Module 2: Foundations in Chemistry V2 Flashcards
1
Q
Define the term relative atomic mass
A
Average mass of an atom ✓
relative to 1/12th the mass ✓
of one atom of Carbon-12 ✓
2
Q
Write the formula for an ammonia and an ammonium ion
A
3
Q
Write the formula for a nitrate ion and nitric acid
A
4
Q
Write the formula for a hydroxide ion and sodium hydroxide
A
5
Q
Write the formula for a hydrogen carbonate ion and sodium hydrogen carbonate
A
6
Q
Write the formula for a silver ion and silver carbonate
A
7
Q
Write the formula for a carbonate ion and sodium carbonate
A
8
Q
Write the formula for a sulfate ion and copper (II) sulfate
A
9
Q
Write the formula for a dichromate ion and potassium dichromate (VI)
A
10
Q
Write the formula for a phosphate ion and phosphoric acid
A
11
Q
In terms of protons, neutrons and electrons, how is an atom of Eu-151 different from an atom of Eu-153
A
153-Eu has two more neutrons ✓
12
Q
In terms of protons, neutrons and electrons, how is an atom of Eu-151 similar from an atom of Eu-153?
A
Both have 63 protons and 63 electrons ✓
13
Q
Define the term relative isotopic mass
A
Mass of the isotope compared to 1/12th the mass of a C-12 atom ✓
14
Q
Which isotope is used as the standard against which the masses of the two rubidium isotopes are measured?
A
C-12 ✓
15
Q
A sample of carbon was found to contain 95% of C-12 and 5% of C-13. How could this information be obtained experimentally
A
Mass spectrometry ✓
16
Q
State three differences between isotopes of the same element
A
Different numbers of neutrons ✓
Different mass numbers ✓
Different physical properties ✓
17
Q
State a similarity between isotopes of the same element
A
Same number of protons and electrons ✓
18
Q
Define what is meant by the weighted mean mass
A
Average mass taking into account the relative abundancies of the isotopes ✓
19
Q
Define empirical formula
A
The simplest whole number ratio of atoms present in a compound ✓
20
Q
Define molecular formula
A
The actual number of atoms of each element in a molecule ✓
21
Q
Suggest 5 reasons why the yield of a reaction will not be 100%
A
The reaction may not go to completion ✓
The reactants may not be pure ✓
Other reactions may occur which lead to by-products ✓
Products may be left behind in the reaction vessel ✓
Separation and purification of the product may result in loss of some of the product ✓
22
Q
Cracking involves breaking down larger hydrocarbons into smaller more useful hydrocarbons. State and explain the percentage atom economy for these reactions
A
100% ✓
All products are useful ✓
No waste products made ✓
23
Q
Addition reactions have 100% atom economy. Explain why?
A
Have only one product ✓
No waste products made ✓
24
Q
What term is given to the .xH2O part of the formula of hydrated salts
A
Number of water of crystalisation ✓
25
Define the term hydrated salt
Contains water of crystalisation ✓
26
Define the term anhydrous salt
Contains no water of crystalisation ✓
27
What is meant by the term standard solution?
A solution of precisely known concentration ✓
28
A student carries out an experiment to find the value of x in the formula of hydrated strontium chloride, SrCl2*xH2O. The student’s method is outlined below.
Step 1 Weigh an empty crucible.Add SrCl2*xH2O to the crucible and reweigh.
Step 2 Heat the crucible and contents for 10 minutes.Allow to cool and reweigh.
Step 3 Heat the crucible and residue for another 5 minutes.Allow to cool and weigh the crucible and residue.Repeat step 3 a further two times.
Suggest why the student takes four readings of the mass of the crucible and residue
To make sure all the water had been removed ✓
29
A student carries out an experiment to find the value of x in the formula of hydrated strontium chloride, SrCl2*xH2O. The student’s method is outlined below.
Step 1 Weigh an empty crucible.Add SrCl2*xH2O to the crucible and reweigh.
Step 2 Heat the crucible and contents for 10 minutes.Allow to cool and reweigh.
Step 3 Heat the crucible and residue for another 5 minutes.Allow to cool and weigh the crucible and residue.Repeat step 3 a further two times.
Suggest two modifications to the method that would reduce the percentage uncertainty in the mass of the residue.
Use balance that weighs to more decimal places ✓
Use a larger mass (of hydrated strontium chloride) ✓
30
Define the term acid
Proton donor ✓
31
Define the term base
Proton acceptor ✓
32
Define what is meant by a "strong acid"
Proton donor ✓
That completely dissociates ✓
33
Define what is meant by a "weak acid"
Proton donor ✓
That partially dissociates ✓
34
State one important use for ammonium sulfate
Fertiliser ✓
35
Apart from the manufacture of a ammonium sulfate state two other large-scale uses of ammonia
Manufacture of explosives ✓
Refrigeration ✓
36
State a hazard with hydrogen gas
Explosive ✓
37
How do you test for the presence of hydrogen gas
Light a splint over the test tube ✓
Squeaky pop ✓
38
Describe how you can test for carbon dioxide gas
Bubble gas through limewater ✓
White precipitate ✓
39
Describe how you can test for the presence of oxygen gas
Light a splint, and blow it out ✓
Insert lit splint into test tube ✓
Lit splint re-ignites ✓
40
What is meant by the term oxidation
Loss of electrons ✓
41
What is meant by the term reduction
Gain in electrons ✓
42
The substance that is reduced in a redox reaction is known as a…
Oxidising agent ✓
43
The substance that is oxidised in a redox reaction is known as a…
Reducing agent ✓
44
Define the term redox reaction
Where one substance is reduced, and another is oxidised ✓
45
Using oxidation numbers, explain why the reaction involves both oxidation and reduction
2Sr(NO3)2(s) → 2SrO(s) + 4NO2(g) + O2(g)
O goes from -2 to 0 ✓
N goes from +5 to +4 ✓
N is reduced and O is oxidised ✓
46
Define the term alkali
Releases OH- ions into solution ✓
47
Define the term salt
When the H+ in an acid is replaced by a metal ion ✓
48
Define what is meant by a concordant
Titre values that are within 0.1 cm3 of each other ✓
49
The student follows the method below.
1. Dissolve a weighed mass of A in 100 cm3 of distilled water and make the solution up to 250 cm3 in a beaker.
2. Add the solution of A to a burette.
3. Titrate the solution of A with a standard solution of sodium hydroxide, NaOH.
The student realised that the solution of A had not been prepared correctly.How should the student have made up the solution?
Use a volumetric flask ✓
50
Several students titrate 25.00 cm3 of the same solution of sodium hydroxide, NaOH(aq) with hydrochloric acid, HCl (aq). One student obtains a larger titre than the other students.
Which procedure explains the larger titre?
A The burette readings are taken from the top of the meniscus instead of the bottom of the meniscus.
B The conical flask is rinsed with water before carrying out the titration.
C An air bubble is released from the jet of the burette during the titration.
D The pipette is rinsed with water before filling with NaOH(aq).
C An air bubble is released from the jet of the burette during the titration. ✓
## Footnote
Air displaced water in burette, So a greater measured volume is needed.
51
Several students titrate 25.00 cm3 of the same solution of sodium hydroxide, NaOH(aq) with hydrochloric acid, HCl (aq). One student obtains a smaller titre than the other students.
Which procedure explains the smaller titre?
A The burette readings are taken from the top of the meniscus instead of the bottom of the meniscus.
B The conical flask is rinsed with water before carrying out the titration.
C An air bubble is released from the jet of the burette during the titration.
D The pipette is rinsed with water before filling with NaOH(aq).
D The pipette is rinsed with water before filling with NaOH(aq) ✓
## Footnote
NaOH is more dilute now, same volume, less moles of NaOH reacting in conical flask, less acid required to neutralise it.
52
State and explain the effect, if any, on the titre value if readings from the burette are taken from the top of the meniscus
No effect ✓
Burettes measure by difference ✓
53
54
55
56
57
58
Calculate the percentage by mass of Ca in CaO
59
Calculate the number of moles of NaOH in 0.2g of the solid
60
61
In a reaction, 0.14 g of calcium oxide is reacted with excess hydrochloric acid.
Calculate the mass of calcium chloride produced.
62
In a reaction, 0.212 g of calcium oxide is reacted with excess hydrochloric acid.
Calculate the mass of calcium chloride produced.
63
64
One coin has a mass of 5.00 g and contains 84.0% of copper, by mass.
Calculate the number of copper atoms in one coin. Give your answer in standard form and to three significant figures.
65
66
67
68
Rearrange the Ideal gas equation, pV = nRT to make
a) p, the subject
b) V, the subject
c) n, the subject
d) R, the subject
e) T, the subject
69
70
71
72
73
Define what is meant by:
Solute
Solvent
Solution
Solute – Substance we dissolve in a solvent ✓
Solvent – Substance we dissolve a solute in ✓
Solution – Substance dissolved in a solvent ✓
74
75
76
77
Describe how you would make a 250 cm3 standard solution.
Dissolve ____ g of the solid in distilled water in a beaker ✓
Transfer the solution to a 250 cm3 volumetric flask ✓
Ensure that all solution is washed out of beaker ✓
Make solution up to 250 cm3 with distilled water ✓
Ensure thorough mixing by inverting the flask several times ✓
78
79
In the manufacture of titanium, what mass of titanium can theoretically be formed when 1.00 kg of titanium chloride reacts with 0.100 kg of magnesium?
## Footnote
Calculations must be done with the limiting reagent i.e. Mg in this case
80
Ethanol can be oxidised to form ethanal. The theoretical yield was 8.80 g.
The ethanal produced was weighed and found the actual yield was 2.10 g.
Calculate the percentage yield.
81
Ethanol can be oxidised to form ethanal.
The percentage yield of the reaction completed in a laboratory is 42.3%. A student is trying to synthesis ethanol, from ethanol. He calculates the theoretical yield as 3.40g
Calculate the amount of ethanal, in g, that the student would produce.
82
The reaction below shows how ethanol can be oxidised by to form ethanal.
9.2g of ethanol was reacted with an oxidising agent (excess) to produce 1.0 g of ethanal.
10.
Calculate the percentage yield.
83
Bromomethane is reacted with sodium hydroxide to make methanol.
Calculate the percentage atom economy for this reaction.
84
85
Calculate the empirical formula of the compound formed when 5.4 g of Aluminium reacts with 4.8 g of oxygen
86
87
Write a balanced equation, including observations for the reaction between magnesium and hydrochloric acid.
88
Write a balanced equation, including observations for the reaction between calcium oxide and hydrochloric acid.
89
Write a balanced equation, including observations for the reaction between sodium hydroxide and nitric acid.
90
Write a balanced equation, including observations for the reaction between sulfuric acid and calcium carbonate.
91
Write a balanced equation, including observations for the reaction between sulfuric acid and aluminium oxide.
92
Write a balanced equation, including observations for the reaction between hydrochloric acid and gallium oxide.
93
Write a balanced equation, including observations for the reaction between hydrochloric acid and huntite.
## Footnote
I balanced found the salts, and products for reacting a metal carbonate and acid first.
Then balanced according to metals, then C, then I balanced water.
94
Write a balanced equation, including observations for the reaction between ethanoic acid and aluminium oxide.
95
Write a balanced equation, including observations for the reaction between propanoic acid and magnesium hydroxide.
96
Write a balanced equation, including observations for the reaction between ammonium hydroxide and nitric acid
97
Calculate the mean titre value.
## Footnote
op manoever - not using concordant results
og manoever - you used the concordant results
98
What is meant by "concordant" results
Titre values that are within 0.1 cm3 of each other ✓
99
Explain why a rough titration is done?
to estimate the titre value ✓
100
A student is doing a titration and notices that the droplets hanging off the tip. He is worried this might affect his results. What should he do?
Nothing ✓
101
102
Define what is meant by a strong acid
Proton donor ✓
That fully dissociates ✓
103
Define what is meant by a weak acid
Proton donor ✓
That partially dissociates ✓
104
105
106
107
Suggest two ways in which percentage error can be reduced in the burette in a titration
Use a more dilute solution of acid/ alkali ✓
Larger volume measured for neutralisation, less uncertainty ✓
Use a burette that measures to a greater number of decimal places ✓
more precise value, less uncertainty ✓
| If given the choice of one, first is preferred.
108
Suggest how uncertainty can be reduced in the mass balance in the practical to determine the empircla formula of a compound
Use a greater mass✓
Larger mass measured, less uncertainty ✓
Use a mass balance that measures to a greater number of decimal places ✓
more precise value, less uncertainty ✓
109
Suggest how measurements of the volume of solution could be made consistent?
Measure from the bottom of the meniscus every time ✓
110
A student is titrating sodium hydrogencarbonate solution with the hydrochloric acid. He made a standard solution of sodium hydrogen carbonate in a volumetric flask. Instead of filling the volumetric flask with distilled water to the graduation mark, he filled it to the top of the flask. State and explain what effect this would have on the titre values.
Titre value would be less ✓
NaHCO3 solution would be more dilute OR less concentrated ✓
## Footnote
NaHCO3 more dilute, less moles in given volume, less moles in conical flask, less volume of acid from burette (titre value) required to neutralise it
111
A student is titrating sodium hydrogencarbonate solution with the hydrochloric acid. He made a standard solution of sodium hydrogen carbonate in a volumetric flask. He washed the pipette out with distilled water instead of sodium hydrogen carbonate solution, then proceeded to add the solution to the conical flask. State and explain the effect on the titre value.
Titre value would be less ✓
NaHCO3 solution would be more dilute OR less concentrated ✓
## Footnote
NaHCO3 more dilute, less moles in given volume, less moles in conical flask, less volume of acid from burette (titre value) required to neutralise it
112
A student is titrating sodium hydrogencarbonate solution with the hydrochloric acid. He made a standard solution of sodium hydrogen carbonate in a volumetric flask. He washed the burette out with distilled water instead of hydrochloric acid solution, then proceeded to run the titration. State and explain the effect on the titre value
Titre value would be larger ✓
hydrochloric acid solution would be more dilute OR less concentrated ✓
## Footnote
HCl more dilute, less moles in given volume, less moles in burette, more volume of acid from burette (titre value) required to neutralise the same moles of NaHCO3
113
In a practical, involving the reaction between sulfuric acid and copper carbonate, a student was trying to determine the percentage by mass of copper. The instructions are outlined and results are as shown.
Suggest why the student conducted the experiment while wearing a lab coat and goggles.
Sulfuric acid is corrosive ✓
114
In a practical, involving the reaction between sulfuric acid and copper carbonate, a student was trying to determine the percentage by mass of copper. The instructions are outlined and results are as shown.
Another student, conducted the same experiment and recorded a result of the gas produced as being 50 cm3 less. The volume is much less than expected, suggest two reasons why?
Gas escaped ✓
reaction did not go to completion ✓
115
In a practical, involving the reaction between sulfuric acid and copper carbonate, a student was trying to determine the percentage by mass of copper. The instructions are outlined and results are as shown.
Explain why it is not necessary to calculate the percentage error of the volume of sulfuric acid.
Sulfuric acid is in excess ✓
116
A student conducted an experiment to determine the relative atomic mass of magnesium. The instructions are outlined and results are as shown.
A student at a college in Burnley conducted the same experiment, but with same amount, in moles, of barium instead. The amount of hydrogen gas produced was the same. Suggest why?
1:1 ratio in stoichiometry ✓
same number of moles of Ba leads to same number of moles of hydrogen ✓
117
A student conducted an experiment to determine the relative atomic mass of magnesium. The instructions are outlined and results are as shown.
Another student repeated the practical but did the same reaction with calcium. The reaction was much faster. Suggest why?
Calcium is more reactive than barium ✓
118
A student conducted an experiment to determine the relative atomic mass of magnesium. The instructions are outlined and results are as shown.
State what effect the use of ethanoic acid, instead of sulfuric acid, would have on the rate of reaction, if any.
Rate is slower ✓
Ethanoic acid is a weaker acid ✓
119
A student conducted an experiment to determine the relative atomic mass of magnesium. The instructions are outlined and results are as shown.
A student at a college in Blackburn used 2000 cm3 of sulfuric acid. Explain why this had no effect on the volume of hydrogen gas produced.
No effect ✓
Sulfuric acid is in excess ✓
120
A student conducted an experiment to determine the relative atomic mass of magnesium. The instructions are outlined and results are as shown.
Suggest another piece of equipment that could be used to measure the volume of hydrogen gas more accurately
Gas syringe ✓
121
A student conducts a practical to determine the emprical formula of magnesium oxide. Method and results are shown below.
Explain why the crucible needed to be raised slightly and periodically.
To allow more oxygen into the crucible ✓
122
A student conducts a practical to determine the emprical formula of magnesium oxide. Method and results are shown below.
Another student conducted the same experiment.
They found the mass of magnesium oxide to be much less than those above.
Suggest what the student could do to receive more valid results.
Heat to constant mass ✓
123
What is meant by the term ionic bonding?
Strong electrostatic attraction between oppositely charged ions ✓
124
What is meant by the term covalent bonding?
A shared pair of electrons ✓
125
What is meant by the term dative covalent bonding?
A shared pair of electrons where both electrons are donated from one atom ✓
126
What is meant by the term metallic bonding?
Electrostatic attraction between positive metal ions and delocalised electrons ✓
127
Describe and explain what states of ionic compounds conduct electricity in i.e. NaCl, KI, MgO, NH4Cl
They can conduct electricity when molten or dissolved only ✓
Ions are free to move and carry a charge ✓
They cannot conduct electricity in the solid state ✓
Ions are fixed in position ✓
128
Describe and explain what states of metallic compounds conduct electricity in i.e. Al, Na, Mg, Ca
Both solid and liquid can conduct electricity ✓
Contain delocalised electrons which are mobile ✓
129
Describe and explain what states covalent compounds (not graphite/ graphene) can conduct electricity in i.e. S8, Br2, N2, O2
Cannot conduct electricity in any state ✓
No mobile charged particles ✓
Electrons are not mobile ✓
130
Describe and explain the melting and boiling points of ionic compounds i.e. NaCl, KI, MgO
Very high melting and boiling points ✓
Giant ionic lattice, Strong electrostatic attraction between oppositely charged ions ✓
Requires large amount of energy to overcome ✓
131
Describe and explain the melting and boiling points of simple covalent compound i.e. Cl2, Br2, N2, O2
Low melting and boiling points ✓
____ has a simple covalent lattice structure with induced dipole-dipole interactions between molecules ✓
Which are weak and require little energy to overcome ✓
132
Describe and explain the melting and boiling points of metals i.e. Na, Li, Mg, Fe
High melting and boiling points ✓
____ has metallic bonding, the electrostatic attraction between positive metal ions and delocalised electrons ✓
Metallic bonds are strong ✓
Large amount of energy is needed to overcome ✓
133
State the name of the structure that exists in ionic compounds
Giant ionic lattice ✓
134
Describe and explain the density of ice compared with water
Ice is less dense than water ✓
Hydrogen bonds hold H2O molecules further apart in ice ✓
135
State and explain two anomalous properties of H2O that depend on its intermolecular forces.
Ice is less dense than water ✓
Hydrogen bonds hold H2O molecules further apart ✓
Ice has a higher melting/ boiling point than expected ✓
Hydrogen bonds are relatively strong and must be broken ✓
136
Explain why HF has a higher boiling point than F2
HF has hydrogen bonds between molecules ✓
F2 has induced dipole-dipole interactions between molecules ✓
Hydrogen bonds are stronger and requires more energy to overcome ✓
137
Explain why HCl has a higher boiling point than H2
HCl has permanent dipole-dipole interactions between molecules ✓
H2 has induced dipole-dipole interactions between molecules ✓
Permanent dipole-dipole interactions are stronger and requires more energy to overcome ✓
138
Explain why H2O has a higher boiling point than H2S
H2O has hydrogen bonds between molecules ✓
H2S has permanent dipole-dipole interactions between molecules ✓
Hydrogen bonds are stronger and require more energy to overcome ✓
139
Explain why the O–H bonds in water and the N–H bonds in ammonia have dipoles.
Oxygen/ nitrogen is more electronegative ✓
Molecule has atoms with different electronegativities ✓
More electronegative atom attracts bonded electron pair more ✓
140
Describe and explain the intermolecular bonding in CH4, Br2 and F2
Induced dipole-dipole interactions ✓
Electrons move causing an uneven distribution in electrons ✓
Leading to induced dipoles in neighbouring molecules ✓
141
Explain why ammonium chloride has a higher boiling point compared to ammonia
Ammonium chloride has a giant ionic lattice structure ✓
Ammonia has a simple covalent lattice structure ✓
Ammonium chloride has ionic bonds, the strong electrostatic attraction between oppositely charged ions ✓
Ammonia has hydrogen bonding between molecules ✓
Ionic bonds are stronger than hydrogen bonds and require more energy to overcome ✓
142
State the four rules in electron configuration
1. Electrons are filled up the lowest energy sub-shells first. ✓
2. Electrons fill orbitals with the same energy singly before they start pairing. ✓
3. For ions, we remove/add electrons from the highest shell. ✓
4. 4s electrons are added first and removed first. ✓
143
How many electrons are in the fourth shell of ceasium
144
A shell contains 50 electrons. Which shell is this?
145
Show the electron configuration of a calcium ion, using an energy level diagram
## Footnote
Remember the charge is now 2+
146
Show the electron configuration of a bromine ions, using an energy level diagram
## Footnote
Remember the charge is now -1.
147
Show the electron configuration of potassium, using an energy level diagram
148
Write the electron configuration of a calcium ion, using an electrons in the box diagram
## Footnote
Remember the charge is now 2+
149
Write the electron configuration of a bromide ion, using an electrons in the box diagram
## Footnote
Remember the charge is now 1-
150
Write the electron configuration of potassium, using an electrons in the box diagram
151
Write the electron configuration of a calcium ion.
## Footnote
Remember the charge is now 2+
152
Write the electron configuration of potassium.
153
Write the electron configuration of a bromide ion.
## Footnote
Remember the charge is now 1-
154
Using a diagram, explain the conductivity of Mg.
Electrostatic attraction between positive metal ions and delocalised electrons ✓
Conducts electricity as delocalised electrons can move and carry a charge ✓
155
State and explain the bond angle around the C atom, bonded to the oxygen
Trigonal planar ✓
120 degrees ✓
There are 4 bonded pairs and 0 lone pairs around the C atom ✓
There are 3 bonding regions/ 3 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
156
State and explain the bond angle around the O atom bonded to H, in ethanoic acid. Include a bond angle in your answer.
Non-linear ✓
104.5 degrees ✓
There are 2 bonded pairs and 2 lone pairs around the O atom ✓
There are 2 bonding regions/ 4 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
157
State and explain the shape around the N atom in ammonium ions. Include a bond angle in your answer.
Tetrahedral ✓
109 degrees ✓
There are 4 bonded pairs and 0 lone pairs around the N atom ✓
There are 4 bonding regions/ 4 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bondedbonding pairs ✓
158
State and explain the shape around the N atom in ammonia. Include a bond angle in your answer.
Pyramidal ✓
107 degrees ✓
There are 3 bonded pairs and 1 lone pairs around the N atom ✓
There are 3 bonding regions/ 4 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
159
State and explain the shape around the C atom in carbon dioxide. Include a bond angle in your answer.
Linear ✓
180 degrees ✓
There are 4 bonded pairs and 0 lone pairs around the C atom ✓
There are 2 bonding regions/ 2 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
160
State and explain the shape around the B atom in boron trifluoride. Include a bond angle in your answer.
trigonal planar ✓
120 degrees ✓
There are 3 bonded pairs and 0 lone pairs around the B atom ✓
There are 3 bonding regions/ 3 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
161
State and explain the shape around the S atom in a molecule of sulfur hexafluoride. Include a bond angle in your answer.
Octahedral ✓
90 degrees ✓
There are 6 bondedbonding pairs and 0 lone pairs around the S atom ✓
There are 6 bonding regions/ 6 regions of electron density ✓
All electron pairs repel ✓
Lone pairs repel more than bonded pairs ✓
162
SbCl3 molecules are polar, explain why?
There is a difference in electronegativities ✓
The molecule is not symmetrical, and dipoles do not cancel ✓
163
Which is the most electronegative element in the periodic table?
F ✓
164
CO2 molecules are non-polar, explain why?
There is a difference in electronegativities ✓
The molecule is symmetrical, and dipoles cancel ✓
165
F2 is non-polar, explain why?
No difference in electronegativity ✓
166
Describe the shape of an s-orbital.
Draw it.
Spherical ✓
167
Describe the shape of p-orbitals.
Draw it.
Dumb-bell shaped ✓
168
Define the term orbital
A region that can hold up to two electrons of opposite spins ✓
169
Using your knowledge of bonding and structure Explain the properties of Ca, Cl2 and CaCl2
Calcium has a giant metallic lattice structure ✓
Contains strong metallic bonds, which are the electrostatic attraction between a positive metal ion and negative delocalised electron ✓
Which require large amounts of energy to overcome ✓
Conducts electricity in liquid and solid states as calcium contains mobile delocalised electrons ✓
Chlorine has a simple covalent lattice structure ✓
Low melting point as it has induced dipole-dipole interactions between molecules ✓
Which require little energy to overcome ✓
Does not conduct electricity in any state as it has no mobile charge carriers ✓
CaCl2 has a giant ionic lattice structure ✓
High melting point due to the strong electrostatic attraction between oppositely charged ions which requires a lot of energy to overcome ✓
Conducts electricity when molten as ions are mobile ✓
and does not when solid as ions are fixed in position ✓
170
Which substances have induced dipole-dipole interactions?
Every simple covalent molecule ✓
Graphite/ Graphene ✓
171
Explain how induced dipole-dipole interactions are formed?
Caused by movement of electrons within the shells ✓
This causes an instantaneous dipole across the molecule. ✓
The instantaneous dipole in one molecule induces a dipole in a neighbouring molecule which in turn induce a dipole on their neighbouring molecules. ✓
The small induced dipoles attract one another causing a weak intermolecular force. ✓
172
Explain why HF has a higher melting and boiling point compared to CO2
HF has hydrogen bonding between molecules ✓
CO2 has only induced dipole-dipole interactions between molecules ✓
Hydrogen bonding is stronger and requires more energy to overcome ✓
173
Explain why PH3 has a higher melting and boiling point compared to F2
PH3 has permanent dipole-dipole interactions between molecules ✓
F2 has only induced dipole-dipole interactions between molecules ✓
Permanent dipole-dipole interactions are stronger and requires more energy to overcome ✓
174
Explain why NH3 has a higher melting and boiling point compared to PH3
NH3 has hydrogen bonding between molecules ✓
PH3 has permanent dipole-dipole interactions between molecules ✓
Hydrogen bonding is stronger and requires more energy to overcome ✓
175
Explain why MgCl2 has a higher melting and boiling point compared to SCl6
MgCl2 has a giant ionic lattice structure ✓
with the strong electrostatic attraction between oppositely charged ions ✓
SCl6 has a simple covalent structure ✓
With induced dipole-dipole interactions which are weak and require little energy to overcome ✓
176
Explain why Mg has a higher melting and boiling compared to PH3
Mg has metallic bonding and a giant metallic lattice structure ✓
Metallic bonding is strong and a large amount of energy is needed to overcome these bonds ✓
PH3 has a simple covalent structure and induced dipole-dipole interactions between molecules ✓
Which are weak and require little energy to overcome ✓
177
The electronegativities of H, C, N and F, and the boiling points of methane, ammonia and hydrogen fluoride are shown below.
Electronegativity H: 2.1 C: 2.5 N: 3.0 F: 4.0
Compound and their boiling point in degrees celcius
CH4 -161.5
NH3 -33.3
HF 19.5
Explain the difference in the boiling points of CH4, NH3 and HF
Boiling point/Energy increases with increased electronegativity difference ✓
HF AND NH3 have hydrogen bonding between molecules ✓
CH4 has induced dipole-dipole interactions between ✓ molecules
HF has stronger hydrogen bonding than NH3 ✓
hydrogen bonding is stronger than London forces ✓
178
Which substance(s) experience(s) induced dipole-dipole interactions (London forces)?
1. Ethanol (C2H5OH)
2. SiO2
3. H2O
Ethanol ✓
H2O ✓
| SiO2 giant covalent lattice structure, others are simple covalent lattic
179
Which element has atoms with the greatest number of singly occupied orbitals?
A: C
B: Cl
C: Ca
D: Ga
A: C ✓
## Footnote
If in doubt, draw the electrons in a box diagram (not normal electron configuration) out
180
Which compound has polar molecules?
A OCl2
B BCl3
C CCl4
D SCl6
A ✓
## Footnote
Dipoles do not cancel
181
student conducted one titration and recorded their results. They made a conclusion based on this result.
Describe what the student should do next to obtain reliable results.
Repeat the titration to obtain concordant results ✓
Calculate a mean ✓
182
Write a balanced equation, including observations for the reaction between calcium carbonate and ethanoic acid.
183
Explain the difference in boiling points between ethanol and ethylamine.
Ethanol has a higher boiling point compared to ethylamine ✓
O atom is more electronegative and has more lone pairs than N atom ✓
Ethanol has stronger hydrogen bonding between molecules ✓ which requires more energy to overcome ✓
184
Explain why HI has a higher melting and boiling point compared to HCl
HI has more electrons and is larger than HCl ✓
HI and HCl have permanent dipole-dipole interactions ✓ between molecules ✓
HI has stronger induced dipole-dipole interactions between molecules ✓ and requires more energy to overcome ✓
185
Explain why NaCl has a higher melting point compared to HCl
NaCl has Ionic bonding ✓ the strong electrostatic attraction between opposite charged ions. ✓
HCl has permanent dipole-dipole interactions ✓ between molecules. ✓
Ionic bonding is stronger than permanent dipole-dipole interactions. ✓
186
A student rinses the pipette with water before filling with base. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value would decrease ✓
Base would be more dilute ✓
187
A student rinses the burette with water before filling with acid. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value would increase ✓
Acid would be more dilute ✓
188
A student rinses the conical flask with water before adding 25cm3 of base using a pipette. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value would stay the same ✓
Same number of moles of base in the conical flask ✓
189
A student takes all titration readings from the top of the meniscus. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value would stay the same ✓
Same titre value measured from a difference between the initial and final reading. ✓
Burette measures by difference ✓
190
An air bubble is released from the burette during a titration. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value will increase ✓
Air bubble will displace volume of acid ✓
191
A student over-fills the volumetric flask when preparing the standard solution of an acid for titration. The acid was placed in the burette.
Would the titre increase, decrease or stay the same? Explain your answer.
Increase ✓
Acid would be more dilute ✓
192
A student dissolves 2.0g of base in 100cm3 of water in a beaker before transferring the solution to a volumetric flask and adding more water to the gradation line. Some solution is spilled whilst transferring solution from the beaker to the volumetric flask. The student also notices some crystals of the base are left in the beaker which were not dissolved. The student also realises that they did not rinse the beaker with water and add the washings to the volumetric flask. Would the titre increase, decrease or stay the same? Explain your answer.
Titre value will decrease ✓
Base would be more dilute ✓
193
Student A dissolves 2.50g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made. 25cm3 of X(OH)3 is transferred to a conical flask using a measuring cylinder. Hydrochloric acid is added from a burette until the base is neutralised.
What should the student do next to increase the accuracy of the results?
Repeat the titration until they obtain concordant results ✓
Calculate a mean of the concordant titre values ✓
194
Student A dissolves 2.50g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made. 25cm3 of X(OH)3 is transferred to a conical flask using a measuring cylinder. Hydrochloric acid is added from a burette until the base is neutralised.
How could student A increase the accuracy of the results when making up and measuring the solution of X(OH)3.
Make a standard solution using a 250cm3 volumetric flask ✓
Use a mass balance which measures to a greater number of decimal places ✓
Use a pipette, which measures 25cm3 ✓
195
Student A dissolves 2.50g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made. 25cm3 of X(OH)3 is transferred to a conical flask using a measuring cylinder. Hydrochloric acid is added from a burette until the base is neutralised.
Other than using different equipment, how else could student A increase the accuracy of the results?
Use a larger volume of X(OH)3 or a larger mass of base dissolved ✓
Make acid less concentrated ✓
196
Student B dissolves 3.05g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made using a volumetric flask.
25cm3 of X(OH)3 is transferred to a conical flask using a pipette.
0.50moldm-3 hydrochloric acid is added from a burette until the base is neutralised.
The student’s results are shown below.
Complete the results table and calculate the mean titre to two decimal places.
197
Student B dissolves 3.05g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made using a volumetric flask.
25cm3 of X(OH)3 is transferred to a conical flask using a pipette.
0.50moldm-3 hydrochloric acid is added from a burette until the base is neutralised.
The student’s results are shown below.
Find the identity of X.
198
Student B dissolves 3.05g of base X(OH)3 in a beaker and adds more water to until a 250cm3 solution of X(OH)3 is made using a volumetric flask.
25cm3 of X(OH)3 is transferred to a conical flask using a pipette.
0.50moldm-3 hydrochloric acid is added from a burette until the base is neutralised.
The student’s results are shown below.
Calculate the percentage error in the pipette and in the mean titre. The pipette has an error of ±0.06cm3 and the error in a burette reading is ±0.05cm3. Which instrument contributes a greater percentage error?
199
Draw the dot-and-cross diagrams for magnesium chloride and calcium phosphide.
200
Draw the dot-and-cross diagrams for lithium nitride and aluminium oxide
201
Draw the dot-and-cross diagrams for nitrogen and carbon dioxide
202
Draw the dot-and-cross diagrams for boron tribromide and sulfur hexachloride
203
Draw the dot-and-cross diagram for potassium hydroxide
## Footnote
Be sure to use a different electron for the K electron i.e. triangle or square.
Note: Triangle (gained electron as its negative charge) can be on the O atom or between the O and H atom in the bond.
204
Draw the dot-and-cross diagrams for LiAlH4 (which contains the AlH4- ion)
## Footnote
Be sure to use a different electron for the Li electron i.e. triangle or square.
205
Draw the dot-and-cross diagram for H3O+
206
Draw the dot-and-cross diagrams for nitrate ions.
## Footnote
Be sure to use a different electron for the gained electron i.e. triangle or square.
Note: Triangle (gained electron as its negative charge) can be on the O atom or between the O and N atom in the bond.
207
Draw the dot-and-cross diagrams for ammonium chloride
## Footnote
Be sure to use a different electron for the gained electron i.e. triangle or square and it matches with the H that lost it be become H+
208
Draw the displayed formula for a molecule of H3O+
209
Draw the displayed formula for an ammonium ion.
210
Using a labelled diagram, Explain the bonding in Sodium.
Electrostatic attraction between positively charged ions and delocalised electrons ✓
211
State and explain the shape of a molecule of H3O+. Include a bond angle in your answer.
Pyramidal 107o. ✓
There are 3 bonded pairs and 1 lone pairs around the oxygen atom. ✓
There are 4 regions of electron density. ✓
All electron pairs repel each other and get as far away as possible from each other. ✓
Lone pairs repel more strongly than bonded pairs of electrons. ✓
212
State and explain the shape around the S atom in sulfuric acid.
Include a bond angle in your answer
Tetrahedral, 109.5 o ✓
There are 6 bonded pairs and 0 lone pairs around the S atom. ✓
There are 4 bonding regions of electorn density/ 4 bonding regions. ✓
All electron pairs repel each other and get as far away as possible from each other. ✓
213
State and explain the shape around the N atom in a nitrate ion.
Include a bond angle in your answer
Trigonal planar, 120 o ✓
There are 4 bonded pairs and 0 lone pairs around the N atom. ✓
There are 3 bonding regions of electorn density/ 3 bonding regions. ✓
All electron pairs repel each other and get as far away as possible from each other. ✓
214
State and explain the shape around H-O-N angle in a nitric acid.
Include a bond angle in your answer
Non-linear, 104.5 o ✓
There are 2 bonded pairs and 2 lone pairs around the O atom. ✓
There are 4 regions of electron density. ✓
All electron pairs repel each other and get as far away as possible from each other. ✓
Lone pairs repel more strongly than bonded pairs of electrons. ✓
215
Using a labelled diagram, show the hydrogen bonding between two molecules of water.
216
Using a labelled diagram, show the hydrogen bonding between two molecules of ammonia.
217
Using a labelled diagram, show the hydrogen bonding between two molecules of HF.
218
Using a labelled diagram, show the hydrogen bonding between two molecules of ethanol.
219
Explain why bromine has a greater boiling point than chlorine
Bromine is larger and has more electrons than chlorine ✓
Bromine has stronger induced dipole-dipole interactions ✓ between molecules ✓
… which requires more to overcome ✓
220
Explain why pure water does not conduct electricity.
There are no mobile charged particles i.e. mobile ions or delocalised electrons. ✓
221
Describe and explain the melting point of magnesium chloride, ammonium chloride and potassium phosphide.
Very high. ✓
magnesium chloride, ammonium chloride and potassium phosphide have a giant ionic lattice structure. ✓
There are lots of ionic bonds ✓, with the strong electrostatic attraction between oppositely charged ions. ✓
Which requires a large amount of energy to overcome. ✓
222
Describe and explain the conductivity of magnesium chloride, ammonium chloride and potassium phosphide, in solid, molten and aqueous states.
Conducts electricity when molten or dissolved – Ions are mobile. ✓
Does not conduct electricity when solid – Ions are fixed in position. ✓
223
Describe and explain the solubility of KCl.
Soluble in water, as water is polar. ✓
Water molecules pull the ions away from the lattice and cause it to dissolve. ✓
224
Describe and explain the melting point of Mg.
High ✓ as Mg has a giant metallic lattice structure. ✓
There are lots of strong metallic bonds, with the electrostatic attraction between positively charged metal ions and delocalised electrons. ✓
Which requires a large amount of energy to overcome. ✓
225
Using a labelled diagram, describe and explain the conductivity of Na, in solid and molten states.
Conducts electricity when solid or molten – Na has mobile delocalised electrons. ✓
There are electrostatic attractions between positively charged metal ions and delocalised electrons. ✓
226
Explain why diomand, Si, B and silicon dioxide all have high melting points
Has a giant covalent lattice structure. ✓
Has a very high melting and boiling point as the structure has lots of strong covalent bonds between carbon atoms. ✓
Which requires a large amount of energy to overcome ✓
227
Explain why graphite has a high boiling point
Has a giant covalent lattice structure ✓
High melting and boiling point – lots of strong covalent bonds between carbon atoms. ✓
Which requires are larger amount of energy to overcome ✓
228
Explain why graphite is used in pencil leads and as a lubricant
Weak induced dipole-dipole interactions between layers allow layers to slide over each other ✓
229
Explain why bromine, chlorine, diomand, fluorine and Si do not conduct electricity
They do not conduct electricity as they have no mobile charged particles. ✓
230
Ammonia, and ammonium nitrate, are compounds of nitrogen. Explain why ammonium nitrate has much higher melting and boiling points compared to ammonia.
Ammonia has hydrogen bonding between molecules. ✓
Ammonium nitrate has ionic bonding, which is the strong electrostatic forces of attraction between oppositely charged ions. ✓
Ionic bonding is stronger and requires more energy to overcome. ✓
231
Two students are discussing the oxidation numbers in ammonium nitrate.
One student claims that the two nitrogen atoms have the same oxidation number. The other student disagrees and claims that the nitrogen atoms have different oxidation numbers.
Explain with reasons which student is correct.
The other student, as nitrogen in ammonium ions is -3. ✓
and nitrogen in nitrate ions is +5. ✓
## Footnote
Remember, when H is bonding to a metal or in ammonium, its oxidation state is -1
232
Give chemical explanations for the following statements.
A carton of milk expands on freezing.
Hydrogen bonds in ice hold H2O molecules further apart
233
State the ideal gas equation and units for each symbol.
234
Complete the following to show what is oxidised and reduced.
235
Complete the following to show what is oxidised and reduced.
236
Complete the following to show what is oxidised and reduced.
237
Complete the following to show what is oxidised and reduced.
238
Complete the following to show what is oxidised and reduced.
239
Complete the following to show what is oxidised and reduced.
240
Complete the following to show what is oxidised and reduced.
241
Complete the following to show what is oxidised and reduced.
242
Assign oxidation numbers and say whether the following reaction is a redox reaction or not.
Mg has changed from 0 in Mg to +2 in MgCl2, so it has been oxidised ✓
H has changed from +1 in HCl to 0 in H2, so it has been reduced ✓
Therefore, this is a redox reaction ✓
243
There is no change in oxidation states ✓
Therefore, this is a not a redox reaction ✓
244
What is the formula of iron(III) sulfate?
245
Find the formula of a phosphate(III) ion given that it has a charge of 3–
246
Give the formulas of the following compounds and ions.
a) Copper(II) sulfate
b) Iron(II) oxide
c) Nitrate(III) ions with an overall charge of -1
d) Chromate(IV) ions with an overall charge of -2
247
Give the formulas of the following compounds and ions.
a) Bromate(III) ions with an overall charge of -1
b) Sulfate(VI) ions with an overall charge of -2
c) Potassium nitrate(V)
d) Chlorate(III) ions with an overall charge of -1
248
Give the formulas of the following compounds and ions.
a) Silicon(IV) oxide
b) Chlorate(VII) ion with an overall charge of -1
c) potassium chlorate (VII)
d) Magnesium bromate (VII)
249
Systematically name the following:
a) Iron(II) sulfate (VI) ✓
b) Iron(III) sulfate (VI) ✓
c) Nitrate(V) ions ✓
d) Bromate(V) ions ✓
250
Systematically name the following:
a) Phosphate(III) ions ✓
b) Iodate(V) ions ✓
c) Iodate(I) ions ✓
d) Bromate (III) ions ✓
251
Systematically name the following:
a) Chlorate(V) ions ✓
b) Sodium chlorate(V) ✓
c) Chlorate(III) ions✓
d) Sodium bromate(I) ✓
252
Systematically name the following:
a) Sodium bromate(V) ✓
b) Calcium chlorate(III) ✓
