school of chemistry university of kwazulu...

School of Chemistry

UNIVERSITY OF KWAZULU-NATAL, WESTVILLE

JUNE 2009 EXAMINATIONS

CHEM 110/CHEM 195: GENERAL PRINCIPLES OF CHEMISTRY

DURATION: 3 HOURS TOTAL MARKS: 100

Moderation Board: Prof J Field Dr M Bala

Dr W van Zyl

University of KwaZulu-Natal

Internal Examiners: Dr P Coombes Ms G Dawson

Prof G Kruger Dr G Maguire

Dr B Moodley Dr Nyamori

Ms L Pillay

IMPORTANT: COMPLETE THIS SECTION IMMEDIATELY

SURNAME AND INITIALS (Optional):___________________________

SIGNATURE: _____________________________________________

Seat

Number:

Student Number:

NOTE: This paper consists of 36 pages, including a periodic table and data sheet. Please

check that you have them all. The use of calculators is permitted. Answer ALL questions on

the question paper and where necessary over the page.

Students are requested, in their own interests, to write legibly.

Question 1 2 3 4 5 6 7 8 9 Total

Internal Marks

External Marks

2School of Chemistry, University of KwaZulu-Natal, Westville campus

CHEM 110W/CHEM 195: General Principles of Chemistry

June 2009 Examination

Section A: Multiple Choice Questions

Answer this section by filling in the correct letter on your MCQ sheet.

1. Oxygen is:

A) a compound.

B) a mixture.

C) an element.

D) always combined with hydrogen. (1)

2. Ethylene glycol has a density of 1.11 g mL-1. What is the volume occupied by

30.0 g of ethylene glycol?

A) 31.1 mL

B) 0.037 mL

C) 30.0 mL

D) 27.0 mL (1)

3. Choose the INCORRECT statement.

A) An anion is a positive ion.

B) A molecule is a group of bonded atoms that exist as an entity.

C) Ionic compounds result from combinations of metals and nonmetals.

D) The molecular formula is the listing of the atoms in an actual

molecule. (1)




Multiple Choice Questions (Continued)

4. Which of the following contains the greatest number of atoms?

A) 1 mole of S8 molecules.

B) 2 moles of P4 molecules.

C) 4 moles of chlorine molecules.

D) All of these contain the same number of atoms. (1)

5. Choose the INCORRECT name/formula combination

A) Sr(IO3)2 strontium iodate

B) NaClO sodium hypochlorite

C) NaClO4 sodium chlorate

D) NaClO2 sodium chlorite (1)

6. Determine the mass percent of H in NH4Cl.

A) 67 %

B) 7.5 %

C) 3.6 %

D) 14 % (1)





7. What is the formula of sodium bicarbonate?

A) Na2CO3

B) NaHCO3

C) NaCO2

D) NaHSO4 (1)

8. Which of these compounds is insoluble in water?

A) Fe(OH)2

B) MnCl2

C) Ni(NO3)2

D) K2S (1)

9. The oxidation state for the element in the reactant that is oxidised in the equation

below is:

4 Al (s) + 3 O2 (g) → 2 Al2O3 (s)

A) 0

B) - 2

C) + 2

D) + 3 (1)





10. Identify a neutralization reaction:

A) AgNO3 (aq) + HCl (aq) → AgCl (s) + HNO3 (aq)

B) 2 HI (aq) + H2O2 (aq) → I2 (s) +2H2O (l)

C) NaOH (aq) + HCl (aq) → NaCl (aq) + H2O(l)

D) MgO (s) + H2O (l) → Mg(OH)2 (s) (1)

11. Identify the reducing agent in the following reaction:

NO3- (aq) + S8 (s) → NO2 (g) + SO4

2- (aq)

A) NO3-

B) S8

C) NO2

D) SO42- (1)

12. A light photon of wavelength 500 nm, when compared to light of wavelength

600 nm, has:

A) a higher frequency.

B) lower energy.

C) a greater velocity.

D) a shorter wavelength. (1)





13. Which electron configuration is impossible?

A) 1s22s

22p

63s

2

B) 1s22s

22p

63s

23p

6

C) 1s22s

22p

62d

2

D) 1s22s

22p

53s

1 (1)

14. A bismuth atom has one more electron than a lead atom. Into which energy sub–

level does this added electron go?

A) 5p

B) 6p

C) 6s

D) 7s (1)

15. The electron configuration of the two outer sublevels of vanadium, element

number 23, is:

A) 3d24s

3

B) 4s24p

3

C) 3d34s

2

D) 3d44s

1 (1)





16. Which contains both covalent and ionic bonds?

A) NH4NO3

B) BaCl2

C) NF3

D) CH2O (1)

17. How many resonance forms are possible for the Lewis formula of SO2?

A) 3

B) 0

C) 4

D) 2 (1)

18. Which one of the following Lewis structures is INCORRECT?

A) N N

B)

C)

D) (1)

H N O

Be HH

C O





19. The Lewis dot structure represents one of the resonance

forms of the thiocyanate ion. The formal charges on the N, C and S atoms,

respectively, are:

N C

S

A) - 1, 0, 0

B) + 2, 0, - 1

C) 0, 0, - 1

D) - 2, 0, - 1 (1)

20. Which one of the following molecules is paramagnetic?

A) N2

B) F2

C) NO

D) HF (1)

21. Which one of the following is an exothermic process?

A) Ice melting.

B) Water evaporating.

C) Boiling soup.

D) Condensation of water vapor. (1)





22. Standard conditions for thermochemistry are,

(a) P = 1 atm

(b) some common temperature, usually 298 K

(c) V = 1 L

A) a only.

B) b only.

C) a and c.

D) a and b. (1)

23. The standard enthalpy formation of NiSO4 (s) at 25 °C is - 872.9 kJ mol-1. The

chemical equation to which this value applies is:

A) ½ Ni (s) + ½ S (s) + ½ O2 (g) → ½ NiSO4 (s)

B) Ni (s) + S (s) + 4 O (g) → NiSO4 (s)

C) Ni (s) + ⅛ S8 (s) + 2 O2 (g) → NiSO4 (s)

D) NiSO4 (s) → Ni (s) + S (s) + 4 O (g) (1)





24. A pressure of 3.00 atm is the same as a pressure of __________ Torr.

A) 253

B) 33.7

C) 2280

D) 33775 (1)

25. Which one of the following statements about gases is FALSE?

A) Gases are highly compressible.

B) Distances between molecules of gas are very large compared to bond

distances within molecules.

C) Gases expand spontaneously to fill the container they are placed in.

D) All gases are colorless and odorless at room temperature. (1)

26. Which statement about ideal behavior of gases is FALSE?

A) At low densities all gases have similar properties.

B) Gas ideality assumes that there are no interactions between gas particles.

C) All particles in the ideal gas behave independently of each other.

D) Low pressures and high temperatures typically cause deviations from the

ideal gas behavior. (1)





27. The van der Waals equation for real gases recognizes that,

A) gas particles have non-zero volumes and interact with each other.

B) the non-zero volumes of gas particles effectively decrease the amount of

"empty space" between them.

C) the molecular attractions between particles of gas decreases the pressure

exerted by the gas.

D) all of the above statements are true. (1)

28. When the reversible reaction

N2 (g) + O2 (g) ⇌ 2 NO (g)

has reached a state of equilibrium,

A) no further reaction occurs.

B) the total amounts, in moles, of reactants and products are equal.

C) the concentration of each substance in the system is constant.

D) the product [N2(g)]e × [O2(g)]e = [NO(g)]e2. (1)





29. Given the reaction:

NO (g) + CO (g) ⇌ ½ N2 (g) + CO2 (g) ΔH = - 374 kJ

The conditions of temperature and pressure that favour the formation of CO2 (g)

are,

A) high T and high P.

B) high T and low P.

C) low T and high P.

D) low T and low P. (1)

30. Given the reaction:

2 NH3 (g) ⇌ N2 (g) + 3 H2 (g) Kc = 230 at 300 °C

which statement most accurately describes the system at equilibrium?

A) More moles of products are present than moles of reactants.

B) More moles of reactants are present than moles of products.

C) Products only are present.

D) Reactants only are present. (1)





31. The equilibrium constant, Kc, of a system at equilibrium, can be altered by

changing:

A) the catalyst.

B) the effective concentrations of the reactants or products.

C) the pressure.

D) the temperature. (1)

32. In which one of the following gas-phase equilibria would Kp = Kc?

A) H2 (g) + I2 (g) ⇌ 2 HI (g)

B) 4 HCl (g) + O2 (g) ⇌ 2 Cl2 (g) + 2 H2O (g)

C) N2 (g) + 3 H2(g) ⇌ 2 NH3 (g)

D) 2 NO (g) + O2 (g) ⇌ 2 NO2 (g) (1)





33. Identify the Brǿnsted-Lowry acids and bases in the equation;

NH4+ (aq) + OH- (aq) ⇌ NH3 (aq) + H2O (l)

A) acids NH4+ (aq), OH- (aq); bases NH3 (aq), H2O (l)

B) acids OH- (aq), H2O (l); bases NH3 (aq), NH4+ (aq)

C) acids NH4+ (aq), H2O (l); bases OH- (aq), NH3 (aq)

D) acids NH4+ (aq), NH3 (aq); bases OH- (aq), H2O (l) (1)





34. Given the following four weak acids:

HF Ka = 6.8 x 10-4

HOCl Ka = 3.8 x 10-8

HNO2 Ka = 4.5 x 10-4

HCN Ka = 4.9 x 10-10

Which is the STRONGEST of these acids?

A) HF

B) HOCl

C) HNO2

D) HCN (1)

[34]




SECTION B: Full Questions

Answer this section on the question paper.

Question 1

1.1 Determine the number of neutrons and electrons in 55Mn2+. (1)

1.2 Naturally occurring chlorine is 75.53 % 35Cl, which has an isotopic mass of 34.969

u, and 24.47 % 37Cl, which has an isotopic mass of 36.966 u. Calculate the

average atomic mass of naturally occurring chlorine (show all the steps of the

calculation). (2)




Question 1 (Continued)

1.3 Calculate the molar mass of sucrose, C12H22O11 (table sugar). (1)

1.4 Calculate the mass, in grams, of 0.433 mol of calcium nitrate [Ca(NO3)2]. The

molar mass of calcium nitrate is 164.1 g mol-1. (1)

[5]




Question 2

2.1 Ascorbic acid (vitamin C) contains 40.92 % C, 4.58 % H, and 54.50 % O by mass.

Calculate the empirical formula of ascorbic acid. (3)

2.2 Balance the following equation by inspection: (1)

NH4NO3 → N2 + O2 + H2O





2.3 How many grams of water are produced in the combustion of 1.00 g of glucose

(C6H12O6)? The molar mass of glucose is 180.0 g mol-1. (3)

C6H12O6 (s) + 6 O2 (g) → 6 CO2 (g) + 6 H2O (l)

2.4 Calculate the percentage yield of a reaction that gives a yield of 2.01 g whilst its

theoretical yield is 2.50 g. (1)

2.5 What mass of potassium hydrogen phthalate, KHP, must be dissolved to make

250.0 mL of a 0.1000 M solution? MMKHP = 204.2 g mol-1. (2)

[10]




Question 3

3. The primary standard, potassium permanganate, KMnO4, was used to standardise

a solution of sodium oxalate, Na2C2O4. The balanced REDOX reaction involved is:

2 MnO4- (aq)+5 C2O4

2- (aq)+16 H+ (aq) →2 Mn2+ (aq)+10 CO2 (g)+8 H2O (l)

If 20.00 mL of the sodium oxalate solution required 20.68 mL of a 0.09628 M

KMnO4 solution to reach the endpoint, what is the molarity of the sodium oxalate

solution? Show all your calculations. (3)

[3]




Question 4

4.1 Study the reaction between chromium chloride and sodium hydroxide and answer

the questions that follow.

CrCl3 (aq) + 3 NaOH (aq) → Cr(OH)3 + 3 NaCl

i) Using the rules for solubility, predict the solubility of the products.

(1)

ii) Write the net ionic reaction for this equation. (1)





4.2 The following reaction occurs in a basic medium:

Cl2 (g) → ClO- (aq) + Cl- (aq)

Balance the redox equation using the half reaction method. Show your work.

(4)

[6]




Question 5

5.1 The frequency of a spectral line of lithium is 4.47 x 1014 s-1. Calculate the energy

of one photon of this light. (1)

5.2 What is the wavelength of light emitted, in nm, when the electron in a hydrogen

atom undergoes a transition from energy level n = 2 to level n = 1?

(3)





5.3 Indicate whether energy is emitted or absorbed when the following electronic

transitions occur: (2)

i) From n = 3 to n = 5

ii) An electron adds to a H+ ion and ends up in the n = 1 energy level.

5.4 If the orbital angular quantum number is l = 3, what are the possible ml values?

(1)





5.5 How many orbitals are there in each of the following sub-levels?

(1)

i) 3d

ii) 4f

5.6 For each of the following write the condensed electronic configuration:

(2)

i) Rb

ii) Bi

[10]




Question 6

6.1 Draw Lewis dot structures for the following compounds.

i) KCl (1)

ii) C2H4 (2)

6.2 Use Lewis structures and formal charges to draw stable resonance structures

for [NO3]¯. (3)





6.3 i) Draw three plausible Lewis structures representing the polyatomic anion,

CNS-, with nitrogen as the central atom. (3)

ii) Decide, on the basis of the formal charges, which one of the structures is the

most plausible. (1)




H

C


HH H H

6.4 The drawn compounds are shown in order of increasing boiling point from (a) to

(d): H

HH

OH

H

CH

O HC

OH

C

H

CHH C H H

HH

(a) (b) (c) (d)

Lowest Highest

Explain fully the above order of boiling points. (4)

[14]




Question 7

7.1 A quantity of 200.0 mL of 0.862 M HCl is mixed with 200.0 mL of 0.862 M KOH in

a constant-pressure calorimeter of negligible heat capacity. The initial

temperature of HCl and KOH solutions is the same at 20.48 °C.

For the process,

H+ (aq) + OH- (aq) → H2O (l)

the heat of neutralization is – 56.2 kJ mol-1.

What is the final temperature of the mixed solution? (Note: The density and

specific heat capacity of the solution are 1.00 g mL-1 and 4.18 J g-1 °C-1

respectively). (4)




Question 7 (continued)

7.2. For the following combustion reaction,

C3H6O (l) + 4 O2 (g) → 3 CO2 (g) + 3 H2O (l) ∆H° = - 1790 kJ mol-1

Calculate the enthalpy of formation (∆H°f) of acetone (C3H6O) using the following

information: the enthalpies of formation for CO2 (g) is - 393.5 kJ mol-1 and for

water is - 285.8 kJ mol-1. (2)

[6]




Question 8

8.1 i) What is the density, in g L-1, of a sample of Cl2 gas at 1124 Torr and

24.0 °C? (Note: The molar mass of Cl2 is 70.90 g mol-1).

(2)

ii) What is the mass of Cl2, in grams, if the volume is 9.22 L? (1)





8.2 Ammonium sulfate is prepared by reacting ammonia with sulfuric acid as shown in

the following reaction,

2 NH3 (g) + H2SO4 (aq) → (NH4)2SO4 (aq)

if 1774 moles of NH3 reacts with sulphuric acid at 315.15 K and under 15.6 atm,

calculate the volume of ammonia required for this reaction in litres. (1)

[4]




Question 9

9.1 For the reaction,

CO (g) + 3 H2 (g) ⇌ H2O (g) + CH4 (g) KC = 190 at 1000 K

if the initial concentrations are [CO (g)] = 0.025 M, [H2 (g)] = 0.045 M, [H2O (g)]

= 0.025 M, and [CH4 (g)] = 0.046 M:

i) Calculate the reaction quotient, Qc, for this reaction. (1)

ii) In which direction will the system move to attain equilibrium? (1)

9.2 Calculate the pOH of a 7.0 x 10-3 M aqueous solution of Ca(OH)2 at 25.0 °C,

assuming complete dissociation. (1)





9.3 What is the [H3O+]e of a solution of formic acid with an initial concentration

0.250 M? Ka = 1.8 x 10-4.

(5)

[8]

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