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CfE Higher Chemistry
Unit 1
Chemical Change and Structure
Homework Booklet
Subsections:
1. Controlling the Rate of Reaction
2. Periodicity
3. Structure and Bonding in Elements and Compounds
Please do not write in this booklet. Answer questions in your homework jotter.
1
1. Controlling the Rate of Reaction
1. Two identical samples of zinc were added to an excess of two solutions of sulphuric
acid, concentrations 2 mol l-1 and 1 mol l-1 respectively.
Which of the following would have been the same for the two samples?
A. The total mass lost
B. The total time for the reaction
C. The initial rate of the reaction
D. The average rate of evolution of gas
2. 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 which molecules need for successful collisions
C. provides energy so that more molecules have successful collisions
D. form bonds with reacting molecules
3. Which of the following graphs of rate of reaction against temperature would apply
to the neutralisation of dilute hydrochloric acid with zinc?
A.
B.
Rate
Rate
Temperature
Temperature
C.
D.
Rate
Rate
Temperature
4.
Temperature
A small increase in temperature results in a large increase in rate of reaction.
The main reason for this is that
A. more collisions are taking place
B. the enthalpy change is lowered
C. the activation energy is lowered
D. many more particles have energy greater than the activation energy
2
5.
The following results were obtained in the reaction between marble chips and
dilute hydrochloric acid.
time/minutes
Total volume of carbon
dioxide production/cm3
0
0
2
52
4
68
6
78
8
82
10
84
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
6.
Number of
molecules
EA
X
Kinetic energy
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
7.
The graph below shows the variation of concentration of a reactant with time
as a reaction proceeds.
concentration/ mol/l
0.20
0.15
0.10
0.05
0.00
0
10
20 30
40
50
time/s
What is the average rate during the first 20s?
A.
B.
C.
D.
0.0025 mol
0.0036 mol
0.0075 mol
0.0090 mol
l-1s-1
l-1s-1
l-1s-1
l-1s-1
3
8.
T1
Number
of molecules
T2
EA
Kinetic energy of the molecules
Which of the following is the correct interpretation of the above energy distribution
diagram for a reaction as the temperature decreases from T2 to T1 ?
A
B
C
D
Activation energy
(EA)
remains the same
decreases
decreases
remains the same
Number of successful
collisions
increases
decreases
increases
decreases
9. The continuous use of large extractor fans greatly reduces the possibility of an
explosion in a flour mill. This is mainly because;
A.
B.
C.
D.
a build up in the concentration of oxygen is prevented
local temperature rises are prevented by the movement of the air
particles of flour suspended in the air are removed
the slow accumulation of carbon monoxide is prevented
10. When copper carbonate reacts with excess acid, carbon dioxide is produced. The
curves shown were obtained under two different conditions.
Volume of
CO2
P
Q
Time
The change from P to Q could be brought about by;
A.
B.
C.
D.
increasing the concentration of the acid
decreasing the mass of copper carbonate
decreasing the particle size of the copper carbonate
adding a catalyst
4
11.
A student carried out an experiment 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 is shown.
Results
Experiment Volume
Of KI(aq)
/cm3
1
25
Volume of Volume of
H2O(aq)
H2O2(aq)
/cm3
/cm3
0
5
Volume of Volume of
H2SO4(aq) Na2S2O3(aq)
/cm3
/cm3
10
10
Rate
/s-1
0.043
2
3
a) Describe how the concentration of the potassium iodide solution was changed
during this series of experiments
b) Calculate the reaction time, in seconds, for the first experiment.
(1)
(1)
12. Nitrogen dioxide gas can be prepared in different ways.
It is manufactured industrially as part of the Ostwald process. In the first stage of
the process, nitrogen monoxide is produced by passing ammonia and oxygen over
a platinum catalyst.
NH3(g) + O2(g)
NO(g) + H2O(g)
a)
b)
Balance the above equation.
Platinum metal is a heterogeneous catalyst for this reaction.
What is meant by a heterogeneous catalyst?
(1)
(1)
13. Catalytic converters in car exhaust systems convert poisonous gases into less
harmful gases.
a)
Two less harmful gases are formed when nitrogen monoxide reacts with carbon
monoxide.
b)
Name the two gases produced.
(1)
The catalyst is made up of the metals platinum, palladium and rhodium.
Explain what happens to molecules in the exhaust gas during their conversion
to less harmful gases.
You may wish to draw a labelled diagram.
(2)
5
14. The rate of carbon dioxide production was measured in three laboratory
experiments carried out at the same temperature and using excess calcium
carbonate.
Experiment
A
B
C
Acid
of 0.10 mol l-1 sulphuric acid
3
40cm of 0.10 mol l-1 sulphuric acid
40cm3 of 0.10 mol l-1 hydrochloric
acid
Calcium carbonate
1g lumps
1g powder
1g lumps
40cm3
total volumee of gas collected/ml
The curve obtained for Experiment A is shown.
120
100
80
60
40
20
0
0
10
20
30
40
50
60
time/s
a) Use the graph to calculate the average reaction rate in mls-1 , between 10 and 20s.
b) Make a rough copy of the graph in your jotter.
Draw curves on the graph to show the results that could be obtained for
experiments B and C. Label each curve clearly.
c) Draw a labelled diagram of the assembled apparatus which could be used to
carry out this experiment.
6
15. A pupil made the following observations on dripping taps. 45cm of water was
collected from Tap A in 3 minutes. 340 cm3 of water was collected from Tap B in 20
minutes. By calculating the average rate of loss of water from each tap, find out
which tap was dripping faster.
16. A farmer records the weight of his pigs every Monday. Here is part of the record for
one of the pigs.
Date
Jan 1
Jan 8
Jan 15
Weight (kg)
75.85
76.50
79.10
Date
Jan 22
Jan 29
Feb 5
Weight (kg)
83.70
86.90
92.10
Calculate the average rate of weight gain per week:i) during the first week
ii) during the first three weeks iii) over the five week period
17. A pupil was attempting to measure the rate of a chemical reaction which produced a
gas. After six seconds 8cm3 of gas had been collected. After ten seconds the total
volume of gas collected was 14cm3.Calculate the average rate of the reaction during
this time interval (from six to ten seconds).
18. The graph opposite shows the
volume of hydrogen gas
released when a 10cm strip of
magnesium (mass = 0.1g) was
added to 30cm3 of 1 mol l-1
hydrochloric acid.
(a) Calculate the average rate
of reaction
i) over the first 15 seconds
ii) between 20 and 30 seconds
(b) How long did it take for the reaction to stop?
(c) The graph shows that the rate of reaction changes as the reaction proceeds. Explain
why it changes in this way.
7
19. Marble chips (calcium carbonate),
reacted with excess dilute
hydrochloric acid.
The rate of reaction was followed
by recording the mass of the
container and the reaction mixture
over a period of time. The results
of the experiment are shown in
the graph opposite.
(a) Write a balanced equation for the
reaction.
(b) Give a reason for the loss of mass of the container.
(c) Calculate the average rate of reaction over the first five minutes.
(d) Why does the average rate of reaction decrease as the reaction proceeds?
20. The results shown opposite were
obtained when 0.42 g of
powdered chalk was added to 20
cm3 hydrochloric acid,
concentration 2 mol l-1 (an
excess of the acid).
(a) Sketch the graph and add a solid line to the graph to show what would happen if
0.42g of chalk lumps was used instead of powdered chalk.
(b) Add a dotted line to the graph to show what would happen if 20 cm3 of 3 mol l1hydrochloric acid was used instead of 2 mol l-1 hydrochloric acid.
8
21. The overall rate of a chemical reaction is often taken as the reciprocal of time
(1/time). Graphs of reaction rate against concentration and reaction rate against
temperature are shown below.
Graph 1
Graph 2
a) From graph 1
i) Calculate the time taken for the reaction when the concentration is 0.4 mol l -1.
ii) Explain why the rate increases as the concentration increases.
b) From graph 2
i) Find the temperature rise needed to double the reaction rate .
ii) Explain why a small temperature rise can give a large increase in reaction rate.
22. A pupil was investigating the effect of temperature on the rate of a chemical
reaction and obtained the following data.
Temperature (oC)
15
25
33
37
44
Time for reaction to
finish (s)
154
66.7
40
30.3
22.2
Relative rate ( .......... )
(a) Copy the table and calculate the relative rate of reaction at each temperature
and add them to the table, putting the correct units in the brackets.
(b) Plot a graph of relative rate against temperature.
(c) Predict what the relative rate of the reaction will be at 50 0C.
9
(d) Use the graph to estimate the time for the reaction to finish at 40 0C.
23. (a) Explain why,
i.)
decreasing the particle size increases the rate of a chemical reaction.
ii)
increasing the concentration speeds up a chemical reaction.
(b) Give an everyday example of a reaction speeded up by
i) decreasing the particle size
ii) increasing the concentration
24. When hydrochloric acid is added to a solution of sodium thiosulphate the following
reaction takes place.
2HCl(aq)
+
Na2S2O3(aq) 
2NaCl(aq) + SO2(g)
+ S(s) + H20(l)
Solid sulphur forms in the solution.
In one set of experiments the effect of varying the concentration of sodium
thiosulphate was studied. Some of the volumes of solutions used are shown.
Volume of 0.05 mol l-1
Na2S2O3(aq)/ cm3
Volume of water/ cm3
Volume of 0.1 mol l-1
HCl(aq)/ cm3
Reaction time /s
200
0
5
20
160
25
120
33
80
50
40
100
(a) Copy and then complete the table to show the volumes of water and acid that
would have been used.
(b) Describe how the reaction time could have been measured.
(c) Describe how the relative rate of reaction would be obtained from each of the
results.
10
25. Hydrogen peroxide can be used to clean contact lenses. In this process, the enzyme
catalase is added to break down hydrogen peroxide. The equation for the reaction
is:
2H2O2

2H2O
+
O2
The rate of oxygen production was measured in three laboratory experiments
using the same volume of hydrogen peroxide at the same temperature.
Experiment
Concentration of H2O2/ mol l-1
Catalyst used
A
0.2
yes
B
0.4
yes
C
0.2
no
The curve obtained for experiment A is shown.
(a) Calculate the average rate of the reaction over the first 40 s.
(b) Copy the graph and add curves to the graph to show the results of experiments B and
C. Label each curve clearly.
(c) Draw a labelled diagram of assembled lab apparatus which could be used to carry
out these experiments.
11
26. The graph shows the concentrations of reactant and product as equilibrium is
established in a reaction.
(a)
Calculate the average rate of reaction over the first 10 s.
(b)
The equilibrium constant for a reaction is given the symbol K
In this reaction K is given by:
K= equilibrium concentration of product
equilibrium concentration of reactant
Calculate the value of K for the reaction.
(c)
The reaction is repeated using a homogeneous catalyst.
(i) What is meant by a homogeneous catalyst?
(ii) What effect would the introduction of the catalyst have on the value of K?
12
27. What is meant by the term “activation energy”?
28. The decomposition of an aqueous solution of hydrogen peroxide into oxygen and
water can be catalysed by iodide ions, I- (aq), or by solid manganese (IV) oxide,
MnO2(s).
For each of these catalysts, state whether the catalysis is homogeneous or
heterogeneous. Give a reason for your answer.
29. The diagram shows the distribution of kinetic energies in a sample of gas at 20°C.
number of
molecules
kinetic energy of the molecules (kJ)
a) Copy the diagram and add another line to show the kinetic energy distribution of
the molecules at 30°C.
1
b) Draw a line to represent the activation energy of a reaction which is slow at 20°C.
1
c) With reference to the completed diagram explain why an increase of temperature
of 10°C can lead to a large increase in reaction rate.
1
13
30.
Potential
energy
(kJ)
reaction pathway
a) i) Is this an endothermic or an exothermic reaction?
ii) Explain your answer.
b) What is the value of:
i) the activation energy for the forward reaction?
ii) the enthalpy change for the reaction?
iii) the energy of the activated complex?
c) The reaction shown can be speeded up by the use of a suitable catalyst.
What effect does a catalyst have on
i) the enthalpy change for the reaction?
ii) the activation energy for the forward reaction
d) i) What is meant by the term ‘activated complex’?
ii) Referring to the above diagram, at what energy would the activated complex be formed?
14
31.
2HI(g)
H2(g) + I2(g)
The reaction above is reversible. The activation energy for the forward reaction is 80 kJ
and the reverse reaction is 50 kJ.
a) Copy and complete the graph below to show how the potential energy varies as the
reaction proceeds.
160
140
Potential
energy
(kJ)
120
100
80
60
reactants
40
20
0
reaction pathway
b) Is the forward reaction is exothermic or endothermic.
c) Gold and platinum both catalyse the reaction. For the forward reaction E A
using gold is 30 kJ, while EA using platinum is 40 kJ.
i) using different dotted lines add this information to the graph.
ii) which is the better catalyst for the reaction? Explain your choice.
d) The gold and platinum catalysts are used in the solid state. Are the catalysts
heterogeneous or homogeneous catalysts? Explain your choice.
32. An advice leaflet given to motorists when catalytic converters were first used states:
“Cars fitted with catalytic converters must be run on unleaded petrol only.”
(a) Outline the reasons for fitting catalytic converters, naming the substances reacting
and what happens to them.
(b) Describe in terms of adsorption how this heterogeneous catalysts work, and state the
effect this has on the activation energy for the reaction.
(c) Describe how a substance poisons a catalyst.
15
2. Periodicity
1. Which equation represents the first ionisation energy of a diatomic element, X2?
A. ½X2(s)
X+(g)
B. ½X2(g)
X-(g)
C. X(g)
X+(g)
D. X(s)
X-(g)
2. As the atomic number of the alkali metal increases
A. the first ionisation energy decreases
B. the atomic size decreases
C. the density decreases
D. the melting point increases
3. Which of the following has the least attraction for bonding electrons?
A. carbon
B. nitrogen
C. phosphorus
D. Silicon
4. The spike graph shows the variation in successive ionisation energies of
an element, Z.
Ionisation
energy/
KJ mol-1
1st 2nd 3rd 4th 5th 6th
Electron removed
In which group of the Periodic Table is element Z?
A. 1
B. 3
C. 4
D. 6
16
4. a) Which arrow (A) or (B) indicates correctly a decrease in
atomic size?
b) Explain why atomic size decreases in this way.
6. a) Define the first ionization energy of an element.
b) Which arrow (A) or (B) indicates correctly a decrease in the
first ionization energy of elements?
c) Give two reasons why the ionisation energy decreases in
this way
d) Explain why the ionisation energy is an endothermic process.
7.
The graph below shows the first ionisation energies of successive elements with
increasing atomic number.
Elements A, B and C belong to the same group of the Periodic Table. Identify the
group
8. a) Explain why a Na+ ion is larger than a Mg2+ ion.
b) Explain a calcium, Ca2+ ion is larger than a Mg2+ ion.
c) Explain why a Si4+ ion is smaller than a P3- ion.
17
9. a) Write an equation corresponding to the first ionisation energy of Aluminium
b) Explain why the first ionisation energy of fluorine is greater than the first
ionisation energy of lithium.
c) Explain why the first ionisation energy of sodium is less than the first ionisation
energy of lithium.
d) Explain why second ionisation energy of sodium is very much higher than the first
ionisation energy of sodium.
10. a) Using the ionisation energies in the data booklet calculate the energy required
For;
Al(g)
Al3+ (g) + 3e --
b) Explain why the third ionisation energy of magnesium (7750 kJ mol-1) is so much
greater than the third ionisation of aluminium (2760 kJ mol-1).
11. Ionisation energies can be found by applying an increasing voltage across test
samples of gases until the gases ionise.
The results below were obtained from experiments using hydrogen atoms and then
helium atoms.
Element
Voltage at which an atom
of gas ionises/V
hydrogen
13.6
no further
change
helium
24.6
54.5
a) Why are there two results for helium but only one for hydrogen?
b) (i) Write an equation which would represent the first ionisation energy of helium
gas.
(ii) Why is the first ionisation of helium higher than that of hydrogen?
c) The ionisation energy, I.E. , can be found from:
I.E. = voltage x 1.6 x 10-19 J
Calculate a value for the first ionisation energy of helium.
18
Bonding and Structure in Elements and Compounds
1. Which of the following chlorides is likely to have the least ionic character?
A. BeCl2
B. CaCl2
C. LiCl
D. CsCl
2. Which chloride is most likely to be soluble in tetrachloromethane, CCl4?
A. barium chloride
B. caesium chloride
C. calcium chloride
D. phosphorus chloride
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. Which of the following occurs when crude oil is distilled?
A. covalent bonds break and form again
B. Van der Waals’ bonds break and form again
C. covalent bonds break and Van der Waals’ bonds form
D. Van der Waals’ bonds break and covalent bonds form
5. A substance melts at 1074oC and boils at 1740oC. The passage of an electric
current through the molten substance results in electrolysis.
What type of structure is presenting the substance?
A. ionic
B. metallic
C. covalent molecular
D. covalent network
6. In which of the substances, in the solid state, would Van der Waals’ attractions be a
significant “intermolecular force”?
A. sodium chloride
B. carbon dioxide
C. magnesium
D. ice
19
7.
Hydrogen gas has a boiling point of –253oC.
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.
8. The elements in the second row of the Periodic Table are shown below.
Li
a)
b)
c)
9.
Be
B
C
N
O
F
Ne
Why does the atomic size decrease crossing the period from lithium to neon?
Diamond and graphite are forms of carbon that exist as network solids.
Name a form of carbon that exists as discrete molecules.
Use the electronegativity values to explain why nitrogen chloride contains pure
covalent bonds.
Compared to other gases made up of molecules of similar molecular masses,
ammonia has a relatively high boiling point.
N
H
H
H
In terms of the intermolecular bonding present, explain clearly why ammonia
has a relatively high boiling point.
10.
The formulae for three oxides of sodium, carbon and silicon are Na2O, CO2 and
SiO2.
Copy and complete the table for CO2 and SiO2 to show both the bonding and
Structure of the three oxides at room temperature.
Oxide
Na2O
CO2
SiO2
Bonding and structure
ionic lattice
20
11.
All types of bonding involve electrostatic attraction between positively charged
particles and negatively charged particles.
Copy and complete the table showing the three types of strong bonding force.
Positively charged
particles
Type of bonding
Negatively charged
particles
shared electrons
ionic
positive nucleus
12.
The melting and boiling points and electrical conductivities of four substances
are given below in the table.
Substance
A
B
C
D
Melting
point/oC
92
1050
773
1883
Boiling point/
oC
190
2500
1407
2503
Solid conducts
electricity?
no
yes
no
no
Melt conducts
electricity?
no
yes
yes
no
Copy and complete the table below by adding the appropriate letter for each type
of bonding and structure.
Substance
Bonding and structure at
room temperature
covalent molecular
covalent network
ionic
metallic
13. a) Predict the type of bonding (non-polar covalent, polar covalent or ionic) between
the atoms of the following elements.
i) magnesium and sulphur
ii) oxygen and phosphorus
iii) nitrogen and nitrogen
iv) fluorine and oxygen
b) Justify, by reference to difference in electronegativity values, your answers to i)
and iv).
21
14. a) Graph 1 shows the boiling points of the Group 7
elements.
Why do the boiling points increase down
Group 7?
b) Graph 2 shows the melting points of
elements from lithium to neon across
the second period. Give reasons for
the high melting points of boron and
carbon.
15.
The Periodic Table below has been divided into four sections - A, B, C and D.
a) State the type of structure in each of the four sections A, B, C and D.
b) In which section(s) will London Dispersion forces between the particles be
significant?
c) Using elements in the above table as examples, explain briefly the difference
between a covalent molecular substance and a covalent network substance.
22
16. There are many types of attractive force, some are weak and some are strong.
A
positively charged ions and
negatively charged ions
B
temporary dipole and
induced dipole
D
permanent dipole and
permanent dipole
E
positively charged nuclei and
shared electrons
C
positively charged nuclei and
delocalised electrons
Identify the statement(s) referring to;
a) London dispersion forces
b) the forces between oxygen and hydrogen atoms in water
c) the intermolecular forces in HCl gas.
d) ionic bonds
17.
The covalent bond in hydrogen chloride gas is polar and the molecule is polar.
The covalent bonds in silicon tetrachloride are also polar. Explain why the silicon
tetrachloride molecule is non-polar.
18.
There are many types of bonding force between atoms and molecules.
A
permanent dipole to
permanent dipole interactions
D
ionic bonds
B
non-polar covalent bonds
C
E
F
metallic bonds
hydrogen bonds
London dispersion forces
a) Identify the three forces present in hydrogen fluoride.
b) Identify the force(s) present in
i) methane ii) sodium chloride iii) hydrogen bromide
c) Identify the bond(s) and/or force(s) of attraction
i) responsible for the low boiling point of argon.
ii) that can exist between molecules.
iii) that allows electrons a lot of free movement.
23
iv) neon v) oxygen
19.
Which of the compounds below have:
Identify the substance(s) where the intermolecular forces are
a) van der Waals’ forces only.
a) London dispersion forces
b) hydrogen bonds
c) permanent dipole to dipole interactions
20.
Elements and compounds show a variety of structures.
A
B
Cl2
D
C
Na
E
SiO2
NaCl
F
NH3
C(diamond)
Identify the substance(s)
a) with a tetrahedral arrangement of bonds in a covalent network.
b) which can conduct electricity because of delocalised electrons.
c) with discrete covalent molecules
21.
Many of the properties of water arise from the presence of polar O - H bonds
which make the water molecules polar.
Carbon dioxide contains polar C = O bonds but its molecules are non-polar.
a) Explain this difference in overall polarity with the aid of diagrams of each
molecule, showing polarities.
b) Water is unusual in that in the solid form (ice) is less dense than the liquid
form. Explain why water behaves in this way.
24
22. Both bonded and non-bonded pairs of electrons surrounding atoms repel each other
and this determines the shape of a molecule. The following procedure is used to find
the total number of pairs of electrons around a central atom.
(i) Note the number of electrons in the outer energy level (shell) of the central
atom.
(ii) Note the number of other atoms present --- each atom provides one electron for
bonding.
(iii) Add (i) and (ii) to give the total number of electrons.
(iv) Divide this number by two to give the number of electron pairs - both bonded and
non-bonded pairs.
Example:- with ammonia, NH3, N is the central atom.
(i) 2,5
= 5 outer electrons
(ii) 3H
3 x 1 = 3 electrons
(iii) Total
= 8 electrons
(iv) 8 electrons gives four pairs.
Since NH3 only has 3 bonds there are 3 bonded pairs and one non-bonded pair. The 4
pairs of electrons repel each other, giving the pyramid shape of the ammonia
molecule as shown in the first row of the table.
Copy and complete the table.
Formula
Outer
electrons in
central atom
Total number
of electrons
Bonded pairs
Non-bonded
pairs
NH3
5
8
3
1
CCl4
4
4
0
H2O
6
8
2
2
10
5
PF5
25
Molecular
shape
23.
The boiling points of compounds depend on the intermolecular forces.
Molecular
mass
Boiling point (0C)
Name
Formula
butane
CH3CH2CH2CH3
- 0.5
propanone
CH3COCH3
56
a) Copy the table and calculate the molecular mass for each compound.
b) Explain why the boiling points are different describing the effect of the London
dispersion forces present and explain how these forces arise.
24.
The Group 5 hydrides are covalent compounds.
Compound
NH3
PH3
AsH3
Enthalpy of formation/kJ mol-1
-46
+6
+172
Boiling point/K
240
185
218
a) What is the trend in stability of the group 5 hydrides?
b) Explain why the boiling point of NH3 is higher than the boiling point of PH3 and
AsH3.
25.The table below shows the boiling point, molecular mass and structure of the
simplest alkanol, methanol, the simplest alkanoic acid, methanoic acid and the ester
methyl methanoate which forms when the acid and the alkanol react together in a
condensation reaction.
a) Using molecular mass as the only criterion, use the boiling points of methanol and
methanoic acid to predict the boiling point of methyl methanoate.
b) The boiling point of the ester is in fact 320C. Explain in terms of the intermolecular
forces why this value is so different from your prediction.
26
26.
The bar chart shows the melting points of chlorides of elements 3 to 20 (with no
bars for 10, 15 and 18).
a) Copy and complete this statement describing the pattern for these melting
points related to the Group number.
In general as the Group number increases the melting point of the chloride
................
b) Explain why no values are given for elements 10 and 18.
c) From the bar chart, state which of the chlorides has the weakest forces
between the
molecules.
c) Predict a value for the melting point of the chloride of element 15.
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Revision of Nat 5 Enthalpy calculations
1.) A student used the simple laboratory apparatus shown to determine the enthalpy
of combustion of methanol.
(a) (i) What measurements are needed to calculate the energy released by the
burning methanol?
(ii) The student found that burning 0.370 g of methanol produces 3.86 kJ of energy.
Use this result to calculate the enthalpy of combustion of methanol.
CH3OH
GFM = 32.
(b) A more accurate value can be obtained using a bomb calorimeter
One reason for the more accurate value is that less heat is lost to the
surroundings than in the simple laboratory method.
Give one other reason for the value being more accurate in the bomb
calorimeter method.
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2.)
The energy changes taking place during chemical reactions have many everyday uses.
(a) Some portable cold packs make use of the temperature drop that takes place when
the chemicals in the pack dissolve in water.
Name the type of reaction that results in a fall in temperature.
(b) Flameless heaters are used by mountain climbers to heat food and drinks. The
chemical reaction in a flameless heater releases 45 kJ of energy.
If 200 g of water is heated using this heater, calculate the rise in temperature of
the water, in °C.
3.)
The enthalpies of combustion of some alcohols are shown in the table.
(a) Using this data, predict the enthalpy of combustion of butan-1-ol, in kJ mol–1.
(b) A value for the enthalpy of combustion of butan-2-ol, C4H9OH, can be
determined experimentally using the apparatus shown.
Mass of butan-2-ol burned =1·0 g
Temperature rise of water =40 °C
Use these results to calculate the enthalpy of combustion of butan-2-ol, in kJ mol–1.
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4.) Ethanol, C2H5OH, can be used as a fuel in some camping stoves.
(a) The enthalpy of combustion of ethanol is -1367 kJ mol-1.
Using this value, calculate the number of moles of ethanol required to raise the
temperature of 500 g of water from 18 °C to 100 °C.
(b) Suggest two reasons why less energy is obtained from burning ethanol in the
camping stove than is predicted from its enthalpy of combustion.
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Moles and Solutions Calculations
1. Calculate the number of moles of sodium chloride in 500cm3 of 0.5 mol l-1
solution.
2. What is the concentration of a solution if 1 mole of sodium chloride is dissolved in
0.5 litres of solution.
3. What volume of 0.4 mol l-1 solution contains 0.5 moles of sodium hydroxide.
4. How many moles are present in 400cm3 of 0.2 mol l-1 sulphuric acid.
5. Calculate the volume of 1.25 mol l-1 potassium chloride solution that contains 0.2
moles.
6. Calculate the concentration of a solution containing 0.05 moles in 150cm 3 of
solution.
7. find by calculation which of the following contains the greatest number of moles
of solute:
8.
a) 100 cm3 of 0.3 mol l-1 hydrochloric acid
b) 50 cm3 of 0.7 mol l-1 sodium chloride
c) 40 cm3 of 0.8 mol l-1 sulphuric acid
d) 80 cm3 of 0.5 mol l-1 sodium phosphate
9. Calculate the concentration of a citric acid solution given that 1.5 moles of citric
acid are dissolved in water and made up to a final volume of 3 litres.
10. Calculate the concentration of a solution of sodium hydroxide, given that it
contains 12g of NaOH in 500cm3 of solution.
11. Calculate the volume of a solution produced if 53g of sodium carbonate Na 2CO3 is
used to make a solution with a concentration of 0.25 mol l-1
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