Questions Booklet - HCC Applied...

Level 3 Applied Science 2017-2018 Unit 1 (Chemistry) 1

Level 3 Applied Science

UNIT 1:

Principles & Applications

of Science I

CHEMISTRY SECTION

Questions

Booklet

Name: ………………………………………………………………………………………..

Teacher: ……………………………………………………………………………………..


1. State the relative mass and charge for protons, neutrons and electrons.

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2. Where is most of the mass in an atom located?

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3. How many electrons are found in the first and second shells of an atom?

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4. In which part of an atom are neutrons found?

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5. State what is meant by the terms:

(i) atomic number

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(ii) mass number

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6. What is the elemental symbol for fluorine?

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7. Calculate the number of protons, neutrons and electrons in the neutral element that has an atomic

number of 34 and mass number of 79 for one of its pure isotopes.

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8. Si has isotopes with a mass number of 28 and 29, yet the relative atomic mass is 28.1. What does this

tell you about the relative amounts of each isotope in a sample of Si found in nature?

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9. Complete the following table:

Pure

Isotope

Atomic no.

(protons)

Mass no. (protons

+ neutrons) No. of electrons No. of neutrons

R.A.M in periodic

table

16O

32S

10B

24Mg

14N

15N

12C

13C

14C

10. State what is meant by the term relative atomic mass.

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11. What is the difference between two isotopes of the same element?

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12. Explain why 14N and 35Cl are not isotopes, but 14N and 15N are isotopes.

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13. Explain why 14N and 15N have the same chemical properties.

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14. What is the standard used when weighing atomic masses?

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15. Calculate the relative atomic mass, Ar, of the following.

(a) A sample of boron containing 18.7% 10B and 81.3% 11B. Give your answer to 3 significant figures.

Answer = ........................................

(b) A sample of magnesium containing 78.6% 24Mg, 10.1% 25Mg and 11.3% 26Mg. Give your answer to 4

significant figures.

Answer = ........................................

(c) A sample of lead containing 1.55% 204Pb, 23.6% 206Pb, 22.6% 207Pb and 52.3% 208Pb. Give your answer

to 4 significant figures.

Answer = ........................................

(d) A sample of iron containing 5.8% 54Fe, 91.6% 56Fe, 2.2% 57Fe and 0.33% 58Fe. Give your answer to 3

significant figures.

Answer = ........................................

(e) A sample of zinc containing 48.9% 64Zn, 27.8% 66Zn, 4.1% 67Zn, 18.6% 68Zn and 0.62% 70Zn. Give

your answer to 3 significant figures.

Answer = ........................................

16. Explain why Mg and Ca have similar chemical properties.

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17. Identify the element in:

(i) Group 3, Period 3 ………………..

(ii) Group 8, Period 1 ………………..

(iii) Group 7, Period 5 ………………..


18. Explain how the elements are arranged in the Periodic Table

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19. Explain what is meant by the term Periodicity. Use examples in your answer.

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20. Compare and explain the atomic radius of F and Cl.

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21. Compare and explain the atomic radius of Li and F.

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22. Explain what is meant by ‘cation’ and ‘anion’.

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23. Explain the octet rule.

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24. Predict the charges on the ions in the table below

Mg2+ F N Al P Br Sr O Cs K

Li Cl S Ca Rb Ba I Cu Fe Mn

25. Draw diagrams to show the formation of the following ions:

i) Mg to Mg2+

ii) K to K+

iii) O to O2-

26. Complete the following table

Atom/Ion Atomic no. Mass no. No. of protons No. of neutrons No. of electrons

14N3-

16O2-

23Na+

27Al3+


27. Compare, and explain, the radius of Ca with its cation.

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28. Compare, and explain, the radius of P with its anion.

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29. Explain what is meant by the term isoelectronic.

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30. Identify any ions that are isoelectronic with Ca.

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31. Explain what is meant by the term first ionisation energy.

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32. Explain why energy is required to remove electrons from atoms. Is this shown by a positive or

negative sign?

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33. Write equations to show the following ionisation energies

(i) first for Li: ………………………………………………………………………………………………..

(ii) third for Ga: ……………………………………………………………………………………………...

(iii) ninth for Na: …………………………………………………………………………………………….

(iv) tenth for Al: ……………………………………………………………………………………………..

(v) sixth for C: ……………………………………………………………………………………………….


34. Explain why the first ionisation energy for Be is greater than the first for Li.

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35. Explain why the first ionisation energy for Na is less than the first for Li.

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36. Explain what is meant by first electron affinity.

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37. Write equations for the first and second electron affinities for:

(i) O ………………………………………………………………………………………………………….

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(ii) N …………………………………………………………………………………………………………

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(iii) P ………………………………………………………………………………………………………...

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(iv) S …………………………………………………………………………………………………………

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38. Explain why energy is released for the first electron affinity but energy is required for the second

electron affinity. Is this shown by a positive or negative sign?

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39. Explain why less energy is released for the first electron affinity of Br than the first for Cl.

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40. Explain why less energy is released for the first electron affinity of Se than the first for Br.

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41. State, and explain, whether the following show ionic, covalent or metallic bonding:

(i) Ca ………………..

(ii) O2 ………………..

(iii) NaBr ………………..

(iv) NH3 ………………..

(v) Mg ………………..

(vi) MgCl2 ………………..

(vii) CO2 ………………..


(viii) CH4 ………………..

(iv) AlCl3 ………………..

42. Complete the table below

Name of ionic

compound

Sodium fluoride Magnesium

fluoride

Aluminium

chloride

Magnesium

phosphide

Cross-over

working

Formula of

ionic compound

43. Draw dot and cross diagrams for the following and also calculate the relative formula mass:

(i) sodium fluoride, NaF

R.F.M ……………

(ii) magnesium chloride, MgCl2

R.F.M ……………

(iii) aluminium oxide, Al2O3

R.F.M ……………

(iv) Barium oxide, BaO

R.F.M ……………

(v) Calcium sulphide, CaS

R.F.M ……………

(vi) Barium nitride, Ba3N2

R.F.M ……………


44. What two factors affect the strength of an ionic bond? Explain your answer.

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45. Draw dot & cross diagrams for the following and calculate the relative molecular mass

(i) hydrogen, H2

R.M.M ……………

(ii) chlorine, Cl2

R.M.M ……………

(iii) methane, CH4

R.M.M ……………

(iv) oxygen, O2

R.M.M ……………

(v) nitrogen, N2

R.M.M ……………

(vi) H3O+ (formed from the reaction of H2O with H+)

R.F.M ……………


46. State what is meant by a dative covalent bond, using a dot & cross diagram of the ammonium ion as

your example.

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47. Draw a 3D diagram of methane. State the name of the shape and the bond angle.

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48. Draw dot & cross diagrams for Cl2 and CO2 . Use these diagrams to compare the strength and length

of the bonds in Cl2 and in CO2.

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49. State what is meant by the term metallic bonding, using a labelled diagram to illustrate your answer.

You should also give examples of substances which have metallic bonding.

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50. Explain the following properties of metals:

(i) conduction of heat and electricity

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(ii) malleability and ductility.

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51. State what is meant by the term electronegativity. Explain the trends in electronegativity values in

the periodic Table.

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52. Show any dipoles on the examples below and state whether the bond is polar or non-polar:

H – F

C – Cl

O – H

S – F

N – H

H – Cl

C – O

N – O

Cl – F

Br – Br

53. Explain what factors affect the electronegativity of an element.

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54. Explain how polar and non-polar molecules arise.

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55. State what is meant by the term intermolecular force. List the three types of intermolecular forces in

order of increasing strength. How do intermolecular forces compare with ionic, covalent and metallic

bonds?

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56. Explain how temporary dipole-induced dipole forces arise.

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57. Explain why the boiling point of the noble gases increases as you go down the group 8.

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58. Draw diagrams showing hydrogen bonding between:

(i) neighbouring NH3 molecules

(ii) neighbouring HF molecules

(iii) one molecule of water and one molecule of ammonia.

59. Identify the intermolecular forces present in the following

(i) O2 …………………………………………………………………………………………………………

(ii) N2 ………………………………………………………………………………………………………..

(iii) F2 ………………………………………………………………………………………………………..

(iv) HF ……………………………………………………………………………………………………….

(v) CH3CH2NH2 .………………………………………………………………………………………….....

(vi) CH3CH2OH ……………………………………………………………………………………………..

(vii) HCl …………………………………………………………………………………………………….

(viii) H2S …………………………………………………………………………………………………….


60. Explain why ICl has a higher boiling point than Cl2.

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61. Explain why H2O has a higher boiling point than H2.

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62. Explain why H2O has a higher boiling point than HCl.

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63. Explain why O2 has a lower melting point than Li.

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64. Explain why Cl2 has a lower melting point than Si.

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65. Explain why Al has a higher melting point than Mg, which has a higher melting point than Na.

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66. Explain why S has a higher melting point than P, which has a higher melting point than Cl, which has

a higher melting point than Ar.

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67. Explain why MgO has a higher boiling point than SO2.

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68. Explain the trend in boiling points down Groups 1, 2 and 7.

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69. Explain what is meant by the term orbital. What are the different types of orbital and how many of

each are there?

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70. Calculate the number of electrons in shell 4, showing your full working.

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71. Draw an S and a P orbital. State their shape.

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72. Draw an electron in box diagram and also write out the electron configuration for Ca and Ca2+.


73. Draw electron in box diagrams and below each write out the electron configuration for:

(i) H

(ii) Be

(iii) O

(iv) K+

(v) Br

(vi) Br-


74. Write out electron configurations only for the following atoms and ions:

He

N

F

Al3+

P3-

75. Write out electron configurations for the following, proving that they do belong to the s or p block.

s-block

Li

Na

Ca

p-block

B

P

Cl

76. Prove that the cations Na+ and Mg2+ are isoelectronic by writing out their electronic configurations.

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77. To which block in the Periodic Table do the following elements belong

(i) Mg ………………..

(ii) Si ………………..

(iii) I ………………..

(iv) Mn ………………..

(v) Ag ………………..


78. Balance the following equations.

1. Mg + N2 → Mg3N2

2. Ba + HCl → BaCl2 + H2

3. C + O2 → CO2

4. C2H6 + O2 → CO2 + H2O

5. Fe + Cl2 → FeCl3

6. Mg + HNO3 → Mg(NO3)2 + H2

7. CuCO3 + HCl → CuCl2 + + CO2 + H2O

79. Construct balanced equations for reaction of the following elements with oxygen. Explain any

properties of the oxides produced.

(i) Na ………………………………………………………………………………………………………...

(ii) Al ………………………………………………………………………………………………………...

(iii) C ………………………………………………………………………………………………………...

(iv) Si ………………………………………………………………………………………………………..

80. Explain why Na2O has a higher melting point than CO2.

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81. Construct balanced equations for the reactions of Na and Al with:

(i) oxygen

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(ii) water

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(iii) hydrochloric acid

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(iv) sulphuric acid.

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82. The reaction between lithium metal and hydrochloric acid is a redox reaction. Write an equation for

this reaction and prove that it is a redox reaction.

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83. The element strontium forms a nitrate, Sr(NO3)2, which decomposes on heating as shown below.

2 Sr(NO3)2 (s) → 2 SrO (s) + 4 NO2 (g) + O2 (g)

Using oxidation numbers, explain why the reaction involves both oxidation and reduction.

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84. Assign oxidation numbers and say whether or not these are REDOX reactions.

i)

Mg + 2 HCl → MgCl2 + H2

ii)

2 Ca + O2 → 2 CaO

iii)

SrO + H2O → Sr(OH)2

iv)

Mg + H2O → MgO + H2

v)

BaO + 2 HCl → BaCl2 + H2O

vi)

Ca + 2 HCl → CaCl2 + H2

85. Prove that the reaction below is a redox reaction.

Fe2O3 + 3 CO → 2 Fe + 3 CO2

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86. Complete the table below. What property of transition metals makes them useful as catalysts?

Formula Name Oxidation no. of

transition element

Copper(I) chloride

CuCl2

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87. Explain why Rb is more reactive than K but Ca is less reactive than K.

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88. Magnesium and copper(II) sulfate, CuSO4, take part in a displacement reaction. Write the equation for

this reaction. Is the reaction a redox reaction? Why does the reverse reaction not take place?

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89. Iron and copper(II) sulfate, CuSO4, take part in a displacement reaction. Write the equation for this

reaction. Is the reaction a redox reaction? Why does the reverse reaction not take place?

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90. Write the equation for the reaction of bromide with chlorine, Cl2. Why does this reaction take place?

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91. Calculate the molar mass of the following:

1. CuO 2. Cl2 3. HF 4. CaCO3

5. NH4NO3 6. Na2SO4 7. Al2(SO4)3 8. Zn(OH)2

92. Calculate the mass, in grams, for the following.

1. 2 mol of Na 2. 0.5 mol of Zn 3. 10 mol of Pb 4. 0.25 mol of SO2

5. 5 mol of Na2CO3 6. 0.2 mol of Br2 7. 0.01 mol of Ag2O 8. 0.02 mol of KMnO4

93. Calculate the amount, in mol, for the following.

1. 2.3 g of Na 2. 58.5 g of NaCl 3. 50.0 g of CaCO3 4. 40.0 g of CH4

5. 0.98 g of H2SO4 6. 8.0 g of Br2 7. 160.0 g of NaOH 8. 303.0 g of KNO3

94. Find the amount, in moles, of solute dissolved in water to form the following solutions.

1. 1 dm3 of 0.2 mol dm-3 H2SO4 2. 100 cm3 of 1 mol dm-3 KCl 3. 25 cm3 of 2.5 mol dm-3 NaCl

4. 22.5 cm3 of 0.2 mol dm-3 HCl 5. 2 dm3 of 2 mol dm-3 H2SO4 6. 5 cm3 of 5 mol dm-3 HNO3


95. Find the concentration of the following solutions in mol dm-3.

1. 1 mol of HCl in 1 dm3 2. 5 mol of HNO3 in 250 cm3 3. 0.1 mol of HNO3 in 25 cm3

4. 2.25 x 10-3 mol in 15 cm3 5. 4 mol in 20 dm3 6. 8.8 x 10-2 mol in 200 cm3

96. 12.7 g of Cu reacts with excess O2 to produce CuO.

2 Cu + O2 → 2 CuO

(i) Calculate the amount, in mol, of Cu used in the reaction.

(ii) Calculate the amount, in mol, of CuO produced in the reaction.

(iii) Calculate the mass, in grams, of CuO produced in the reaction.

97. 6.0 g of Mg reacts with an aqueous solution of CuSO4 to form MgSO4 and Cu.

CuSO4 + Mg → MgSO4 + Cu

(i) Calculate the amount, in mol, of Mg used in the reaction.

(ii) Calculate the amount, in mol, of Cu produced in the reaction.

(iii) Calculate the mass, in grams, of Cu produced in the reaction.


98. In a reaction, SrCO3 is used to neutralise 25 cm3 of a 0.15 mol dm-3 HCl solution.

SrCO3 + 2 HCl → SrCl2 + H2O + CO2

(i) Calculate the amount, in mol, of HCl used in the reaction.

(ii) Calculate the mass, in grams, of SrCO3 used to neutralise the HCl in the reaction.

99. In an experiment 0.50 g of K metal was added to 500 cm3 of water making a solution of KOH.

2 K + 2 H2O →2 KOH + H2

(i) Calculate the amount, in mol, of K used in the reaction.

(ii) Calculate the concentration, in mol dm-3, of the KOH solution formed.

100. An excess amount of Ba is treated with 200 cm3 of 0.05 mol dm-3 HCl to form an aqueous solution of

barium chloride.

Ba + 2 HCl → BaCl2 + H2

(i) Calculate the amount, in mol, of HCl used in the reaction.

(ii) Calculate the concentration, in mol dm-3, of the BaCl2 solution formed.


101. 1 Kg of MgO is used to neutralise a 0.04 mol dm-3 solution of HCl.

MgO + 2 HCl → MgCl2 + 2 H2O

(i) Calculate the amount, in mol, of MgO used in the reaction.

(ii) Calculate the volume, dm3, of the HCl solution used.

102. 50 kg of pure sulfuric acid was accidently released into a lake when a storage vessel leaked. It was

proposed that CaCO3 was used to neutralise it.

Calculate the mass of CaCO3 required to neutralise the 50 kg of H2SO4.

H2SO4 + CaCO3 → CaSO4 + H2O + CO2

103. A student dissolved 2.794 g of an acid HY in 100 cm3 of water. Once HY was fully dissolved, the

student then added a further 150 cm3 of water to make 250 cm3 of a solution of HY in water. 25cm3 of this

HY solution was then titrated against a 0.0614 mol dm-3 of Na2CO3 solution, requiring 23.45 cm3 of the

Na2CO3 solution for complete neutralisation.

2 HY + Na2CO3 → 2 NaY + H2O + CO2

(i) Calculate the amount, in mol, of Na2CO3 used in the titration.

(ii) Calculate the amount, in mol, of HY used in the titration.


(iii) Calculate the amount, in mol, of HY in the original 250 cm3 solution.

(iv) Calculate the molar mass of HY.

104. A chemist reacts 1.00g of Cu with excess oxygen. After purification, 1.00g of CuO is isolated from

the reaction.

2 Cu + O2 → 2 CuO

(i) Calculate the amount, in moles, of Cu reacted.

(ii) Determine the theoretical moles of CuO.

(iii) Calculate the actual moles of CuO made.

(iv) determine the percentage yield of the reaction.

105. A chemist reacts 1.00g of methanol, CH3OH, with an oxidant. After purification, 0.98g of methanol,

CH2O, is isolated from the reaction. Determine the percentage yield.

CH3OH + [O] → CH2O + H2O


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