Download 3. Chemical changes and Structure Unit Questions

Department of
Chemistry
Exam Questions for CfE Higher
Chemistry
Chemical Changes and Structure
 Periodicity/Structure and Bonding
 Controlling rate/ Chemical Analysis
2
Contents
3.1 Periodicity
Page 4
3.2 Structure and Bonding
Page 14
3.3 Controlling Rate
Page 24
3.4 Chemical Analysis
Page 37
3
3.1
Periodicity
Summary
The periodic table
o groups are columns in the periodic table ↕
o periods are rows in the periodic table ↔
o The groups in the periodic table are:
o Group 1- alkali metals
o Group 2- rare earth metals
o Group 7- halogens
o Group 8- Noble gases
o Chemical formula can be written:
o From the name if there is a prefix. Di, tri, tetra etc.
o Using S,V,S,D,F
o The state symbols are:
o (s)-solid
o (l)- liquid
o (g)- gas
o (aq)- in solution (dissolved)
o The first 20 elements can be separated into the following:
o Metallic-left of the thick black line
o Monatomic- noble gases
o covalent network- B,C and Si
o covalent molecular (including diatomic)- everything else (H,N,O,F down)
Trends in the first 20 elements:
Reactivity:
o Atoms in the same groups react in similar ways because they have the same number of
outer electrons
o they get more reactive because it’s easier for the outer electrons to react, this is due to
more smaller shells shielding the outer electrons
Melting and Boiling Points:
o In groups 1 to 4 the melting and boiling points decrease going down the group
o In groups 5 to 8(0) the melting and boiling points increase going down the group
o The trend across the period is the melting and boiling points increase from group 1 to 4
and decrease after that
Density:
o Moving down a group the density increases
o Going across a period that the density increases to a maximum in the middle then
decreases
4
Atomic Size (atomic radius):
o The atomic size increases going down a group because there are more electron shells
o Atomic size decreases moving across a period because there are more protons attracting
the same size of shell
Ionisation:
o The ionisation energy is the energy required to remove one mole of electrons from one
mole of gaseous atoms
o You can have second and third ionisation energies, but note that if there is a full outer
shell the atom will not want to lose another electron making the energy very high.
o The general formula for ionisation energy is given in the databook- where E is replaced
with an element
o Ionisation energy decreases going down a group because the outer electrons are more
shielded from the nucleus, making it easier to remove
o Ionisation energy increases moving across a period because the atomic size is decreasing
(see atomic size above), so the electron is closer to the nucleus, this makes the electron
harder to remove
Electronegativity:
o Electronegativity is the attraction atoms have for electrons
o Electronegativity decreases going down a group because the electrons are further from
the nucleus
o Electronegativity increases moving across a period because the atoms want a full outer
shell
5
Questions on Periodicity
1. An element (melting point above 3000 °C) forms an oxide which is a gas at room
temperature. Which type of bonding is likely to be present in the element?
A Metallic
B Polar covalent
C Non-polar covalent
D Ionic
2. What type of bonding and structure is found in a fullerene?
A Ionic lattice
B Metallic lattice
C Covalent network
D Covalent molecular
3. At room temperature, a solid substance was shown to have a lattice consisting of positively
charged ions and delocalised outer electrons. The substance could be
A graphite
B sodium
C mercury
D phosphorus
4. The two hydrogen atoms in a molecule of hydrogen are held together by
A a hydrogen bond
B a polar covalent bond
C a non-polar covalent bond
D a van der Waals’ force.
5. Which of the following does not contain covalent bonds?
A Hydrogen gas
B Helium gas
C Nitrogen gas
D Solid sulphur
6. Which of the following elements exists as discrete molecules?
A Boron
B Carbon (diamond)
C Silicon
D Sulfur
7. Which type of bonding is never found in elements?
A Metallic
B London dispersion forces
C Polar covalent
D Non-polar covalent
6
8. The diagram shows the melting points of successive elements across a period in the Periodic
Table.
Which of the following is a correct reason for the low melting point of element Y?
A It has weak ionic bonds.
B It has weak covalent bonds.
C It has weakly-held outer electrons.
D It has weak forces between molecules.
9. As the relative atomic mass in the halogens increases
A the boiling point increases
B the density decreases
C the first ionisation energy increases
D the atomic size decreases.
10. Element X was found to have the following properties.
(i) It does not conduct electricity when solid.
(ii) It forms a gaseous oxide.
(iii) It is a solid at room temperature.
Element X could be
A magnesium
B silicon
C nitrogen
D sulphur.
11. Hydrogen will form a non-polar covalent bond with an element which has an
electronegativity value of
A 0·9
B 1·5
C 2·2
D 2·5
7
12. Which line in the table is likely to be correct for the element francium?
13. Which of the following elements is most likely to have a covalent network structure?
14. Which of the following elements has the greatest attraction for bonding electrons?
A Lithium
B Chlorine
C Sodium
D Bromine
15. Which of the following statements is true?
A The potassium ion is larger than the potassium atom.
B The chloride ion is smaller than the chlorine atom.
C The sodium atom is larger than the sodium ion.
D The oxygen atom is larger than the oxide ion.
16. Which of the following atoms has the least attraction for bonding electrons?
A Carbon
B Nitrogen
C Phosphorus
D Silicon
8
17. As the atomic number of the alkali metals increases
A the first ionisation energy decreases
B the atomic size decreases
C the density decreases
D the melting point increases.
18. Which equation represents the first ionisation energy of a diatomic element, X 2?
A ½ X2(s)  X+(g)
B ½ X2(g)  X–(g)
C X(g)  X+(g)
D X(s)  X–(g)
19. Which of the following equations represents the first ionisation energy of fluorine?
A F–(g) → F(g) + e–
B F–(g) → ½ F2(g) + e–
C F(g) → F+(g) + e–
D ½ F2(g) → F+(g) + e–
20. Which of the following reactions refers to the third ionisation energy of aluminium?
A Al(s) → Al3+(g) + 3e–
B Al(g) → Al3+(g) + 3e–
C Al2+(g) → Al3+(g) + e–
D Al3+(g) → Al4+(g) + e–
21. The table shows the first three ionization energies of aluminium.
Using this information, what is the enthalpy change, in kJ mol–1, for the following reaction?
Al3+(g) + 2e– → Al+(g)
A +2176
B –2176
C +4590
D –4590
9
22. The elements lithium, boron and nitrogen are in the second period of the Periodic Table.
Complete the table below to show both the bonding and structure of these three elements
at room temperature. (2)
Name of Element
Bonding
Lithium
Structure
Lattice
Boron
Nitrogen
Covalent
23. Hydrogen gas has a boiling point of –253 °C. Explain clearly why hydrogen is a gas at room
temperature. In your answer you should name the intermolecular forces involved and
indicate how they arise. (2)
24. (a) Atoms of different elements have different attractions for bonded electrons. What term
is used as a measure of the attraction an atom involved in a bond has for the electrons
of the bond? (1)
(b) Atoms of different elements are different sizes. What is the trend in atomic size across
the period from sodium to argon? (1)
(c) Atoms of different elements have different ionisation energies. Explain clearly why the
first ionisation energy of potassium is less than the first ionisation energy of sodium. (2)
25. The Periodic Table allows chemists to make predictions about the properties of elements.
(a) The elements lithium to neon make up the second period of the Periodic Table.
(i) Name an element from the second period that exists as a covalent network. (1)
(ii) Why do the atoms decrease in size from lithium to neon? (1)
(iii) Which element in the second period is the strongest reducing agent? (1)
(b) On descending Group 1 from lithium to caesium, the electronegativity of the elements
decreases. Explain clearly why the electronegativity of elements decreases as you go
down the group. (2)
10
26. (a) Lithium starts the second period of the Periodic Table.
What is the trend in electronegativity values across this period from Li to F? (1)
(b) Graph 1 shows the first four ionisation energies for aluminium.
Why is the fourth ionisation energy of aluminium so much higher than the third ionisation
energy? (1)
(c) Graph 2 shows the boiling points of the elements in Group 7 of the Periodic Table.
Why do the boiling points increase down Group 7? (1)
11
27. The following answer was taken from a student’s examination paper. The answer is
incorrect. Give the correct explanation.
28. The elements from sodium to argon make up the third period of the Periodic Table.
(a) On crossing the third period from left to right there is a general increase in the first
ionisation energy of the elements.
(i) Why does the first ionisation energy increase across the period? (1)
(ii) Write an equation corresponding to the first ionisation energy of chlorine. (1)
(b) The electronegativities of elements in the third period are listed on page 10 of the
databook. Why is no value provided for the noble gas, argon? (1)
29. Attempts have been made to make foods healthier by using alternatives to traditional
cooking ingredients.
(a) An alternative to common salt contains potassium ions and chloride ions.
(i) Write an ion-electron equation for the first ionisation energy of potassium. (1)
(ii) Explain clearly why the first ionisation energy of potassium is smaller than that of
chlorine. (3)
12
30. (a) The first ionisation energy of an element is defined as the energy required to remove
one mole of electrons from one mole of atoms in the gaseous state.
The graph shows the first ionisation energies of the Group 1 elements.
(i) Clearly explain why the first ionisation energy decreases down this group. (2)
(b) The ability of an atom to form a negative ion is measured by its Electron Affinity.
The Electron Affinity is defined as the energy change when one mole of gaseous atoms
of an element combines with one mole of electrons to form gaseous negative ions.
Write the equation, showing state symbols, that represents the Electron Affinity of
chlorine. (1)
13
3.2
Structure and Bonding
Summary
Bonding in molecules (intramolecular bonding):
o There is a bonding is a continuum as shown:
o pure covalent polar covalentionic
o This occurs because of differences in electronegativity,
o if there is a small difference the molecule is neutral over all
o if there is a small difference the molecule has a slight positive and a slight
negative making the molecule polar
o is there is a large enough difference the electron from one atom can be ripped
away making ions and ionic bonding
o a rough guide for bonding is:
o <0.5 =pure covalent
o >0.5 but <1.6 =polar covalent
o >1.6= ionic
Covalent Bonding:
o Covalent bonding is usually between two non-metal atoms-you must check the properties
to be sure
o A covalent bond is made when electrons are held between two atoms nuclei-sharing the
electrons
o Polar covalent bonding occurs when an atom has a greater electronegativity, it pulls the
electrons closer to its nucleus, making that atom more negative and the other more
positive
Ionic bonding:
o Ionic bonding usually occurs between a metal and a non-metal atom
o An ionic bond is made when the electronegativity is greater than the ionisation energy.
The non-metal atom takes the electron from the metal atom making a negative and a
positive ion, these are attracted (electrostatically)
Metallic bonding:
o Metallic bonding occurs between metals, the atoms are held together because electrons
are free to move throughout the compound
Bonding between atoms and molecules (intermolecular bonding):
o London dispersion forces (LDF)
o LDF occur when an atom or molecule forms a temporary dipole (eg. O2), this
causes a v.v.v. weak attraction between molecules.
o Permanent dipole interactions
o When polar covalent molecules are together the δ+ and δ- of different molecules
attract. This gives a slight increase to melting point and boiling point of a
substance
o Hydrogen bonding
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o Hydrogen bonding is a strong intermolecular force but not as strong as a covalent
bond
o It only occurs between very polar molecules which contain:
 O-H bonds
 N-H bonds
 F-H bonds
o Because this bonding is so strong it gives substances an unusually high melting and
boiling point and a unusually low density (which is why ice floats on water)
Properties of bonding
o Intramolecular bonding is most significant when looking at properties, but intermolecular
forces do play a role
Intermolecular
forces possible?
Mpt/bpt
Density
Solubility
conductivity
Covalent network
No
High
High
No
No
Covalent
molecular
LDF
Low
Low
No
No
Slightly
higher
Slightly
higher
Slightly
Slightly
Unusually
high
Unusually
high
Yes
Yes
High
High
Very
Only when
molten or in
solution
Ionic
pure
Pola
r
Yes, permanent
dipole
interactions
Hydr Hydrogen
ogen bonded
bond
ed
No
Metallic
No
High
V high
No
Yes
Monatomic
LDF only
V low
V low
No
No
Intermolecular forces on properties:
Temperature:
o The extra attraction between the molecules increases the melting point and boiling
point.
this more evident with hydrogen bonding, looking at the table:
15
we can see that the three hydrogen bonded molecules have odd boiling points because of
this attraction.
Density:
o When hydrogen bonding occurs the molecules density is decreased. At low temperatures
the hydrogen bonds have a greater effect than at higher temperature’s, so when water
freezes the molecules are pushed further away from each other and this makes them less
dense.
Solubilityo Polar molecules can only dissolve in polar solvents and vice versa
16
Questions on Structure and Bonding
1. Which of the following compounds contains both a halide ion and a transition metal ion?
A Iron oxide
B Silver bromide
C Potassium permanganate
D Copper iodate
2. Particles with the same electron arrangement are said to be isoelectronic. Which of the
following compounds contains ions which are isoelectronic?
A Na2S
B MgCl2
C KBr
D CaCl2
3. Which property of a chloride would prove that it contained ionic bonding?
A It conducts electricity when molten.
B It is soluble in a polar solvent.
C It is a solid at room temperature.
D It has a high boiling point.
4. Which of the following chlorides is likely to have least ionic character?
A BeCl2
B CaCl2
C LiCl
D CsCl
5. Which of the following compounds has the greatest ionic character?
A Caesium fluoride
B Caesium iodide
C Sodium fluoride
D Sodium iodide
6. Which of the following chlorides is most likely to be soluble in tetrachloromethane, CCl4?
A Barium chloride
B Caesium chloride
C Calcium chloride
D Phosphorus chloride
7. Which of the following compounds exists as discrete molecules?
A Sulphur dioxide
B Silicon dioxide
C Aluminium oxide
D Iron(II) oxide
17
8. Which of the following compounds has polar molecules?
A CO2
B NH3
C CCl4
D CH4
9. When two atoms form a non-polar covalent bond, the two atoms must have
A the same atomic size
B the same electronegativity
C the same ionisation energy
D the same number of outer electrons.
10. In which of the following molecules will the chlorine atom carry a partial positive
Charge (δ+)?
A Cl−Br
B Cl−Cl
C Cl−F
D Cl−I
11. Which line in the table represents the solid in which only London dispersion forces are
overcome when the substance melts?
12. Which of the following structures is never found in compounds?
A Ionic
B Monatomic
C Covalent network
D Covalent molecular
13. Atoms of nitrogen and element X form a bond in which the electrons are shared equally.
Element X could be
A carbon
B oxygen
C chlorine
D phosphorus.
18
14. Which line in the table shows the correct entries for tetrafluoroethene? 1013B
15. In which of the following compounds would hydrogen bonding not occur?
16. Coniceine is a deadly poison extracted from the plant hemlock.
Which of the following would be the best solvent for coniceine?
A Propanoic acid
B Propan-l-ol
C Heptane
D Water
19
17. The structures for molecules of four liquids are shown below. Which liquid will be the most
viscous?
18. The shapes of some common molecules are shown. Each molecule contains at least one
polar covalent bond. Which of the following molecules is non-polar?
19. Some covalent compounds are made up of molecules that contain polar bonds but the
molecules are overall non-polar. Which of the following covalent compounds is made up of
non-polar molecules?
A Ammonia
B Water
C Carbon tetrachloride
D Hydrogen fluoride
20. A positively charged particle with electron arrangement 2, 8 could be
A a neon atom
B a fluoride ion
C a sodium atom
D an aluminium ion.
20
21. Compared to other gases made up of molecules of similar molecular masses, ammonia has a
relatively high boiling point.
In terms of the intermolecular bonding present, explain clearly why ammonia has a
relatively high boiling point. (2)
22. The formulae for three oxides of sodium, carbon and silicon are Na 2O, CO2 and SiO2.
Complete the table for CO2 and SiO2 to show both the bonding and structure of the three
oxides at room temperature. (2)
23. Hydrogen cyanide, HCN, is highly toxic. Information about hydrogen cyanide is given in the
table.
Although hydrogen cyanide has a similar molecular mass to nitrogen, it has a much higher
boiling point. This is due to the permanent dipole–permanent dipole attractions in liquid
hydrogen cyanide.
What is meant by permanent dipole–permanent dipole attractions?
Explain how they arise in liquid hydrogen cyanide. (2)
21
24. A student writes the following two statements. Both are incorrect. In each case explain the
mistake in the student’s reasoning.
(a) All ionic compounds are solids at room temperature. Many covalent compounds are
gases at room temperature. This proves that ionic bonds are stronger than covalent
bonds. (1)
(b) The formula for magnesium chloride is MgCl2 because, in solid magnesium chloride,
each magnesium ion is bonded to two chloride ions. (1)
25. Electronegativity values can be used to predict the type of bonding present in substances.
The type of bonding between two elements can be predicted using the diagram below.
(a) Using the information in the diagram, state the highest average electronegativity found
in ionic compounds. (1)
(b) The diagram can be used to predict the bonding in tin iodide.
Electronegativity of tin = 1·8
Electronegativity of iodine = 2·6
Average electronegativity = 2·2
Difference in electronegativity = 0·8
Predict the type of bonding in tin iodide. (1)
(c) The electronegativities of arsenic and chlorine are shown below.
Electronegativity of arsenic = 2·2
Electronegativity of chlorine = 3·0
22
Place a small cross on the diagram to show the position of arsenic chloride. Show
calculations clearly. (2)
26. The structures below show molecules that contain chlorine atoms.
The compounds shown above are not very soluble in water. Trichloromethane is around ten
times more soluble in water than tetrachloromethane.
Explain clearly why trichloromethane is more soluble in water than tetrachloromethane.
Your answer should include the names of the intermolecular forces involved. (3)
23
3.3
Controlling the Rate
Summary
Rate:
o To increase the rate of reaction you could:
o Increase the temperature-particles have more energy leading to more collisions
with the activation energy
o Increase the concentration-more particles in a fixed space to collide with
o Decrease particle size- larger surface area
o Use a catalyst
o The average rate of reaction is calculated using the equation:
𝑐ℎ𝑎𝑛𝑔𝑒
𝑎𝑣𝑒𝑟𝑎𝑔𝑒 𝑟𝑎𝑡𝑒 =
ℎ𝑜𝑤 𝑙𝑜𝑛𝑔 𝑖𝑡 𝑡𝑜𝑜𝑘 𝑡𝑜 𝑐ℎ𝑎𝑛𝑔𝑒
Notes: the change could be change in gas produced of change of mass
Depending on what changes affects the units
o The relative rate can be calculated using the equation:
1
𝑟𝑒𝑙𝑎𝑡𝑖𝑣𝑒 𝑟𝑎𝑡𝑒 =
𝑡
Notes: The relative rates units is in time -1
You are often asked to work out time by rearranging the above equation
Reaction profiles:
o The activated complex is a transition state between reactants and products
o It is found at the highest point on the reaction profile
o The activation energy is the difference between the starting energy and the highest point
on the reaction profile
o If the final energy is lower than the initial energy, then energy has been given out and
the reaction is exothermic (∆H=-ve)
o If the final energy is higher than the initial energy, then energy has been taken in and the
reaction is endothermic (∆H=+ve)
Catalysts:
o Catalysts speed up a reaction but can be removed unchanged at the end
o there are two types of catalyst
o heterogeneous-catalyst and reactants are in different states
o homogeneous- catalyst and reactants are in the same state
o catalysts lower the activation energy of a reaction by providing a different pathway for
the reaction
o the new pathway includes:
o absorption
o reaction
o de-absorption
o the effect on the energy profile is that it reduces how high the curve goes up
Temperature and kinetic energy:
24
o temperature is the average amount of thermal energy in a substance
o A small increase in temperature results in more particles having greater than or equal to
the activation energy, therefore more successful collisions
25
Rate Questions
1. The graph shows the variation of concentration of a reactant with time as a reaction
proceeds.
What is the
A
B
C
D
average reaction rate during the first 20 s?
0.0025 mol l–1 s–1
0.0050 mol l–1 s–1
0.0075 mol l–1 s–1
0.0150 mol l–1 s–1
2. The graph shows how the rate of a reaction varies with the concentration of one of the
reactants.
What was the reaction time, in seconds, when the concentration of the reactant was
0.50 mol l–1?
A
0.2
B
0.5
C
2.0
D
5.0
26
3. The following results were obtained in the reaction between marble chips and dilute
hydrochloric acid.
What is the average rate of production of carbon dioxide, in cm3 min–1, between 2 and 8
minutes?
A
5
B
26
C
30
D
41
4. Excess zinc was added to 100 cm3 of hydrochloric acid, concentration 1 mol l–1. Graph I refers
to this reaction.
Graph
A
B
C
D
II could be for
excess zinc reacting with 100 cm3 of hydrochloric acid, concentration 2 mol l–1
excess zinc reacting with 100 cm3 of sulphuric acid, concentration 1 mol l–1
excess zinc reacting with 100 cm3 of ethanoic acid, concentration 1 mol l–1
excess magnesium reacting with 100 cm3 of hydrochloric acid, concentration 1 mol l–1.
5. Which of the following is not a correct statement about the effect of a catalyst?
The catalyst:
A
provides an alternative route to the products
B
lowers the energy that molecules need for successful collisions
C
provides energy so that more molecules have successful collisions
D
forms bonds with reacting molecules.
6. In which of the following will both changes result in an increase in the rate of a chemical
reaction?
A
A decrease in activation energy and an increase in the frequency of collisions.
B
An increase in activation energy and a decrease in particle size.
C
An increase in temperature and an increase in the particle size.
D
An increase in concentration and a decrease in the surface area of the reactant
27
7.
In area X
A
molecules always form an activated complex
B
no molecules have the energy to form an activated complex
C
collisions between molecules are always successful in forming products
D
all molecules have the energy to form an activated complex.
8. For any chemical, its temperature is a measure of
A
the average kinetic energy of the particles that react
B
the average kinetic energy of all the particles
C
the activation energy
D
the minimum kinetic energy required before reaction occurs.
9.
Which line in the table is correct for a reaction as the temperature decreases from T2 to T1?
28
10. The potential energy diagram below refers to the reversible reaction involving reactants R
and products P.
What is the enthalpy change, in kJ mol–1, for the reverse reaction P  R?
A
+ 30
B
+ 10
C
–10
D
–40
11.
When a catalyst is used, the activation energy of the forward reaction is reduced to
35 kJ mol–1. What is the activation energy of the catalysed reverse reaction?
A
30 kJ mol–1
B
35 kJ mol–1
C
65 kJ mol–1
D
190 kJ mol–1
29
12. A reaction was carried out with and without a catalyst as shown in the energy diagram.
What is the enthalpy change, in kJ mol–1, for the catalysed reaction?
A–100
B –50
C +50
D +100
13. The following potential diagram is for a reaction carried out with and without a catalyst.
The activation energy for the catalyzed reaction is
A 30 kJ mol–1
B 80 kJ mol–1
C 100 kJ mol–1
D 130 kJ mol–1
30
14.
The enthalpy change for the forward reaction can be represented by
A x
B y
C x+y
D x−y
15. A potential energy diagram can be used to show the activation energy (EA) and the enthalpy
change (ΔH) for a reaction. Which of the following combinations of EA and ΔH could never be
obtained for a reaction?
A EA = 50 kJ mol–1 and ΔH = –100 kJ mol–1
B EA = 50 kJ mol–1 and ΔH = +100 kJ mol–1
C EA = 100 kJ mol–1 and ΔH = +50 kJ mol–1
D EA = 100 kJ mol–1 and ΔH = –50 kJ mol–1
31
16. Temperature has a very significant effect on the rate of a chemical reaction.
(a) The reaction shown below can be used to investigate the effect of temperature on
reaction rate.
5(COOH)2(aq) + 6H+(aq) + 2MnO4–(aq) → 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
The instructions for such an investigation are shown below.
(i) What colour change indicates that the reaction is over? (1)
(ii) With each of the experiments, the temperature of the solution was measured both
during heating and at the end of the reaction. When plotting graphs of the reaction rate
against temperature, it is the temperature measured at the end of reaction, rather than
the temperature measured while heating, that is used. Give a reason for this. (1)
(b) The graph shows the distribution of kinetic energy for molecules in a reaction mixture at
a given temperature.
Why does a small increase in temperature produce a large increase in reaction rate? (1)
32
17. A student carried out the Prescribed Practical Activity (PPA) to find the effect of
concentration on the rate of the reaction between hydrogen peroxide solution and an
acidified solution of iodide ions.
H2O2(aq) + 2H+(aq) + 2I–(aq) → 2H2O(l) + I2(aq)
During the investigation, only the concentration of the iodide ions was changed.
Part of the student’s results sheet for this PPA is shown.
(a) Describe how the concentration of the potassium iodide solution was changed during this
series of experiments. (1)
(b) Calculate the reaction time, in seconds, for the first experiment. (1) 8B3
18. The following answer was taken from a student’s examination paper. The answer is
incorrect. Give the correct explanation.
19. Hydrogen peroxide decomposes as shown:
H2O2(aq) → H2O(l) + ½O2(g)
The reaction can be catalysed by iron(III) nitrate solution. What type of catalyst is iron(III)
nitrate solution in this reaction? (1
33
20. A student carried out three experiments involving the reaction of excess magnesium ribbon
with dilute acids. The rate of hydrogen production was measured in each of the three
experiments.
Experiment
Acid
1
100cm3 of 0.10 mol l-1 sulphuric acid
2
50cm3 of 0.20 mol l-1 sulphuric acid
3
100cm3 of 0.10 mol l-1 hydrochloric acid
The equation for Experiment 1 is shown.
Mg(s) + H2SO4(aq)  MgSO4(aq) + H2(g)
The curve obtained for Experiment 1 is drawn on the graph.
Draw curves on the graph to show the results obtained for Experiment 2 and Experiment 3.
Label each curve clearly. (2)
21. The reaction of oxalic acid with an acidified solution of potassium permanganate was
studied to determine the effect of temperature changes on reaction rate.
5(COOH)2(aq) + 6H+(aq) + 2MnO4–(aq) → 2Mn2+(aq) + 10CO2(g) + 8H2O(l)
The reaction was carried out at several temperatures between 40 °C and 60 °C. The end of
the reaction was indicated by a colour change from purple to colourless.
(a) (i) State two factors that should be kept the same in these experiments. (1)
(ii) Why is it difficult to measure an accurate value for the reaction time when the
reaction is carried out at room temperature? (1)
(b) Sketch a graph to show how the rate varied with increasing temperature. (1)
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22. An experiment was carried out to determine the rate of the reaction between hydrochloric
acid and calcium carbonate chips. The rate of this reaction was followed by measuring the
volume of gas released over a certain time.
(a) Describe a different way of measuring volume in order to follow the rate of this reaction. (1)
(b) What other variable could be measured to follow the rate of this reaction? (1)
23. Enzymes are biological catalysts.
(a) Name the four elements present in all enzymes. (1)
(b) The enzyme catalase, found in potatoes, can catalyse the decomposition of hydrogen
peroxide.
2H2O2(aq) → 2H2O(l) + O2(g)
A student carried out the Prescribed Practical Activity (PPA) to determine the effect of
pH on enzyme activity. Describe how the activity of the enzyme was measured in this
PPA. (1)
(c) A student wrote the following incorrect statement.
When the temperature is increased, enzyme-catalysed reactions will always speed up
because more molecules have kinetic energy greater than the activation energy.
Explain the mistake in the student’s reasoning. (1)
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24. Chloromethane, CH3Cl, can be produced by reacting methanol solution with dilute
hydrochloric acid using a solution of zinc chloride as a catalyst.
(a) What type of catalysis is taking place? (1)
(b) The graph shows how the concentration of the hydrochloric acid changed over a period
of time when the reaction was carried out at 20 °C.
(i) Calculate the average rate, in mol l–1 min–1, in the first 400 minutes. (1)
(ii) On the graph above, sketch a curve to show how the concentration of hydrochloric acid
would change over time if the reaction is repeated at 30 °C.(1)
36
3.4 Chemical Analysis
Summary
Chromatography
o Chromatography is a technique used to splitting up or analysing mixtures of compounds
o There are always two phases in chromatography, the mobile phase and the stationary
phase
o If you want to remove a polar substance you would use a polar solvent eg. ethanol
o If you want to remove a non-polar substance you would use a non-polar solvent eg.
hexane
o Other types of chromatography include gas-liquid chromatography and high pressure
liquid chromatography (HPLC)
Volumetric Titrations
o Titrations are used when you have a known concentration of a substance and you want to
work out the concentration of another substance
o The substances could be acid/alkali or oxidising/reducing agents
o The apparatus for titrations are:
o Pipette
o Burette
o An indicator
o The N5 indicator is universal indicator but this has no clear end point (a clear colour
change)
o Indicator to note are:
o Methyl orange- red in acidyellow in neutral and alkaline
o Phenolphthalein - colourless in acid and neutral pink in alkaline
o Potassium permanganate- purplecolourless when reduced
o Iodine- brown  colourless when oxidised
o When carrying out titrations the results should be reproduced
o You should have at least three concurrent results (roughly the same) and these are
averaged
o They may be outliers-these should be ignored
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Chemical Analysis Questions
1. An organic chemist is attempting to synthesise a fragrance compound by the following
chemical reaction.
compound X + compound Y → fragrance compound
After one hour, a sample is removed and compared with pure samples of compounds X and Y
using thin-layer chromatography. Which of the following chromatograms shows that the
reaction has produced a pure sample of the fragrance compound?
2. The alcohol content of wine was analysed by four students. Each student carried out the
experiment three times.
The most reproducible results were obtained by
A Student A
B Student B
C Student C
D Student D
3. 45 cm3 of a solution could be most accurately measured out using a
A 50 cm3 beaker
B 50 cm3 burette
C 50 cm3 pipette
D 50 cm3 measuring cylinder.
4. Aluminium carbonate can be produced by the following reaction.
2AlCl3(aq) + 3K2CO3(aq) → Al2(CO3)3(s) + 6KCl(aq)
The most suitable method for obtaining a sample of the aluminium carbonate is
A collection over water
B distillation
C evaporation
D filtration.
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5. Seaweeds are a rich source of iodine in the form of iodide ions. The mass of iodine in a
seaweed can be found using the procedure outlined below.
(a) Step 1
The seaweed is dried in an oven and ground into a fine powder. Hydrogen peroxide
solution is then added to oxidise the iodide ions to iodine molecules.
The ion-electron equation for the reduction reaction is shown.
H2O2(aq) + 2H+(aq) + 2e– → 2H2O(l)
Write a balanced redox equation for the reaction of hydrogen peroxide with iodide ions.
(1)
(b) Step 2
Using starch solution as an indicator, the iodine solution is then titrated with sodium
thiosulphate solution to find the mass of iodine in the sample. The balanced equation
for the reaction is shown.
2Na2S2O3(aq) + I2(aq) → 2NaI(aq) + Na2S4O6(aq)
In an analysis of seaweed, 14.9cm3 of 0.00500 mol l–1 sodium thiosulphate solution was
required to reach the end-point. Calculate the mass of iodine present in the seaweed
sample. Show your working clearly. (3)
6. Oxalic acid is found in rhubarb. The number of moles of oxalic acid in a carton of rhubarb
juice can be found by titrating samples of the juice with a solution of potassium
permanganate, a powerful oxidising agent.
The equation for the overall reaction is:
5(COOH)2(aq) + 6H+(aq) + 2MnO4 –(aq) → 2Mn2+ (aq) + 10CO2(aq) + 8H2O(l)
(a) Write the ion-electron equation for the reduction reaction. (1)
(b) Why is an indicator not required to detect the end-point of the titration? (1)
(c) In an investigation using a 500 cm3 carton of rhubarb juice, separate 25.0cm3 samples
were measured out. Three samples were then titrated with 0.040 mol l–1 potassium
permanganate solution, giving the following results.
Average volume of potassium permanganate solution used = 26.9cm3.
(i) Why was the first titration result not included in calculating the average volume of
potassium permanganate solution used? (1)
(ii) Calculate the number of moles of oxalic acid in the 500 cm 3 carton of rhubarb juice.
Show your working clearly. (2)
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7. The number of moles of carbon monoxide in a sample of air can be measured as follows.
Step 1 The carbon monoxide reacts with iodine(V) oxide, producing iodine.
5CO(g) + I2O5(s) → I2(s) + 5CO2(g)
Step 2 The iodine is then dissolved in potassium iodide solution and titrated
against sodium thiosulphate solution.
I2(aq) + 2S2O32–(aq) → S4O62–(aq) + 2I–(aq)
(a) Write the ion-electron equation for the oxidation reaction in Step 2. (1)
(b) Name a chemical that can be used to indicate when all of the iodine has been
removed in the reaction taking place in Step 2. (1)
(c) If 50.4cm3 of 0.10 mol l–1 sodium thiosulphate solution was used in a titration,
calculate the number of moles of carbon monoxide in the sample of air.
Show your working clearly. (2)
8. A major problem for the developed world is the pollution of rivers and streams by nitrite and
nitrate ions. The concentration of nitrite ions, NO2 –(aq), in water can be determined by
titrating samples against acidified permanganate solution.
(a) Suggest two points of good practice that should be followed to ensure that an accurate
end-point is achieved in a titration.
(b) An average of 21·6cm3 of 0·0150 mol l–1 acidified permanganate solution was required to
react completely with the nitrite ions in a 25·0 cm3 sample of river water. The equation
for the reaction taking place is:
2MnO4– (aq) + 5NO2–(aq) + 6H+(aq) → 2Mn2+(aq) + 5NO3–(aq) + 3H2O(l)
(i) Calculate the nitrite ion concentration, in mol l–1, in the river water. Show your
working clearly. (2)
(ii) During the reaction the nitrite ion is oxidised to the nitrate ion. Complete the ionelectron equation for the oxidation of the nitrite ions.
NO2–(aq) → NO3–(aq)
(1)
9. (a) The concentration of chromate ions in water can be measured by titrating with a solution
of iron(II) sulphate solution. To prepare the iron(II) sulphate solution used in this
titration, iron(II) sulphate crystals were weighed accurately into a dry beaker. Describe
how these crystals should be dissolved and then transferred to a standard flask in order
to produce a solution of accurately known concentration. (2)
(b) A 50·0cm3 sample of contaminated water containing chromate ions was titrated and
found to require 27·4 cm3 of 0·0200 mol l–1 iron(II) sulphate solution to reach the endpoint.
The redox equation for the reaction is:
2+
3Fe (aq) + CrO42–(aq) + 8H+(aq) → 3Fe3+(aq) + Cr3+(aq) + 4H2O(l)
Calculate the chromate ion concentration, in mol l–1, present in the sample of water.
Show your working clearly. (2)
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10. Hydrogen sulfide, H2S, can cause an unpleasant smell in water supplies. The concentration
of hydrogen sulfide can be measured by titrating with a chlorine standard solution. The
equation for the reaction taking place is
4Cl2(aq) + H2S(aq) + 4H2O(l) → SO42−(aq) + 10H+(aq) + 8Cl−(aq)
50·0 cm3 samples of water were titrated using a 0∙010 mol l−1 chlorine solution.
(a) Name an appropriate piece of apparatus which could be used to measure out the water
samples. (1)
(b) What is meant by the term standard solution? (1)
(c) An average of 29·4 cm3 of 0∙010 mol l−1 chlorine solution was required to react
completely with a 50·0 cm3 sample of water. Calculate the hydrogen sulfide
concentration, in mol l−1, present in the water sample. Show your working clearly. (3)
11. Zinc is an essential element for the body and is found in a variety of foods.
(a) The mass of zinc in four 100 g samples taken from a cheese spread was measured.
Calculate the average mass of Zn, in mg, in 100 g of this cheese spread. (1)
(b) The recommended daily allowance of zinc is 9·5 mg for an adult male. 100 g of peanuts
contains 3·3 mg of zinc. Calculate the mass of peanuts which would provide the
recommended daily allowance of zinc. (2)
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12. Solutions containing iodine are used to treat foot rot in sheep. The concentration of iodine
in a solution can be determined by titrating with a solution of thiosulfate ions.
(a) Write an ion-electron equation for the reaction of the oxidising agent in the titration. (1)
(b) Three 20·0 cm3 samples of a sheep treatment solution were titrated with 0·10 mol l–1
thiosulfate solution. The results are shown below.
(i) Why is the volume of sodium thiosulfate used in the calculation taken to be
18·15 cm3, although this is not the average of the three titres in the table? (1)
(ii) Calculate the concentration of iodine, in mol l–1, in the foot rot treatment solution. Show
your working clearly. (3)
(iii) Describe how to prepare 250 cm3 of a 0·10 mol l−1 standard solution of sodium
thiosulfate, Na2S2O3. Your answer should include the mass, in g, of sodium thiosulfate
required. (3)
42
13. A student carried out an investigation to measure the nitrite level in the school water
supply. A compound, which reacts with the nitrite ions to form a product that absorbs light,
is added to water samples. The higher the concentration of nitrite ions present in a water
sample, the greater the amount of light absorbed.
(a) The student prepared potassium nitrite solutions of known concentration by diluting
samples from a stock solution.
(i) Calculate the mass, in mg, of potassium nitrite, KNO2, needed to make 1 litre of
stock solution with a nitrite ion concentration of 250 mg l–1. (2)
(ii) Describe how the weighed potassium nitrate is dissolved to prepare the stock
solution to ensure that its concentration is accurately known. (2)
(iii) Why should the student use distilled or deionised water rather than tap water when
dissolving the potassium nitrite? (1)
(iv) To prepare a solution with a nitrite ion concentration of 0·05 mg l –1 the student
dilutes the stock solution. Why is this method more accurate than preparing a
solution by weighing out potassium nitrite? (1)
(b) The graph below shows results for five solutions of potassium nitrite and a sample of
distilled water.
The results for four tap water samples are shown below.
What is the concentration of nitrite ions, in mg l–1, in the tap water? (2)
43
14. Soft drinks contain many ingredients. Caffeine is added to some soft drinks. The
concentration of caffeine can be found using chromatography. A chromatogram for a
standard solution containing 50 mg l−1 of caffeine is shown below.
Results from four caffeine standard solutions were used to produce the calibration graph
below.
44
Chromatograms for two soft drinks are shown below.
(a) What is the caffeine content, in mg l−1 of soft drink X? (1)
(b) The caffeine content of the soft drink Y cannot be determined from its chromatogram.
What should be done to the sample of soft drink Y so that the caffeine content could be
reliably calculated? (1)
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15. When forensic scientists analyse illegal drugs, anaesthetics such as lidocaine, benzocaine
and tetracaine are sometimes found to be present. The gas chromatogram below is from an
illegal drug.
(a) The structures of lidocaine, benzocaine and tetracaine are shown below.
Suggest why benzocaine has a shorter retention time than tetracaine. (1)
(b) Why is it difficult to obtain accurate values for the amount of lidocaine present in a
sample containing large amounts of caffeine? (1)
(c) Add a peak to the diagram below to complete the chromatogram for a second sample
that only contains half the amount of tetracaine compared to the first. (1)
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