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
                      atomic radius                   Periodic Law

         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

          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
                                  traditional naming system              molecular formula
                                  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?

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
                               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
                                  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
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
       2HCl         H2     +     Cl2
       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 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.