Exercise 1.1

1. Calculate the energy in kJ.mol-1, corresponding to a) a wavenumber of 1000 cm-1

b) a wavelength of 620 nm

2. A typical microwave oven operates at a frequency of 2.45 x 109 Hz. Calculate the wavelength of this radiation.

3. Light from a sodium street lamp is found to have a frequency of 5.09 x 1014 Hz. Calculate the wavelength of this light in nanometers.

4. Data book information records that the flame colour of potassium is lilac and has a wavelength of 405 nm. Calculate the frequency of this radiation.

5. Use the data book to find the wavelength of light emitted by a sample of copper in a flame and thus calculate its frequency.

6. The wavenumber of radiation absorbed in the infrared region of the spectrum by the C-H bonds in an alkane is given in the data book as around 2900 cm-1. Calculate the wavelength and frequency of this radiation.

Exercise 1.2

1. The bond enthalpy of a Cl-Cl bond is 243 kJ.mol-1. Calculate the maximum wavelength of light that would break one mole of these bonds to form individual chlorine atoms.

2. Chlorinated hydrocarbon molecules contain C-Cl bonds. The energy reuired to break these bonds is 326 kJ.mol-1.a) Calculate the wavelength of light required to break one mole of these bonds.b) By reference to the electromagnetic spectrum suggest why these molecules are

unstable in the upper atmosphere.

3. The red line in the hydrogen spectrum has a wavelength of 656.3 nm. Calculatea) The energy value of one photon of light at this wavelength (in Jmol-1).b) The energy value in kJ.mol-1 for one mole of photons at this wavelength.

Exercise 1.3

1. Using both spectroscopic and orbital box notations write down the electronic configurations for:

a) Oxygen d) Argon g) Mg2+

b) Aluminium e) Li+ h) S2-

c) Phosphorus f) F- i) K+

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2. Which of these is the spectroscopic designation for a lithium atom?A 1s2 2s2 C 1s2 2s1

B 2s1 1s2 D 2s2 1s1

3. Which element has atoms with the spectroscopic designation [Ar] 4s1

A hydrogen C chlorineB lithium D potassium

4. How many electrons are there in the 2p sub-shell of the oxygen atom?

5. Which number would complete this spectroscopic notation for a nitrogen atom?1s2 2s2 2p?

6. Carbon has two unpaired electrons. How many unpaired electrons would boron have?

7. Which element is represented as [Ne] 3s2?A helium C neonB oxygen D magnesium

Exercise 1.4

1. Which block contains the noble gases?A s C dB p D f

2. Which block contains the most reactive metals?A s C dB p D f

3. What one word name is given to the d-block elements?

In each of the next four questions classify the element as being s, p, d or f block.

4. Aluminuim

5. 1s2 2s2 2p6

6. Scandium

7. 1s2 2s2 2p6 3s1

Exercise 1.5

This spectrum was obtained from the atmosphere around the sun.

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1. Write the name of the element which is responsible for the line at 656.3 nm.

2. Write the name of the other element present.

These were obtained from equal sized soil samples that have been treated and the spectra measured.Sample A – from productive farmlandSample B – from an old factory site where insecticides and pesticides were produced.

The grid contains elements whose spectra appear at the end of the exercise.A

Li

B

Hg

C

HeD

Tl

E

Na

F

Ca3. Which element is responsible for the triplet of lines around 650 nm?

4. Which other metal is present in both samples?

5. Use the spectra at the end of the exercise to work out which element is present in sample B but not in A.

6. Mercury is a metal whose salts are well known poisons and thallium salts are used in some countries as a rat poison. Is there evidence of mercury in sample A?

7. Is there evidence of mercury in sample B?

8. Write the name of the element which you think the spectra prove to be the cause of the pollution.

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Exercise 1.6

1. Use the electronegativity values in the data book to predict which of these bonds is the most polar.

A H-Br C Br-ClB N-Cl D N-Br

2. Predict which of these molecules is polar.A CH4 C Br2B CHCl3 D CO2

In each of the following four questions classify the bond as polar or non-polar.

3. N-O

4. N-N

5. H-Cl4

6. C-S

In each of the following four questions classify the bond as polar or non-polar.

7. CO2

8. CBr4

9. Cl2

10.CClBr3

Exercise 1.7

1. Which of these would show the strongest electron pair repulsion?A bonding pair : bonding pairB non-bonding pair : bonding pairC non-bonding pair : non-bonding pair

2. Which of these would result in the smallest bond angle?A bonding pair : bonding pairB non-bonding pair : bonding pairC non-bonding pair : non-bonding pair

Answer the following three questions as TRUE or FALSE

3. Two lone pairs of electrons repel bonding electrons more than one lone pair.

4. The four electron pairs in water take up a tetrahedral arrangement.

5. The water molecule is tetrahedral.

6. Which of these molecules has a non-bonding electron pair on the central atom?A PF3B BF3C BeCl2D H2

7. Draw molecules of the following species showing their shapes:a) PCl3 e) BeF2b) SnCl4 f) H2Sc) BrF4- g) NF3d) PF5

8. What is the likely structure of an antimony(V) chloride molecule?A linearB tetrahedralC trigonal bipyramidalD octahedral

9. Which of these would have a bond angle greater than 109.5 ?A CCl4B NH3C SCl2D BeF2

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10.Sulphur has six outer shell electrons. Draw a diagram to show the SF5- ion structure (the negative ion).

11.What shape describes the arrangement of the electron pairs you have just drawn?

12.Suggest a name for the shape of this molecule.

Exercise 1.8

1. Write down the electronic configurations in both spectroscopic and orbital box notation for the following atoms and ions.a) Cu d) Ni2+

b) Ti3+ e) Fe3+

c) Co2+

2. Zinc invariably forms the 2+ ion and the only ion of scandium is the 3+ ion. Using spectroscopic notation, write down the electronic configuration for both these ions and use them to explain why zinc and scandium are often regarded as not being transition metals.

The filling of the d orbitals follows the ‘aufbau’ principle with two exceptions.

3. What is the symbol of the exception which has the lower atomic number?

4. What is the symbol of the other exception?

Question 5-9: Decide whether the statement is True or False.

5. Copper has a complete 4s sublevel

6. Chromium has the most unpaired electrons.

7. Zinc has the least unpaired electrons.

8. Chromium has an electronic configuration with all five d orbitals half full.

9. Copper has a complete 3d sublevel.

10.What is likely to be the most common charge on transition metal ions?

11.Which copper ion allows copper to fit the definition of a transition metal?

12.For each of the following ions or compounds calculate the oxidation number of the transition metal. (Remember to give your answer as ‘+x’ or ‘-x’.)

a) VO2+ f) K2Cr2O7b) MnO4- g) Na4NiCl6c) CrO42- h) Fe O42-

d) VOCl2 i) K2MnO4e) Cr2O3 j) K3CoF6

Exercise 1.9

1. Write an ion-electron equation for Fe2+ acting as:a) an oxidising agentb) a reducing agent

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2. Work out the oxidation number of Cr in Cr2O72- and decide whether the conversion of Cr2O72- to Cr3+ is oxidation or reduction. Is the Cr2O72- acting as an oxidising agent or a reducing agent in this reaction? Confirm your answer by writing the appropriate ion-electron equation.

3. Work out the oxidation number of chromium in Cr2O72- and in CrO42- and decide whether the conversion of Cr2O72- to CrO42- is oxidation or reduction.

4. The most common oxidation states of iron are +2 and +3. Using orbital box notation, draw out their respective electronic configurations and suggest which of the two ions is the more stable.

Questions 5 and 6 refer to the equation:2Fe2+ + Br2 2Fe3+ + 2Br-

5. Which is the oxidising agent?

6. Which is the reducing agent?

Questions 7 – 9 refer to the equation:Cr2O7

2- + 6I- + 14H+ 2Cr3+ + 3I2 + 7H2O

7. Using your answer to question 2 above, does the oxidation number of chromium decrease or increase in this reaction?

8. What type of reaction does Cr2O72- to Cr3+ cause?

9. The oxidation number of iodine in the iodide ion is –1 and in the iodine molecule is 0. Is the iodide ion, I-, an oxidising agent or a reducing agent?

Questions 10 and 11 refer to the equation:MnO4

- + 5V2+ + 8H+ Mn2+ + 5V3+ + 4H2O

10.Which is the oxidising agent?A MnO4-

B V2+

C Mn2+

D V3+

11.Which is the reducing agent?A MnO4-

B V2+

C Mn2+

D V3+

Questions 12 and 13 refer to the following equation:2FeO4

2- + 3SO32- + 10H+ 3SO4

2- + 2Fe3+ 5H2O

12.Which is the oxidising agent?A FeO42-

B SO32-

C Fe3+

D SO42-

13.Which is the reducing agent?A FeO42-

B SO32-

C Fe3+

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D SO42-

Exercise 1.10

1. Write the correct name for each of the following complexes.

a) [Co(H2O)6]Cl2 c) K4[Fe(CN)6]b) Na[CrF4] d) K3[Fe(C2O4)3]

2. What is the coordination number of the transition metal ion in each of the following complexes?

a) Na[CrF4]b) K3[Fe(C2O4)3]c) K4[Fe(CN)6]

3. Predict the shape of the complex ion in question 2c.

What is the correct structural formula for each of the following compounds?

4. sodium tetrachloroplatinate(II)A Na[PtCl4]B Na2[PtCl4]C Na[Pt2Cl4]D Na4[PtCl4]

5. diaquadicyanocopper(II)A (H2O)2(CN)2CuB [Cu(CN2)(H2O2)]C Cu(CN)(H2O)2D [Cu(CN)2(H2O)2]

6. pentaaquachlorochromium(III) chlorideA [CrCl(H2O)5]Cl2B [Cr(H2O)5]Cl3C [CrCl(H2O)5]Cl3D [CrCl5(H2O)5]Cl

7. tetraamminedichlorocobalt(III) chlorideA [CoCl2(NH3)4]Cl3B [CoCl2(NH3)4]Cl2C [CoCl2(NH3)4]ClD [Co(NH3)4]Cl3

8. Name the following complexes:

a) [CoCl4]2- d) [PtCl6]2-

b) Fe(CN)6]4- e) [Cu(NH3)4]2+

c) [Ni(CN)6]4-

Exercise 1.11

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1. Use the wavelength of the most intense absorption in the visible spectrum above to calculate in kJ.mol-1 (to one decimal place) the crystal field splitting () caused by the chloride ion.

2. What colour would you predict for a solution containing [CrCl6]3- ions?A redB blue-green (cyan)C violet (magenta)D yellow

3. Use the wavelength of the most intense absorption in the visible spectrum above to calculate in kJ.mol-1 (to one decimal place) the crystal field splitting () caused by the water ligand.

4. What colour would you predict for a solution containing [Cr(H2O)6]3+ ions?A redB yellowC violet (magenta)D green

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5. Use the wavelength of the most intense absorption in the visible spectrum above to calculate in kJ.mol-1 (to one decimal place) the crystal field splitting () caused by the ammonia ligand.

6. What colour would you predict for a solution containing [Cr(NH3)6]3+ ions?A redB yellowC violet (magenta)D green

7. The ligands can be placed in order of the crystal field splitting () with the ligand of lowest energy first. Which of the following shows the correct order?A NH3 < H2O < Cl-B NH3 < Cl- < H2OC Cl- < H2O < NH3D Cl- < NH3 < H2O

Exercise 1.12

1. For a reversible reaction at equilibrium, which two options show factors which are;a) equalb) constant, but not equal?

A rate of forward reactionB concentration of reactantsC concentration of productsD rate of reverse reaction

2. When chlorine is dissolved in water the following equilibrium is set up;Cl2 + H2O 2H+ + ClO- + Cl-

The hypochlorite ion, ClO-, is responsible for the bleaching action of this solution.What effect on the bleaching efficiency of a solution of chlorine in water would the following have? Adding;a) dilute nitric acidb) sodium chloride crystalsc) potassium hydroxide solution

3. Reaction 1 : H2(g) + I2(g) 2HI(g)Reaction 2 : 2CO(g) + O2(g) 2CO2(g)Reaction 3 : CH3OH(g) CO(g) + 2H2(g)

a) In which of the above reactions will an increase in pressure;10

i) shift the equilibrium to the right?ii) have no effect on the equilibrium position?

b) In reaction 1, the forward reaction is exothermic. What effect, if any, will anincrease in temperature have in the equilibrium position?

Exercise 1.13

Write an equilibrium expression for the following reactions;

1. 2Fe3+(aq) + 3I-(aq) 2Fe2+(aq) + I3- (aq)2. H3PO4(aq) 2H+(aq) + HPO42-(aq)

3. For the reaction H2 + I2 2HI The equilibrium constant is defined as; K = [HI]2

[H2].[I2]and at 453 ºC it has a value of 50. At 453 ºC which compound is present in greatest concentration?A hydrogenB iodineC hydrogen iodideD all the same concentration

4. The K value for the reaction; PCl5 PCl3 + Cl2is 0.021 at 160 ºC. Which compound is present in greatest concentration at

equilibrium?A phosphorous(V) chlorideB phosphorus(III) chlorideC chlorineD all the same

5. The following equilibrium constants apply at room temperature (25 ºC).Zn(s) + Cu2+(aq) Cu(s) + Zn2+(aq) K = 2 x 1037

Mg(s) + Cu2+(aq) Cu(s) + Mg2+(aq) K = 6 x 1090

Fe(s) + Cu2+(aq) Cu(s) + Fe2+(aq) K = 3 x 1026

Of the metals Zn, Mg and Fe, which removes Cu(II) ions from solution most completely?A ZnB MgC Fe

6. Write an appropriate equilibrium expression for;2SO2(g) + O2(g) 2SO3(g)

7. Which expression is correct for the reaction;Zn(s) + Cu2+

(aq) Zn2+(aq) + Cu(s)

A [Zn2+].[Cu] C [Zn 2+ ]

[Zn].[Cu2+] [Cu2+]

B [Zn].[Cu2+] D [Cu 2+ ] [Zn2+].[Cu] [Zn2+]

8. Which expression is correct for the reaction;NH4NO3(s) 2H2O(g) + N2O(g)

A

11

B

C

D

9. The solubility of halide ions in water decreases in the order AgCl > AgBr > AgI most soluble least soluble

Because of the difference in solubility, the bromide ions will displace chloride ions from solid silver chloride. Calculate the equilibrium constant, K, for the reaction;

AgCl(s) + Br-(aq) AgBr(s) + Cl-

(aq)Given that the equilibrium concentration of chloride and bromide ions are 0.0997 mol.l-1

and 0.0003 mol.l-1 respectively.

Exercise 1.14

1. What is the hydrogen ion concentration in solutions with the following pH values?a) pH = 3b) pH = 12c) pH = 0

2. What is the pH of the following solutions?a) [H+] = 1 x 10-4 mol.l-1b) [OH-] = 1 x 10-5 mol.l-1

3. Calculate the pH of the following solutionsa) Hydrochloric acid, [H+] = 2.7 x 10-4 mol.l-1b) Sodium hydroxide, [H+] = 3.1 x 10-10 mol.l-1c) Hydrochloric acid, [H+] = 2.2 x 10-5 mol.l-1d) 0.034 mol.l-1 sulphuric acide) 0.0022 mol.l-1 barium hydroxide

4. Given that a solution of a strong acid has a pH of 2.25, calculate the [H+]

5. Given that a solution of a strong alkali has a pH of 9.8, calculate the [OH-]

6. What is the pH of a 0.02 mol.l-1 solution of potassium hydroxide?

7. What is the pH of a 0.4 mol.l-1 solution of sulphuric acid?

8. Calculate the pH of a solution of hydrogen chloride with a concentration of 10 g.l -1

1.76g of sodium hydroxide was added to 50 cm3 of 0.025 mol.l-1 nitric acid. Calculate the

pH of the resulting solution.

9. At 100ºC Kw = 50 x 10-14. Calculate the pH of pure water at 100ºC.

10.0.52g of sodium hydroxide were added to a flask and water was added to make a 50 cm3

solution. 25 cm3 of 0.36 mol.l-1 sulphuric acid was added to the flask and finally 60 cm3 of

1.25 mol.l-1 potassium hydroxide was added to the flask. Calculate the pH of;a) the initial solutionb) the solution after sulphuric acid was addedc) the final solution

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Exercise 1.15

1. The terms weak and strong, dilute and concentrated are often confused.a) Explain clearly the difference between a dilute and concentrated acid, then a strong

and weak acid.b) Give one example of a weak acid and one example of a strong acid.c) As a strong acid is diluted state what happens to the concentration of hydrogen

ions and the number of hydrogen ions.

2.Acid pH of 2 mol l-1 solution

A CCl3COOH 0.5B CHCl2COOH 0.9

a) Explain which is the stronger acid.b) Which has the higher conductivity?c) Which reacts faster with magnesium?d) 25cm3 of 2 mol l-1 sodium hydroxide solution is required to neutralise a fixed volume

of acid A. What volume of the same alkali would be required to neutralise the same volume of acid B?

e) Explain your answer to question (d).

3. Explain what happens to the pH of ammonia solution when ammonium is added.

4. Explain what happens to the pH of water when carbon dioxide is bubbled through.

Exercise 1.16

1. When sodium carbonate dissolves in water the pH increases. Explain this change (showing equations)

2. A sodium sulphonate solution has a pH greater than 7.

a) What does this indicate about the sulphonic acid from which it has been made?b) Explain why the sodium sulphonate solution is alkaline.

3. Salts can have a pH less than 7, a pH equal to 7 or a pH greater than 7. Make a table with these headings and place each of the following salts in the correct column.

Sodium sulphite, potassium chloride, lithium carbonateAmmonium sulphate potassium nitrate ammonium nitrateSodium ethanoate lithium sulphate ammonium chloride

4. Solutions of the salt potassium cyanide (KCN) are alkaline.a) What is the formula of the acid from which potassium cyanide is derived?b) Is it a strong or a weak acid?c) Explain why potassium cyanide solution is alkaline.

Exercise 1.17

1. Which is the strongest acid?13

A H3BO3B HFC H2SO3D H3PO4

2. Which is the weakest acid?A H3BO3B HFC H2SO3D H3PO4

3. Which of the following statements would describe a strong acid?A high Ka, high pKa B low Ka, high pKaC high Ka, low pKaD low Ka, low pKa

4. On page 12 of the data book, which is the strongest acid?

5. State the relationship between Ka and the strength of an acid.

6. Which of the following statements would describe a weak acid?A high Ka, high pKa B low Ka, high pKaC high Ka, low pKaD low Ka, low pKa

7. On page 12 of the data book, which is the weakest acid?

8. What is the pH of a solution of 0.02 mol.l-1 propanoic acid?

9. What is the pH of a solution of 0.05 mol.l-1 sulphurous acid?

10.A solution of methanoic acid has a pH of 2.4. What is the molar concentration?

11.Calculate the pH of a 0.2 mol.l-1 solution of Ethanoic acid if Ka = 1.7 x 10-5.

12.0.02 mol.l-1 benzoic acid C6H5COOH, a monoprotic acid, was found to have a pH of 2.94.Calculate the Ka of this weak acid.

Exercise 1.18

1. Using the table on page 37 of the unit calculate the pKa and Ka for the four indicators.

Questions 2 and 3 refer to the following equilibrium for the indicator litmus;HIn(aq) + H2O(l) H3O+(aq) + In-(aq)red blue

2. In which direction will the equilibrium position move of the [H+] is increased?

3. What colour will the solution be?

4. Bromothymol blue has a pKIn value of 7.0.Calculate the ratio for bromothymol blue at pH 9.

5. What colour will the solution be? (you may need to refer to Unit 2 page 41)

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6. Calculate the ratio for bromothymol blue at pH 7.0.

7. What colour will the solution be? (you may need to refer to page 80)

Questions 8 and 9 refer to the table of common indicators on page 80.8. Which indicator is most suitable for the titration of a weak acid with a strong alkali?

9. Which indicator is most suitable for the titration of a strong acid with a weak alkali?Exercise 1.19

1. Calculate the pH of the buffer solution made from 1.0 mol.l-1 methanoic acid and 1.78 mol.l-1 sodium methanoate solution.

2. Calculate the pH of the buffer solution made from ethanoic acid and potassium ethanoate solution, both 1.0 mol.l-1.

3. Calculate the pH of the buffer solution made by adding 2.05g of sodium ethanoate to 1 litreof 0.09 mol.l-1 ethanoic acid.

4. Calculate the pH of a buffer solution containing 0.10 mol.l-1 ethanoic acid and 0.50 mol.l-1 sodium ethanoate.

5. Calculate the composition of methanoic acid and sodium methanoate required to make abuffer solution with a pH of 4.0. Quote your answer as a ratio of salt to 1 (a ratio of 6.31

: 1would be 6.31).

6. Calculate the concentrations of acid and salt solutions required to make ;a) a buffer of pH 6.0 from carbonic acid and sodium hydrogen carbonateb) a buffer of pH 3.1 from chloroethanoic acid and its potassium salt (pKa of

chloroethanoic acid is 2.9)

Exercise 1.20

1. a) Define the enthalpy of formation of butaneb) Write a balanced equation for the enthalpy of formation of butanec) Using the required enthalpies of combustion from the data booklet, calculate the enthalpy of formation of butane.

2. The enthalpy of combustion of carbon monoxide is -284kJ.mol-1. Using this value andrequired enthalpy of combustion from the data booklet, calculate the enthalpy ofpartial combustion of propane to form carbon monoxide and water.

3. Ethyne is used as a fuel in oxyacetylene burners.Use the enthalpies of combustion of ethyne, carbon and hydrogen in the data booklet

tocalculate the enthalpy of formation of ethyne.

4. Use the enthalpies of combustion of carbon, hydrogen and ethanol in the data booklet tocalculate the enthalpy of formation of ethanol.

5. Calculate the enthalpy of formation of carbon disulphide (CS2) given;15

∆H o / kJ.mol -1 C(s) + O2(g) CO2(g) -394 S(s) + O2(g) SO2(g) -297 CS2(s) + 3O2(g) CO2(g) + 2SO2(g) -1072

6. Calculate the enthalpy of formation of benzene using the enthalpies of combustion ofcarbon, hydrogen and benzene in the data booklet.

7. Given the following information; ∆H o / kJ.mol -1

½N2(g) + ½O2(g) NO(g) +90.2 ½N2(g) + O2(g) NO2(g) +33.2Calculate the enthalpy change for the following reaction;

NO(g) + ½O2(g) NO2(g)

8. The molar enthalpy of formation of iron(III) oxide and aluminium oxide are -827 and -1676kJ.mol-1 respectively. Calculate the enthalpy change which takes place in the thermitereaction;

Fe2O3(s) + 2Al(s) 2Fe(s) + Al2O3(s)

Exercise 1.21

1. Cycloalkanes and aromatic hydrocarbons. Consider the sequence shown below.a) Use the data in the table below to calculate ΔSΘ for the conversion of octane to o-

xylene and hydrogen at 298K.

Compound SΘ/J K–1 mol–1

Octane 463

o-xylene 352

Hydrogen 131

2. 1,2-dichloroethane has been used as a solvent for lacquers and oils. One proposed method of production is the addition of hydrogen chloride to ethyne:

C2H2 + 2HCl → CH2ClCH2ClUsing the data given in the table below,a) calculate the standard entropy change, in J K mol–1, for the reactionb) calculate the standard enthalpy change, in kJ mol–1, for the reaction.

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Compound SΘ/J K–1 mol–1 ΔHΘf /kJ mol–1

C2H2 201 227HCl 187 –92.3

CH2ClCH2Cl 208 –166

3. The apparatus shown in the diagram below can be used to find the decomposition temperature of sodium hydrogencarbonate.

The equation for the decomposition is: 2NaHCO3(s) Na2CO3(s) + H2O(g) + CO2(g) ΔHΘ = +129 kJ mol–1

Substance SΘ/J K–1 mol–1NaHCO3(s) 102.1Na2CO3(s) 136.0

H2O(g) 188.7CO2(g) 213.6

Calculate the entropy change for the reaction.

Exercise 1.22

1. Using the data in the table, calculate ∆Go for the reaction;H2(g) + ½O2(g) → H2O(l)

∆Ho (kJ.mol-1) So (J.K-1.mol-1)

H2(g) 0 131O2(g) 0 205H2O(l) -286 70

2. Use the table of standard free energies of formation below to calculate values of ∆Go for the two reactions below and thus predict whether or not the reactions are spontaneous.

substance ∆Gfo (kJ.mol-

1)CO2 -394MgO -569ZnO -318CuO -130All elements 0

a) 2Mg(s) + CO2(g) → 2MgO(s) + C(s)b) 2CuO(s) + C(s) → 2Cu(s) + CO2(g)

3. Calculate the standard free energy change at both 400 K and 1000 K for the reaction; MgCO3(s) → MgO(s) + CO2(g)

∆Hof / kJ.mol-1 -1113 -602 -394So / J.K-1.mol-1 66 27 214

4. Use the data given along with data book values to calculate the temperature at whichthe Haber process becomes feasible.

N2(g) + 3H2 (g) → 2NH3(g)∆Hof / kJ.mol-1 0 0 -46.4So / J.K-1.mol-1 ? ? 193.2

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ureas

Exercise 1.23

1. By experiment the reaction is found to be first order with respect to both bromide and bromate but second order with respect to hydrogen ions.

5Br-(aq) + BrO3-(aq) + 6H+(aq) → 3Br2(aq) + 3H2O(l)

a) Write the rate equation for this reaction.b) What is the overall order of the reaction?

2. The following reaction shows the hydrolysis of urea in the presence of the enzyme urease;

NH2CONH2(aq) + H2O(l) → 2NH3(g) + CO2(g)

The rate equation for the reaction is found by experiment to be;Rate = k.[urea].[urease]

a) What is the overall order of the reaction?b) What is the order with respect to water?c) What is the order with respect to urea?d) What is the order with respect to urease?

3. The following reaction shows the reaction between propanone and bromine in alkaline solution. The balanced equation is;

CH3COCH3(aq) + Br2(aq) + OH-(aq) → CH3COCH2Br(aq) + H2O(l) + Br-(aq)

The experimentally determined rate equations is;Rate = k.[CH3COCH3].[OH-]

Answer True or False:a) The reaction is first order with respect to bromineb) The reaction involves a simple one-step processc) The reaction is second order overalld) The rate determining step involves one molecule of propanone and one molecule of

bromine

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