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quet1. What is the formula for the compound formed by calcium and nitrogen?
A. CaN B. Ca2N C. Ca2N3 D. Ca3N2
2. What is the best description of the carbon-oxygen bond lengths in CO32–
?
A. One short and two long bonds B. One long and two short bonds
C. Three bonds of the same length D. Three bonds of different lengths
3. What is the number of sigma () and pi () bonds and the hybridization of the carbon atom in
O
C HH O
Sigma Pi Hybridization
A. 4 1 sp2
B. 4 1 sp3
C. 3 2 sp3
D. 3 1 sp2
4. Element X is in group 2, and element Y in group 7, of the periodic table. Which ions will be
present in the compound formed when X and Y react together?
A. X+ and Y
– B. X
2+ and Y
– C. X
+ and Y
2– D. X
2– and Y
+
5. Which of the following increase(s) for the bonding between carbon atoms in the sequence of
molecules C2H6, C2H4 and C2H2?
I. Number of bonds II. Length of bonds III. Strength of bonding
A. I only B. I and III only C. III only D. I, II and III
6. Which of the following contain a bond angle of 90°?
I. PC14+
II. PCl5 III. PCl6–
A. I and II only B. I and III only C. II and III only D. I, II and III
7. Which allotropes contain carbon atoms with sp2 hybridization?
I. Diamond II. Graphite III. C60 fullerene
A. I and II only B. I and III only C. II and III only D. I, II and III
8. Based on electronegativity values, which bond is the most polar?
A. B―C B. C―O C. N―O D. O―F
9. What is the Lewis (electron dot) structure for sulfur dioxide?
A. O S O B. O S O C. O S O
D. O S O
10. Which substance is most soluble in water (in mol dm–3
) at 298 K?
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A. CH3CH3 B. CH3OCH3 C. CH3CH2OH D. CH3CH2CH2CH2OH
11. What is the molecular shape and the hybridization of the nitrogen atom in NH3?
Molecular shape Hybridization
A. tetrahedral sp3
B. trigonal planar sp2
C. trigonal pyramidal sp2
D. trigonal pyramidal sp3
12. Which statement about sigma and pi bonds is correct?
A. Sigma bonds are formed only by s orbitals and pi bonds are formed only by p orbitals.
B. Sigma bonds are formed only by p orbitals and pi bonds are formed only by s orbitals.
C. Sigma bonds are formed by either s or p orbitals, pi bonds are formed only by p orbitals.
D. Sigma and pi bonds are formed by either s or p orbitals.
13. According to VSEPR theory, repulsion between electron pairs in a valence shell decreases in
the order
A. lone pair-lone pair > lone pair-bond pair > bond pair-bond pair.
B. bond pair-bond pair > lone pair-bond pair > lone pair-lone pair.
C. lone pair-lone pair > bond pair-bond pair > bond pair-lone pair.
D. bond pair-bond pair > lone pair-lone pair > lone pair-bond pair.
14. Which molecule is linear?
A. SO2 B. CO2 C. H2S D. Cl2O
15. Why is the boiling point of PH3 lower than that of NH3?
A. PH3 is non-polar whereas NH3 is polar.
B. PH3 is not hydrogen bonded whereas NH3 is hydrogen bonded.
C. Van der Waals’ forces are weaker in PH3 than in NH3.
D. The molar mass of PH3 is greater than that of NH3.
16. Which molecule is non-polar?
A. H2CO B. SO3 C. NF3 D. CHCl3
17. Consider the following statements.
I. All carbon-oxygen bond lengths are equal in CO32–
.
II. All carbon-oxygen bond lengths are equal in CH3COOH.
III. All carbon-oxygen bond lengths are equal in CH3COO–.
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Which statements are correct?
A. I and II only B. I and III only C. II and III only D. I, II and III
18. What happens when sodium and oxygen combine together?
A. Each sodium atom gains one electron. B. Each sodium atom loses one electron.
C. Each oxygen atom gains one electron. D. Each oxygen atom loses one electron.
19. Which statement is correct about two elements whose atoms form a covalent bond with
each other?
A. The elements are metals. B. The elements are non-metals.
C. The elements have very low electronegativity values.
D. The elements have very different electronegativity values.
20. In ethanol, C2H5OH (l), there are covalent bonds, hydrogen bonds and van der Waals’
forces. Which bonds or forces are broken when ethanol is vaporized?
A. only hydrogen bonds B. covalent bonds and hydrogen bonds
C. covalent bonds and van der Waals’ forces
D. hydrogen bonds and van der Waals’ forces
21. Which substance has the lowest electrical conductivity?
A. Cu(s) B. Hg(l) C. H2(g) D. LiOH(aq)
22. Which statement best describes the attraction present in metallic bonding?
A. the attraction between nuclei and electrons
B. the attraction between positive ions and electrons
C. the attraction between positive ions and negative ions
D. the attraction between protons and electrons
23. Which statement is correct about multiple bonding between carbon atoms?
A. Double bonds are formed by two π bonds.
B. Double bonds are weaker than single bonds.
C. π bonds are formed by overlap between s orbitals.
D. π bonds are weaker than sigma bonds.
24. When the following bond types are listed in decreasing order of strength (strongest first), what
is the correct order?
A. covalent hydrogen van der Waals’ B. covalent van der Waals’ hydrogen
C. hydrogen covalent van der Waals’ D. van der Waals’ hydrogen covalent
25. Which statement is true for most ionic compounds?
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A. They contain elements of similar electronegativity.
B. They conduct electricity in the solid state.
C. They are coloured. D. They have high melting and boiling points.
26. What is the valence shell electron pair repulsion (VSEPR) theory used to predict?
A. The energy levels in an atom B. The shapes of molecules and ions
C. The electronegativities of elements D. The type of bonding in compounds
27. Which fluoride is the most ionic?
A. NaF B. CsF C. MgF2 D. BaF2
28. Which particles can act as ligands in complex ion formation?
I. C1–
II. NH3 III. H2O
A. I and II only B. I and III only C. II and III only D. I, II and III
29. Which statements correctly describe the NO2– ion?
I. It can be represented by resonance structures.
II. It has two lone pairs of electrons on the N atom.
III. The N atom is sp2 hybridized.
A. I and II only B. I and III only C. II and III only D. I, II and III
30. Which substance is most similar in shape to NH3?
A. GaI3 B. BF3 C. FeCl3 D. PBr3
31. Which statement is a correct description of electron loss in this reaction?
2Al + 3S Al2S3
A. Each aluminium atom loses two electrons.
B. Each aluminium atom loses three electrons.
C. Each sulfur atom loses two electrons.
D. Each sulfur atom loses three electrons.
32. Which molecule has the smallest bond angle?
A. CO2 B. NH3 C. CH4 D. C2H4
33. In which substance is hydrogen bonding present?
A. CH4 B. CH2F2 C. CH3CHO D. CH3OH
34. Which is a correct description of metallic bonding?
A. Positively charged metal ions are attracted to negatively charged ions.
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B. Negatively charged metal ions are attracted to positively charged metal ions.
C. Positively charged metal ions are attracted to delocalized electrons.
D. Negatively charged metal ions are attracted to delocalized electrons.
35. Which is the smallest bond angle in the PF5 molecule?
A. 90 B. 109.5 C. 120 D. 180
36. Which types of hybridization are shown by the carbon atoms in the compound CH2 = CHCH3?
I. sp II. sp2
III. sp3
A. I and II only B. I and III only C. II and III only D. I, II and III
37. What intermolecular forces are present in gaseous hydrogen?
A. Hydrogen bonds B. Covalent bonds
C. Dipole-dipole attractions D.Van der Waals’ forces
38. Which molecule is polar?
A. CO2 B. PF3 C. CH4 D. BF3
39. What are responsible for the high electrical conductivity of metals?
A. Delocalized positive ions B. Delocalized valence electrons
C. Delocalized atoms D. Delocalized negative ions
40. Which compound has the least covalent character?
A. SiO2 B. Na2O C. MgCl2 D. CsF
41. Identify the types of hybridization shown by the carbon atoms in the molecule
CH3CH2CH2COOH
I. sp II. sp2
III. sp3
A. I and II only B. I and III only C. II and III only D. I, II and III
42. When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the
correct order?
A. C2H6, C2H2, C2H4 B. C2H4, C2H2, C2H6
C. C2H2, C2H4, C2H6 D. C2H4, C2H6, C2H2
43. Which compound contains both ionic and covalent bonds?
A. MgCl2 B. HCl C. H2CO D. NH4Cl
44. When the species BF2+, BF3 and BF4
– are arranged in order of increasing F−B−F bond angle,
what is the correct order?
A. BF3, BF4–, BF2
+ B. BF4
–, BF3, BF2
+
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C. BF2+, BF4
–, BF3 D. BF2
+, BF3, BF4
–
45. Which species has a trigonal planar shape?
A. CO32–
B. SO32–
C. NF3 D. PCl3
46. When C2H4, C2H2 and C2H6 are arranged in order of increasing C–C bond length, what is the
correct order?
A. C2H6, C2H2, C2H4 B. C2H4, C2H2, C2H6
C. C2H2, C2H4, C2H6 D. C2H4, C2H6, C2H2
47. Which molecule is square planar in shape?
A. XeO4 B. XeF4 C. SF4 D. SiF4
48. What is the hybridization of nitrogen atoms I, II, III and IV in the following molecules?
H N N H H N N H2 2
I II III IV
I II III IV
A. sp2 sp
2 sp
3 sp
3
B. sp3 sp
3 sp
2 sp
2
C. sp2 sp
2 sp sp
D. sp3 sp
3 sp sp
49. What is the formula for an ionic compound formed between an element, X, from group 2 and an
element, Y, from group 6?
A. XY B. X2Y C. XY2 D. X2Y6
50. In the molecules N2H4, N2H2, and N2, the nitrogen atoms are linked by single, double and triple
bonds, respectively. When these molecules are arranged in increasing order of the lengths of
their nitrogen to nitrogen bonds (shortest bond first) which order is correct?
A. N2H4, N2, N2H2 B. N2H4, N2H2, N2
C. N2H2, N2, N2H4 D. N2, N2H2, N2H4
51. The compounds listed have very similar molar masses. Which has the strongest intermolecular
forces?
A. CH3CHO B. CH3CH2OH C. CH3CH2F D. CH3CH2CH3
52. What is the shape of the CO32–
ion and the approximate O–C–O bond angle?
A. Linear, 180 B. Trigonal planar, 90
C. Trigonal planar, 120 D. Pyramidal, 109
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53. What is the molecular geometry and the Cl–I–Cl bond angle in the ICl4– ion?
A. Square planar 90 B. Square pyramidal 90
C. Tetrahedral 109 D. Trigonal pyramidal 107
54. What is the geometry of the bonds around an atom with sp2 hybridization?
A. 2 bonds at 180 B. 3 bonds at 120
C. 2 bonds at 90, 1 bond at 180 D. 4 bonds at 109
55. Which combination of Hvaporization and boiling point is the result of strong intermolecular
forces?
Hvaporization Boiling Point
A. large high
B. large low
C. small low
D. small high
56. What is the formula of the compound formed when aluminium reacts with oxygen?
A. Al3O2 B. Al2O3 C. AlO2 D. AlO3
57. Which statement is true for compounds containing only covalent bonds?
A. They are held together by electrostatic forces of attraction between oppositely charged
ions.
B. They are made up of metal elements only.
C. They are made up of a metal from the far left of the periodic table and a non-metal from
the far right of the periodic table.
D. They are made up of non-metal elements only.
58. How many electrons are used in the carbon-carbon bond in C2H2?
A. 4 B. 6 C. 10 D. 12
59. Which compound has the highest boiling point?
A. CH3CH2CH3 B. CH3CH2OH C. CH3OCH3 D. CH3CHO
60. What type of solid materials are typically hard, have high melting points and poor electrical
conductivities?
I. Ionic II. Metallic III. Covalent-network
A. I and II only B. I and III only C. II and III only D. I, II and III
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61. How many sigma (σ) and pi (π) bonds are present in the structure of HCN?
σ π
A. 1 3
B. 2 3
C. 2 2
D. 3 1
62. How many lone pairs and bonding pairs of electrons surround xenon in the XeF4 molecule?
Lone pairs Bonding pairs
A. 4 8
B. 0 8
C. 0 4
D. 2 4
63. The boiling points of the hydrides of the group 6 elements are shown below.
400
300
200
100
0
Boiling point / K
H O H S H Se H Te2 2 2 2
(i) Explain the trend in boiling points from H2S to H2Te.
(ii) Explain why the boiling point of water is higher than would be expected from the group
trend.
64. (i) State the shape of the electron distribution around the oxygen atom in the water molecule
and state the shape of the molecule.
(ii) State and explain the value of the HOH bond angle.
65. Explain why the bonds in silicon tetrachloride, SiCl4, are polar, but the molecule is not.
66. The diagrams below represent the structures of iodine, sodium and sodium iodide.
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A B C
(a) (i) Identify which of the structures (A, B and C) correspond to iodine, sodium and
sodium iodide.
(ii) State the type of bonding in each structure.
(b) (i) Sodium and sodium iodide can both conduct electricity when molten, but only
sodium can conduct electricity when solid. Explain this difference in conductivity in
terms of the structures of sodium and sodium iodide.
(ii) Explain the high volatility of iodine compared to sodium and sodium iodide.
67. The boiling points of the hydrides of group 6 elements increase in the order
H2S < H2Se < H2Te < H2O.
Explain the trend in the boiling points in terms of bonding.
68. (i) Draw the Lewis structures for carbon monoxide, carbon dioxide and the carbonate ion.
(ii) Identify the species with the longest carbon-oxygen bond and explain your answer.
69. (i) Draw Lewis (electron dot) structures for CO2 and H2S showing all valence electrons.
(ii) State the shape of each molecule and explain your answer in terms of VSEPR theory.
(iii) State and explain whether each molecule is polar or non-polar.
70. Identify the strongest type of intermolecular force in each of the following compounds.
CH3Cl ...................................................................................................................................
CH4 .......................................................................................................................................
CH3OH .................................................................................................................................
71. (i) List the following substances in order of increasing boiling point (lowest first).
CH3CHO C2H6 CH3COOH C2H5OH
(ii) State whether each compound is polar or non-polar, and explain the order of boiling
points in (i).
72. (a) An important compound of nitrogen is ammonia, NH3. The chemistry of ammonia is
influenced by its polarity and its ability to form hydrogen bonds. Polarity can be
explained in terms of electronegativity.
(i) Explain the term electronegativity.
10
(ii) Draw a diagram to show hydrogen bonding between two molecules of NH3.
The diagram should include any dipoles and/or lone pairs of electrons
(iii) State the H–N–H bond angle in an ammonia molecule.
(iv) Explain why the ammonia molecule is polar.
(b) Ammonia reacts with hydrogen ions forming ammonium ions, NH4+.
(i) State the H–N–H bond angle in an ammonium ion.
(ii) Explain why the H–N–H bond angle of NH3 is different from the H–N–H bond
angle of NH4+; referring to both species in your answer.
73. In 1954 Linus Pauling was awarded the Chemistry Nobel Prize for his work on the nature of the
chemical bond. Covalent bonds are one example of intramolecular bonding.
Explain the formation of the following.
(i) σ bonding
(ii) π bonding
(iii) double bonds
(iv) triple bonds
74. Atomic orbitals can mix by hybridization to form new orbitals for bonding.
Identify the type of hybridization present in each of the three following molecules.
Deduce and explain their shapes.
(i) OF2
(ii) H2CO
(iii) C2H2
75. Three scientists shared the Chemistry Nobel Prize in 1996 for the discovery of fullerenes.
Fullerenes, like diamond and graphite, are allotropes of the element carbon.
(i) State the structures of and the bonding in diamond and graphite.
(ii) Compare and explain the hardness and electrical conductivity of diamond and graphite.
(iii) Predict and explain how the hardness and electrical conductivity of C60 fullerene would
compare with that of diamond and graphite.
76. State the type of bonding in the compound SiCl4. Draw the Lewis structure for this
compound.
77. Outline the principles of the valence shell electron pair repulsion (VSEPR) theory.
78. (i) Use the VSEPR theory to predict and explain the shape and the bond angle of each of the
molecules SCl2 and C2Cl2
(ii) Deduce whether or not each molecule is polar, giving a reason for your answer.
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79. For the following compounds
PCl3, PCl5, POCl3
(i) Draw a Lewis structure for each molecule in the gas phase.
(Show all non-bonding electron pairs.)
(ii) State the shape of each molecule and predict the bond angles.
(iii) Deduce whether or not each molecule is polar, giving a reason for your answer.
80. (i) Explain the meaning of the term hybridization.
(ii) Discuss the bonding in the molecule CH3CHCH2 with reference to
the formation of σ and π bonds
the length and strength of the carbon-carbon bonds
the types of hybridization shown by the carbon atoms
81. Draw a Lewis structure of a water molecule, name the shape of the molecule and state and
explain why the bond angle is less than the bond angle in a tetrahedral molecule such as
methane. (Total 4 marks)
82. Predict and explain the order of the melting point for propanol, butane and propanone with
reference to their intermolecular forces. (Total 4 marks)
83. (a) Draw the Lewis structures for the compounds XeF4, PF5 and BF4–.
(3)
(b) Use the valance shell electron pair repulsion (VSEPR) theory to predict the shapes of the
three compounds in (a). State and explain the bond angles in each of the three
compounds. (3)
(Total 6 marks)
84. (a) State the meaning of the term hybridization. State the type of hybridization shown by the
nitrogen atoms in N2, N2H2 and N2H4.
(4)
(b) By referring to the N2H2 molecule describe how sigma () and pi () bonds form and
describe how single and double bonds differ. (4)
(Total 8 marks)
85. (i) Explain why the first ionization energy of magnesium is lower than that of fluorine. (2)
(ii) Write an equation to represent the third ionization energy of magnesium. Explain why the
third ionization energy of magnesium is higher than that of fluorine. (3)
(Total 5 marks)
86. The elements sodium, aluminium, silicon, phosphorus and sulfur are in period 3 of the periodic
table.
Describe the metallic bonding present in aluminium and explain why aluminium has a higher
melting point than sodium.
12
87. Draw the Lewis structure of NCl3. Predict, giving a reason, the Cl – N – Cl bond angle in NCl3.
(Total 3 marks)
88. Draw the Lewis structures, state the shapes and predict the bond angles for the following
species.
(i) PCl5
(3)
(ii) SCl2
(3)
(iii) ICl4–
(3)
89. (a) (i) State the meaning of the term hybridization. (1)
(ii) State the type of hybridization around the carbon atoms in C60 fullerene, diamond
and graphite. (3)
(iii) Explain why graphite and C60 fullerene can conduct electricity.
(2)
(b) (i) Compare how atomic orbitals overlap in the formation of sigma () and pi ()
bonds. (2)
(ii) State the number of sigma bonds and pi bonds in H2CC(CH3)CHCH2.
(2)
(Total 10 marks)
90. Arrange the following in decreasing order of bond angle (largest one first), and explain your
reasoning.
NH2–, NH3, NH4
+
(Total 5 marks)
91. (i) Outline the principles of the valence shell electron pair repulsion (VSEPR) theory. (3)
(ii) Use the VSEPR theory to deduce the shape of H3O+ and C2H4. For each species, draw the
Lewis structure, name the shape, and state the value of the bond angle(s). (6)
(iii) Predict and explain whether each species is polar. (2)
(iv) Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or
N-H would be most polar. (2)
(Total 13 marks)
92. Predict and explain which of the following compounds consist of molecules:
NaCl, BF3, CaCl2, N2O, P4O6, FeS and CBr4.
(Total 2 marks)
93. Diamond, graphite and C60 fullerene are three allotropes of carbon.
13
(i) Describe the structure of each allotrope. (3)
(ii) Compare the bonding in diamond and graphite. (2)
(Total 5 marks)
94. State two physical properties associated with metals and explain them at the atomic level. (Total 4 marks)
95. (i) Apply the VSEPR theory to deduce the shape of 2NO , ICl5 and SF4. For each species,
draw the Lewis (electron dot) structure, name the shape, and state the value of the bond
angle(s). (9)
(ii) Discuss the bond angle(s) in SF4.
(1)
(iii) Explain the hybridization involved in the C2H4 molecule.
(4)
(iv) State the hybridization involved in the 2NO ion and comment on the nitrogen-oxygen
bond distances. (2)
(v) Using Table 7 of the Data Booklet, predict and explain which of the bonds O-H, O-N or
N-H would be most polar. (2)
(Total 18 marks)
96. (a) Draw the Lewis structure of methanoic acid, HCOOH. (1)
(b) In methanoic acid, predict the bond angle around the (2)
(i) carbon atom. .....................................................................................................
(ii) oxygen atom bonded to the hydrogen atom. ...................................................
(c) State and explain the relationship between the length and strength of the bonds between
the carbon atom and the two oxygen atoms in methanoic acid. (3)
(Total 6 marks)
97. (a) Explain the meaning of the term hybridization.
(b) State the type of hybridization shown by the carbon atom in the H–C≡N molecule, and
the number of and bonds present in the C≡N bond.
(c) Describe how and bonds form.
98. D
99. C
100. A
101. B
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102. B
103. C
104. C
105. B
106. D
107. C
108. D
109. C
110. A
111. B
112. B
113. B
114. B
115. B
116. B
117. D
118. C
119. B
120. D
121. A
122. D
123. B
124. B
125. D
126. B
127. D
128. B
129. B
130. D
131. C
132. A
133. C
134. D
135. B
136. B
137. D
138. C
139. C
140. D
141. B
142. A
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143. C
144. B
145. B
146. A
147. D
148. B
149. C
150. A
151. B
152. A
153. B
154. D
155. A
156. B
157. B
158. C
159. D
160. (i) as molecules become larger/heavier/have higher Mr values/
number of electrons increases; van der Waals’/London/
dispersion forces increase; 2
(ii) hydrogen bonding between molecules in H2O; this bonding is stronger
(than van der Waals’ forces); 2
Must be an implied comparison with (i)
[4]
161. (i) tetrahedral (accept correct 3-D diagram);
bent/V-shape/angular (accept suitable diagram); 2
(ii) 105° (accept 103 – 106°);
lone pairs repel each other more than bonding pairs; 2
Do not accept repulsion of atoms.
[4]
162. bonds are polar as Cl more electronegative than Si;
Allow ―electronegativities are different‖
molecule is symmetrical, hence polar effects cancel out/OWTTE; 2 [2]
163. (a) (i) A – sodium iodide, B – sodium, C – iodine (three correct [1]); 1
Accept correct formulas.
(ii) A – ionic bonding;
B – metallic bonding;
C – van der Waals’ forces (and covalent bonding); 3
(b) (i) (for Na) (lattice of) positive ions/atoms;
delocalized/free electrons/sea of electrons;
(for NaI) oppositely charged ions/positive and negative ions;
free to move (only) in molten state; 4
(ii) forces between I2 molecules are weak;
ionic/metallic bonding strong(er); 2 [10]
16
164. for H2S, H2Se and H2Te, as size/mass/Mr increases, van der Waals’
forces increase (and b pt. increases);
H2O experiences H–bonding;
H–bonding stronger than van der Waals’/explanation of H–bonding; 3 [3]
165. (i)
C C CO O O
O O
O
2–
OTTWE 3
Award [1] each. Need charge on CO32–
for [1].
Penalize missing lone electron pairs only once.
(ii) CO32–
;
bond order 3
11 /3
11 bonds each compared to double bonds in CO2 and
triple bond in CO;
the fewer the number of bonding electrons, the less tightly nuclei
are held together, the longer the bond; 3 [6]
166. (i) O C O
H S H
;
;
2
Accept dots, crosses, a combination of dots and crosses or a
line to represent a pair of electrons.
(ii) CO2 is linear;
two charge centres or bonds and no lone pairs (around C);
H2S is bent/v-shaped/angular;
two bond pairs, two lone pairs (around S); 4
(iii) CO2 is non-polar, H2S is polar;
bond polarities cancel CO2 but not in H2S; 2
[8]
167. CH3Cl – dipole-dipole attractions;
CH4 – van der Waals’/dispersion/London forces;
CH3OH – hydrogen bond; 3
[3]
168. (i) C2H6 < CH3CHO, < C2H5OH < CH3COOH; 2
Award [2] if all correct, [1] if first and last correct.
(ii) C2H6 non polar;
CH3CHO polar;
C2H5OH polar;
CH3COOH polar;
Award [2] for all four correct, [1] for 3 or 2 correct.
boiling point depends on intermolecular forces;
least energy required for van der Waals’ forces/maximum energy
for hydrogen bonding;
C2H6 van der Waals’ forces only;
17
CH3CHO dipole-dipole;
C2H5OH and CH3COOH hydrogen bonding;
hydrogen bonding is stronger in CH3COOH/greater polarity/
greater molecular mass/greater van der Waals’ forces; 8 [10]
169. (a) (i) (relative) measure of an atom’s attraction for electrons; in a bond; 2
(ii)
+ H
+ H
N
N
H
H
H
H
+
+
+
+
–
–xx
xx
hydrogen bonding
Suitable diagram indicating
dipoles;
lone pairs of electrons;
hydrogen bonding; 3
(iii) 107°; 1
Accept answer in range 107 to 109° .
(iv) molecule is asymmetrical/OWTTE; 1
(b) (i) 109.5°; 1
(ii) NH4+ has four bonding pairs
(around central atom so is a regular tetrahedron);
NH3 has three bonding pairs (of electrons) and one non-bonding pair;
non-bonding pairs (of electrons) exert a greater repulsive force; 3
Accept suitable diagrams.
[11]
170. (i) ―head on‖ overlap of (2) orbitals;
along axial symmetry/along a line drawn through the
2 nuclei/OWTTE; 2
Accept suitable diagram for 2nd mark.
(ii) parallel p orbitals overlap sideways on;
above and below the line drawn through the 2 nuclei/OWTTE; 2
Accept suitable diagram for 2nd mark.
(iii) 1 σ and 1 π/σ and π; 1
(iv) 1 σ and 2 π/σ and π; 1 [6]
171. (i) OF2
sp3;
V-shaped/bent/angular;
2 bonding + 2 non-bonding (electron pairs); 3
(ii) H2CO
sp2;
trigonal planar;
18
2 areas of electron density/negative charge centres; 3
(iii) C2H2
sp;
linear;
2 areas of electron density/negative charge centres; 3
Accept suitable diagrams for shapes.
Allow [2] for ECF if correct explanation given for
incorrect formula, e.g. C2H4.
172. (i)
Diamond giant molecular/macromolecular/3-D covalent bonds only;
Graphite covalent bonds and van der Waals’ forces layer structure;
Award [1] for both shape and bonding in each case. 2
Accept suitable diagrams.
(ii)
Diamond Graphite
poor/non-conductor good conductor
no delocalized electrons delocalized electrons
hard soft
rigid structure layers can slide
Award [1] per row. 4
(iii) softer than diamond/harder than graphite;
as C60 molecules can move over each other;
conducts better than diamond/worse than graphite;
as C60 has less delocalisation (of the unpaired bonding electrons)
than graphite; 4 [10]
173. Si—Cl bonds are covalent; 3
Cl ClSi
Cl
Cl
Accept lines for electron pairs.
Award [1] for covalent bonds and [1] for lone pairs.
[3]
174. find number of electron pairs/charge centres in (valence shell of) central atom;
electron pairs/charge centres (in valence shell) of central atom repel each other;
to positions of minimum energy/repulsion/maximum stability;
pairs forming a double or triple bond act as a single bond;
non-bonding pairs repel more than bonding pairs/OWTTE; 3 max
Do not accept repulsion between bonds or atoms.
Award [1] each for any three points.
[3]
175. (i) SCl2 two bonding pairs, two non-bonding pairs;
angular/bent/non-linear/V-shaped;
Both these marks can be scored from a diagram.
90° < angle < 107°;
C2Cl2 two charge centres around each C;
linear;
Both these marks can be scored from a diagram.
19
angle = 180°; 6
(ii) SCl2 is polar;
C2Cl2 is non-polar; 3
No net dipole movement for C2Cl2 but angular SCl2 has a
resultant dipole / OWTTE;
Mark can be scored from a diagram.
Allow ECF based on the answers given to (i).
[9]
176. (i)
Award [1] for each correct Lewis structure.
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
Cl
P
P
O
P
PCl
PCl
POCl
3
5
3
3
Accept use of dots or crosses to represent electron pairs.
Subtract [1] if non-bonding pair on P in PCl3 is missing.
Subtract [1] if non-bonding pair(s) on Cl or O are missing.
Accept legitimate alternatives for POCl3, e.g. see below.
Cl P Cl
O
Cl
Cl P ClO
Cl
(ii)
PCl3 PCl5 POCl3
trigonal pyramid; trigonal bipyramid; tetrahedral;
20
Accept answers in range
100° to 108°;
90° and 120°; Accept answers in range
100° to 112°;
Allow ECF if based on legitimate chemical structure. 6
(iii)
PCl3 PCl5 POCl3
polar, polarities do not
cancel/OWTTE;
non-polar, polarities
cancel/OWTTE;
polar, polarities do not
cancel/OWTTE;
3
Award [2] for three polarities correct, [1] for two polarities
correct, and [1] for correct reason(s).
Accept argument based on dipole moments.
Allow ECF if based on legitimate chemical structure.
177. (i) combining of atomic orbitals to form new orbitals/OWTTE; 1
(ii) σ : overlap of orbitals between nuclei/end-on overlap;
: overlap above and below line joining nuclei/sideways overlap;
Award [1] if candidate counts bonds (8 σ, 1 π), or
describes all three types of bonds
(i.e. C—H is σ, C—C is σ, C=C is σ and π).
single bonds longer than double;
double bonds stronger than single;
C of CH3 is sp3;
other two C are sp2; 6
Accept suitable diagrams.
178.
OH
H
OH H
Allow a combination of dots, crosses or lines.
bent/V shaped/angular
104.5;
Accept answers in range 104 to 106.
repulsion of the two non-bonding pairs of electrons forces bond angle
to be smaller/non-bonding pairs repel more than bonding pairs; 4 [4]
179. butane < propanone < propanol;
butane has van der Waals’ forces;
Accept vdW, dispersion or London forces or attractions
between temporary dipoles.
propanone has dipole-dipole attractions;
propanol has (the stronger) H-bonding; 4 [4]
180. (a)
21
; lone pairs on Xe required for the mark.
; square brackets and charge required for the mark.
FP
F
F
F
F
x
xxx
xxx
x
x
x
x
x
x
x xxx
xxxx
xx
xx
xx
xx
x
XeF
FF
Fxxx
x
x
xx
x
x
x
x
x
x
xx
xx
x
x
xxx
xxx x
x x
xxx
x
x
xx
x
xx
x
x
xxx
xx
x
x
x
B
F
FF
F
x xxx
Accept any combination of dots, crosses and lines.
Penalise missing fluorine lone pairs once only.
(b) XeF4
Square planar and 90;
PF5
trigonal bipyramid and 90 and 120;
BF4–
Tetrahedral and 109.5/109; 3
Allow clear suitable diagrams instead of name.
No ECF from (a).
[6]
181. (a) hybridization: mixing/merging of atomic orbitals;
N2 sp;
N2H2 sp2;
N2H4 sp3; 4
(b) bonds (result from the) overlapping of orbitals end to end/along inter-nuclear
axis;
bonds (result from the) overlapping of parallel/sideways p orbitals;
(single bonds) bonds only;
(double bonds) have a σ bond and a bond; 4
Suitable clear and labelled diagrams acceptable for all marks.
[8]
182. (i) electron removed from higher energy level/further from nucleus/greater
atomic radius;
increased repulsion by extra inner shell electrons/increased shielding
effect; 2
(ii) Mg2+
(g) Mg3+
(g) + e;
22
(even though) valence electrons in the same shell/main energy level/
Mg2+
has noble gas configuration;
Mg has greater nuclear/core charge/more protons; 3 [5]
183. delocalized electrons;
(attracted) to positive ions;
more delocalized/mobile/outer shell electrons/higher ionic charge; 3 [3]
184.
NCl
Cl
Cl
All electrons must be shown.
Accept molecular structures using lines to represent bonding
and lone electron pairs.
bond angle: 107109
greater repulsion between lone pair and bonding pairs/OWTTE; 3
NOT between electron pairs and atoms.
Award [1 max] if lone pair missed on nitrogen, ECF for bond
angle of 120.
[3]
185. (i)
;PClx
x x
x xCl Cl
Cl
ClCl
Cl
ClCl
ClP
trigonal bipyramidal;
90;
120;
180; 3
Award [1] for 2 correct bond angles.
(ii)
S
Cl
Clxx
xx
xx
;Cl
Cl
S
Bent/angular/V-shaped;
100107; 3
(iii)
;
Cl Cl
ClCl
I
square planar; 3
90;
23
No ECF allowed.
Penalize once only [1] mark for missing lone pairs.
Accept structures using lines to represent bonding and lone
electron pairs.
[9]
186. (a) (i) mixing/combining of atomic orbitals/OWTTE; 1
(ii) C60 fullerene: sp2;
graphite : sp2;
diamond: sp3; 3
(iii) each carbon atom is bound to 3 other carbon atoms/ bonding;
leading to delocalized electrons; 2
(b) (i) sigma/ bonds are formed by orbitals overlapping end to end/
along the internuclear axis/along line directly between nuclei;
Accept suitable diagram.
pi/ bonds are formed by p orbitals overlapping sideways; 2
Accept suitable diagram.
(ii) 12 sigma bonds;
2 pi bonds; 2 [10]
187. NH4+ > NH3 > NH2
–;
NH4+ has four bonded electron pairs (and no lone electron pairs);
NH3 has three bonded electron pairs and one electron lone pair;
NH2– has two bonded electron pairs and two electron lone pairs;
Accept correct Lewis structures with lone electron pairs clearly
shown.
lone pair-lone pair > lone pair-bonded pair > bonded pair-bonded pair/
lone pairs of electrons repel more than bonding pairs of electrons/OWTTE; 5
Do not accept repulsion between atoms.
[5]
188. (i) Find number of electron pairs/charge centres in (valence shell of)
central atom;
electron pairs/charge centres (in valence shell) of central atom repel
each other;
Any one of the following:
to positions of minimum energy/repulsion/maximum stability;
pairs forming a double or triple bond act as a single bond;
non-bonding pairs repel more than bonding pairs/OWTTE; 3 max
Do not accept repulsion between bonds or atoms.
24
(ii) 6
Species Lewis (electron-dot)
structure Shape Bond angle(s)
H3O+ O
H HH
– +
;
Trigonal/triangular
pyramidal;
Allow values in the
range 106° to 109.5°;
C2H4
H;
H
H H
Trigonal/triangular
planar;
Allow values of
approximately 120°;
Accept crosses and dots for electrons in Lewis structures also.
As the Lewis structures were asked for, and not 3D
representations, do not penalize incorrectly drawn geometries.
Do not accept structure of hydronium cation without lone pair
on oxygen.
No penalty for missing charge.
(iii) H3O+: is polar and explanation either using a diagram or in words,
involving the net dipole moment;
e.g. the three individual O-H bond dipole moments add as
vectors to give a net dipole moment.
C2H4: is non-polar and explanation either using a diagram or in words,
involving no net dipole moment; 2
e.g. the vector sum of the individual bond dipole moments is
zero.
For simple answers such as bond polarities do not cancel for
H3O+ and do cancel for C2H4, Award [1], only for the last two
marking points.
(iv) O-H is most polar;
O-H has greatest difference between electronegativities/calculation
showing values of 1.4, 0.5 and 0.9 respectively; 2 [13]
189. BF3, N2O, P4O6 and CBr4;
Non-metals only/small difference in electronegativity values of the elements; 2 [2]
190. (i) 3
Allotrope Structure
Diamond 3D array/network involving tetrahedral
carbons/each carbon atom joined to four others;
Graphite layer structure involving trigonal (triangular) planar
carbons/with each carbon atom joined to three
others/with hexagonal (six-membered) rings of
carbon atoms;
C60 fullerene truncated icosahedrons;
Accept carbon atoms form a ‗ball‘ with 32 faces, of
which 12 are pentagons and 20 are hexagons,
exactly like a soccer ball. Do not accept soccer ball
alone.
(ii) Diamond: covalent bonds (only);
Graphite: covalent bonds and the separated layers held together by
(weak) London/van der Waals’/dispersion forces; 2
25
[5]
191. Electrical conductivity:
Bonding electrons are delocalised;
Current flow occurs without displacement of atoms within the metal/
able to flow within the metal;
Malleability:
Can be hammered into thin sheets;
atoms capable of slipping with respect to one another; 4 [4]
192. (i)
Species Lewis (electron-dot)
structure Shape Bond angle(s)
NO2–
O
N
O;
Bent/V-
shaped/angular; 109.5° < θ < 120°;
ICl5
Cl
I ;
Cl
Cl
Cl
Cl
Square pyramidal;
|Inplane Cl-I-out-of-plane Cl| < 90°;
Allow corresponding correct
statement for other correctly identified
bond angles.
SF4
F
SF
FF
;
See-saw;
|Equatorial F-S-Equatorial F| < 120°;
Allow corresponding correct
statement for axial-equatorial and
axial-axial F-S-F angles.
9
Accept crosses and dots for electrons in the Lewis structures
also.
If all ideal bond angles are given, penalize once only.
As the Lewis structures were asked for, and not 3D
representations, do not penalize incorrectly drawn geometries.
(ii) (equatorial F-S-equatorial F) less than 120° since non-bonding electron pairs
(exert greater repulsive forces and thus) compress the bond angles/OWTTE; 1
(iii) orbital diagram representation of carbon ground-state going to carbon
excited-state electron configuration;
mixing of orbitals to give three new entirely equivalent hybrid orbitals, sp2,
on each carbon;
sp2 orbitals trigonal (triangular) planar in shape;
unhybridized orbitals overlap to give π-bond; 4
(iv) sp2;
both N-O bond lengths equal, (intermediate between double and single
bonds) due to resonance/delocalisation; 2
(v) O-H is most polar;
O-H has greatest difference between electronegativities/calculation
showing values of 1.4, 0.5 and 0.9 respectively; 2 [18]
193. (a)
26
H
H
C
O
O 1
No mark without lone electron pairs.
Correct shape not necessary.
Do not award mark if dots/crosses and bond lines are shown.
Accept lone pairs represented as straight lines.
(b) O − C − O = 120°/H − C − O = 120°;
C − O − H = 109°/<109°; 2
No mark for 109.5°
Accept answer in range 100–109°
(c) length: C = O < C − O; strength: C = O > C – O;
greater number of electrons between nuclei pull atoms together and require greater
energy to break;
Or
double bonds are shorter/single bonds are longer;
double bonds are stronger/single bonds are weaker; 3
Accept stronger attraction between nuclei and (bonding)
electrons.
[6]
194. (a) mixing/joining together/combining/merging of atomic orbitals to
form molecular orbitals/new orbitals/orbitals of equal energy; 1
Accept specific example such as mixing of s and p orbitals.
(b) sp;
Do not award mark if sp2 or sp
3 is also stated.
one sigma and two pi (bonds); 2
(c) ( bond formed by) end-on/axial overlap;
electrons/electron density between the two (carbon) atoms/OWTTE;
(π bond formed by) sideways/parallel overlap;
electrons/electron density above and below bond/OWTTE; 4
Marks can be scored from a suitable diagram.
Do not award 2nd
and 4th
marks if electrons are not mentioned.
[7]
