.
16.45 / (32.0 + 10) = _____
g.
15.0 − 0.0134 = _____
h.
7.18 x 103 / 1.51 x 104 = _____
6.
Perform the following unit conversions. Use scientific notation where appropriate.
a.
3.01 cm = _____ m = _____ mm = _____ nm
b.
50.0 cm3 = _____ dm3 = _____ mL = _____ L = _____ m3
c.
282 oz = _____ lb = _____ g = _____ mg = _____ µg
7.
Write the symbol (with mass and atomic number) for an atom with:
a.
6 protons and 6 neutrons
b.
13 protons and 14 neutrons
c.
19 protons and 20 neutrons
8.
How many protons, electrons and neutrons are in each of the following ions?
a. 23Na+
b. 81Brc. 39K+
d. 40Ca2+
e.
115
In3+
f. 127Te2-
9.
An element has two isotopes with masses of 62.9396 amu and 64.9278 amu. 30.83% of the atoms are the heavier
isotope. Calculate the atomic mass of this element and give its name and symbol.
10.
Naturally occurring rubidium is 72.17% 85Rb (atomic mass = 84.912 amu). The remaining atoms are 87Rb (atomic mass
= 86.909 amu). Calculate the average atomic mass of rubidium.
11.
Determine the molar mass of the following compounds:
a. phosphorus trichloride
b. potassium carbonate
c. nickel (II) phosphate
d. sodium thiosulfate Na2S2O3
12.
a.
b.
A 1.836 g sample of coal contains 1.5484 g of carbon. Calculate the mass percentage of carbon in the coal.
A fertilizer is advertised to contain 15.8% nitrogen by mass. Calculate the mass (in g) of N in 4.15 kg of fertilizer.
13.
a.
b.
Calculate the quantity (in mol) of HCl in 0.730 g HCl.
Calculate mol Na2SO4 in 284 g of the compound.
14.
Calculate the mass (in g) of each of the following:
a. 0.40 mol of CO2
b. 0.20 mol of K2SO4
c. 3.00 mol of NH3
15.
Calculate the mass (in g) of a single atom, molecule, ion, or formula unit of each of the following:
a. Br
b. BrF3
c. BrO4–
d. KBrO4
16.
Calculate the number of atoms, molecules, or formula units in each of the following samples:
a. 16.8424 g of magnesium
b. 13.65 g of aluminum oxide
c. 25.0 g of xenon hexafluoride
17.
Consider a 15.10 mL sample of ethanol (CH3CH2OH, d= 0.789 g/mL).
a.
Calculate the mass (in g) of this sample.
b.
Calculate the mol of ethanol in this sample.
c.
Calculate the number of ethanol molecules in this sample.
18.
Name the following elements and list their group and period numbers. Classify each as a metal or non-metal.
a. P
b. Sb
c. Ag
d. Zn
e. Ba
19.
Pewter is an alloy (metal mixture) with the following composition by mass: 6.00% antimony, 1.50% copper, 92.5% tin.
Calculate the mass (in g) of each element in a 47.7 g pewter figurine.
20.
Provide the missing name or chemical formula for each of the following:
barium bromide
aluminum oxide
copper (II) nitrate
Cl2O
N2O4
NF3
SnBr2
sulfurous acid
tetraphosphorus hexaoxide
sodium hypoiodite
HClO2 (aq)
potassium permanganate
SrSO3
CuCN
iron (III) hydroxide
21.
The bitter tasting compound quinine is a component of tonic water and is used as a protection against malaria. Quinine
contains only C, H, O, and N. When a sample of mass 0.487 g was burned, 1.321 g carbon dioxide, 0.325 g of water,
and 0.0421 g of N2 were produced. The molar mass of quinine is 324 g/mol. Determine the empirical and molecular
formulas of quinine.
22.
A compound containing only Na, Al, and F is 32.79% Na, 13.02 % Al, and 54.19% F by mass. Determine the empirical
formula of this compound.
23.
One mole of N2 contains how many N2 molecules? N atoms?
24.
Consider one mole of iron (III) sulfate, Fe2(SO4)3. Calculate the mol SO42– in this sample. Calculate the number of O
atoms in this sample.
25.
Balance the following chemical equations:
a.
Na (s) + H2O (l)  NaOH (aq) + H2 (g)
b.
C2H4 (g) + O2 (g)  CO2 (g) + H2O (l)
c.
Al (s) + HCl (aq)  AlCl3 (aq) + H2 (g)
d.
Mg3N2 (s) + HCl (aq)  MgCl2 (aq) + NH4Cl (aq)
e.
Fe(OH)3 (s) + H2SO4 (aq)  Fe2(SO4)3 (aq) + H2O (l)
26.
Iron in the form of fine wire burns in oxygen to form iron (III) oxide. How many mol of O2 are required to produce 5.21
mol of iron (III) oxide?
27.
Butane, C4H10 burns with the oxygen in air to give carbon dioxide and water:
2 C4H10 (g) + 13 O2 (g)  8 CO2 (g) + 10 H2O (g)
What type of reaction is this? What mass of CO2 can be produced from the reaction of 0.15 mol C4H10 with excess O2?
28.
Ethanol, C2H5OH burns in air to give carbon dioxide and water. This combustion reaction is describe by the following
equation: C2H5OH (l) + 3 O2 (g)  2 CO2 (g) + 3 H2O (l)
a.
Consider a 88.8 mL sample of ethanol (d = 0.789 g/mL). Calculate the mol C in this sample. Calculate the number
of oxygen atoms in this sample.
b.
What is the amount of water (in mol) produced from the complete reaction of 0.25 mol ethanol?
c.
What is the mass of water (in grams) produced from the complete combustion of 55.48 g of ethanol?
29.
Hematite, Fe2O3, is an important ore of iron. Iron metal is obtained by reacting hematite with carbon monoxide (CO) in
a blast furnace. The balanced chemical equation for this reaction is:
Fe2O3 (s) + 3 CO (g)  2 Fe (s) + 3 CO2 (g)
Calculate the mass (in g) of iron that can be produced from the complete reaction of 1.00 kg of Fe2O3.
30.
Chlorine can be produced in the laboratory by heating hydrochloric acid with manganese (IV) oxide. How many grams
of HCl are required to react with 5.00 g of MnO2 according to this equation?
4 HCl (aq) + MnO2 (s)  2 H2O (l) + MnCl2 (aq) + Cl2 (g)
31.
Zinc sulfide ore is an important source of zinc metal. The first step in the processing of the ore consists of heating the
sulfide with oxygen to give zinc oxide, ZnO and sulfur dioxide, SO2. Calculate the mass (in kg) of O2 (g) required to
react completely with 5.00 x 103 g of ZnS.
32.
Nitric acid, HNO3, is manufactured by the Ostwald process, in which nitrogen dioxide, NO2 reacts with water:
3 NO2 (g) + H2O (l)  2 HNO3 (aq) + NO (g)
How many grams of nitrogen dioxide (with excess water) would be required to produce 5.00 g of nitric acid?
33.
Methanol, CH3OH, is prepared industrially from the gas-phase catalytic reaction:
CO (g) + 2 H2 (g)  CH3OH (g)
In a laboratory test, a reaction vessel was filled with 35.4 g CO and 10.2 g H2. Calculate the mass (in g) of methanol
that can be produced in this reaction. Which reactant remains unconsumed at the end of this reaction? What mass (in g)
of it remain?
34.
Carbon disulfide (CS2) burns in oxygen. Complete combustion gives the reaction:
CS2 (g) + 3 O2 (g)  CO2 (g) + 2 SO2 (g)
Calculate the grams of sulfur dioxide produced when a mixture of 15.0 g of carbon disulfide and 35.0 g of oxygen
reacts. Which reactant remains unconsumed at the end of the combustion? How many grams remain?
35.
Calculate the molarity of the following solutions:
a.
0.0834 mol of Na2SO4 dissolved in enough water to form 650.0 mL of solution
b.
1.45 g of NaCl dissolved in enough water to form 250.0 mL of solution
36.
Calculate the mass (in g) of solute present in each of the following solutions:
a.
1.00 x 102 mL of 0.150 M Na2SO4
b.
2.50 x 102 mL of 0.0500 M KBrO3
37.
A sample of 0.0341 mol of iron (III) chloride was dissolved in water to give 25.0 mL of solution. Calculate the
molarity of this solution. What is the molarity of chloride ions in this solution?
38.
An experiment calls for 0.0353 g of potassium hydroxide, KOH. What volume (in mL) of 0.0176 M KOH is required?
39.
You wish to prepare 0.12 M HNO3 from a stock solution of nitric acid that is 15.8 M. What volume (in mL) of the
stock solution is required to prepare 1.00 L of the 0.12 M HNO3?
40.
Calculate the volume of 0.250 M HNO3 that will react completely with 42.4 mL of 0.150 M Na2CO3 according to the
following reaction: 2 HNO3 (aq) + Na2CO3 (aq)  2 NaNO3 (aq) + H2O (l) + CO2 (g)
41.
A 3.33 g sample of iron ore is treated resulting in a solution of iron (II) sulfate. The resulting solution is then reacted
with 0.150 M K2Cr2O7 (see balanced equation given below). 41.4 mL of the potassium dichromate solution are required
for complete reaction. Calculate the mass percent iron in the ore sample.
6 FeSO4 (aq) + K2Cr2O7 (aq) + 7 H2SO4 (aq)3 Fe2(SO4)3 (aq) + Cr2(SO4)3 (aq) + 7 H2O (l) + K2SO4 (aq)
42.
Bone is dissolved in hydrochloric acid resulting in 50.0 mL of a solution containing calcium chloride, CaCl2. To
precipitate the calcium ion from this solution, an excess of potassium oxalate is added. 1.437 g of calcium oxalate is
collected. Calculate the molar concentration of the CaCl2 solution.
Chem 171:
Review - Exam 1 Answers
1.
8.4 x 103 kg/m3
2.
9.26 cm3
3.
a. Na
4.
a. 5
b. 2
5.
a. 216
b. .037
6.
b. Cl
c. Ra
c. 2
d. Ne
d. 3
c. .224
e. 1
d. .5
f. 5
e. 13.51
g. 6
f. .39
h. 2
g. 15.0
h. .475
7
a. .0301 m, 30.1 mm, 3.01 x 10 nm
b. .0500 dm3, 50.0 mL, .0500 L, 5.00 x 10—5 m3
c. 17.6 lb, 7.98 x 103 g, 7.98 x 106 mg, 7.98 x 109 µg
27
12
Al
39
7.
C
a. 6
8.
a. 11 p, 10 e, 12 n
b. 35 p, 36 e, 46 n
c. 19 p, 18 e, 20 n
d. 20 p, 18 e, 20 n
e. 49 p, 46 e, 66 n
f. 52 p, 54 e, 75 n
b. 13
K
c. 19
9.
Cu, atomic number 29, 63.55 amu
10.
85.47 amu
11.
a. 137.3 g/mol
b. 138.2 g/mol
c. 366.1 g/mol
d. 158.1 g/mol
12.
a. 84.34%
b. 656 g N
13.
a. 0.0200 mol HCl
b. 2.00 mol Na2SO4
14.
a. 18 g CO2
b. 35 g K2SO4
—22
g
—22
b. 2.273 x 10
c. 51.1 g NH3
15.
a. 1.327 x 10
g
16.
a. 4.172 x 1023 atoms
b. 8.059 x 1022 formula units
17.
a. 11.9 g
b. .258 moles
c. 2.390 x 10—22 g
c. 6.14 x 1022 molecules
c. 1.55 x 1023 molecules
d. 3.039 x 10–22 g
18.
a. phosphorus, group 5A (or group 15), period 3, non-metal
b. antimony, group 5A (or group 15), period 5, metal (or metalloid)
c. silver, group 11, period 5, metal
d. zinc, group 12, period 4, metal
e. barium, group 2A (or group 2), period 6, metal
19.
2.86 g Sb, 0.716 g Cu, 44.1 g Sn
20.
BaBr2
Al2O3
Cu(NO3)2
dichlorine oxide
dinitrogen tetraoxide
nitrogen trifluoride
tin (II) bromide
H2SO3 (aq)
P4 O 6
NaIO
chlorous acid
KMnO4
strontium sulfite
copper (I) cyanide
Fe(OH)3
21.
C10H12NO, C20H24N2O2
22.
Na3AlF6
23.
6.022 x 1023 N2 molecules, 1.20 x 1024 N atoms
24.
3 mol SO42— ions, 7.226 x 1024 O atoms
25.
a. 2 Na (s) + 2 H2O (l)  2 NaOH (aq) + H2 (g)
b. C2H4 (g) + 3 O2 (g)  2 CO2 (g) + 2 H2O (l)
c. 2 Al (s) + 6 HCl (aq)  2 AlCl3 (aq) + 3 H2 (g)
d. Mg3N2 (s) + 8 HCl (aq)  3 MgCl2 (aq) + 2 NH4Cl (aq)
e. 2 Fe(OH)3 (s) + 3 H2SO4 (aq)  Fe2(SO4)3 (aq) + 6 H2O (l)
26.
7.82 mol O2
27.
26 g CO2
28.
a. 3.04 mol C, 9.16 x 1023 O atoms
29.
699 g Fe
30.
8.39 g HCl
31.
2.46 kg O2
32.
5.48 g NO2
33.
40.4 g CH3OH, H2, 5.12 g H2
34.
25.2 g SO2, O2, 16 g O2
35. a. .128 M
b. .0992 M
36. a. 2.13 g
b. 2.09 g
—
37. 1.36 M FeCl3 (aq), 4.09 M Cl (aq)
38. 35.7 mL
39. 7.6 mL
40. 50.9 mL
41. 62.5 %
42. 0.224 M
b. 0.75 mol H2O
c. 65.10 g H2O