Download Precision, accuracy and significant figures

Worksheet 1.1
Precision, accuracy and
significant figures
INTRODUCTION
For a quantity to have an exact value, it must either be defined or obtained by counting. All
measured quantities have an inherent uncertainty because all instruments used to make
measurements have limitations, and the people operating the instruments have varying skills.
The accuracy of a measurement is an expression of how close the measured value is to the ‘correct’
or ‘true’ value. The precision of a set of measurements refers to how closely the individual
measurements agree with one another. Thus, precision is a measure of the reproducibility or
consistency of a result.
The precision of a measurement is sometimes expressed as an uncertainty using a plus/minus (±)
notation to indicate the possible range of the last digit. An alternative method is to indicate the
certainty of the measurement by the use of significant figures.
To clarify the number of significant figures in a measurement, the value may be written in standard
form. A number written in standard form is expressed as a number greater than 1 but less than 10
multiplied by 10x, where x is an integer.
When a calculation involves multiplication and division, the result should have the same number of
significant figures as the factor with the least number of significant figures. For addition and
subtraction calculations, the result should have the same number of decimal places as the number
used with the fewest decimal places. In most calculations you will need to round off numbers to
obtain the correct number of significant figures.
No.
Question
Answer
1
Which of the following quantities would have an
inherent uncertainty?
A The number of pages in this book
B Your measured height (in cm)
C The number of mL in 6.0 L
D A volume of liquid measured using a pipette
2
Shooting at targets may be used as an analogy to show the ideas of precision and accuracy in
measurements. Label each of the shooting targets shown as representing one of the following
situations.
N for neither accuracy nor precision
B for both precision and accuracy
P for precision, but inaccuracy
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Worksheet 1.1
Precision, accuracy and
significant figures
No.
Question
Answer
3
Why are measurements in experiments often
repeated several times and the results averaged?
4
State the number of significant figures in each of
the following measured quantities.
a A temperature reported as 26.1°C
b A burette reading of 32.34 mL
c A mass reading of 0.0471 g
d A time recorded as 6.000 s
5
Express each of the following numbers in
standard form, ensuring you use the correct
number of significant figures.
a 140.7
b 5005
c 980.0
d 0.0075
6
Round each of the following numbers to three
significant figures, and express in standard form.
a 7.8001
b 600.5
c 98.345
d 0.000600
7
Express the number 6000 in standard form to
show that it contains:
a 1 significant figure
b 4 significant figures.
8
Calculate each of the following and express the
answers to the correct number of significant
figures.
a 5.6 × 120
b 0.0045 × 67.1
c 0.046 ÷ 0.023
d 63 × 7.06
9
Perform the following calculations and round off
the answers to the correct number of significant
figures.
a 3.256 + 45.2 – 3.815
b 12.13 + 342.0 + 4.108
10
A dozen eggs have a mass of 722 g. What is the
average mass of the eggs?
Page 2
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Worksheet 1.2
Calculations involving gases
and solutions
NAME:
CLASS:
INTRODUCTION
The following formulas are essential! You must know how and when to use them.
Formula
What it means
m
mass of sample (g)
n=
Amount (mol) =
M
molar mass (g mol−1 )
N
number of particles
n=
Amount (mol) =
NA
Avogadros number (mol−1 )
For solutions
n=c×V
Amount (mol) = concentration (mol L–1) × volume (L)
For gases at standard temperatures and pressures
V
volume of gas (L)
n=
Amount (mol) =
VM
molar volume (L mol−1 )
where molar volume, VM, is
22.4 L mol–1 at STP (0°C and 1.0 atm)
24.5 L mol–1 at SLC (25°C and 1.0 atm)
For gases
pV = nRT
Pressure (kPa) × volume (L)
and
= amount (mol) × gas constant (8.31 J K–1 mol–1) × temperature (K)
m
pV =
RT
M
For a fixed amount of a gas
pV
= k or
pressure of gas × volume of gas
T
= a constant
temperature of gas (K)
p 2V2
p1V1
=
T1
T2
Unit of
concentration
% m/m
% m/v
% v/v
ppm
ppb
What it means
Mass of solute (in g) in 100 g of solution
Mass of solute (in g) in 100 mL of solution
Volume of solute (in mL) in 100 mL of solution
Mass of solute (in g) in 106 g of solution
(equivalent to mg L–1 for dilute solutions)
Mass of solute (in g) in 109 g of solution
(equivalent to μg L–1 for dilute solutions)
For dilution of a solution: c1 × V1 = c2 × V2
where c1 = initial concentration, V1 = initial volume, c2 = final concentration, V2 = final volume.
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Worksheet 1.2
Calculations involving gases
and solutions
No.
Question
Answer
1
a What amount (in mol) of oxygen is
present in 5.00 g of glucose, C6H12O6?
b How many atoms are present in 8.36 g
of hydrogen peroxide, H2O2?
2
What volume of water should be added to
35.0 mL of 0.30 M H2SO4 in order to
produce a 0.090 M solution?
3
What pressure does 0.49 mol of SO3 exert
in a sealed 3.0 L vessel at 54ºC?
4
The heaviest known atom has a mass of
about 4 × 10–22 g. What would be the
mass of one mole of these atoms?
5
2.5 g of a gas initially occupying a
volume of 600 mL, at 260 K, is heated to
325 K at constant pressure. What would
its new volume be?
6
What volume will 62.0 g of carbon
dioxide gas occupy at a temperature of
124°C and 210 kPa?
Page 2
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Worksheet 1.2
Calculations involving gases
and solutions
No.
Question
Answer
7
What volume of water must be added to
1.0 L of a solution containing 70.2 g of
NaCl to produce a solution of 0.67 M
NaCl?
8
2.06 g of a hydrocarbon occupies 16 L at
27ºC and 20 kPa. Find the molar mass of
this compound, and so identify the
hydrocarbon.
9
a What volume of 5.00% m/v cloudy
ammonia cleaning solution is needed
to make 250 mL of a 1.50% m/v
solution?
b What mass of ammonia is present in
150 mL of the 1.50% solution?
10
3.0 g of carbon dioxide occupies 687 mL
at 143 800 Pa. What volume does it
occupy at a pressure of 199 kPa,
assuming temperature is constant?
11
0.778 g of one of the halogens (Group 17)
was found to occupy a volume of 122 mL
at a pressure of 99.8 kPa and a
temperature of 26°C. Which halogen was
it?
12
A sample of water from a waterway is
found to contain 600 ppm mercury. What
is this concentration in ppb?
Page 3
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Worksheet 1.2
Calculations involving gases
and solutions
No.
Question
Answer
13
3.5 g of Pb(NO3)2 is added to 60 mL of
distilled water in a beaker and stirred to
dissolve the solid. 10 mL of this solution
is then transferred to another beaker and
mixed with 20 mL of distilled water.
What are the concentrations (in mol L–1)
of these two solutions?
14
30 mL of a sodium carbonate solution is
made up to a total volume of 300 mL
with distilled water. The resultant
solution has a sodium carbonate
concentration of 0.108 M. What mass of
sodium carbonate was present in the
original solution?
15
50 mL of a 2.0% m/v glucose solution is
mixed with 50 mL of a 6.0% m/v glucose
solution. The solution is made up to a
total of 300 mL with distilled water. What
is the concentration of the final solution?
16
A sample of N2 gas collected at 25°C and
750 mmHg pressure occupies 190 mL.
What volume will it occupy at STP?
17
50 mL of a 16% m/v silver nitrate
solution is added to an equal volume of
distilled water. What is the concentration
of the dilute solution in ppm?
18
A solution of silver nitrate (AgNO3) is
made by dissolving 2.33 g of solid in
398 mL of distilled water. What is the
concentration of this solution in
a % m/v?
b M?
Page 4
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Worksheet 1.3
Stoichiometry
NAME:
CLASS:
INTRODUCTION
Stoichiometric problems can be tackled by remembering five basic steps.
Step 1: Write a balanced chemical equation for the reaction.
Step 2: List all data given, including relevant units. Remember to also write down the symbol of the
unknown quantity.
Step 3: Convert the data given to moles, using the relevant formulas:
N
pV
V
m
n=
n=c×V n=
n=
n=
M
NA
RT
VM
Step 4: Use the chemical equation to determine the mole ratio of the unknown quantity to the
known quantity. This ratio enables calculation of the number of moles of the unknown
quantity.
Step 5: Finally, convert this number of mole back into the relevant units of the unknown.
In all calculation questions, take care to show all steps in your working, including reacting ratios
where relevant. Also, take care to add the correct unit to your answer (e.g. mol, L) and give your
answer to the correct number of significant figures.
No.
Question
Answer
1
SO2 in the atmosphere contributes to
acid rain. The equation for formation of
the acid is represented by the equation:
2SO2(g) + O2(g) + 2H2O(l)
→ H2SO4(aq)
What mass of sulfuric acid will form
from 50.0 L of sulfur dioxide at SLC?
2
Phosphoric acid can be generated by the
oxidation of phosphorus with nitric acid
according to the following equation:
P(s) + 5HNO3(aq)
→ H3PO4(aq) + H2O(l) + 5NO2(g)
If sufficient reactants are available to
produce 1.00 kg of phosphoric acid
(H3PO4), what mass of nitrogen dioxide
will also be generated in the reaction?
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Worksheet 1.3
Stoichiometry
No.
Question
Answer
3
23.8 mL of sulfuric acid is just
neutralised by 29.9 mL of 2.86 M
sodium hydrogen carbonate solution,
according to the equation:
2NaHCO3(aq) + H2SO4(aq)
→ Na2CO3(aq) + CO2(g) + H2O(l)
Determine the concentration of the
sulfuric acid solution used.
4
In a car accident, the impact triggers
ignition of a detonator cap in the air bag,
which causes sodium azide (NaN3) to
decompose explosively.
a Write the equation to show the
decomposition of sodium azide into
solid sodium and nitrogen gas.
b If the bag contained 75 g of sodium
azide, what volume of gas would
form at 30ºC and 101.3 kPa?
5
Sodium thiosulfate reacts with bromine
in alkaline solution according to the
equation:
Na2S2O3(aq) + 4Br2(l) + 10NaOH(aq)
→ 2Na2SO4(aq) + 8NaBr(aq) + 5H2O(l)
In order to completely react with
10.0 mL of bromine of density
3.12 g mL–1, 239 mL of NaOH was
added along with excess Na2S2O3.
Determine the concentration of the
sodium hydroxide solution required.
6
Ammonium sulfate, an important
fertiliser, can be prepared by the reaction
of ammonia with sulfuric acid according
to the equation:
2NH3(g) + H2SO4(l) → (NH4)2SO4(aq)
Calculate the volume of NH3 needed to
react with 19.56 g of H2SO4 at 87°C and
2.99 atm.
Page 2
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Worksheet 1.3
Stoichiometry
No.
Question
Answer
7
When magnesium carbonate is heated
strongly, it decomposes to magnesium
oxide and carbon dioxide gas.
What volume of carbon dioxide would
be produced at STP when 100 g of
magnesium oxide was generated in this
reaction?
8
When solid calcium carbonate reacts
with nitric acid solution, neutralisation
takes place.
a Write the equation for this reaction.
b If 10.0 g of calcium carbonate reacts
with 100 mL of 0.500 M nitric acid,
what volume of carbon dioxide is
formed at SLC?
9
Arsenic undergoes oxidation by a hot,
concentrated solution of sodium
hydroxide to produce sodium arsenate
and hydrogen gas according to the
equation:
2As(s) + 6NaOH(aq)
→ 2Na3AsO3(s) + 3H2(g)
6.57 g of arsenic is reacted with 250 mL
of 0.779 M sodium hydroxide solution.
Calculate the mass of hydrogen gas
evolved in the process.
10
8.00 g of barium hydroxide is dissolved
in 120 mL of 1.886 M hydrochloric acid
to produce barium chloride solution and
water.
a Write an equation to represent this
neutralisation reaction
b Determine the concentration of the
barium chloride solution that results.
Page 3
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Worksheet 2.1
Determining the molecular
formula of a gas
NAME:
CLASS:
INTRODUCTION
Empirical formula: The formula showing the smallest whole-number ratio of atoms in a substance.
Molecular formula: The formula showing the actual number of atoms of each element present in a
molecule of a compound. For example, the molecular formula of glucose is C6H12O6 and its
empirical formula is CH2O.
No.
Question
1
A compound used in ceramics
contains, by mass, 22.8% Na,
21.5% B and 55.7% O. What is its
empirical formula?
2
A sweet-smelling organic
compound contains 0.0556 mol of
carbon, 0.112 g of hydrogen
atoms and 9.57 × 1021 oxygen
atoms. Calculate its empirical
formula.
3
If a compound contains 75.7%
arsenic and 24.3% oxygen by
mass, and has a molar mass of
395.6 g mol–1, what is its
molecular formula?
Answer
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Worksheet 2.1
Determining the molecular
formula of a gas
No.
Question
4
A compound has the following
composition by mass: carbon
40.0%, hydrogen 6.67% and
oxygen 53.3%. If the compound
has a molecular mass of
approximately 60, determine both
its empirical and molecular
formulas.
Answer
The apparatus shown below may be used to determine the empirical and molecular formulas of an
unknown compound of general formula CxHyOz.
The questions that follow refer to this apparatus.
No.
Question
5
What is the purpose of the
CaCl2(s)?
6
What is the purpose of the
NaOH(aq)?
Answer
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Worksheet 2.1
Determining the molecular
formula of a gas
No.
Question
7
14.34 g of a hydrocarbon was
burned in excess oxygen. The
mass of the NaOH(aq) increased
by 43.54 g, while the mass of
CaCl2(s) increased by 22.26 g.
What is the empirical formula of
the hydrocarbon?
Answer
A method used to determine the relative molecular mass of a volatile liquid is shown below.
Steam is passed through the outer jacket surrounding a graduated syringe that contains a small,
measured volume of air. When the temperature has stabilised, a weighed sample of a few mL of
liquid is injected into the graduated syringe using a small hypodermic syringe. The liquid vaporises
and the final volume of air plus sample is recorded.
The questions which follow refer to this apparatus.
No.
Question
8
Why would this method be
unsuitable for a liquid with a
boiling point above 90°C?
9
Using the apparatus shown, 0.16 g
of liquid was vaporised to produce
a volume of 46 mL, at a
temperature of 100°C and a
pressure of 1.0 atm. Determine the
molar mass of the liquid.
Answer
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Worksheet 2.1
Determining the molecular
formula of a gas
No.
Question
10
In one experiment an 18.99 g
sample of the compound was
burnt in excess oxygen. When the
gases evolved were passed
through anhydrous CaCl2, its mass
increased by 11.38 g. The
remaining gases, when bubbled
through a NaOH solution,
increased its mass by 27.83 g.
In a separate experiment, a 6.21 g
sample of the compound was
vaporised. The vapour occupied
2.17 L at 200°C and 1.25 × 105 Pa.
Calculate the molecular formula
of the compound.
11
Which instrumental method is
routinely used to determine the
molar mass of an unknown
organic compound?
Answer
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Worksheet 2.2
Gravimetric analysis of a
fertiliser
NAME:
CLASS:
INTRODUCTION
The purity of a sample of ammonium sulfate fertiliser was determined by gravimetric analysis in
which the sulfate was precipitated as barium sulfate. The following steps were used in this
procedure:
1
A sample of about 5 g of the fertiliser was dried in an oven for an hour at 110ºC.
2
The dried fertiliser was ground into fine powder using a mortar and pestle.
3
1.00 g of the dried, ground fertiliser was weighed out into a 250 mL beaker.
4
150 mL of distilled water and 2 mL of 10 M HCl were added to the fertiliser with stirring.
5
The solution was heated to nearly boiling.
6
The fertiliser solution was maintained at that temperature while the precipitating agent was
prepared. If the fertiliser was pure, it was calculated that 37.9 mL of the 0.200 M BaCl2 stock
solution would be required. 45 mL of the stock solution was placed in a 100 mL beaker and
heated to nearly boiling.
7
The fertiliser solution was stirred vigorously as the hot barium chloride solution was added
slowly.
8
Heating of the beaker continued while the precipitate was allowed to settle. A few drops of
the supernatant liquid were tested for complete precipitation by adding them to a few drops of
barium chloride solution on a watchglass. No cloudiness was observed.
9
The fertiliser–barium chloride mixture was covered and left to ‘digest’ just below boiling
point for one hour. (This digestion process is necessary for experiments involving barium
sulfate because the crystals formed are initially too small to filter. Digestion allows larger
crystals to form.)
10 The hot solution was filtered under vacuum filtration into a pre-weighed sintered glass
crucible.
11 The precipitate was washed several times with hot water.
12 The crucible and precipitate were dried to constant mass in an oven at 60ºC over two
successive days; 1.63 g of precipitate was obtained.
No.
Question
1
Why was it necessary to dry and
grind the fertiliser before
weighing?
2
If the fertiliser was 100%
ammonium sulfate, what amount
(in mol) of ammonium sulfate
would be present in 1.00 g?
Answer
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Worksheet 2.2
Gravimetric analysis of a
fertiliser
No.
Question
3
This procedure assumes that there
were no insoluble components in
the fertiliser. If there had been,
what extra step would have been
necessary?
4
Step 6 states that 37.9 mL of
barium chloride was needed.
Show how this volume was
calculated.
5
If 37.9 mL was the maximum
volume needed, why was 45 mL
used?
6
How was it shown that complete
precipitation had taken place?
7
Why was the precipitate washed
in hot water?
8
Define the term ‘weighed to
constant mass’.
9
Use the final mass of precipitate
to calculate the % m/m of
ammonium sulfate in the fertiliser.
10
Where could errors have occurred
in this experiment?
Answer
Page 2
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Worksheet 2.3
Gravimetric analysis problems
NAME:
CLASS:
INTRODUCTION
Analysis by precipitation sounds very simple, but, in practice, obtaining accurate results by this
method requires planning and careful experimentation. The method involves considering such
variables as the degree of solubility, the size of the precipitate particles formed and the possibility
that other ions present might interfere.
Gravimetric analysis involves precipitating one of the ions of interest and weighing the precipitate.
The amount of precipitate tells us the amount of the ion of interest that must have been present in
the original solution. Sometimes the positive ion is precipitated, sometimes the negative ion.
No.
Question
1
Gravimetric analysis cannot be
used for solutions of sodium
nitrate. Why is this?
2
Silver nitrate is added to a
solution containing a mixture of
sodium chloride and potassium
chloride.
a Write a balanced equation for
the reaction that occurs
between:
i silver nitrate and sodium
chloride
ii silver nitrate and potassium
chloride.
b Explain why this procedure can
be used for an analysis of the
chloride ion concentration, but
not the potassium ion
concentration.
3
Most precipitates are not 100%
insoluble. What is the main
precaution taken during a
gravimetric procedure to ensure
that this does not affect the
accuracy too much?
Answer
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Worksheet 2.3
Gravimetric analysis problems
No.
Question
4
Use a solubility table to suggest a
compound that could be added to
each of the following solutions to
produce a precipitate.
a Na2CO3
b MgI2
c NaOH
5
The sulfate content of fertiliser
can be found by adding barium
nitrate, Ba(NO3)2, to precipitate
the sulfate as barium sulfate,
BaSO4.
a Write an ionic equation for the
reaction between barium nitrate
and sulfate ions.
b A student suggests that the
sulfate content could also be
found by boiling the water
from a solution of the fertiliser
and weighing the solid left
behind. Why does this
procedure not work?
6
When barium sulfate is collected
as a precipitate, care must be
taken to use a filter paper that has
very fine pores. Explain why this
is needed.
7
a Write a balanced equation for
the reaction between silver
nitrate and lithium chloride
solutions.
b A solution contains 0.02 mole
of lithium chloride. A student
wishing to precipitate this
lithium chloride adds 25 mL of
2 M silver nitrate solution. Has
the student added an excess of
silver nitrate?
Answer
Page 2
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Worksheet 2.3
Gravimetric analysis problems
No.
Question
8
An organic compound has a
formula C8H15Cl3. The purity of a
sample of this compound is to be
checked by the addition of silver
nitrate to precipitate the chlorine
atoms as silver chloride. It is
difficult to write a complete,
balanced equation for this reaction
but fortunately only a partial
equation is required.
a Write a partial equation for this
process.
b A 2.0 g sample leads to 0.478 g
of precipitate. Calculate the %
purity by mass of the organic
compound in the sample.
9
A precipitate of Fe2O3 is obtained
from an iron(III) chloride, FeCl3,
solution. The mass of precipitate
obtained is 0.644 g.
a Write a partial equation for this
reaction.
b If the mass of sample used was
2.0 g, determine the percentage
by mass of iron(III) chloride in
the sample.
10
As households use more grey
water on their gardens, the
phosphorus content of the water
becomes more relevant.
Phosphorus can be precipitated as
Mg2P2O7. During an analysis of a
water sample, the Mg2P2O7
precipitate is found to weigh
0.744 g. What will be the effect on
the calculated phosphorus content
if:
a the precipitate is not
completely dry?
b some of the precipitate passed
through the filter paper?
Answer
Page 3
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Worksheet 3.1
Revision of acids and bases
NAME:
CLASS:
INTRODUCTION
These reactions are typical of acids:
acid + metal → salt + hydrogen gas (a redox reaction)
acid + metal hydroxide → salt + water
acid + metal oxide → salt + water
acid + metal carbonate → salt + water + carbon dioxide
Acids are proton donors; bases are proton acceptors. Acid–base reactions involve the transfer of
protons. Acidic species may be described as monoprotic (donate one proton), diprotic (donate two
protons), triprotic (donate three protons) or amphiprotic (both donate and accept protons).
Acid–base strength is a measure of how readily protons are donated or accepted.
The pH scale is used as a measure of the acidity or basicity of a solution. The scale is usually
applied over the range 0 to 14 (but does extend beyond these values). pH = –log10[H3O+]. pH is a
logarithmic scale, so a difference of one unit on the pH scale means a 10-fold difference in the
hydrogen ion concentration. For dilute solutions at 25°C, Kw = [H3O+] × [OH–] = 10–14.
No.
Question
1
Write balanced chemical
equations for the reactions
occurring when the following
chemicals are mixed.
a Solid potassium hydrogen
carbonate and sulfuric acid
b Iron(III) oxide and nitric acid
c Calcium hydroxide solution
and hydrochloric acid
2
Write balanced equations to
illustrate the following reactions.
a Dissociation of Ca(OH)2 in
aqueous solution
b Successive ionisations of
H2SO4
c Neutralisation of Ba(OH)2
solution with HCl solution
Answer
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Worksheet 3.1
Revision of acids and bases
No.
Question
3
Write the formulas of:
a the conjugate bases of HS– and
OH–
b the conjugate acids of HSO4–
and H2PO4–.
4
For each of the following identify
the Lowry–Brønsted conjugate
acid–base pairs involved in the
reaction.
a NH3(g) + HCl(g) → NH4Cl(s)
b K2O(s) + H2O(l) → 2KOH(aq)
c Na2CO3(s) + H2SO4(aq)
→ Na2SO4(aq) + H2O(l) + CO2(g)
5
HSO4– is an amphiprotic ion.
Write chemical equations to show
this ion acting as:
a an acid
b a base.
6
Give two reasons why two acid
solutions of equal concentration
could have different pH values.
7
Give an explanation for each of
the following observations.
a The electrical conductivity of a
1.0 M solution of methanoic
acid (HCOOH) is less than that
of a 1.0 M solution of
hydrochloric acid (HCl).
b Ethanoic acid (CH3COOH) is
monoprotic, even though it
contains four hydrogen atoms.
8
List the following 1.0 M solutions
in order of decreasing pH. Give
reasons for your order.
NaOH, H2O, NH3, CH3COOH,
H2SO4, HNO3
Answer
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Worksheet 3.1
Revision of acids and bases
No.
Question
9
Calculate the pH of each solution.
a 0.050 M HNO3
b 0.50 M Ba(OH)2
10
20.0 mL of a solution of pH 3.0 is
diluted to produce a total volume
of 200.0 mL. What is the pH of
the resulting solution?
11
What volume of water must be
added to 50.0 mL of a
hydrochloric acid solution of pH
2.0 to increase the pH to 2.5?
12
Calculate the pH of a solution
formed when 20.00 mL of
0.00100 M HCl is mixed with
20.00 mL of 0.00100 M Ba(OH)2.
Answer
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Worksheet 3.2
Acid–base titrations
NAME:
CLASS:
INTRODUCTION
Probably the most used analytical procedure of old is the acid–base titration. The calculations
involved use the same stoichiometry procedures encountered earlier in your studies. There are
certain practical steps necessary, however, to ensure accurate results. These involve the choice of
indicators, choice of liquids for rinsing glassware and the correct preparation of standard solutions.
No.
Question
1
Explain the steps required to
prepare 250 mL of a standard
solution of 0.0500 M sodium
carbonate using anhydrous
Na2CO3.
2
20 mL of 0.10 M nitric acid,
HNO3, sits in a flask. Several
drops of indicator are added to the
flask. When sodium hydroxide
solution, NaOH, of an unknown
concentration is added dropwise
into the flask, there is a colour
change after 5 drops of sodium
hydroxide has been added. What
conclusion can you draw about the
concentration of the sodium
hydroxide?
3
20 mL of 0.1 M nitric acid sits in
a flask. Indicator is added. When
sodium hydroxide is added, a
colour change occurs after exactly
5 mL of sodium hydroxide has
been added. Without calculating
any mole quantities, what must
the concentration of the sodium
hydroxide be?
Can you use this ratio technique to
solve all titration calculations?
Answer
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Worksheet 3.2
Acid–base titrations
No.
Question
4
a Write a balanced equation for
the reaction between solutions
of sodium carbonate, Na2CO3,
and nitric acid.
b What volume of 0.20 M
sodium carbonate is required to
neutralise 40 mL of 0.10 M
nitric acid?
5
a Write a balanced equation for
the reaction between solutions
of ammonia, NH3, and
hydrochloric acid, HCl.
b 0.1 M ammonia in a flask is to
be titrated with 0.1 M
hydrochloric acid. Ammonia is
a weak base. The approximate
pH of the ammonia solution
before the titration begins is 11.
i What do you think the pH of
the solution in the flask will
be at the equivalence point?
(Look at the products of the
reaction.)
ii After significant extra
hydrochloric acid has been
added, what will the
approximate pH be?
iii Sketch the pH curve for this
titration.
6
A 25.0 mL sample of ethanoic
acid is diluted to 100.0 mL. A
20.00 mL aliquot is then titrated
with 0.114 M sodium hydroxide.
The titre required is 15.45 mL.
Calculate the concentration of the
original ethanoic acid solution.
Answer
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Worksheet 3.2
Acid–base titrations
No.
Question
7
Suggest two reasons why an
indirect (back) titration might be
performed, rather than a direct
titration.
8
a Draw a pH curve for the
addition of 0.1 M ethanoic
acid, CH3COOH, to 0.1 M
sodium hydroxide, NaOH.
b What will be the approximate
pH at the equivalence point?
c Use your curve to explain why
an indicator such as methyl red
would be a poor choice for this
titration.
9
10.0 mL of hydrochloric acid is
added to a 250 mL volumetric
flask. The flask is made up to the
mark and 20.0 mL aliquots of this
solution are added to conical
flasks for titration with 0.100 M
sodium carbonate from a burette.
a State which liquid should be
used to rinse each of the
following items.
i Volumetric flask
ii Pipette
iii Burette
iv Conical flasks
b The average titre of sodium
carbonate is 18.3 mL.
Calculate the concentration of
the original HCl used.
Answer
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Worksheet 3.2
Acid–base titrations
No.
Question
10
A 0.10 M magnesium hydroxide,
Mg(OH)2, solution is used to find
the concentration of a solution of
nitric acid. An aliquot of 25.0 mL
of magnesium hydroxide is added
to a flask. The volume of nitric
acid required to neutralise it is
13.5 mL.
Calculate the concentration of the
nitric acid solution.
Answer
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Worksheet 3.3
A back titration
NAME:
CLASS:
INTRODUCTION
Marble is a metamorphic rock that is almost pure calcium carbonate. The following experiment was
conducted to determine the calcium carbonate content of a marble sample. An accurately weighed
sample of crushed marble was added to a measured volume of recently standardised HCl solution.
The solution was heated to drive off the evolved carbon dioxide. The remaining solution was
titrated with a recently standardised NaOH solution, using a methyl red indicator. The results
obtained are shown below.
Mass of marble sample: 1.740 g
Volume of HCl solution added: 40.00 mL
Concentration of the standardised HCl solution: 1.020 M
Concentration of the standardised NaOH solution: 0.275 M
Average titre of NaOH: 25.56 mL
No.
Question
1
Write an equation for the reaction
of HCl with:
a CaCO3(s)
b NaOH(aq)
2
Calculate:
a the amount (in mol) of HCl
added initially
b the amount (in mol) of NaOH
used in the titration
c the amount (in mol) of
unreacted HCl
d the amount (in mol) of HCl
reacting with the CaCO3
e the amount (in mol) of CaCO3
in the marble sample
f the mass (in g) of CaCO3 in the
marble sample
g the percentage by mass of
CaCO3 in the marble sample.
Answer
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Worksheet 3.3
A back titration
No.
Question
3
The sodium hydroxide solution
used was recently standardised.
Why is sodium hydroxide
unsuitable as a primary standard?
4
Why was it necessary to drive off
the evolved carbon dioxide before
performing the titration? How
would the result be affected if this
step was omitted?
5
Why was it necessary to use a
back titration for this analysis,
rather than a direct titration of the
marble with HCl solution?
6
How would each of the following
errors, if made during the analysis,
alter the calculated value for the
percentage CaCO3?
a The 40.0 mL pipette used to
deliver the HCl was rinsed only
with water prior to its use.
b The burette was rinsed only
with water prior to its use.
c The volumetric flask was
rinsed only with water prior to
its use.
7
An alternative method of analysis
involves reacting the crushed
marble with excess HCl and
collecting the evolved carbon
dioxide. In one such experiment,
95.0 mL of gas was collected at a
pressure of 765 mmHg at 23°C
when 0.411 g of marble was
reacted. Determine the percentage
by mass of CaCO3 in the marble
sample, based on this data.
8
Suggest reasons why the value
obtained using this method is
smaller than that obtained using
the back titration.
Answer
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Worksheet 4.1
Oxidation numbers and redox
equations
NAME:
CLASS:
INTRODUCTION
The concept of oxidation numbers (ON) or oxidation states was applied to determine whether or not
electrons had moved from one species to another in a chemical reaction. An oxidation–reduction
(redox) reaction is one in which one or more atoms change oxidation numbers. Oxidation occurs
when an atom’s oxidation state becomes more positive, indicating that electrons have been lost.
Reduction occurs when an atom’s oxidation state becomes less positive, indicating that electrons
have been gained.
OXIDATION
…–5 –4 –3 –2 –1 0 +1 +2 +3 + 4 +5….
REDUCTION
The oxidation numbers assigned to atoms in covalent compounds are hypothetical charges, and the
atoms do not really have these charges within the compound, since they are only sharing electrons.
For example, in the compound CO2 we say that carbon has a +4 ON and oxygen is in a –2 ON, but
the atoms do not really have these charges as electrons are not transferred in covalent compounds.
Recall that the oxidant is the reactant being reduced, so its ON will decrease. The reductant is the
reactant being oxidised, so its ON will increase.
Rules for assigning oxidation numbers
1
The oxidation number of an element is zero.
2
For a monatomic ion, the oxidation number is the charge on the ion.
3
The oxidation number of combined hydrogen is usually +1.
4
The oxidation number of combined oxygen is usually –2.
5
The sum of all oxidation numbers of atoms in a compound is zero.
6
The sum of all oxidation numbers of atoms in an ion is equal to the charge on that ion.
No.
Question
1
Write the formula of a substance in
which nitrogen has the following
oxidation numbers:
a +3
b +5
c 0
d –3
Answer
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Worksheet 4.1
Oxidation numbers and redox
equations
No.
Question
2
What is the oxidation number of the
atom in bold type in each of the
following?
a H2SO4
b K2Cr2O7
c CO2
d H2O2
3
For the following equations,
determine which are redox
processes. For those that are redox
reactions, identify the oxidant and
the reductant.
a 2OH–(aq) + Cr2O72– (aq)
→ 2CrO42–(aq) + H2O(l)
b I2O5(s) + 5CO(g)
→ I2(s) + 5CO2(g)
c PBr3(l) + 3H2O(l)
→H3PO3(aq) + 3HBr(aq)
d 2Hg2+(aq) + N2H4(aq)
→ 2Hg(l) + N2(g) + 4H+(aq)
e 3H2S(g) + 2H+(aq) + 2NO3–(aq)
→ 3S(s) + 2NO(g) + 4H2O(l)
f 3NO2(g) + H2O(l)
→ 2HNO3(aq) + NO(g)
4
The highest positive oxidation
number possible for any atom is
equal to the number of electrons in
its valence shell. For example,
nitrogen has a maximum oxidation
number of +5. What is the maximum
oxidation number of:
a oxygen?
b magnesium?
c chlorine?
5
Occasionally you will find that an
atom in a compound has an
oxidation number of zero. What is
the oxidation number of each atom
in glucose, C6H12O6?
Answer
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Worksheet 4.1
Oxidation numbers and redox
equations
Rules for balancing redox half-equations in acidic media
1
Write the species undergoing oxidation or reduction on the left and its conjugate on the right,
leaving out any spectator ions.
2
Balance for atoms other than hydrogen and oxygen.
3
Balance for oxygen by adding water, H2O.
4
Balance for hydrogen ions, H+.
5
Balance for charge by adding electrons to the side with the greater positive charge.
No.
Question
6
For each of the following, write
the two half-equations, then add
them together to get the overall
redox equation. Make sure you
leave out any spectator ions.
a Acidified potassium
dichromate solution (K2Cr2O7)
is used to oxidise ethanol to
ethanoic acid. The dichromate
ion is converted to
chromium(III) ions.
b The bromate ion, BrO3– may be
used in acidic solution to
oxidise iodide ions to iodine.
The reduction product is the
bromide ion.
c Purple acidified potassium
permanganate solution
(KMnO4) is decolourised by
acidified iron(II) sulfate
solution, producing Mn2+ and
Fe3+.
d Hydrogen peroxide (H2O2) is
added to acidified potassium
permanganate and oxygen is
evolved.
e Hydrogen sulfide (H2S) gas is
bubbled into a solution of
acidified potassium
dichromate, producing a
deposit of sulfur.
Answer
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Worksheet 4.2
Recovering silver from
solution
NAME:
CLASS:
INTRODUCTION
Given the high price of silver, it is desirable to have an efficient method for the recovery of ‘waste’
silver from aqueous solutions containing Ag+ ions. This recovery can be achieved by using a redox
reaction in which Ag+(aq) is reduced to Ag(s). This worksheet concerns the analysis and extraction
of silver from a 0.1 M solution, and involves both gravimetric and volumetric techniques.
PART 1: VOLUMETRIC DETERMINATION OF SILVER ION CONCENTRATION
A 20.00 mL aliquot of 0.100 M NaCl solution was pipetted into a 250 mL conical flask.
Approximately 1 mL of 0.1 M K2CrO4 solution was also added to the flask. The flask contents
were then titrated with a solution of approximately 0.1 M AgNO3 until the first permanent redbrown colour appeared. This red-brown colour indicated that the reaction between Ag+(aq) and Cl–
(aq) was complete (the CrO4– ion reacts with the Ag+ ion to form a red-brown precipitate. This
reaction occurs only when there is no Cl– ion present). The titration was completed to obtain
concordant titres; an average titre of 19.32 mL was required.
No.
Question
1
Name the liquid that should be
used for the final rinse of each of
the following pieces of equipment
prior to their use in the titration.
a Burette
b Pipette
c Conical flask
2
Write an ionic equation for the
precipitation reaction between:
a Ag+ and Cl– ions
b Ag+ and CrO4– ions.
3
Using the titration results,
determine the concentration of the
approximately 0.1 M AgNO3
solution.
4
Suggest two sources of error in
this determination.
Answer
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Worksheet 4.2
Recovering silver from
solution
PART 2: RECOVERY OF SILVER FROM SOLUTION
A 20.00 mL aliquot of the approximately 0.1 M AgNO3 solution was pipetted into a large test tube.
A spiral of copper wire was inserted into the test tube so that most of the spiral was in the solution.
The solution was stirred occasionally using the copper spiral. The test tube and spiral were allowed
to stand overnight. The spiral was removed after shaking to dislodge deposited silver, and any
remaining silver washed into the test tube with distilled water. The deposited silver was collected,
dried and weighed. The mass of silver recovered was 0.206 g.
No.
Question
Answer
5
How could you ensure that the recovered
silver was dry?
6
a Write an ionic equation for the
reaction between Cu and Ag+ ions.
b Name the oxidant in this reaction.
7
Determine the percentage of silver
recovered from the solution.
8
Suggest two reasons to account for the
less than 100% recovery of silver from
the solution.
9
Most methods for the recovery of silver
from solution involve the reduction of
Ag+(aq) to Ag(s). Another reductant used
is the dithionite ion, S2O42–, which is
oxidised to the sulfite ion, SO32–. Write a
half-equation for the oxidation of the
dithionite ion.
10
In an alternative process, an electric
current is forced through a solution
containing silver ions. Silver deposits at
the negative electrode. At the positive
electrode, water is oxidised to oxygen gas
and hydrogen ions. Write a half-equation
for this oxidation.
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Worksheet 4.3
Redox titrations
NAME:
CLASS:
INTRODUCTION
Titrations are not restricted to acid–base reactions. They can also be applied to redox reactions. The
principles of rinsing and careful searching for an endpoint apply to redox titrations, but instead of
acid–base indicators, the colour change that occurs when elements change oxidation state is often
used.
No.
Question
Answer
1
The presence of iodine, I2, can be
detected by the use of starch as an
indicator. The starch forms a blue colour
when iodine is present.
a Write a balanced half-equation for the
reduction of iodine to iodide ions.
b Explain how the starch acts as an
indicator in reactions involving iodine.
2
The half-equation for the reaction of
vitamin C, ascorbic acid, is shown.
C6H4O2(OH)4(aq)
→ C6H4O4(OH)2(aq) + 2e– + 2H+(aq)
a Write a balanced overall equation for
the reaction between ascorbic acid and
iodine.
b 25.0 mL of 0.10 M iodine reacts with
17.8 mL of ascorbic acid. Calculate the
concentration of the ascorbic acid.
3
The now obsolete puff bag used the
reaction of potassium dichromate,
K2Cr2O7, and ethanol.
a Write balanced half-equations for the
oxidation and reduction reactions.
b Write a balanced overall equation for
this reaction.
c A 25.0 mL sample of whiskey is
diluted to 250.0 mL and 20.0 mL
aliquots of the diluted whiskey are
titrated against 0.200 M potassium
dichromate. The average titre required
is 19.45 mL. Calculate the ethanol
concentration (in M) of the whiskey.
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Worksheet 4.3
Redox titrations
No.
Question
Answer
4
The concentration of an iron solution,
Fe2+, can be found by titration with
permanganate, MnO4–, ions. The Fe2+ is
oxidised to Fe3+ and the purple
permanganate is reduced to colourless
Mn2+ in acid conditions.
a Write balanced half-equations for the
oxidation and reduction reactions.
b Write an overall equation for this
reaction.
c How would you judge that the
endpoint has been reached in this
reaction?
d 20.0 mL aliquots of an iron solution
are titrated with 0.110 M KMnO4. The
average titre required is 12.8 mL.
Calculate the concentration of the iron
solution.
5
A rusty nail with a mass of 2.87 g is
added to a beaker containing 100.0 mL of
1.08 M hydrochloric acid. After the
reaction stops, the excess acid is
neutralised by the addition of 1.00 M
sodium hydroxide. The volume of sodium
hydroxide required is 45.9 mL.
a Write a balanced equation for the
reaction between hydrochloric acid
and iron.
b Write balanced half-equations to show
that this is a redox reaction.
c Calculate the % by mass of iron in the
nail.
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Worksheet 5.1
Analysis of iron in iron ore
NAME:
CLASS:
INTRODUCTION
Iron ores often contain a mixture of oxides and contain both Fe2+ and Fe3+ ions. The iron content of
these ores may be determined by a variety of analytical techniques. This worksheet concerns three
of these techniques.
PART 1: SPECTROSCOPIC ANALYSIS
The iron content of the ore sample was analysed using UV-visible spectrophotometry. The iron(III)
ion, Fe3+, reacts with the thiocyanate ion, SCN–, to form a complex ion with an intense red colour.
This complex ion may be detected in a spectrophotometer set at a wavelength of about 580 nm. A
0.100 g sample of the ore was dissolved in concentrated hydrochloric acid. The extract was filtered,
treated to ensure all the iron present was converted to Fe3+, then 20.0 mL of potassium thiocyanate
solution was added. The volume was then made up to 100.0 mL with deionised water. Four
standard solutions of iron(III) were similarly treated and their absorbances measured to generate the
calibration curve shown below.
No.
Question
1
How would it have been
determined that 580 nm was an
appropriate wavelength for this
analysis?
2
Why was it necessary to construct
a calibration curve for this
determination?
Answer
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Worksheet 5.1
Analysis of iron in iron ore
No.
Question
3
The absorbance of the iron ore
solution was found to be 0.400.
Calculate the percentage by mass
of iron in the ore sample.
Answer
PART 2: VOLUMETRIC ANALYSIS
In a second experiment the iron content of the ore was determined volumetrically. A 0.268 g sample
of ore was dissolved in acidic solution, filtered, and the filtrate treated to convert all the iron present
to Fe2+. This solution was titrated with a standardised, acidified 0.0335 M potassium permanganate
(KMnO4) solution. The titration required 19.75 mL of the permanganate solution to reach the light
pink-purple endpoint.
No.
Question
4
Write half-equations and a
balanced redox equation for the
titration reaction, given that the
products of the titration reaction
include Fe2+(aq) and Mn2+(aq).
5
Calculate the percentage by mass
of iron in the ore sample, based on
this volumetric analysis.
6
If some of the iron was present in
the solution as Fe3+ prior to the
titration with permanganate
solution, how would this have
affected the value determined for
percentage of iron in the ore?
Answer
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Worksheet 5.1
Analysis of iron in iron ore
PART 3: GRAVIMETRIC ANALYSIS
In a third experiment the iron content of the ore was determined gravimetrically. A 1.01 g sample of
the ore was dissolved in concentrated hydrochloric acid. The extract was filtered to remove any
insoluble material. Excess sodium hydroxide solution was added, and the precipitate collected and
heated to convert it to solid iron(III) oxide. 1.08 g of Fe2O3 was obtained.
No.
Question
7
Write an ionic equation for the
precipitation reaction between the
following:
a Iron(II) ion and hydroxide ion
b Iron(III) ion and hydroxide ion
8
Suggest why the iron hydroxide
precipitate was not simply
collected, dried and weighed in
this determination.
9
Calculate the percentage by mass
of iron in the ore sample, based on
this gravimetric analysis.
10
Suggest a possible reason why the
value obtained using gravimetric
analysis is significantly larger
than the other determined values.
Answer
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Worksheet 5.1
Analysis of iron in iron ore
CONCLUSION
No.
Question
11
Which method of analysis used
(spectroscopic, volumetric or
gravimetric):
a is the most expensive?
b is the most prone to error?
c requires the least specialised
equipment to perform?
d is likely to be the most
accurate?
12
Suggest another method of
analysis of the iron content of the
iron ore.
Answer
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Worksheet 5.2
Colorimetric determination of
manganese in steel
NAME:
CLASS:
INTRODUCTION
A 2.890 g sample of steel was analysed for its manganese content in the following manner.
The sample was dissolved in concentrated nitric acid, producing manganese(II) ions.
Potassium persulfate, K2S2O8, was added to remove the carbon in the steel. This oxidised the carbon
to carbon dioxide.
The virtually colourless manganese(II) ions were converted to deep purple permanganate (MnO4–)
ions by boiling with a solution containing periodate (IO4–) ions. In this reaction the periodate ions
are converted to IO3– ions.
Phosphoric acid was added in order to convert the yellow iron(III) ions to a colourless complex (so
the iron ions would not interfere with the colorimetric MnO4– ion analysis).
The solution was diluted with distilled water to 1.00 L in a volumetric flask.
A series of standards was prepared, containing MnO4– ranging from 1.00 to 8.00 ppm.
The colorimeter was set to measure absorbance at 520 nm, and calibrated using a solvent blank. The
absorbances of the standard solutions and the steel sample were taken and are given in the table
below.
Concentration of MnO4– (ppm)
0.00
1.00
2.00
3.00
5.00
6.00
8.00
Steel sample
Absorbance
0.000
0.076
0.151
0.227
0.381
0.455
0.610
0.195
No
Question
Answer
1
a Write a half-equation for the
conversion of Mn2+ to MnO4–
in acidic aqueous solution.
b Write a half-equation for the
conversion of IO4– to IO3– in
acidic aqueous solution.
c Add the two half-equations
together to obtain the overall
equation for step 3.
2
Plot a graph of concentration of MnO4– (x-axis) against absorbance (y-axis). Draw a straight
line of best fit through the points.
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Worksheet 5.2
Colorimetric determination of
manganese in steel
No
Question
3
Determine the concentration of
MnO4– in the diluted steel
solution.
4
Using your answer to question 3,
find the % m/m of manganese in
the steel sample.
5
After step 1, the solution
contained a high concentration of
iron(III) ions. Where did they
come from?
6
If the standards were prepared
using anhydrous potassium
permanganate, calculate the mass
of solute needed to prepare
100.0 mL of an 8.0 ppm MnO4–
solution.
7
As well as water, what else
should be added to the standards
and the solvent blank and why?
8
Why was 520 nm a more
appropriate wavelength than
400 nm for this analysis?
9
During calibration, the
absorbance was set to zero using
a solvent blank. Would your
calculated percentage of Mn
(m/m) have been affected had
this not been done? Explain.
10
Suggest an alternative method for
the determination of the
manganese content of the steel
sample.
Answer
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Worksheet 5.3
Analysing mass spectra
NAME:
CLASS:
INTRODUCTION
As early as the 1920s, mass spectrometry was used to identify the isotopes of elements, and their
relative isotopic masses (RIM) and abundances. Armed with this information, accurate relative
atomic masses (Ar) for elements could be determined. A flowchart of the steps in the operation of
the mass spectrometer is shown below.
No.
Question
Answer
1
Name the processes occurring in
stages 1, 2 and 4.
2
Describe the processes occurring in:
a stage 2
b stage 4.
3
The mass spectrum data obtained for a sample containing a mixture of two monatomic gases
is shown below. Label each peak on the mass spectrum using the ZA X n notation. You can
refer to the periodic table.
Page 1
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Worksheet 5.3
Analysing mass spectra
No.
Question
4
Element X has two naturally
occurring isotopes with relative
isotopic masses of 120.90 and
122.90. If the relative atomic mass
of the element is 121.75, determine
the percentage abundance of the
lighter isotope.
Answer
The modern mass spectrometer is a powerful and accurate tool that is used extensively for the
detection and identification of a wide range of compounds, rather than for RIM determinations.
When a molecular substance is placed in the mass spectrometer, the mass spectrum shows a peak
corresponding to the molecular ion, allowing the molecular mass to be determined. The bombarding
electrons in the mass spectrometer also cause molecules to fragment, giving rise to a fragmentation
pattern of lower molecular mass ions. The masses of the units broken off the molecule often give
clues as to the structure of the molecule.
No.
Question
Answer
5
The molecular structure and simplified mass spectrum of 2-pentanol are shown below. Label
the peaks with the probable formula for the ions producing peaks at the following mass
numbers.
a 88
b 87
c 73
d 45
Page 2
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Worksheet 5.3
Analysing mass spectra
No.
Question
Answer
6
A simplified mass spectrum and the molecular structures of two possible compounds
forming the spectrum are shown below. Based on an inspection of the mass spectrum, which
compound is most likely to have produced the spectrum shown? Explain your choice.
7
Analysis of fragments in the mass
spectrum gives clues to the
structure of an unknown
substance. Explain how a mass
spectrum may be used to clearly
identify an unknown substance.
Page 3
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Worksheet 5.4
Spectroscopic analysis of
organic compounds I
NAME:
CLASS:
INTRODUCTION
Spectroscopy is used extensively in the analysis of organic compounds. Infrared spectroscopy is
used to identify particular functional groups in compounds. Mass spectroscopy is used to determine
the relative molecular mass of compounds, and is widely used to identify the structure of molecules.
The fragmentation pattern obtained when molecules are broken is used to help determine their
molecular structure. Nuclear magnetic resonance (NMR) spectroscopy is used for determining the
precise structure of organic compounds. Absorption of radiation by 1H nuclei in different chemical
environments provides information about the arrangement of hydrogen atoms within the compound.
The questions that follow provide practice in the interpretation of spectroscopic data used for the
identification of selected organic compounds. (You will learn more about these compounds and
their functional groups in chapter 7 of the coursebook.)
No.
Question
Answer
1
An unknown compound is thought to be either
pentanal or 2,2-dimethylpropanal. The
structures of these two compounds are shown
below.
NMR spectroscopic analysis is used to help
identify the compound. The 1H NMR spectrum
for each compound is produced.
a How many peaks would you expect to see
on the 1H spectrum of:
i pentanal?
ii 2,2-dimethylpropanal?
b What is the expected ratio of areas for the
peaks on the 1H spectrum of:
i pentanal?
ii 2,2-dimethylpropanal?
c Given that the 1H NMR spectrum shows two
peaks with the integration ratio 155:17.3,
what is the identity of the unknown?
d How could infrared spectroscopy be used to
confirm the identity of the unknown?
Page 1
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Worksheet 5.4
Spectroscopic analysis of
organic compounds I
No.
Question
Answer
2
Two possible structures for an unknown
organic compound are shown below. Both have
the molecular formula C3H6O2.
The correct structure is to be determined using
spectroscopic analysis.
a The infrared spectrum of the unknown
compound shows a distinct absorption peak
at 1710 cm–1, but no distinct peak in the
2500–3000 cm–1 range.
i Briefly explain why the compound
absorbs radiation in the infrared section
of the electromagnetic spectrum.
ii On the basis of the infrared data
provided, which of the two structures is
likely to be correct? Explain your choice.
b The 1H nuclear magnetic resonance
spectrum of the compound shows two peaks,
one at a chemical shift of 1.9 ppm, the other
at 3.8 ppm. The peak areas are in the ratio
1:1.
i Which type of electromagnetic radiation
is absorbed by atoms in molecules to
produce an NMR spectrum?
ii Based on the NMR spectrum, which of
the two structures is most likely to be
correct? Explain your choice.
In more detailed NMR analysis, more information can be obtained. In high resolution NMR,
absorption peaks are split into closely spaced peaks, called doublets, triplets etc. This splitting is the
result of the influence of neighbouring nuclei on the nucleus causing the peak. Neighbouring nuclei
are also tiny bar magnets and so they influence the effect of the applied magnetic field on the
nucleus under investigation. Valuable information about these neighbouring groups can be obtained
by analysing the splitting patterns of each signal in the NMR spectrum.
Page 2
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Worksheet 5.4
Spectroscopic analysis of
organic compounds I
No.
Question
Answer
3
An unknown compound is thought to be either
compound I or II (molecular structures are
shown below). The 1H NMR spectrum of the
unknown compound shows two peaks with
integration ratio 2:3. Which compound, I or II,
is the unknown?
4
The high resolution 1H NMR spectrum of the
unknown compound in question 3 is shown
below. The numbers 2 and 3 represent relative
peak areas. What is the relationship between
the number of splits in a peak and the number
of 1H nuclei adjacent to the 1H nucleus causing
the peak?
Page 3
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Worksheet 5.4
Spectroscopic analysis of
organic compounds I
No.
Question
Answer
5
An unknown compound is thought to be either
compound I or II, (molecular structures are
shown). The high resolution 1H NMR spectrum
of the unknown compound is also shown.
a How many peaks would be expected for
i compound I?
ii compound II?
b What is the expected ratio of areas for the
peaks for
i compound I?
ii compound II?
c What is the expected splitting of the peaks
for
i compound I?
ii compound II?
d What other information from the NMR
spectrum could be used to distinguish
between compounds I and II?
Page 4
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