Chemistry Study Guide (DOC) by mikesanye

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									Chemistry                                                             Name
                                                                      Period
Chemistry Study Guide
Answer the following questions on a separate sheet of paper.
                                                     Chapter 10 States of Matter
1. State the 5 basic assumptions of the kinetic molecular theory.
   a. Gases consist of large numbers of tiny particle that are far apart relative to their size.
   b. Collisions between gas particles and particles and container walls are elastic collisions.
   c. Gas particles are in continuous, rapid, random motion.
   d. There are no forces of attraction between gas particles.
   e. The temperature of a gas is dependent on its average kinetic energy.
2. What is the relationship between the temperature, speed and kinetic energy of gas molecules?
   They are dependent on each other, as kinetic energy increases so do temp and speed
3. Explain how the attractive forces between particles in a liquid are related to the equilibrium vapor pressure of that liquid.
   The stronger the attractive forces (intermolecular bonds), the high the vapor pressure. The liquid particles with high attractive
   forces require more energy (more pressure) to break the bonds between the molecules.
4. What is the relationship between atmospheric pressure and the boiling point of a liquid?
   In order for a substance to boil, the atmospheric pressure must equal the vapor pressure of the liquid.
5. Water has a high heat of fusion (6.009 KJ/mol) and a high heat of vaporization (40.79 KJ/mol). Explain what this means in
   terms of attraction between particles.
   Heat of fusion is the amount of energy it takes to change 1 mole from solid to liquid. Heat of vaporization is the amount of
   energy needed to change 1 mole of a liquid to a gas. H 2O has a high heat of fusion and vaporization because water is polar and
   has strong H-bonds holding the molecules together in the solid or liquid state.
6. Using the values given in #5, calculate the amount of energy needed to melt 7.95x105 g of ice. Is this an exothermic or
   endothermic reaction? Endothermic, 2.65x105kJ
7. Explain why the water molecule is polar. In your explanation, include why the water molecule is bent?
   The oxygen on water has 2 unshared pairs of electrons, which repels the covalent bonds more than normal creating the bent
   angle of 105º. The electrons in the covalent bond are shared unequally creating a dipole between the two elements.
8. From your knowledge of intermolecular forces, arrange the following in order of increasing surface tension (least to most):
   Water, hexane, ethanol, ethanal            hexane<ethanal<ethanol<water
9. Describe how the intermolecular forces in water allow for each of the following properties of water:
   a. low vapor pressure Requires more E to overcome the NP molecules in air
   b. high heat of vaporization     A lot of E is needed to overcome the strong H-bonds between molecules
   c. solid H2O is less dense than liquid H2O
        When water freezes into ice, it increases its surface area by forming rigid, honeycomb shapes.
   d. high boiling point for a molecule of its mass
        A lot of energy is needed to overcome the strong intermolecular forces (H-bonds) between the molecules.
                                                          Chapter 11 Gases
1. Explain what happens to the pressure inside a balloon when you blow into it.
   As you increase the number of particles into the balloon, the pressure increases.
2. Why can’t an ideal gas be liquefied? Ideal gas particles do not attract to each other and change state.
3. What would be the number of moles and grams of carbon dioxide gas contained in an 855 mL container at 35˚C and 1860
   mmHg of pressure? 0.083mol, 3.6g
4. If a 2.75 L container of gas at 24.0˚C and 95.2 kPa was compressed to 1820 mL and warmed to 40.0˚C, what would be the new
   pressure? 152 kPa
5. A mixture of 3 gases, A,B and C, is at a total pressure of 6.11 atm. The partial pressure of A is 168 kPa and B is 3.89 atm. What
   is the partial pressure of gas C? 0.56 atm
6. Ammonia and ethanol are released at the same time across a room. Which will you smell first? ammonia
7. What is the volume of 8.00 grams of oxygen at STP? 5.60 L
                                                    Chapter 16 Thermochemistry
1. How much heat, in calories, does 32.0 g of water absorb when it is heated from 25ºC to 80ºC? How many joules is this?
   1800 cal, 7400 J
2. The temperature of a piece of copper with a mass of 95.4 g changes from 25.0ºC to 48.0ºC when the metal absorbs 849 J of
   heat. What is the specific heat of copper? Is this an exothermic or endothermic reaction? 0.387J/g x ºC, endothermic
3. Will the specific heat of 50.0 g of a substance be the same as or greater than the specific heat of 10.0 g of the same substance?
   The same (specific heat is a physical property)
4. If 28.2 g of CaCl2 are dissolved in 125 mL of H2O at 25.0˚C and the final temp. reaches 42.0˚C, calculate the ΔH in kJ and
   ΔHsoln in kJ/mol. Also diagram this process. 8.88 kJ, 34.9 kJ/mol
5. Use the following equation to calculate the amount of energy released when 725 g of CaO are dissolved in water:
   CaO(s) + H2O(l)                  Ca(OH)2(s) + 65.2 kJ        843kJ
                                                              Chapter 12 Solutions
1.   Define the following terms: (look up on own)
     a. solubility           b. saturated solution        c. unsaturated solution     d. Henry’s law
2.   How does each of the following affect the solubility of (a) a solid dissolved in a liquid, and (b) a gas dissolved in a liquid.
     a. an increase in temperature (a) more collisions between particles causing an increase in dissolving particles
                                       (b) decreases solubility, as T increases more dissolved gas escapes
     b. shaking, agitation (a) more collisions between particles causing an increase in dissolving particles
                             (b) increases solubility
     c. an increase in pressure        (a) no change      (b) increase solubility
     d. an increase in pressure with a decrease in temperature (a) no change          (b) increases solubility (more than c.)
3.   Differentiate between the following:
     a. a dilute unsaturated solution and a dilute saturated solution
          A dilute unsaturated solution contains less solute than a dilute saturated solution. A dilute saturated solution contains more
          solute but more solute could still be added.
     b. a concentrated saturated solution and a concentrated unsaturated solution
          A concentrated saturated solution contains more solute than a concentrated unsaturated solution. A concentrated
          unsaturated solution has less solute but, more solute can still be added.
4.   Calculate the following (SHOW ALL WORK):
     a. The molarity of a solution containing 42.6g of sodium hydroxide in 3.00L of water. 0.357M
     b. The number of moles of solute present in 680mL of a 0.25M Na2SO4 solution. 0.17 mol
     c. The number of grams of KBr present in 500mL of a 0.100 M solution. 6 g
5.   Describe how to prepare the following solutions. Include calculations, a description of the procedure, and specific equipment.
     a. 400.0 mL of 0.15 M solution of copper (II) sulfate from 0.75M stock solution.
          Using M1V1 = M2V2, calculate the volume of stock needed to make your dilute solution. Answer is 80 mL of stock and add
          320 mL of water (V2-V1) to make your dilution. Use a 500 mL graduated cylinder for the water and a 100 mL graduated
          cylinder to measure the amount of stock needed.
     b. 50.0 mL of a 0.20 M solution of potassium nitrate from a 4.0 M stock solution.
          Using M1V1 = M2V2, calculate the volume of stock needed to make your dilute solution. Answer is 2.5 mL of stock and add
          47.5 mL of water to make your dilution. Use a 100 mL graduated cylinder for the water and a 10 mL graduated cylinder to
          measure the amount of stock needed.
6.   Calculate the following (SHOW ALL WORK):
     a. The maximum amount of KCl that can be dissolved in 45.0g of water at 20 oC (solubility of KCl is 34g/100g H2O). 15.3 g
     b. A student dissolved 52.0g of sodium fluoride in 125g of water (solubility of NaF is 12g/100g H 2O at 25oC). Is the solution
     saturated or unsaturated? 15 g; saturated
                                           Chapters 17-18 Rates of Reaction and Equilibrium
1.   Define the following terms: (look up on own)
     a. activation energy              c. chemical equilibrium              e. Le Chatelier’s principle
     b. catalyst                       d. collision theory                  f. reaction mechanism
2.   Draw energy diagrams and label the reactants, products, activated complex, activation energy, and energy change for:
     a. an exothermic reaction         b. and endothermic reaction          c. a reaction with and without the presence of a catalyst




3.   Use the collision theory to explain how and why each of the following affects the rate of a reaction:
     a. an increase in temperature increase in collisions, increases the kinetic energy which increases the rxn rate
     b. a decrease in the concentration of a reactant less concentration, decreases the number of collisions, decreasing the rate
     c. particle size a smaller particle has a larger surface area allowing an increase in particle collision which increases the rate
     d. use of a catalyst a catalyst speeds up a rxn by using less starting energy, it doesn’t get used up unless an inhibitor is present
4.   Predict how each of the following changes will affect the equilibrium position of the given reaction:
                                                  PCl5(g)            PCl3(g) + Cl2(g) + energy
     a. increase the pressure          shift left         d. add more PCl5 to the reaction vessel        shift right
     b. decrease the temperature       shift right        e. increase the volume of the container        shift right
     c. decrease the amount of Cl2 shift right
5.   Write the equilibrium constant expression for each of the following reactions: (NOTE – you must first balance the equations!)
     a. 4 NH3 (g) + 5 O2 (g)           4 NO (g) + 6 H2O (g) [NO] 4[H2O] 6 c. COCl2 (aq)                   CO (g) + Cl2 (g) [CO][Cl2]
                                                              [NH3] 4[O2] 5                                                 [COCl]
     b. N2O4 (g)              2 NO2 (g) [NO2] 2                               d. CaO (s) + CO2 (g)          CaCO3(s)        [CaCO3]
                                        [N2O4]                                                                           [CaO][CO2]
6.   For each of the following reactions, determine whether the reactants or the products are favored:
     a. H2Cr2O7 (aq) + 4 NaOH (aq)                2 Na2CrO4 (aq) + 3 H2O(l)             Keq = 5.62 x 1049 products
      b. N2 (g) + O2 (g)                 2NO (g)                                       Keq = 1.2 x 10-4 reactants
7.    Calculate Keq for the following general equations, given the concentration provided:
      a. 2 A + B                2C      [A] = 0.6M         [B] = 0.1M         [C] = 0.4 M Keq = 4
      b. A + 3B                 C       [A] = 1.2M         [B] = 0.4M         [C] = 0.2 M Keq = 3
      c. For the reactions above, are the products or the reactants favored? products
                                       Chapter 13-15 Acids and Bases, Neutralization and Titration
1.    Differentiate between strong electrolytes, weak electrolytes, and nonelectrolytes. Give two examples of each and explain why
      you placed them in each category.
      Strong electrolytes –dissociates/ ionizes completely, strong conductor of electricity (H2SO4, NaOH)
      Weak electrolytes – partially ionizes, poor conductor (tap H2O, NH3)
       Nonelectrolytes – doesn’t ionize (sucrose, distilled water)
2.    Predict the solubility of each of the following substances in: (1)water and (2) heptane
      a. sodium iodide                  c. hydrogen bromide         e. ethanol         g. tetrachloromethane
      b. nickel (solid)                 d. calcium carbonate        f. benzene
      (1) sodium iodide, hydrogen bromide, calcium carbonate, ethanol
      (2) benzene, tetra chloromethane
3.    Why do you believe each of the above is or isn’t soluble in the solvents mentioned? “Like dissolves like” - water is polar and
      dissolves P covalent and Ionic cmpds; heptane is NP and only dissolves NP covalent cmpds
4.    List at least three properties of acids and three properties of bases.
      Acids – taste sour, conducts electricity, pH 1-6, ionizes in water
      Bases – taste bitter, slippery, conducts electricity, pH 8-14, ionic cmpds dissociate in water
5.    Differentiate between Arrhenius, Bronsted-Lowry, and Lewis acids and bases. Give an example of each.
      Arrehenius acid – produces H+(H2SO4) Arrhenius base – produces OH- (H2O)
      Bronsted-Lowry acid – proton donor( HCl) Bronsted-Lowry base – proton acceptor (H2O)
      Lewis acid – accepts a electron pair (BF3) Lewis base – donates an electron pair (NH3)
6.    For the following reaction, label the acid, base, conjugate acid, and conjugate base. List the conjugate acid-base pairs.
                                        CH3COOH + H2O                 CH3COO- + H3O+
                                             Acid      Base          Conj. base Conj. acid
7.    Calculate the following:
      a. [OH] when [H+] of a solution is .0024 M 4.2x10-12M                   d. pOH of a .025M solution of HCl            12.398
      b. [H+] when [KOH] is 1.6 x 10-12                0.0063M                e. pH of a 2.20M solution of Ca(OH)2         14.34
      c. pH of a .025M solution of HCl                 1.600
8.    Show the balanced equation for the self-ionization of water. What is the pH at 25oC? H2O → H+ + OH-, pH = 7
9.    Differentiate between the following and give an example of each:
      a. strong acids and weak acids – strong acids(HCl or H2SO4) ionize completely while weak acids partially ionize(CH3COOH)
      b. strong bases and weak bases – strong bases (NaOH) dissociate completely while weak bases partially ionize (NH3)
10.   Explain how it is possible to have:
      a. a concentrated weak acid - a lot of solute but stays more as a compound rather than as ions
      b. a dilute strong acid – low concentration but ionizes completely
11.   Predict the products for the following acid-base reactions & write balanced equations.
      a. hydrochloric acid + sodium hydroxide              HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
      b. carbonic acid + calcium hydroxide                 H2CO3(aq) + Ca(OH)2(aq) → CaCO3(aq) + 2 H2O(l)
      c. ammonia + sulfuric acid                           NH3(aq) + H2SO4(aq) → (NH4)2SO4(aq)
      d. ethanoic acid + potassium hydroxide               CH3COOH(aq) + KOH(aq) → KCH3COO(aq) + H2O(l)
12.   Calculate the following:
      a. The number of moles of NaOH needed to neutralize 3.5 moles of H2CO3 7.0 mol NaOH
      b. The number of moles of HCl needed to neutralize 6.0 moles of Mg(OH) 2 12 mol HCl
13.   Calculate the following:
      a. [HCl ] if 2.0 mL was needed to neutralize 40.0 mL of a 0 .20M NaOH solution 4.0 M HCl
      b. Number of moles of Mg(OH)2 present in 50.0 mL of solution, if 75.0 mL of a 0.1 M HCl solution was needed 0.004 mol
      c. The molarity of a NaOH solution, if 20.0 mL of the solution was neutralized by 28.0 mL of a 1.0 M H 3PO4 solution 4.2 M
                                                  Chapter 19 Oxidation—Reduction Reactions
1.    Define the following terms:       (look up on own)
      a. oxidation b. reduction         c. half-reaction d. oxidation number           e. oxidizing agent        f. reducing agent
2.    Assign oxidation states for all atoms in each of the following compounds:
      a. KMnO4 b. Na2S2O3               c. SF4             d. H2O             e. I2              f. NaH          g. CH3OH
        +1,+7,-2         +1,+2,-2         +4,-1              +1,-2               0                  +1,-1          -2,+1,-2,+1
3.    For the following reactions, label (1) substance that is oxidized (2) substance that is reduced (3) oxidizing agent and (4)
      reducing agent
      a. Cu(s)        +        2Ag+(aq)           2Ag(s) +           Cu2+(aq)
            (1) (4)            (2) (3)
     b.    SiCl4 (l) +        2Mg(s)            2MgCl2 (s)      +        Si(s)
          Si(2) (3)           (1) (4)
     c. CH4 (g) +             2O2 (g)           CO2 (g) +      2H2O(g)
          (1) (4)             (2) (3)
4.   Balance the following chemical equations using the half-reaction method. SHOW WORK.
     a. As2O3 (s) + NO3- (aq)           H3AsO4 (aq) + NO(g)              (in acidic solution)
     b. CH3OH(aq) + Cr2O72-             CH2O(aq) + Cr3+ (aq)             (in acidic solution)
     c. CN-(aq) + MnO4 -(aq)            CNO-(aq) + MnO2 (s)              (in basic solution)
     d. NO2-(aq) + Al(s)                NH3 (g) + AlO2-(aq)              (in basic solution)
                         Chapter 22 Hydrocarbon Compounds, Functional Groups and Organic Reactions
1.   Define the following terms: (look up on own)
     a. hydrocarbon                   c. structural isomer     e. saturated and unsaturated compounds
     b. aromatic compounds            d. geometric isomer      f. organic compound
2.   Name the following:
     a. C3H8                          propane                  e. C3H7CH=CH2                1-pentene
     b. CH3(CH2)6CH3                  octane                   f. C6H6                      benzene
     c. C2H5CH(CH3)C6H13              3-methylnonane           g. Δ                         cyclopropane
     d. CH3C(triple bond)CCH2CH3 2-pentyne
3.   Draw the structural formulas for each of the following:
     a. 3-ethyl-4- propyloctane       c. 2,2-dimethylbutane              e. 1,3-dimethylbenzene          g. cyclobutene
             C-C                           C                                          CH3
        C-C-C-C-C-C-C-C                 C-C-C-C
                C-C-C                      C                                          CH3
     b. 2-hexyne                      d. 4-ethyl-2-heptene               f. 3-phenylhexane               h. toluene
                                                 C-C                              C-C-C-C-C-C                     CH3
          C-C≡C-C-C-C                    C-C=C-C-C-C-C

4.   Draw the structural formulas for all the possible structural isomers of C5H10.
                                                  C                             C                   C
     C=C-C-C-C              C-C=C-C-C          C=C-C-C                    C=C-C-C                 C-C=C-C
     1-pentene              2-pentene          2-methyl-1-butene          3-methyl-1 butene       2-methyl-2-butene

5.   Calculate the mass of carbon dioxide formed when 6.0L of ethane gas undergoes complete combustion. Show the balanced
     equation. (Reminder: molar volume of a gas 22.4L) 2 C2H6 + 7 O2 → 4 CO2 +6 H2O, 24 g
6.   Give the general formulas for each of the following:
     a. aldehyde RCHO c. ether ROR                     e. ester RCOOR         g. amine RNH2
     b. alcohol ROH         d. carboxylic acid RCOOH f. halocarbon RX         h. ketone RCOR
7.   Draw the structural formulas for each of the following:
     a. benzoic acid        d. ethylhexyl ether        g. 2,2-dichloroheptane           j. 1,2-dichlorocyclopentane
                 COOH                                        Cl
                            C-C-O-C-C-C-C-C-C              C-C-C-C-C-C-C                                Cl
                                                             Cl                                        Cl
     b. 1-pentanol          e. phenol                  h. 1,2-dibromobenzene           k. propanoic acid
                                     OH                             Br                         O
     C-C-C-C-C-OH                                                                         C-C-C-OH
                                                                    Br
     c. butanal             f. methylamine             i. butyl propanoate             l. butanone
             O                                                O                            O
     C-C-C-CH                  C-N                       C-C-C-O-C-C-C-C               C-C-C-C

8.   Name the following compounds:
     a. C6H5OH phenol                        d. CH3COOC2H5 ethyl ethanoate g. CHCl3 trichloromethane
     b. CH3CH(OH)C2H5 2-butanol              e. CH3COOH ethanoic acid           h. CH3CH2OCH2CH3 diethyl ether
     c. C2H5CHO propanal                     f. C2H5NH2 ethylamine              i. CH3COCH2CH3 2-butanone
9.   Diagram the following reactions using structural formulas. Name all products AND indicate the type of reaction occurring.
     a. benzene + bromine                    + Br2                 + HBr
                                                                  Br

     b. propene + hydrofluoric acid
                                      C-C=C + HF                C-C-C-F
    c. 1-propanol + acetic acid                      O                   O
                                   C-C-C-OH +      C-C-OH              C-C-O-C-C-C + H2O

    d. ethanol + benzoic acid                            COOH                 COOC-C
                                   C-C-OH +


10. Draw the structural formulas and name the products formed from the oxidation of the following:
    a. ethane (in the presence of oxygen) c. acetone                  e. ethanal                g. tertiary-butyl alcohol
    b. ethane (in the absence of oxygen)   d. benzaldehyde            f. isopropyl alcohol
11. Diagram the addition polymerization reaction for cis-2-butene. Use structural formulas.
        C         C                  C CCCC C
    n      C=C                      -C-C-C-C-C-C-

12. Use structural formulas to diagram a condensation polymerization reaction for the following:
                                                              O O O O
    a. HOOCCH2COOH + HOCH (C6H5) OH                         -C-C-C-C-C-C-C-C- + H2O



                                                               O     O
    b.   HOOCCH2CH2COOH + H2NCH2NH2                           -C-C-C-C-N-C-N-      + H2O

    What TYPE of bond was formed in each case? covalent bond
    What is the specific name of the bond formed in reaction (b) peptide bond
13. Place in order from lowest to highest boiling points:        water, ethane, ethanal, ethanol, heptanes
         Ethane < heptane < ethanal < ethanol < water
                                                         Chapter 23 Biochemistry
1. Identify each of the following as monosaccharides, disaccharides, or polysaccharides.
    a. sucrose D            b. starch P      c. fructose M       d. maltose D       e. cellulose P    f. glucose M
2. Identify the storage form of glucose in animals and in plants. Identify the structural form of glucose in plants . starch
3. Match the monomer to the polymer (monomers can be used more than once):
    POLYMER: 1. Cellulose glucose            3. DNA nucleotide            5. starch glucose           7. RNA nucleotide
                  2. protein amino acid      4. lipid fatty acid + glycerol 6. Peptide amino acid 8. Glycogen glucose
    MONOMER: a) nucleotide           b) amino acid     c) glucose         d) fatty acid + glycerol
4. Name the bond formed between glucose molecules and the molecule produced in this reaction. Covalent bond/ glycogen
5. Name the bond formed between amino acids. Draw an example. Label the functional groups involved in this reaction.
              O        O       O             -peptide bond forms between the carbonyl and the amine
         -N-C-C-N-C-C-N-C-C- + H2O -functional groups are: amine, amide, carboxylic acid
            R       R       R
6. For each of the four major types of biological compounds, describe the polarity and solubility.
    Carbs – P, dissolves in polar covalent + ionic cmpds
    Lipids – P head but very NP tail, dissolves in NP molecules
    Protein – P, soluble in polar covalent + ionic cmpds
    Nucleic Acids - P, dissolves in polar covalent + ionic cmpds
7. For each of the four major types of biological compounds, identify the elements that are found in them, and write the general
    formula for a monosaccharide.
    Carbs – CHO             Lipids – CHO     Protein – CHON               Nucleic Acids - CHONPS
    Monosaccharides – C6H12O6

								
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