G&OC 2010 S1 Final Examination.docx

ID Number_______________

Faculty of Health and Environmental Sciences

School of Applied Sciences

General and Organic Chemistry

775205

Final Examination

June 2010

Time Allowed: 2 Hours plus 5 minutes reading time.

Instructions:

1. Answer all questions in sections A and B and 5 of the 8 questions in section C.

2. Section A is Multichoice. Answer the questions on the answer sheet at the end of Section A. Leave this sheet attached to the exam paper

3. Answer sections B and C in the spaces provided. DO NOT WRITE IN PENCIL

4. Calculators may be used.

5. Make sure you write your ID Number at the top of this page and on the Multichoice answer sheet.

6. Total number of pages including this page = 16

General and Organic Chemistry Examination S1 2010

Section A (Each question is worth 1 mark)

Multiple ChoiceIdentify the choice that best completes the statement or answers the question and circle the choice on the answer sheet provided at the end of this section.

1. Which one of the following is a chemical change?

a. Distillation of an organic product.b. Making ice cubes.c. Melting Lead.

d. Digesting food.

2. In which one of the following are the lengths given in the correct order?

a. cm < m < km < mmb. km < m < cm < mmc. mm < cm < m < kmd. cm < mm < m < km

3. Which one of the following is true?

a. 0.075 kg = 75 mgb. 0.075 kg = 750 mgc. 0.075 kg = 7500 mgd. 0.075 kg = 75000 mg

4. The atomic nucleus described by the symbol 35Cl contains: 17

a. 35 neutrons and 18 protons.b. 18 neutrons and 35 protons.c. 17 protons and 18 neutrons.d. 18 protons and 17neutrons.

5. How many orbitals are there in the 3d subshell?

a. 1 b. 2 c. 4 d. 5

6. Sulphur has an Atomic Number of 16. Its electron configuration is:

a. 1s22s22p43s23p6

b. 1s22s22p63s23p4

c. 1s22s22p43s43p4


d.. 1s22s22p23s23p63d2

7. Atoms of Uranium 234U decay by α particle emission. What is the Mass Number of the atom formed as a result? 92

a. 230 b. 231 c. 232 d. 234

8. Atoms of Lead 212Pb decay by β particle emission forming atoms of Bismuth.

What is the Atomic Number of Bi?

a. 83 b. 82 c. 81 d. 80

9. For the reaction 3FeS(s) + 2H3PO4(aq) → Fe3(PO4)2(s) + 3H2S(g), how many moles of H3PO4 are required to produce 1 mole of H2S?

a. 0.67 b. 1.5 c. 2 d. 3

10. If 20 g of a compound is dissolved in water to make 250 mL of a 0.5 mol L-1 solution, the Molar Mass of the compound in g mol-1 is:

a. 10 b. 40 c. 80 d. 160

11. The rates of all chemical reactions increase with a rise in temperature. Which one of the following is correct?

This statement is: a. True, since the activation energies for all reactions increase with a rise in temperature. b. False, since the rate of an exothermic reaction decreases with a rise in temperature. c. False, since the rate of an endothermic reaction decreases with a rise in temperature. d. True, since a greater proportion of molecular collisions will result in reaction.

12. A particular reaction carried out at 60oC takes 5 minutes. Approximately how long will it take at 30oC ?

a. 10 minutes b. 20 minutes c. 40 minutes d. 21/2 minutes

13. In which one of the following reactions does an element show the greatest increase in oxidation number?

a. MnO4- + 8H+ + 5Fe2+ → Mn2+ + 5Fe3+ + 4H2O

b. 2Al + 6H+ → 2Al3+ + 3H2

c. SO2 + 2H2S → 3S + 2H2O d Cl2 + Sn2+ → 2Cl- + Sn4+


14. Which one of the following will occur if the temperature of a gas is increased from 30oC to 60oC?

a. The volume will doubleb. The volume will increase by a small amountc. The volume will decrease by a small amountd. The volume will be halved

15. Which property of water is due to it being a polar molecule?

a. Its boiling point. b. Its ability as a solvent c. Its high specific heat d. All of the above.

16. Which one of the following acids has the weakest conjugate base?

a. H3PO4 b. H2SO4 c. H2CO3 d. CH3COOH

17. Which one of the following cannot act as a base?

a. H2O b. HCO3- c. H2CO3 d. H2PO4

-

18. In an aqueous solution the [H3O+] is 2.0 x 10-6 mol L-1. The [OH-] in mol L-1 is:

a. 5.0 x 10-9 b. 1.0 x 10-7 c. 2.0 x 10-6 d. 1.0 x 108

19. The addition of acid to a solution at pH = 5 changes the pH to pH = 1. The concentration of H3O+ has increased by:

a. 4 b. 10 c. 100 d. 10,000

20. Which one of the following would not form a buffer solution?

a. H3PO4 and NaH2PO4 b. H2CO3 and NaHCO3 c. HCl and NaCl d. NH3 and NH4Cl


21. Which one of the following organic molecules is the most polar?

a. Propanone.b. Propanoic acid.c. Propanal.d. 1-Propanol.

22. Which one of the following hydrocarbons would have geometric (cis-trans) isomers?

a. CH3–CH2–CH=CH–CH3 b. CH3–CH2–CH2–CH=CH2

c. CH3–C≡C–CH2–CH2–CH3 d. CH3–CH2–CH2–CH2–CH3

23. What type of alcohol is formed by the addition of H2O to 1-butene?

a. Primary.b. Secondary.c. Tertiary.d. None of these, the product is not an alcohol.

24. In which class of compounds has the carbon been most highly oxidized?a. Carboxylic acids.b. Aldehydes.c. Ketones. d. Alcohols

25. How many polar covalent bonds are found in a carboxylic acid?

a. 1 b. 2 c. 3 d. 4

_____________________End of Section A ________________________


Multichoice Answer Sheet ID Number_______________Circle the answer you have chosen that best answers the question.Leave this sheet attached to the question paper

1. a b c d

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24. a b c d

25. a b c d


Section B (30 marks)Write your answers in the spaces provided on this question paper. Show working in calculations.

1. Name the following compounds.

(a) BaSO3 ____________________________________

(b) Ca(HCO3)2 ____________________________________

(c) NH4SCN ____________________________________

(d) CH3COOK ________________________________ ____ (4 marks)

2. Give the formulas of the following compounds

(a) Sodium phosphate ______________

(b) Mercuric nitrate ______________

(c) Ferric sulphate ______________

(d) Silver chromate ______________ (4 marks)

3. (i) Write a balanced molecular equation for the reaction between potassium iodide and lead nitrate solutions.

+ → + (2 marks)

(ii) Rewrite the equation in (i) as an ionic equation leaving out the spectator ions.

→ (1 mark)

(iii) What is the colour of the precipitate? ____________________ (1 mark)

4. Write the Lewis electron dot structures for the following molecules using lines for bonds and dots for non bonding electrons.

(i) Dichloromethane CH2Cl2

(2 marks)

(ii) Methanoic acid HCOOH

(2 marks)


5. (a) State the Bronsted-Lowry definition of an acid and a base.

An acid is a ______________________ A base is a _______________________ (1 mark)

(b) Write the equation for the reaction with water of the dihydrogen phosphate ion, H2PO4-

when it is acting

(i) as an acid: H2PO4- + H2O →

(ii) as a base: H2PO4- + H2O →

(3 marks)

(c) Substances which can act as both an acid and a base are called __________________

(1 mark)

6. For each of the following equations state whether the vanadium species undergoes oxidation reduction or neither. Justify your answer.

(a) V2O5 + 2H+ → 2VO2+ + H2O

_________________ because________________________________________________

(b) 2VO2+ + Zn + 4H+ → 2V3+ + Zn2+ + 2H2O

_________________ because _______________________________________________

(3 marks)7. Name the following organic compounds:

(a) CH3–CH2–CH2–CH–CH3 (b) CH3–CH2–CH2–C–CH3

│ || OH O _________________________ ________________________

(c) CH3–CH2–CH2–CH2–C–H (d) CH3–CH2–CH2– CH2–C–OH || || O O _________________________ _________________________

(e) CH3–CH2–C–O–CH2–CH3

|| O

______________________________________ (6 marks)

_____________________End of Section B_______________________


Section C (Longer answer questions)

Answer FIVE of the eight questions only. (All worth 9 marks each)

1. The graph below shows the titration curves for the titrations of 20.0 mL aliquots of 0.100 mol L-1 HCl and 0.100 mol L-1 propanoic acid with dilute NaOH.

(a) Explain why:

(i) The initial pH for propanoic acid is higher than for HCl. _____________________

_________________________________________________________________

_________________________________________________________________

. (ii) Both titrations require the same volume of NaOH to reach the equivalence point.

________________________________________________________________

________________________________________________________________

(iii) Both titrations have the same pH values after the equivalence point..

________________________________________________________________

(b) From the equivalence point on the graph calculate the concentration of the NaOH.

nHCl = x = mol

nNaOH in ________mL = because _______________

(Equiv Point)

[NaOH] = ____________________mol = __________ mol L-1


L contd

1. (c) Show on the diagram how to determine the approximate pKa of propanoic acid and so calculate an approximate value for Ka.

pKa ___________ therefore Ka = ______________________

2. (a) Calculate the pH of a 0.1 mol L-1 solution of acetic acid given Ka = 1.75 x 10-5

and for a weak acid [H3O+] = √KaMa

(b) Morphine is a naturally occurring weak base used to relieve pain. Calculate the pH of a 0.1 mol L-1 solution of morphine given Kb = 1.60 x 10-6 and for a weak base [OH-] = √KbMb.

(c) Blood contains a carbonate buffer of H2CO3 and HCO3- .

Given Ka for H2CO3 = 4.47 x 10-7 = ##### Calculate the pH of a buffer of 0.05 mol L-1 H2CO3 and 0.1 mol L-1 HCO3

- .


3. 0.350g of a steel sample was analysed for its manganese content by dissolving it in acid

and oxidising the Mn2+ to MnO4- and making the solution up to 100 mL.

The absorbance of this solution was measured at 525nm together with the absorbances of a series of standard KMnO4 solutions.

The following results were obtained:

[KMnO4] / mol L-1 Absorbance 0 0.000 0.002 0.074 0.004 0.152 0.006 0.224 0.008 0.300 Steel Solution 0.180

(a) On the graph paper below plot a suitable graph and determine the concentration of MnO4

- in the steel solution In mol L-1.

(b) From (a) calculate the moles of Mn (= moles of MnO4

-) in 100 mL and the mass of Mn (54.9 g mol-1) in 100 mL (= mass of Mn in the steel sample).

(c) Calculate the % Mn (54.9g mol-1) in the steel sample.

(d) What is the general name given to: (i) the graph you have plotted? _______________________________________


(ii) the wavelength of 525 nm which was used? _________________

4. (a) Given the following table of standard reduction potentials:

Eo / Volts

Fe3+ + e- → Fe2+ +0.77

I2 + 2e- → 2I- +0.54 Cu2+ + 2e- → Cu +0.34 Sn2+ + 2e - → Sn -0.14

Write a balanced ionic equation for the reaction between the following pairs of atoms or Ions. If no reaction occurs write “no reaction” and state why..

(i) Iron (III) ions and iodide ions._______________________________________

(ii) Tin (II) ions and copper. ___________________________________________

(iii) Iron (II) ions and copper.___________________________________________

(b)

DIAGRAM

For the cell pictured above:

(i) Which electrode is the cathode? ____________ because _________________

_______________________________________________________________

(ii) What is the direction of the electron flow when the switch is closed? _________

_________________________because _______________________________

(iii) Write the equation for the cell reaction. + → +

(iv) If the solutions are 1 mol L-1 calculate the cell voltage


5. (a) Pure hydrogen was collected in a 5.0 litre bulb which had previously been evacuated. At 298K the pressure in the bulb was 55 kPa. (i) Calculate the moles of hydrogen in the bulb. ( PV = nRT. R = 8.314 J K-1 mol -1.)

(ii) Calculate the mass of hydrogen in the bulb. (H = 1 g mol-1)

(b) The first step in the manufacture of nitric acid is the oxidation of ammonia to nitric oxide. The reaction is an equilibrium: 4NH3(g) + 5O2(g) ⇌ 4NO(g) + 6H2O(g) + heat

(i) Write the equilibrium constant expression for this reaction.

K = [ ] [ ]

[ ] [ ] (ii) Why is the reaction carried out at a reduced pressure? ______________________

__________________________________________________________________

(iii) What is the effect on the yield of nitric oxide if the temperature is increased?

_______________ because ___________________________________________

(iv) Bearing in mind your answer to (iii), why do you think the reaction is carried out at 850oC rather than 100oC?

__________________________________________________________________

(v) The ammonia and oxygen are passed over a platinum gauze catalyst.

What effect would the catalyst have on the concentration of nitric oxide?

__________________________ because ________________________________

__________________________________________________________________

Why is a catalyst used on a reaction which is an equilibrium?

__________________________________________________________________


6. (a) The Specific Heat of iron is 0.44 J g-1 oC-1 and of water is 4.18 J g-1 oC-1.

A 10 g nail heated to 500oC was placed in 250 mL (= 250 g) of water and the final temperature of the system was 25oC.

(i) Calculate the amount of heat (Q) lost by the nail. (Q = SH x m x ∆T).

(ii) The answer to (i) will also be the heat gained by the water. Calculate the temperature of the water before the nail was placed in it.

(b) Iron reacts rather slowly with oxygen at room temperature (rusting). A student thought he could speed up the reaction and make the iron burn by

heating a nail and dropping it in to a gas jar of oxygen but the nail did not burn. Suggest an improvement to the experiment to test if iron will burn and explain

why your experiment is better.

______________________________________________________________

______________________________________________________________

(c) A bottle of radioactive material was purchased by a radiotherapy laboratory butwas mislaid in a storeroom. When the storeroom was cleaned sometime later, the bottle was found. Analysis with a Geiger Counter showed the bottle contained 1.6 g of the radioactive isotope but the label on the bottle showed the mass was 50 g at the time of delivery. How long was the bottle mislaid for if the half-life of the radioisotope is 30 days?


7. Complete the following equations and name the products.

(a) CH3–CH2–CH=CH2

HBr

_____________________

(b) CH3–CH2–CH2–CH2–OH

Cr2O7

2−

H+

_____________________

(c) CH3–CH2–C–OH + CH3–OH conc

|| H2SO4

O

________________________________

(d) CH3

│

CH3 –C–CH2 –CH3 conc

│ HCl OH

_____________________________

(e) CH3 –C–CH3 NaBH4 ||

O

_____________________


8. (a) A liquid organic compound has the Molecular Formula C3H6O.

(i) If it forms an orange precipitate with Brady’s Reagent draw and name two possible structural formulas for the compound.

(ii) Describe ,giving the reagent and observation, a chemical test which would distinguish between the two possible compounds. ________________________________________________________________

_

_________________________________________________________________

(b) An unknown organic liquid was tested as described below. State the inference from each test.

(i) A small amount of the liquid burned with a non-smokey flame. ______________

(ii) The liquid was soluble in water. ______________

(iii) A colourless gas was formed when the liquid was added to sodium hydrogen carbonate solution. ____________________________

(iv) A sweet smell was detected when some of the liquid was heated with ethanol and 2 drops of concentrated H2SO4. __________________________________

(v) Give three physical properties which be used to identify the name of the unknown liquid. __________________________________________________________

______________________End of Section C _______________________________

Reminder – Answer FIVE questions only from Section C


17 17 17Cl 17 17Cl

92 92 92U 92 92U

82 82 82Pb 82Pb


G&OC 2010 S1 Final Examination.docx

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