# Final Exam Review 271 Answer Key

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```					                                 Final Exam Review - 271
Unit 1 Atomic Structure Terms:
mass number        neutron          sublevels     average atomic mass
atomic number      electrons        orbitals      valence electrons
proton             ion              isotopes      kernel electrons

1. 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 #            Neutrons       Electrons
Ar              18              38                 20              18
Rb              37              72                 35              37
Si              14              29                 15              14
Na              11              22                 11              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

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. What were some problems with Dalton’s Atomic Theory? Solid mass, no isotopes

7. What did Rutherford’s experiment prove? The existence of a positively charged nucleus.
8. How many electrons are in an s orbital? 2
9. Which sublevels are in the 4th energy level? s, p, d, f

1s     2s     2p       3s      3p      4s
10. Al

11. 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                                    c. I 1s22s22p63s23p64s23d104p65s24d105p5
VS = 4 VE = 2 KE = 28                                     VS = 5 VE = 7 KE = 46
b.    Ba 1s22s22p63s23p64s23d104p65s24d105p66s2                 d. Cd 1s22s22p63s23p64s23d104p65s24d10
VS = 6 VE = 2 KE = 54                                     VS = 5 VE = 2 KE = 46

Unit 2 Periodic Table Terms:
metals/nonmetals                transition elements
groups/periods                  ionization energy

1. What were Mendeleev’s contributions to the Periodic Table? Initial
organization/Predicted undiscovered elements.
2. Who arranged the periodic table to its current format?   Henry Moseley

3. Write the symbol of 2:

a. metals K, Fe                  d. noble gases He. Ne         g. alkaline earth
metals Mg, Ca
b. nonmetals S, O                e. transition metals
Ag, Mo                     h. inner transition
c. halogens F, Br                f. alkali metals Rb, Li        metals U, Ce

4. Write complete electron configurations and 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. Fe [Ar]4s23d6 VS=4 VE=2 KE= 24

b. P [Ne]3s23p3 VS=3 VE=5 KE= 10                 d. Ba [Xe]6s2 VS=6 VE=2 KE= 54

5. What are the trends for atomic radius and ionization energy? Atomic radius
increases top to bottom and decreases left to right. Ionization energy is the
opposite.

6. Which is larger? a. S or O               b. Na or K            c. Ca or Ca2+
Unit 3 Chemical Bonding Terms:
double bond                     cation           covalent bond
molecular geometry              anion            ionic bond

1. Draw dot diagrams for Ca, S, and K1+           Ca               S            [ K ]1+

2. 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.

3. 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.

4. 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.

12. Draw a structure and identify the molecular geometry of the following:

a.   NH3                                      c.    CH4

b.   H2O                                      d.    CCl4

Unit 4 Study Terms:         chemical formula                       empirical formula
ternary compound                       binary compound
stock naming system                    diatomic molecule
binary ionic                           binary acid

13. 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. Hydrochloric acid HCl           h. Fe2(CO3)3 iron(III)carbonate o. MgCl2 magnesium chloride
b. MgSO3 magnesium sulfite          i. acetic acid HC2H3O2          p.   copper(I)oxide Cu2O
c. ammonium oxalate (NH4)2C2O4      j. NH4Cl ammonium chloride q. CaC2O4 calcium oxalate
d. sodium chloride NaCl             k. AgOH silver hydroxide        r. strontium bromide SrBr2
e. iron(II)oxide FeO                l. aluminum sulfide Al2S3       s. AlCl3 aluminum chloride
f. Sr(NO3)2 strontium nitrate       m. H2SO4 sulfuric acid          t. Titanium(IV)oxide TiO2
g. nitric acid HNO3                 n. Rubidium oxide Rb2O
14. What is the molecular and empirical formula for the following structural formulas?

C6H6
C2H4O2         CH2O                              CH

Unit 5 Measurement Terms:          scientific notation        rules for calculating
qualitative/quantitative   accuracy/precision

1. What is the difference between quantitative and qualitative? Quantitative uses numbers to
describe a measurement, while qualitative is a general description without numbers.

2. What is the difference between accuracy and precision? Accuracy is a measure of how close
the answer is to the accepted value, while precision is being able to get a consistent result with
each trial.

3. Express these in scientific notation:

a.   400             4.00 x 102
b.   .0404           4.04 x 10-2
c.   600 x 103       6.00 x 105
d.   25.3 x 10-4     2.53 x 10-3

4. Calculate:
a.     (6.0 X 104)(2.0 X 105)= 1.20 X 1010
b.     (6.0 x 106)/(2.0 x 104)= 3.00 x 102

Unit 6 The Mole Terms:       Avogadro’s number          percentage composition
formula mass               molar mass
Molarity                   atomic mass
empirical/molecular formula

1. Calculate the formula mass of the following compounds:
a. N2O5       108g b. CO2          44 g       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. Determine the mass of 213 L of NO2?
213 L NO2 x 1 mol NO2 x 46 g NO2                    = 437.4 g NO2
22.4 L NO2 1 mol NO2

e. Calculate the mass of 2.26 x 1024 molecules of PCl3 at STP?
24                                                                                    2
2.26 x 10 molecules PCl3 x        1 mol PCl3           x 137.5 g PCl3     = 516.2 g = 5.16 x 10 g PCl3
23
6.02 x 10 molecules        1 mol PCl3

3. Determine the percentage composition by mass for the following compounds.

a. CO 43% C                     b. NaCl             39.3% Na        2.04% H
c. H2SO4
57% O                                         60.7% Cl        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       2(CH3) = C2H6 = molecular
12 g C
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                2(CH3O) = C2H6O2 = molecular
1gH
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     yields               endothermic
products      exothermic           coefficient
perfect mixing ratio               limiting reactant
types of reactions

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    Balanced

d.      2Sb              +                  3H2O                       Sb2O3           +          3H2

e.      2Re              +                  3Br2                       2ReBr3

4. Which type of reactions are these?

a. 2K(s) +       Cl2(g)                 2KCl(s)                         synthesis

b. Fe2O3(s)       +       2Al(s)              Al2O3(s)          + 2Fe(s)        single replacement

c. HNO3 (aq)      +       NaOH (aq)                   H2O (l)      + NaNO3 (aq)         double replacement

d. PbO2 (s)               Pb (s)    +       O2 (g)              decomposition
5. 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

6. 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

7. 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
23
6.02 x 10 molecules HCl       2 mol HCl      1 mol Cl2

8. Determine the number of molecules of sulfur trioxide that is produced when 64 L of oxygen
gas at STP reacts with excess sulfur according to this balanced chemical equation.
S     +      3O2          2SO3
23                           24
64 L O2 x 1 mol O2 x 2 mol SO3 x 6.02 x 10 molecules SO3 = 1.15 x 10                  molecules S
22.4 L O2  3 mol O2        1 mol SO3

NOTE: Review all labs. At least 1 question related to each lab will appear on the final.

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