Exam for Grade 12 SCH 4U1

This is an exam to use a practice. Do not assume that all minor points are covered or all types of sample problems are addressed. If this is all you study you probably will be missing something.

Someday answers may appear but don't hold your breath! Multiple choice answers at the bottom of file. For the sections: come to class with your question and it can be dealt with.

1. Which of the following is false for a reaction pathway in the presence of a catalyst?
   a) The catalyst speeds up the rate of reaction for both the forward and the reverse process.
   b) the catalyst enables the reaction to proceed rapidly at a lower temperature.
   c) the catalyst involves the formation or an activated complex.
   d) the catalyst lowers the ΔH of the reaction.
   e) the catalyst is not used up as the reaction proceeds.
   
2. The K_sp of CdS is 3.6 x 10^-29 aat 18^oC. The concentration of cadmium ion in a satrurate solution of cadmium sulfide in moles/litre is
   a) 3.6 x 10^-29 b) 1.4 x 10^-5 c) 1.4 x 10^-6
   d) 6.0 x 10^-15 e) 6.0 x 10^-14
   
3. One hundred millilitres of 0.2 mol/L HCl are placed in a flask How many mL of 0.1 mol/L NaOH are required to bring the solution to a pH of 7?
   a) 5 mL b) 35 mL c) 50 mL d) 100 mL e) 200 mL
4. If n = 5, l = 2 then there are how many possible values of m_l ?
   a) 2 b) 3 c) 4 d) 5 e) 7
   
5. Which of these atoms has a last shell ending as 4f³? tricky
   a) vandium, b) neodymium c) arsenic d) indium e) praseodymium
   
6. This configuration represents the element 1s₂ 2s₂ 2p₆ 3s₂ 3p₆4s₂ 3d₁₀4p₆5s₂4d₅
   a) Mn b) Kr c) Ru d) Tc e) Mo
   
7. Which of these functional groups represents an aldehyde unit?
   Cann't draw in html but you should get the idea.
   
8. When you add solid ammonium nitrate (NH₄NO₃) to water, the solution feels cold to your hand. Which statement best describes this observation?
   a) NH₄NO₃(s) ---> NH₄NO₃(aq) + 6.8 kJ
   b) NH₄NO₃(s) ---> NH₄NO₃(aq) delta H = + 6.8 kJ
   c) The reaction is exothermic
   d) Heat is released from the system, so it feels cooler.
   e) The boiling point is raised so it takes more heat to keep it at room temperature, hence it feels cooler.
   
9. Use the information in the following equations to determine delta H for the combustion of acetylene.
   2C_(s) + H_{2 (g)} ---> C₂H_{2 (g)} Δ H = + 54.2 kJ
   C_(s) + O_{2 (g)} --> CO_{2 (g)} Δ H = -94.1 kJ
   H_{2 (g)} + 1/2O_{2 (g)} --> H₂O_(g) Δ H = - 57.8 kJ
   The Δ H for the burning of acetlylene, in kJ/mol is
   a) -97.7 b) -191.8, c) -206.1, d) -300.2, e) -258.0.
   
10. The molecular enthalpy of a substance is comprisee of which of the following forms of energy?
    1 Chemical potential energy, 2) nuclear energy, 3) potential energy, 4) vibrational kinetic energy
    5) translational and rotational kinetic energy
    a) 1 & 3, b) 3 & 4, c) 3 & 5, d) 1, 2, 4, & 5, e) all of the above
    
11. According to collision theory the rate of a chemical reaction is not dependent on
    a) whether or not chemical bonds are broken in the reaction
    b) the number of successful collisions between molecules per unit of time.
    c) The kinetic energy of the colliding molecules.
    d) the relative amount of surface area in contact with reactants.
    e) the mass of the colliding molecules.
    
12. Reaction rates are also affected by concentration, collision geometry, and the presence of a catalyst. Which of the following staements is false?
    a) Increasing the concentration of reacting particles increases the chance for collisions.
    b) Poor collision geometry slows the rate of reaction
    c) A catalyst increases the kinetic energy of the reacting molecules
    d) The slowest reaction involved in a reaction mechanism is the rate determing step
    e) A catalyst lowers the activation enerrgy requirement.
    
13. Which of the following will not be obseved in a system at equilibrium?
    a) constant colour. b) constant temperature. c) constant pressure, d) constant volume
    e) a constant flow of reactants entering the system and of products leaving it.
    
14. In the following reaction
    2CH₃OH(l) + 3O₂(g) ----> 4H₂O(l) + 2CO₂(g) + 1452.8 kJ
    If this reaction is reversed then the heat of reaction Δ H becomes
    a) +1452.8 kJ b) - 1452.8 kJ c) can't determine, depends on the temperature of reaction d) stays the same
    
15. Which of these molecules is unsaturated?
    a) 2-methylpropanol, b) 3-cyclohexylbutanoic acid, c) 3,3dichloro-2-methyl -1- heptene, d) 2-pentanone
    
16. Which of the following is a weak acid?
    a) HNO₃ , b) H₂SO₄, c) NH₄OH , d) H₃PO₄
    
17. In a titration an indicator is used to tell
    a) the equivalence point, b) the pH , c) the end point, d) the colour change, e) strong acids from weak acids.
    
18. You have an electrolytic cell made with chrome and lead and their appropriate solutions. The theoretical voltage for this pair of half cells is
    a) +0.74 volts b) - 0.74 volts, c) +0.40 volts d) +0.87 volts , e) +0.61 volts, f) -0.61 volts
    
19. In order to determine the rate law exponents one must
    a) Consult the correct table in the text, b) theoretically calculate using the formula, c) apply the correct Order of Reaction formula , d) can only be determined experimentally.
    
20. If the value of K (an equilibrium constant) is very small then
    a) reaction proceeds mostly to completion, b) give you a 2^nd or 3^rd order reaction, c) hardly any products are formed, d) the reaction is reversible
    
21. Which of the following isn't a property of a van der Waals forces in the solid state?
    a) low melting point, b) can evaporate at room temperature, c) hard crystals, d) occur in nonpolar molecules.
    
22. Which of the following has the greatest bond strength?
    a) covalent bond, b) Van der Waals forces, c) Hydrogen bond, d) dipole forces
    
23. How many sublevels are there in the principle energy level n = 4?
    a) 1, b) 2, c) 3, d) 4, e) 5.
    
24. Consider the following reactionat equilibrium at 1000^o C
    2CO_(g) + O_{2 (g)} ----> 2CO_{2 (g)} Δ H^o = -135.2 kcal
    Which of the following changes would not result in a larger yield of CO₂?
    a) decrease the volume, b) decrease the temperature, c) adding a catalyst, d) increase the partial pressure of CO_(g)
    e) increasing the partial pressure of O₂.
25. Phosphoric acid is a weak acid. If 1.23 g of phosphoric acid are dissolved in enough water to make 1.00 L of solution, the concentration of the solution will be
    a) 7.97 mol/L, b) 0.797 mol/L, c) 0.126 mol/L , d) 0.0126 mol/L, e) 0.00126 mol/L.
    

Written Responses

1. A quantity of 100 mL of 0.500 M HCl is mixed with 100 mL of 0.500 M NaOH in a constant pressure calorimeter having a heat capacity of 335 J/C^o. The initial temperature of the HCl and NaOH is the same 22.50^oC and the final temperature of the mixed solution is 24.90 ^oC . Calculate the heat change for the neutralization reaction. Assume that the densities and specific heats are the same as for water.

   Ans: --> -56.2 kJ/mol

   p 232
2. Describe the hybridization state of phosphorus in phosphorus pentabromide (PBr₅)
3. From the following equations and enthalpy changes,
   (a) C (graphite) + O₂(g) ---> CO₂(g) /\H^o = -395.5 kJ
4. List the different ways to write the four quantum numbers that designate an electron in a 3p orbital p 298
5. Use the VSEPR model to predict the geometry of the following molecules and ions: a) AsH₃, b) OF₂,
   c) AlCl₄^-, d) I₃^-, e) C₂H₄ p 401
6. What type(s) of intermolecular forces exist between the following pairs: a) HBr and H₂S,
   (b) Cl₂ and CBr₄, (c) I₂ and NO₃^-, (d) NH₃ and C₆H₆ P 447
7. Which of the following can form hydrogen bonds with water? CH₃OCH₃, CH₄, F^-, HCOOH and Na⁺
   p 450
8. Write the reaction rate expressions for the following reations in terms of the disappearance of the reactants and the appearance of products:
   (a) I^-(aq) + OCl^-(aq) -------> Cl^-(aq) + OI^-(aq)
   (b) 3O₂(g) --------> 2O₃(g)
   (c) $NH₃(g) + 5O₂(g) ---------> 4NO(g) + 6H₂O(g) p 549
9. The rate of the reaction A + 2B --------> C has been observed at 25^oC. from the following data, determine the rate law for the reaction and calculate the rate constant. p 552 table of data to come.
10. The conversion of cyclopropane to propene in the gas phase is a first order reaction with a rate constant of 6.7 x 10^-4s^-1at 500^oC. Calculate the half-life of this reaction.
    

    Ans; 1000 s or about 17 minutes p 557

    
    
11. Write the equilibrium constant expression K_c and K_p if applicable, for each of the following heterogeneous systems:
    (a) (NH₄)₂Se(s) <======> 2NH₃(g) + H₂Se(g)
    (b) AgCl(s) <======> Ag⁺(aq) + Cl^-(aq)
    (c) P₄(s) + 6Cl₂(g) <======> 4PCl₃(l) p 605
12. In a calorimetry experiment why is it important to know the heat capacity of the calorimeter?
13. Define these terms; enthalpy, and enthalpy of reaction.
14. Photosynthesis can be represented by the reaction
    6CO₂ + H₂O ---------> C₆H₁₂O₆ + 6O₂, where C₆H₁₂O₆ is glucose. How would you determine experimentally the /\H^o value for this reaction?
15. What hybrid orbitals are used by carbon atoms in the following species? (a) CO, (b) CO₂, (c) CN^- ion.
16. How would you distinguish between a sigma bond and a pi bond?
17. Define a dipole moment? Give an example of a bond with a dipole moment. Give and example and explain why even though a bond may have a dipole moment the molecula itself may be none polar.
18. Why are metals good conductors of heat and electricity?
19. Describe and give examples of, the following types of crystals: (a) ionic crystals, (b) covalent crystals, (c) molecular crystals, (d) metallic crystals
20. Explain the difference between physical equilibrium and chemical equilibrium. Give TWO example of each.
21. Define homogeneous equilibrium and heterogeneous equilibrium. Give TWO examples of each.
22. What do the symbols K_c and K^p represent?
23. Define reaction quotient. How does it differ from the equilibrium constant?
24. Explain or define Le Chatelier's principle. How can this principle help us maximize the yields of reactions?
25. Does the addition of a catalyst have any effects on the position of an equilibrium? Explain your answer.
26. Write the equation relating K^a for a weak acid and K^b for its conjugate base.
27. Define salt hydrolysis. Categorize salts according to how they affect the pH of a solution>
28. Define the term " common ion effect ".Explain this effect in terms of Le Chatelier's principle.
29. A photon has a wavelength of 624 nm. Calculate the energy of this photon in joules.
30. What is an energy level? Wxpalin the difference between a ground state and an excited state.
31. From the standard enthalpies of formation, calculate /\H^o for the reaction
    C₆H₁₂(l) + 9O₂ (g) -----> 6CO₂(g) + 6H₂O(l)
32. What is meant by the order of a reaction?
33. Sketch a potential energy versus reaction progress plot for the following reactions:
    a) S(s) + O₂(g) --------> SO₂(g) /\H_o = -296.06 kL
    b) Cl₂(g) ------> Cl(g) + Cl(g) /\H₂ = 242.7 kJ.
34. Calculate the pH of a buffer solution prepared by adding 20.5 g of CH₃COOH and 17.8 g of CH₃COONa to enough water to make 5.00 x 10² mL of solution.
35. The pH of a bicarbonate -carbonic acid buffer is 8.00. Calculate the ratio of the concentration of carbonic acid to that of the bicarbonate ion.
36. Calculate the solubilty in g/L of AgBr in (a) pure water and (b) in 0.0010 M NaBr.
37. The molar solubility of Pb(IO₃)₂ in a 0.10 M NaIO₃ solution is 2.4 x 10^-11 mol/L.
    What is the K_sp for Pb(IO₃)₂?
38. For these complete redox reactions below, (i) break down each reaction into its half-reactions; (ii) identify the oxidizing reagent; (iii) identify the reduceing reagent.
    
    • 2Sr + O₂ -----> 2SrO
    • 3Mg + N₂ ----> Mg₃N₂
    • Cl₂ + 2NaBr -----> 2NaCl + Br₂
    • 2CCl₄ + CrO₄^-2 ------> 2COCl₂ +CrO₂Cl₂ + 2Cl^-1
39. Give the oxidation state number for the element that is underlined
    • NaHCO₃
    • H₃AsO₃
    • K₂Cr₂O₇
40. You are given a strip of zinc and a strip of silver. Outline how you would make an galvanic cell from these two metals, appropriate ionizing solutions and the necessary lab ware. Write the equation for the oxidation, reduction reactions and the overall reaction. What is the theoretical voltage of this system ( A diagram would be most helpful in your answer.)
41. Balance these redox equations by the ion-electron method.
    • Mn⁺² + H₂O₂ -----> MnO₂ in basic solution
    • S₂O₃^-2 + I₂ ------> I^-1 + S₄O₆^-2 in acid solution
    • Cu (s) + HNO₃(aq) ---------> Cu(NO₃)₂(aq) + H₂O(l) + NO₂(g)
    • Zn(s) + NO₃(aq) ----------> Zn⁺²(aq) + NH₄⁺¹(aq) + H₂O(l)
42. Draw & Label a diagram of a standard hydrogen electrode. What is the voltage of this electrode?
43. Arrange the following species in order of increaseing strength as oxidizing agents: MnO₄^-1 (in acid solution), Sn⁺², Al⁺³, Co⁺³ and Ag⁺¹. Assume species are in standard states.
44. Outline what corrosion is; give an example with balanced chemical equations.
45. Anode? Cathode? Define these terms and why are they relative to the material under discussion?
46. A galvanic cell consists of a Mg electrode in a 1.0 M Mg(NO₃)₂ solution and a Cu electrode in a 1.0 M solution of CU(NO₃)₂ solution. Calculate the standard emf of this electrochemical cell at 25^oC.

Chem II Final Exam #3

1. Complete and balance the following redox equation:

ClO₃^-1 + Cr(OH)₃ -----> CrO₄^-2 + ClO₂^-1 (In basic soln.)

Determine the electrochemical potential of the reaction at a pH of 8.2 when the concentration of all other ionic species is at 0.1 M.

2. A two gram sample of iron ore was dissolved in acid and all iron present was reduced to Fe²⁺ The resultant solution was titrated with a 0.5 M soln of permanganate. The volume required to reach the endpoint was 35.71 ml. During the titration Fe²⁺ was oxidized to Fe³⁺ and MnO₄^1- was reduced to Mn²⁺. Determine the percentage of iron, as the metal, present in the sample.

3. A buffer solution is made up of equal volumes of 0.1 M acetic acid and 0.5M sodium acetate. What is the pH of the solution?

4. In an experiment a sample of sodium chlorate was 95% decomposed in 48 minutes. Approximately how long would this decomposition have taken if the same sample had been heated at a temperature 20 degrees celcius higher?

5. Hydrogen chloride, HCl, and ammonia, NH₃, escape from bottles of their solutions and react to form the white glaze often seen on glass in chemistry laboratories.

a) Calculate the free energy change, /\H & /\G, for this reaction.
b) At what temperature will /\G for the reaction be equal to zero?

6. The equilibrium constant K for the reaction

is 0.0211 at a certain temperature. What are the equilibrium concentrations of PCl₅, PCl₃, and Cl₂ starting with a concentration of PCl₅ of 1.00M?

7. At what temperature does the decomposition of dinitrogen trioxide become spontaneous. The unbalanced reaction is given below. (All substances are in the gaseous state.)

N₂O₃ ------> NO + NO₂

8. Given the following relationship: log N ₀ - log N = 0.301 * t/t 1/2 . .
What information can be gained from the slope and the y-intercept of the line obtained when one plots log N vs. t?

9. The data in the table below apply to the following reaction:

Determine the rate law for this reaction, find your m & n's, then the overall reaction order, then the value of k

[A]	[B]	[C]	Rate
0.2	0.4	0.1	1
0.4	0.4	0.2	8
0.2	0.2	0.2	1
0.4	0.4	0.1	4

Multiple Choice Answers

Answers to select questions

• # 1 Calculate the heat of the solution using Q = m/\tc. do this for each individual solution.
  Q_sol = 2.01x10³J. Find Q for the calorimeter. Q_cal = 804 J
  Add the two together. Q = 2.81x10³J or 2.81kJ
  Both HCl and NaOH have the same concentration 0.5 mol/L which is in 100 mL of solution which translates into 0.05 moles
  Therefore the heat of neutralization per mole is 2.81 kJ/ 0.05 mol or 56.2 kJ/mol.
  

• #10 Half-lives Page 752
  Use the formula 18.10 page 753 with k = 6.7x10^-4 and solve for t_1/2.
  

More Practice Exam Questions

Final Examination

December 11, 1996

Choose the best answer and be sure to mark only one choice for each question. Each question is worth 4 points. There are a total of 27 questions.