Download Final Exam Review 271 Answer Key 2015

Final Exam Review – 271 KEY
Unit 3 Atomic Structure Terms:
mass number
atomic number
proton
neutron
electrons
isotopes
1.
average atomic mass
ion
energy levels
sublevels
orbitals
valence electrons
kernel electrons
quantum numbers
Atomic Theory/History
nuclear decay eq. (alpha/beta)
half life
What is the mass, charge and location of an electron, proton and neutron?
Mass
Charge
Location
proton
1 amu
positive
nucleus
neutron
1 amu
neutral
nucleus
electron
0 amu
negative
outside nucleus
2. Assume these are all neutral atoms.
Element
Atomic #
Mass #
Ar
18
38
Rb
37
72
Si
14
29
Na
11
22
Neutrons
20
35
15
11
Electrons
18
37
14
11
3. List the number of protons, neutrons and electrons for the following ions.
Element
Protons
Neutrons
Electrons
12
13
10
8
9
10
4. Neon has 2 isotopes. Neon-20 has a mass of 19.992 u and neon-22 has a mass of 21.991 u. In any
sample of 100 neon atoms, 90 will be neon-20 and 10 will be neon-22. What is the average atomic
mass of neon?
Mass
% Abundance
Neon - 20
19.992 x
90
=
17.9928
Neon – 22
21.991 x
10
=
+ 2.1991
20.1919 u = 20.192 u
5. What is the average atomic mass of silicon if 92.21 % of its atoms have mass 27.97 u, 4.70% have mass
28.976 u, and 3.09 % have mass 29.974 u?
Mass
% Abundance
Si 27.97 x
92.21 =
25.79
Si 28.976 x
4.70 =
1.362
Si 29.974 x
3.09 = +
.9262
28.08 u
6. Write electron configurations for the following elements using the Diagonal Rule. Identify the energy
level of the valence shell, the number of valence electrons and the number of kernel electrons.
a. Zn 1s22s22p63s23p64s23d10
VS = 4 VE = 2 KE = 28
c. O2- 1s22s22p6
VS = 2 VE = 8 KE = 2
b. Ag 1s22s22p63s23p64s23d104p65s14d10
VS = 5 VE = 2 KE = 45
d. Cr 1s22s22p63s23p64s13d5
VS = 4 VE = 2 KE = 22
7. Phosphorus-32 has a half-life of 14.3 days. How many milligrams of phosphorus-32 remain
after 57.2 days if you start with 4.0 mg of the isotope? 57.3/14.3 = 4 half lives
# half lives
1
2
3
4
8.
Starting amount = 4.0 mg
2.0 mg
1.0 mg
0.5 mg
.25 mg left after 4 half lives
Write an equation for Alpha Emission of the following element:
4
234
Uranium-238 238
92 Po
2 He +
90 Th
Unit 4 Periodic Table Terms:
metals/nonmetals
transition elements
groups/periods
noble gas notation
Periodic Law/History
atomic radius
ionization energy
electronegativity
1. Write the symbol of 2:
a. metals K, Fe
d. noble gases He. Ne
b. nonmetals S, O
e. transition metals
Ag, Mo
f. alkali metals Rb, Li
c. halogens F, Br
g. alkaline earth
metals Mg, Ca
h. inner transition
metals U, Ce
2. Write noble gas notation for the following elements using the Periodic Table. Identify the energy
level of the valence shell, the number of valence electrons and the number of kernel electrons.
a. Li [He]2s1 VS=2 VE=1 KE= 2
c. F1- [He]2s2p6 VS=2 VE=8 KE= 1
b. Fe3+ [Ar]3d5 VS=4 VE=8 KE= 18
d. Sn [Kr]5s24d105p2 VS=5 VE=4 KE= 46
3. Of the elements Ga, Br, and Ca, which has the highest electronegativity? Br
4. For each of the following pairs, circle the atom or ion having the larger radius.
a.
S
or
b.
Ca or
O
Ca2+
c.
Na1+
or K1+
d.
Na or K
e.
S2–
f.
F or
or
O2–
F1–
Unit 5 Chemical Bonding Terms:
cation
covalent bond/properties of
anion
ionic bond/properties of
double bond triple bond
molecular geometry
polar/nonpolar
1. Describe how the ionic bond forms between an atom of potassium and an atom of
chlorine.
Potassium transfers 1 electron to the chlorine in order to bond. Potassium
becomes a 1+ ion and chlorine (chloride) becomes a 1- ion.
2. Describe how a covalent bond forms between an atom of hydrogen and an atom of
chlorine.
Hydrogen and chlorine share electrons to form a bond. There is an unequal sharing
(polar covalent bond) since chlorine has a higher electronegativity than hydrogen but
they are still sharing. No ions are formed.
3. Why do atoms form bonds? [Hint: Octet Rule]
Atoms form bonds in order to become more stable in their electron configuration like
the nearest noble gas. Metals and nonmetals achieve this through ionic bonding while
nonmetals bond by covalent bonding.
4. Draw a structure and identify the molecular geometry of the following:
a.
NH3
pyramidal
c.
CH4
b.
H2O
bent
d.
CO2
Unit 6 The Mole Terms:
Avogadro’s number
formula mass
molar mass
atomic mass
molecular mass
empirical/molecular formula
1. Calculate the formula mass of the following compounds.
a. N2O5 108g
b. CO2
44 g
tetrahedral
linear
molarity
percentage composition
c. H3PO4
98 g
2. Do the following mole conversion problems:
a. A sample of neon has a volume of 11.2 L. How many moles does this represent?
11.2 L Ne x 1 mol Ne = 0.5 mol =
5.00 x 10-1 mol Ne
22.4 L Ne
b. How many atoms are there in 2.75 moles of copper?
2.75 mol Cu x 6.02 x 1023 atoms Cu = 1.66 x 1024 atoms Cu
1 mol Cu
c. What is the mass of 1.51 x 1024 atoms of carbon?
1.51 x 1024 atoms C x
1 mol C
x 12 g C = 30.09 g C = 3.01 x 101 g C
23
6.02 x 10 atoms C 1 mol C
d. Calculate the mass of 2.26 x 1024 molecules of PCl3 at STP?
2.26 x 1024 molecules PCl3 x
1 mol PCl3
x 137.5 g PCl3
PCl3
6.02 x 1023 molecules
1 mol PCl3
= 516.2 g = 5.16 x 102 g
e. What is the volume of 1.51 x 1023 molecules of N2O3 at STP?
1.51 x 1023 molecules N2O3 x 1 mol N2O3
x
22.4 L N2O3 = 5.62 L N2O3
6.02 x 1023 molecules
1 mol N2O3
f. A sample of SO2 has a mass of 38.5 g. What is the volume at STP?
38.5 g SO2 x 1 mol SO2
x 22.4 L SO2 = 13.45 L SO2
64.1 g SO2
1 mol SO2
3. Determine the percentage composition by mass for the following compounds.
a. CO 43% C
57% O
b. NaCl
39.3% Na
60.7% Cl
c. H2SO4
2.04% H
32.7% S
65.2% O
4. What is the empirical formula of a compound, given that a 212.1 g sample of the compound
contains 42.4 g of hydrogen and 169.7 g of carbon? What is the molecular formula of the
compound, given it has a molecular mass of 30.0 u.
42.4 g H x 1 mol H = 42.4 mol H/14.08 = 3 mol H
1gH
empirical = CH3
30.0 u/15.0 = 2
169.7 g C x 1 mol C = 14.08 mol C/14.08 = 1 mol C
f.m. = 15
12 g C
2(CH3) = C2H6 = molecular
5. Calculate the empirical and molecular formula of a substance that is 38.7% C, 9.6% H and 51.6%
O. The molecular mass of the compound is 62 u.
38.7 g C x 1 mol C = 3.225 mol C/3.225 = 1 mol C
12 g C
empirical = CH3O
62 u/31 = 2
9.6 g H x 1 mol H = 9.6 mol H/3.225 = 3 mol H
f.m. = 31
1gH
2(CH3O) = C2H6O2 = molecular
51.6 g O x 1 mol O = 3.225 mol O/3.225 = 1 mol O
16 g O
6. The molecular mass of a compound is 118.0 u and its empirical formula is C2H3O2. What is the
molecular formula of this compound?
118.0 u / 59 = 2
2(C2H3O2) = C4H6O4
7. What is the molarity of 2.3 moles of sodium chloride, NaCl in 0.45 liters of water?
M= mol
2.3 mol
L
0.45 L = 5.11 M solution
8. What is the molarity of 98 grams of sodium hydroxide, NaOH in 2.2 liters of water?
98 g NaOH x 1 mol NaOH = 1.11 M solution
2.2 L 40 g NaOH
9. How many grams of sodium sulfate are needed to make 0.75 L of a 0.25 M sodium sulfate,
Na2SO4.
M = mol
mol = M x L so…
0.75 L x 0.25 mol Na2SO4 x 142.1 g Na2SO4 = 26.64
g
L
1L
1 mol Na2SO4
10. What is the molarity of a solution that contains 2.5 L of solvent and 660 grams of calcium
phosphate, Ca3(PO4)2? 660 g Ca3(PO4)2 x 1 mol Ca3(PO4)2 = 0.851 M solution
2.5 L
310.3 g Ca3(PO4)2
Unit 7 Stoichiometry Terms:
reactants
products
endothermic
exothermic
types of reactions
coefficient
limiting reactant
percent yield
1. Solid sodium reacts with oxygen gas to produce solid sodium oxide. Write a word equation and
a formula equation, using the correct chemical formulas and phase symbols. Balance the
equation. sodium(s) + oxygen (g)
sodium oxide(s)
4Na(s) + O2 (g)
2Na2O (s)
2. Write the word and the formula equation for the decomposition of solid calcium carbonate into
solid calcium oxide and carbon dioxide gas. Balance the equation.
calcium carbonate(s)
calcium oxide(s) + carbon dioxide(g)
CaCO3 (s)
CaO(s) + CO2(g)
3. Balance the following equations:
a.
5O2
+
Sb2S3
Sb2O4
+
3SO2
b.
Cu
+
Cl2
CuCl2
Balanced
c.
CuO
+
H2
Cu
+
H2O
d.
2Sb
+
3H2O
Sb2O3
+
3H2
e.
2Re
+
3Br2
2ReBr3
Balanced
4. How many liters of oxygen gas at STP can be produced by the decomposition of 36.0 grams of
water? The balanced equation is as follows.
2H2O
2H2
+
O2
36.0 g H2O x 1mol H2O
x 1 mol O2 x 22.4 L O2
= 22.4 L O2
18.0 g H2O
2 mol H2O
1 mol O2
5. How many grams of water vapor will be produced when 56.5 L of hydrogen at STP reacts
completely with oxygen according to the balanced chemical equation?
2H2 +
O2
2H2O
56.5 L H2 x 1 mol H2 x 2 mol H2O x 18 g H2O = 45.4 g H2O
22.4 L H2 2 mol H2
1 mol H2O
6. What volume of chlorine gas, measured at STP, can be produced by the decomposition of 2.15 x
1024 molecules of hydrogen chloride gas given the following balanced chemical equation?
2HCl
H2
+
Cl2
24
2.15 x 10 molecules HCl x
1 mol HCl
x 1 mol Cl2 x 22.4 L Cl2 = 40 L Cl2
6.02 x 1023 molecules HCl 2 mol HCl
1 mol Cl2
Unit 8 Solutions and pH Terms:
solubility/solvation solute
solvent
solution
pH scale
acid/base
neutral
insoluble
percent by mass (concentration)
neutralization
miscibility
immiscibility
molarity
saturated solution
unsaturated solution
supersaturated solution
suspension
1. When two substances mix together and one becomes insoluble what does it form? precipitate
2. What does the phrase “like dissolves like” refer to? Polar dissolves in, or is miscible in, other
polar substances, while polar and non polar substances are immiscible in each other.
3. Give 3 ways that solvation of a substance can be increased. Heat the solution, crush the solute to
increase surface area, stir faster.
4. What number on the pH scale is considered neutral? 7
5. Which types of substances have high pH values (8-14) and which has low (0-6)? Bases have high
pH values and turn pH paper blue, while acids have low pH values and turn pH paper red.
6. What happens when an acid dissolves in water? A base? When an acid dissolves in H2O, H+ ions
are released. When bases dissolve in water OH- ions are released.
7. What happens when an acid and a base mix together? Neutralization occurs and water is
produced.
8. What is the mass percent of sodium chloride when 30 g of NaCl are dissolved in 100 g of
solution?
30 g/(30 g + 100 g) x 100 = 23.1% NaCl by mass
Nomenclature:
chemical formula
monatomic ion
diatomic molecule
empirical formula
molecular formula
traditional naming
system
stock naming system
binary molecular
binary ionic
ternary ionic
binary acid/ternary acid
1. Classify the compound. Write the name when given the formula or vice-versa. Identify the name and
number of cations and anions for ionic compounds. Identify the number of atoms of each element for
ternary ionic compounds. Write the alternate name for molecular compounds.
a. plumbic oxide PbO2
h. N2O3 dinitrogen trioxide
b. MgSO3 magnesium sulfite
i. AgOH silver hydroxide
c. ammonium oxalate (NH4)2C2O4
j. H2SO4 sulfuric acid
d. CCl4 carbon tetrachloride
k. ferrous chloride FeCl2
e. Sr(NO3)2 strontium nitrate
l. BeCl2 beryllium chloride
f. nitric acid HNO3
m. CaC2O4 calcium oxalate
g. Fe2(CO3)3 iron(III)carbonate
n. nitrogen dioxide NO2