CHEMISTRY – Second Semester EXAM Review

PART I: Multiple Choice - Complete this part on the Scantron form.  You may write on the pages but only the answers on the Scantron will be counted.  Use your periodic table and calculator as necessary.

1.      Raising the temperature of a gas in a fixed volume container will most likely change the  (a) number of particles in the container.  (b) density of the gas in the container.  (c) pressure exerted by the gas in the container.  (d) size of the gas particles in the container.

2.      Gases do not form liquids, unless the temperature is lowered, because of the  (a) extremely small size of the gas particles.  (b) relatively slow motion of the gas particles.  (c) weak attractive forces between particles.  (d) presence of repulsive forces between particles.

3.      If the amount and temperature of a gas are kept constant, the pressure and volume of the gas are
(a) also unchanging.  (b) directly proportional.  (c) inversely proportional.  (d) independent of one another.

4.      The law of partial pressures of gases was formulated by  (a) Avogadro.  (b) Boyle.  (c) Charles.      (d) Dalton.

5.      The ideal gas law describes the behavior of real gases under  (a) all conditions of temperature and pressure.  (b) relatively high temperature and low-pressure conditions.  (c) low temperatures and high-pressure conditions.  (d) any specified temperature and pressure conditions.

6.      Compared to the particles of a gas, the particles of a liquid  (a) are not in constant motion.
(b) collide with the sides of the container, but not with each other.  (c) can take the shape but not volume of their container.  (d) can move longer distances.

7.      A characteristic that distinguishes solids from liquids is the solid’s  (a) fixed shape.  (b) greater diffusion of particles.  (c) absence of particle movement.  (d) weaker attraction between particles.

8.      If the phases of matter are arranged in order of increasing disorder, the arrangement would be        (a) solid, liquid, gas.  (b) gas, solid, liquid.  (c) gas, liquid, solid.  (d) liquid, solid, gas.

9.      The strength of induced dipole attractions  (a) increases with the size of the molecules.
(b) is stronger in molecules or atoms with a large number of electrons.  (c) is the explanation for the increasingly higher boiling point of the noble gases going down the group.  (d) is characterized by a, b and c.  (e) is characterized by none of the previous answers.

10.  Dipole-dipole forces would most likely be found in compounds made up of atoms of  (a) the same element.  (b) different atomic radius.  (c) different electronegativity.  (d) the same mass.

For questions 11-16 refer to the figure at the right. (This was a phase diagram)

11.  Where do the three states of matter exist
in equilibrium?  (a) only at the origin
(b) only at the triple point  (c) on any
solid line  (d) between any two solid lines

12.  In which region will water exist as a liquid?
(a) a  (b) b  (c) c  (d) in all three

13.  What state(s) of water would be present at
15oC and 20 mm Hg?  (a) solid only
(b) liquid only  (c) gas only
(d) liquid and gas

14.  Where would you find boiling water?                                              
(a) at the intersection of all three solid lines
(b) in area b  (c) at any point along the line separating regions b and c
(d) at any point along the line separating regions a and c

15.  If the temperature of a sample at 0oC and 25 mm Hg is reduced at constant pressure, what would happen to the sample?  (a) The sample would melt.  (b) The sample would freeze to solid ice.
(c) The sample remains in equilibrium between solid and liquid.  (d) The sample would not change.

16.  How could you convert a sample of water at 10oC and 20 mm Hg to vapor without increasing the temperature?  (a) Decrease the pressure to 15 mm Hg.  (b) Decrease the pressure to 5 mm Hg, so conditions for the sample fall below the solid line.  (c) Decrease the pressure to that of the triple point of water.  (d) It can’t be done.

17.  A solution that contains as much solute as can possibly be dissolved under the existing conditions is said to be:  (a) saturated.  (b) supersaturated.  (c) soluble.  (d) concentrated.

18.  Which of the following is not an important factor influencing solubility?  (a) chemical nature of solute  (b) temperature  (c) chemical nature of solvent  (d) volume of solvent

19.  The rate at which a solid solute can be dissolved in a liquid solvent can be increased by 
(a) lowering the temperature of the solvent.  (b) grinding the solute into small pieces. 
(c) increasing the air pressure on the liquid.  (d) lowering the temperature of the solute.

23.  How many ions will 3 formula units of CaCl2 produce when dissolved in water?  (a) none  (b) one  (c) three  (d) six  (e) nine

Use a solubility table to answer questions 24-27

24.  Using the figure at right, what will happen when fairly concentrated sodium chloride and lead (II) nitrate solutions are mixed?  (a) Nothing will happen.  (b) Sodium nitrate will form a precipitate.  (c) Lead (II) chloride will form a precipitate.  (d) Chlorine gas will be produced.

25.  Which cations listed in the figure at right would not tend to form precipitates with sulfate ions?
(a) Group 1A only  (b) Group 1A, silver and lead ions only  (c) Group 1A and ammonium ions only  (d) ammonium, silver and lead ions only

26.  Which of the following ions would always be spectator ions in precipitation reactions? (a) Group 1A only  (b) Group 1A, sulfate and nitrate ions only  (c) Group 1A and ammonium ions only  (d) Group 1A, ammonium and nitrate ions only

27.  What net ionic equation represents the reaction that occurs that occurs when fairly concentrated solutions of barium sulfide and lead nitrate are mixed?  (a) Ba2+(aq) + 2NO3-(aq) à Ba(NO3)2(s)
(b) Pb2+(aq) + S2-(aq) à PbS(s)   (c) Ba2+(aq) + SO42-(aq) à BaSO4(s)
(d) Ba2+(aq) + 2NO3-(aq) + Pb2+(aq) + S2-(aq) à PbS(s) + Ba(NO3)2(s)

28.  In the Bronsted-Lowry definition, an acid is a substance that  (a) donates a proton.  (b) accepts a proton.  (c) produces H3O+ ions in solution.  (d) accepts a pair of electrons.

29.  The reaction of many metals with acids produces the gas  (a) oxygen.  (b) nitrogen.  (c) neon. 
(d) hydrogen.

30.  The relative strengths of acids and bases are determined by the extent to which the acid or base  
(a) react with materials.  (b) eat holes in ones clothing.  (c) ionize in water.  (d) react with each other.

31.  How would you describe the solution of the salt formed from the reaction of a weak acid and a strong base?  (a) neutral  (b) acidic  (c) basic  (d) need more information

32.  An acidic hydrogen is usually bonded to elements that attract the bonding electrons  (a) very weakly.  (b) somewhat weakly.  (c) very strongly.  (d) to the same degree as the hydrogen attracts the electrons.

Use the chart below to answer questions 33-35

Acid                 Reaction                                                           Ka (at 25oC)
hydrocyanic      HCN + H2O çè  H3O+  +  CN-                    Ka = 4 x 10-10
acetic               HC2H3O2  + H2O çè  H3O+  +  C2H3O2-      Ka = 1.8 x 10-5
hydrofluoric      HF + H2O çè  H3O+  +  F-                           Ka = 6.7 x 10-4
phosphoric       H3PO4  + H2O çè  H3O+  +  H2PO4-            Ka = 7.1 x 10-3
carbonic           H2CO3  + H2O çè  H3O+  +  HCO3-            Ka = 4.4 x 10-7

33.  Identify the strongest acid from the list above.  (a) HCN  (b) HF  (c) H3PO4  (d) H2CO3

34.  What would you expect to find after equilibrium is reached in the dissociation of hydrofluoric acid? 
(a) an equal mix of products and reactants  (b) mostly reactants  (c) mostly products  (d) all products

35.  Using the above table which acid with an initial 1.0 M concentration would produce the lowest pH?  (a) HCN  (b) HC2H3O2  (c) HF (d) H3PO4

36.  If the pH of a solution changes from 10 to 11,  (a) hydronium ions increase by a factor of 10. 
(b) hydroxide ions increase by a factor of 10.  (c) hydronium ions decrease by ½.  (d) hydroxide ions double.

37.  The higher the pH of a solution,  (a) the more acidic it is.  (b) the greater [OH-].  (c) the less basic it is.  (d) both answers a and c apply.

38.  How can you determine the concentration of a weak acid in water?  (a) Measure the pH with a pH meter.  (b) Observe the color change of litmus paper.  (c) Determine the extent of electrical conductivity.  (d) Perform an acid-base titration.

39.  Endothermic reactions are reactions that  (a) release heat.  (b) absorb heat.  (c)  do not involve heat.  (d) take place instantaneously.

40.  In a balanced exothermic reaction, where does a heat term appear?  (a) on the product side only 
(b) on the reactant side only  (c) on both sides  (d) on neither side

41.  What principle relates the enthalpy change for a net reaction to the enthalpy changes of a series of summed reactions?  (a) Boyle’s Law  (b) Hess’s Law  (c) Avogadro’s hypothesis  (d) Henry’s Law

42.  What is the symbol of entropy change?  (a) DS  (b) DG  (c) DH  (d) DT

43.  Whenever products have more disorder than reactants,  (a) DH is negative.  (b) DH is positive. 
(c) DS is negative.  (d) DS is positive.

44.  DG is equal to:  (a) DH + TDS.  (b) DH - TDS.  (c) DS + TDH.  (d) DS - TDH

48.  What is true of the spontaneity of the melting of a solid?  (a) It is always spontaneous.  (b) It is spontaneous at high temperatures.  (c) It is spontaneous at low temperatures.  (d) It is probably not spontaneous.

49.  What is happening to the enthalpy and to the entropy of the following system: 
heat + S(g) + O2(g) à SO2(g)   (a) There is an increase in enthalpy and a decrease in entropy.
(b) There is a decrease in enthalpy and an increase in entropy.  (c) There is a decrease in both enthalpy and entropy.  (d) There is an increase in both enthalpy and entropy.

 

PART II - PROBLEM SOLVING or SHORT ANSWER -- Answer the following problems in the space provided. Neatly show your work and clearly indicate your answers.  Answer short answer questions in phrases or sentences or as otherwise indicated.

1)      Write the balanced overall equation and balanced net ionic equation for what occurs when a saturated solution of potassium sulfide (K2S) is mixed with a saturated solution of chromium(III) nitrate (Cr(NO3)3), given that the sulfides of all transition metals are insoluble, and all nitrates are soluble.

a)      overall:

b)      net ionic:

2)      How many grams of Na2CO3 are required to prepare 250. mL of a 0.158 M solution?

3)      Answer the following questions related to gases in this chemical equation:
                  C3H8(g)  +  5 O2(g)  à  4 H2O(l)  +  3 CO2(g)
      [R= 8.31 (L-kPa/mol-K), 0.0821 (L-atm/mol-K), 62.4 (L-mmHg/mol-K), STP = 0oC, 1 atm]

(a)    How many moles of C3H8 are contained in a 25.0-liter container at 23oC at 775 mmHg pressure?

(b)   How many molecules of O2 are required to completely react the moles of C3H8 in part (a)?

(c)    Assuming all of the C3H8 and O2 react in parts (a) and (b), how many liters of CO2 would be produced under STP conditions?

(d)   What is the pressure exerted by 35.7 grams of CO2 at 150.0oC in a rigid container whose volume is 25.0 liters?

4.      How does the pressure of an enclosed, fixed-volume gas change when the gas is heated?  Explain why this change occurs using the kinetic particle theory of gases.

5)      In the acid-base reactions identify the following:

a)      the acid and its conjugate base:  H2PO4-  + OH- ßà HPO42-  + H2O

b)      the base and its conjugate acid:  HF  +  NH3  ßà  F-  +  NH4+ 

6)      Human blood has a pH of 7.45.  What is the [H3O+], [OH-] and pOH of human blood.

7)      The following concentrations were measured for CH3COOH at equilibrium:
[CH3COOH] = 1.24 M, [H3O+] = 0.0047 M and [CH3COO-] = 0.0047 M.
Write a chemical equation for the dissociation of CH3COOH in water.

(a)    Write the Ka expression for this acid.

(b)    Calculate the Ka of this acid?

(c)    If you were to measure the equilibrium concentrations of these same species of a solution of CH3COOH that had a concentration of 0.50 M, would you expect the Ka to be different or the same?  Explain.

8) Given the equation 2 SO2 (g) + O2 (g) à 2 SO3 (g), DH = -197.8 kJ, answer the following questions:

(a)    Is this reaction endo- or exothermic?

(b)   Do the reactants or products have greater internal energy?

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