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EL1: Where do the chemical elements come from?
Learning objectives:
 Describe protons, neutrons and electrons in terms of their mass and relative charge;
 Describe the structure of atoms in terms of electrons and a central nucleus containing
protons and neutrons;
 Understand that knowledge of the structure of the atom developed in terms of a
succession of gradually more sophisticated models, and given information, interpret
these and other examples of such developing models [Activity EL 1.1];
 Explain and use the terms atomic number, mass number, isotope, Avogadro constant,
relative isotopic mass, relative atomic mass, relative formula mass and relative
molecular mass;
 Describe the electron structure of atoms in terms of main energy levels (electron shells)
up to Z = 36;
 Recall that in fusion reactions lighter atoms join to give heavier atoms (under conditions
of high temperature and pressure) and understand that this is how certain elements are
formed [Activity EL 1.2];
 Explain the occurrence of absorption and emission atomic spectra in terms of changes in
electronic energy levels, and compare and contrast the features of these spectra;
[Activity EL 1.3];
 Understand the relationship between the energy emitted or absorbed and the
frequency of the line produced in the spectra, ΔE = hν;
 Recall that the nuclei of some atoms are unstable and that these atoms are radioactive;
 Recall and explain the different properties of α, β and γ radiations;
 Recall that the term half-life refers to the time taken for half the radioactive nuclei in a
sample to decay, and that the half-life is fixed for any given isotope;
 Carry out half-life calculations [Activity EL 1.4];
 Use nuclear symbols to write equations for nuclear processes, including fusion and
radioactive decay;
 Understand how radioactive isotopes can be used as ‘tracers’ in the body and (given
information) for other uses;
 Explain that the half-life of ‘tracers’ must be of an appropriate length to allow detection
but not cause undue damage;
 Understand the use of radioisotopes in the dating of archaeological and geological
material;
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Key definitions:
Compile a glossary by writing your own definitions for the following key terms related to the
learning objectives above.
Key term
proton
neutron
electron
nucleus
atomic number (Z)
mass number (A)
isotope
Avogadro constant
relative isotopic mass
relative atomic mass
relative formula mass
relative molecular mass
fusion
radioactive decay
emissions
Definition
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Key term
lightweight star
red giant
white dwarf
spectroscopy
electromagnetic
spectrum
photosphere
chromosphere
corona
absorption spectrum
absorption lines
excited states
emission spectra
Definition
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Chemical Ideas 2.1
A simple model of the atom
Sub-atomic particles
Atoms have a diameter of 0.2nm to 0.4nm (0.2 x 10-9m to 0.4 x 10-9m so they are too small to
see. They comprise:
 protons which are positively charged and found in the nucleus at the centre of the
atom;
 neutrons which have no charge but are also found in the nucleus;
 electrons which have a much lower mass than protons and neutrons;
Electrons have an equal but opposite (negative) charge to protons and they are arranged
around the nucleus.
particle
relative mass
charge/C
proton
1
+ 1.602 x 10-19
neutron
1
0
electron
1/1836
- 1.602 x 10-19
Nuclear symbols
In a nuclear symbol such as



A
Z
X:
A is the mass number (the number of protons added to the number of neutrons)
Z is the atomic number (the number of protons)
X is the elements chemical symbol
Atoms have equal numbers of protons and electrons so the atomic number also tells you how
many electrons the atom has.
mass number, A
7
3
Li
atomic number, Z
protons and neutrons
7
protons
3
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Development of the ‘model of the atom’
John Dalton is credited with developing the modern
theory of the atom. Following a long series of
experiments, Dalton suggested that:
 different elements have atoms which differ in mass;
 each element is characterised by the mass of its
atoms;
Four postulates describe how atoms behave:
 the atoms in a given element are all of the same kind;
 a compound contains atoms of two or more elements
combined together in fixed proportions;
 an atom retains its identity during a chemical
reaction;
 during a chemical reaction, the atoms in the reacting
substances rearrange to form the products of the
reaction;
Dalton also introduced the idea of chemical symbols. He used a system based on circular
systems. This system is no longer used. The current familiar system based on letters was
devised by the Swedish scientist, Jacob Berzelius in 1811.
The discovery of the electron
In 1897, J. J. Thomson deflected cathode rays with both electric and magnetic fields, and used
his results to measure the ratio of their charge to their mass. He reasoned that, if cathode rays
have mass, then they must be composed of a stream of particles. We now call these particles
electrons.
Thomson assumed
the charge to be
the same as the
smallest charge
observed during
electrolysis, and so
calculated the mass
of an electron to be
nearly 2000 times
smaller than the
mass of a hydrogen
atom.
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J. J. Thomson proposed a model that described atoms as negatively charged electrons
embedded in a sphere of positive charge. It became known as the ‘plum pudding’ model
because the electrons were spread randomly throughout the positive charge like the dried fruit
in a pudding. This model was soon to be shown to be inadequate.
The nucleus and protons
The Geiger-Marsden experiment involved measuring the deflection of alpha particles as they
struck a thin sheet of gold foil less than 1µm thick. Alpha particles were known to be positively
charged. Predictions based on Thomson’s model of the atom suggested that the alpha particle
would be deflected by a few degrees from the straight-ahead direction when they passed
through the gold foil.
Geiger and Marsden made two unexpected observations:
1. The vast majority of the alpha particles were deflected by less than one degree, showing
that they must have passed through essentially empty space.
2. A few (about 1 in 8000) actually bounced back towards the source. Rutherford remarked
that this was ‘almost as incredible as if you had fired a 15-inch shell at a piece of tissue
paper and it has bounced back and hit you’.
In 1911, Rutherford proposed that the atom has at its centre a very small positively charged
nucleus, which contains almost all the mass of the atom. This nucleus is tiny and the rest of the
atom is mostly empty space. (If you magnified an atom to the size of a football stadium, its
nucleus would be about the size of a pea.) So most of the alpha particles passed through empty
space. Only very rarely would one travel close enough to the nucleus to be repelled strongly by
the dense concentration of positive charge. Sufficient electrons surround the nucleus of an
atom to balance the charge of the nucleus and to make the atom neutral overall.
Between 1917 and 1921, Rutherford bombarded six different elements with alpha particles and
found that they all gave out the same positive particle, which was identical to the nucleus of the
hydrogen atom. As this was the first particle found in the nucleus, he called it the ‘proton’ (from
the Greek protos, meaning ‘first’). Rutherford concluded that protons made up the positive part
of the nuclei of all elements. The proton carries a positive charge of exactly the same
magnitude as the negative charge on the electron.
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The number of electrons in an atom equals the number of
protons in its nucleus. The number of protons is called the
atomic number (also called the proton number) of the element
concerned. In 1913, Henry Moseley found that when he
bombarded elements with high-speed electrons, they emitted
X-rays. He observed that the frequency of the emitted X-rays
depended on the element used. Moseley plotted a graph of
atomic number against the square root of the X-ray frequency.
The result was a straight-line plot.
The nucleus and neutrons
Experimental evidence for the neutron was found in 1932, when James Chadwick bombarded
the element beryllium with alpha particles. This bombardment produced highly penetrating
stream of particles, which could pass through many centimetres of solid lead and which was not
deflected by electric or magnetic fields. Chadwick decided that the stream must consist of
particles with almost the same mass as protons but with no charge. Chadwick had detected the
neutrons postulated earlier by Rutherford. Protons and neutrons are collectively known as
nucleons because they are both found in the nucleus.
Neutrons have almost the same mass as protons, so they contribute to the mass of an atom.
The mass number of an atom is defined as the sum of the numbers of protons and neutrons in
the atom’s nucleus e.g. the nucleus of a fluorine atom contains 9 protons and 10 neutrons, so
the atomic number of fluorine is 9 and its mass number is 19 (9 + 10).
Isotopes
Isotopes are atoms of the same element but with different numbers of neutrons. Their proton
numbers are identical but their mass numbers are different. For example: 11 H, 21 H and 31 H are all
isotopes of hydrogen. Each one has one proton and one electron but they have 0, 1 and 2
neutrons respectively.
Isotopes of an element are chemically identical but have different physical properties.
Isotopes are atoms of the same element
but with different numbers of neutrons.
35
17
Cl
35 - 17 = 18 neutrons
Chlorine-35 and chlorine-37 are
chemically identical because they each
have 17 electrons.
Their atomic number is the same but
their mass number is different.
37
17
Cl
37 - 17 = 20 neutrons
Isotopes of an element with more
neutrons have:
o higher mass
o higher density
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Chlorine-35 and chlorine-37 are chemically identical because they each have 17 electrons.
Isotopes of an element with more neutrons have:
higher mass
higher density
higher melting and boiling points
slower rates of diffusion
The relative atomic mass Ar of an element is defined as the mass of one atom of that element
relative to 1/12th the mass of one atom of carbon-12 i.e. 12C = exactly 12. The relative atomic
mass is the average of the masses of the stable isotopes of the element, weighted to take into
account the relative abundance of each isotope. For the example of chlorine given above, the
relative atomic mass is calculated as follows:
75.8
24.2
Ar (Cl) = (
𝑥 35) + (
𝑥 37) = 35.5
100
100
So the relative atomic mass of chlorine is 35.5.
SUMMARY
The atom has at its centre a very small, positively charged nucleus.
Almost all the mass of an atom is concentrated in the nucleus.
The nucleus of an atom contains protons.
The proton has a positive charge.
The charge on a proton is equal in magnitude, and opposite in sign, to the charge on an
electron.
The total positive charge on the nucleus equals the total negative charge of the electrons.
The atomic number of an element equals the number of protons in its nucleus.
The atomic number of an element distinguishes that element from all others.
The neutron has zero charge and almost the same mass as the proton.
Protons and neutrons are called nucleons because they are found in the nucleus of an atom.
The mass number is the sum of the numbers of protons and neutrons in the nucleus of an
atom.
Isotopes of an element have the same number of protons but different numbers of
neutrons; they therefore have different masses.
A nuclide is an isotope with a specified mass number.
The relative atomic mass of an element is the average mass of one atom of the element
relative to 1/12th the mass of one atom of carbon-12.
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Chemical Ideas 2.2
Nuclear reactions
The nuclei of some isotopes are unstable. They can break down or decay by giving off emissions
of particles or rays. This happens spontaneously and is not affected by chemical or physical
processes. Isotopes that decay like this are radioactive and the emissions are called radiation.
There are three main types of radiation:
α radiation which comprises particles identical to helium nuclei (two protons and two neutrons
joined together). When an isotope decays by α-decay its mass number decreases by 4 and its
proton number decreases by 2. A new element is formed.
β radiation which comprises high-energy electrons ejected from the nucleus. When an isotope
decays by β-decay its mass number stays the same but its proton number increases by 1. A new
element is formed.
γ radiation which comprises high-frequency electromagnetic radiation. γ-decay often
accompanies α- and β-decay. The nucleus loses energy but does not change into a different one
unless α- or β-decay is also present.
Nuclear fusion
In nuclear fusion reactions, nuclei fuse together to make the nucleus of a heavier element. The
positive nuclei would normally repel each other. The very high pressures and temperatures
found only in stars are needed to overcome this repulsion.
In the Sun, hydrogen nuclei collide with enough energy to overcome their repulsion and form
helium nuclei. Energy is released in these nuclear fusion reactions. Although the production of
each helium nucleus only releases 4 x 10-12J, the Sun makes 1038 helium nuclei per second. As a
result the Sun’s energy output is enormous. It has enough hydrogen for around 5 billion years
when it will expand and become a red giant. It will then use larger nuclei such as helium nuclei
as its fuel.
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1. Chemists have put forward various models to try to explain atomic
structures. Scientific models often change in the light of new evidence.
The following diagrams show possible structures for atoms based on
different models, A-D.
Complete the flow diagram by filling in the boxes on the right with the
letter corresponding to the appropriate model.
[3]
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2. The current model of the atom describes a nucleus containing protons and
neutrons surrounded by electrons in energy levels.
(a) Complete the following table showing the properties of the sub-atomic
particles.
[3]
(b) Explain why:
 mass numbers of atoms are always whole numbers;
 the relative atomic mass of an element may not be a whole number;
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [3]
3. In the Sun, 32 He is formed when two different isotopes of hydrogen join
together.
(i) What term is used to describe this process?
____________________________________
[1]
(ii) Write a nuclear equation for this process.
[2]
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4. ‘Radiopharmaceutical imaging’ is a type of medical scan. Tiny amounts of
radioactive materials are injected into the blood stream to allow an image
of the various tissues to be obtained. Radioactive materials contain atoms
that spontaneously break down. As the nuclei of atoms break down, they
emit rays and particles called emissions. Three different types of emission
have been identified: α, β and γ.
Complete the table to show the properties of these emissions.
[3]
5. The most widely used radioisotope in radiopharmaceutical imaging is
technetium-99. This is the product of the β-decay of molybdenum-99.
(i) Complete the nuclear equation for the β-decay of molybdenum-99.
[2]
(ii) Molybdenum has several naturally occurring isotopes.
Explain the term isotopes.
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
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6. Beryllium is a Group 2 metal and has several modern day uses. It also
played an important role in the development of ideas about the structure
of the atom.
In 1932, James Chadwick fired α-particles at beryllium metal and found that
particles were emitted that were not deflected by electric fields. Chadwick
had discovered the neutron. Give the mass number and atomic number of
this particle.
mass number = __________; atomic number = __________;
[2]
7. Information on how the composition of the earth’s atmosphere has
changed over thousands of years can be obtained by drilling 3 kilometre
deep ice cores in the Antarctic.
Significant changes in concentrations of lead-207 and lead-208 isotopes
over the last hundred years have been identified from the analysis of polar
ice cores.
Complete the table below to give the atomic structure of the lead-207
isotope.
[1]
8. The mass spectrum of a sample of material showed the abundance of the
two isotopes as 16O, 99.64% and 18O, 0.3600%.
Calculate a value for the relative atomic mass of oxygen based on these
figures.
Give your answer to four significant figures.
relative atomic mass = _______________ [3]
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9. Oxygen is the third most abundant element in the Universe. It is produced
in some stars by the ‘carbon burning process’. This process involves a series
of nuclear fusion reactions.
(i) Complete the equation for the following nuclear fusion process:
12
6
C +
12
6
C  168 O + 2
[2]
(ii) Describe and explain the conditions necessary for nuclear fusion to
occur.
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
10. Gallium has two naturally occurring isotopes: 69Ga and 71Ga.
A chemist determines that the percentage abundances of these isotopes in
gallium are:
69
Ga 60.1%
Ga 39.9%
71
(i) Calculate the relative atomic mass, Ar, of gallium.
Give your answer to three significant figures.
Ar = _______________ [2]
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(ii) Complete the table below to show the atomic structures of the two
gallium isotopes.
[2]
11. There has been a steep rise in demand in recent years for the isotope
helium-3. This isotope is used widely in cold temperature research and
medical imaging.
(a) Complete the following table to show the atomic structures of helium-3
and helium-4.
[1]
(b) Helium-3 is produced by the radioactive decay of hydrogen-3, one of the
isotopes of hydrogen.
State the type of radioactive decay in which hydrogen-3 is converted to
helium-3.
Explain your answer.
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
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(c) Helium-3 is also formed in the Sun by nuclear fusion processes involving
isotopes of hydrogen.
Write the nuclear fusion equation for the production of helium-3 from
two different isotopes of hydrogen.
[1]
(d) Nuclear fusion processes could be a useful energy source. It is very
difficult, however, to create the high temperatures and pressures to
bring about safe fusion on Earth.
Suggest why very high temperatures and pressures are needed for
nuclear fusion.
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
12. Analysis of tooth enamel from ancient human skeletons provides
information on the place of origin of people from the past.
The technique is based on the fact that the concentrations of strontium and
oxygen isotopes in tooth enamel vary with geographical locality.
(a) Give the number of protons, electrons and neutrons in the strontium
isotope, Sr-88.
protons _____
electrons _____
neutrons _____
[1]
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(b) A sample of strontium has four stable isotopes with abundances as
shown in the table below:
Use the above data to calculate a value for the relative atomic mass of
strontium in the sample.
Show your working and give your answer to three significant figures.
relative atomic mass of strontium in the sample = _______________ [3]
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Experimental error and
error analysis
Just how good are those results?
Practical chemistry can involve making all sorts of measurements. You may have to measure
the mass of a chemical, or the temperature or volume of a solution. Whenever you make a
measurement, there is an experimental uncertainty associated with the value you record.
This article looks at:
 the language used by scientists when discussing measurements (see Box 1);
 how to calculate uncertainties
 how to identify the most significant uncertainty;
 how to combine uncertainties;
A typical sixth-form experiment
The experiment described below involves a variety of measurements, all of which have a
degree of uncertainty associated with them. By looking at the procedures that carry the
greatest level of uncertainty, a chemist can start to appreciate the most important stages in the
experiment and concentrate on ways to minimise the level of uncertainty in the final answer.
A student was given the task of finding the concentration of a solution of hydrochloric acid,
which was thought to be around 0.15mol dm-3.
First he weighed out (using a 2 decimal place top pan balance) exactly 1.04g of sodium
hydroxide into a weighing boat. The sodium hydroxide was carefully transferred into a clean
beaker and distilled water was added. The weighing boat was washed through with distilled
water and the rinsings were transferred to the beaker. Once all of the solid had dissolved, the
solution was transferred into a 250cm3 (Grade B) volumetric flask. The beaker was washed
several times and each time the rinsings were transferred to the volumetric flask. After allowing
time for cooling to room temperature (sodium hydroxide dissolves exothermically), the solution
was made up to the calibration mark on the volumetric flask with distilled water. The flask was
stoppered and shaken to ensure thorough mixing.
Next, the student took 25.00cm3 samples of this ‘standard solution’ of sodium hydroxide (using
a Grade B pipette) and titrated them against the unknown hydrochloric acid (which was in a
50cm3 Grade B burette) using a suitable indicator. Several titrations were carried out until
concordant results were obtained. These results are shown in Table 1.
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Box 1: Sorting out the language
The term experimental error is often used in this context but isn’t a particularly helpful phrase
as it suggests that you may have done something wrong. A more appropriate term is
experimental uncertainty (or measurement uncertainty). It is important to remember that
even results obtained by the most experienced and careful operator will have an experimental
uncertainty associated with them.
Uncertainties in results can be caused by:
 imperfections in the measuring device
 imperfections in the experimental procedure
 judgements made by the operator
If you repeat a series of measurements and obtain values which are close together, your
results are said to be precise. Since one person obtained the results, the procedure used can
be described as repeatable. If the same procedure is carried out by a number of different
people and the results are still close together, the procedure is said to be reproducible.
If your results are also close to the true value, then they are said to be accurate.
A systematic error causes a bias in your measurement in one direction (but always in the same
direction). Systematic errors can be taken into account in your calculation, if you become
aware of them. For example, if a group of students were all carrying out the same experiment
using 25cm3 pipettes, but one student accidentally used a 20cm3 pipette, his titres would
consistently be lower than those of the rest of the group. Any calculations done using this
student’s results would differ from the rest of the group. However, the student might realise
what he has done and apply a correcting factor, enabling his results to be compared with
those of his friends.
Random errors occur in all experimental measurements (however careful or experienced the
operator may be) and are beyond the control of the operator. One can reduce the effect of
random errors by carrying out many repeat experiments, although this may not always be
practical. The average value from a set of repeat measurements is generally a better estimate
of the true value of the quantity.
Table 1
Mass of sodium hydroxide used = 1.04g
Titration 1
Titration 2
Titration 3
Titration 4
Initial burette reading/cm3
0.00
19.05
0.05
18.75
Final burette reading/cm3
19.05
37.65
18.75
37.40
Titre/cm3
19.05
18.60
18.70
18.65
Average titre of hydrochloric acid = 18.65cm3
(The first titration result was carried out quickly to get a rough idea of the titre. This result is
discarded and the average titre is calculated using results 2, 3 and 4).
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“Since you asked, there’s a 1% uncertainty in
alcohol content, 2% uncertainty in the volume
measurement and 0% uncertainty in the price .
. .”
Box 2: Calculating uncertainties
The uncertainty in a measurement depends on the
precision of the equipment being used to make the
measurement. The table below gives typical
measurement uncertainties for some common
laboratory equipment, taken from manufacturers’
literature.
Equipment
Size/class
Top pan balance 3 decimal place
Top pan balance 2 decimal place
Top pan balance 1 decimal place
Pipette
25cm3 Grade A
Pipette
25cm3 Grade B
Measuring cylinder 25cm3
Burette
50cm3 Grade A
Burette
50cm3 Grade B
Volumetric flask
250cm3 Grade B
Uncertainty value
0.0008g
0.008g
0.08g
0.02cm3
0.04cm3
0.5cm3
0.06cm3 (when used
to deliver 25cm3)
0.08cm3 (when used
to deliver 25cm3)
0.2cm3
Finding the percentage uncertainty (commonly called
the percentage error) in a measurement allows you to
compare uncertainties and decide which stage of your
procedure is likely to introduce the greatest
uncertainty.
The percentage uncertainty is calculated using the
formula:
Percentage uncertainty = uncertainty x 100
result
Calculating percentage uncertainties
for each measurement
The information given in Box 2 is used to
calculate these uncertainties:
Uncertainty in measurement using = 0.008 x 100 = 0.77%
2 decimal place top pan balance
1.04
Uncertainty in measurement using = 0.08 x 100 = 0.43%
Grade B burette
18.65
Uncertainty in measurement using = 0.2 x 100 = 0.08%
Grade B volumetric flask
250
Uncertainty in measurement using = 0.04 x 100 = 0.16%
Grade B pipette
25
Calculating the concentration of
sodium hydroxide
amount of sodium =
mass of sodium hydroxide
hydroxide used
molar mass of sodium hydroxide
= 1.04 = 0.026 moles
40
concentration of = amount of sodium hydroxide (mol)
sodium hydroxide
volume of solution (dm3)
= 0.026 = 0.104mol dm3
0.250
Combining uncertainties
When adding or subtracting measurements, the
maximum uncertainty is the sum of the uncertainties
associated with each individual measurement. For
example, if two temperatures are measured as 21.4°C
± 0.05°C and 31.7°C ± 0.05°C, the difference in
temperature is 10.3°C ± 0.1°C.
When multiplying or dividing measured quantities,
the maximum percentage uncertainty is the sum of the
percentage uncertainties for each of the quantities.
“I’ve double checked and I’m 100% certain we’re 0.1% uncertain”
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Calculating the concentration of the
acid
Which measurement has the greatest
uncertainty?
The equation for the reaction between sodium
hydroxide and hydrochloric acid is:
In this experiment, the percentage uncertainty in
measuring the mass of sodium hydroxide (0.77%)
is the greatest.
HCl (aq) + NaOH(aq)  NaCl(aq) + H2O(l)
Since hydrochloric acid and sodium hydroxide
react in a 1:1 mole ratio,
moles of NaOH = moles of HCl
so
conc. of NaOH x vol. of NaOH = conc. of HCl x vol. of HCl
therefore
conc. of HCl = concentration of NaOH x volume of NaOH
volume of HCl
= 0.104 x 25 = 0.1394mol dm-3
18.65
The percentage uncertainty on this value is given
by the sum of the percentage uncertainties for
each measurement (see Box 2).
% uncertainty in concentration of NaOH
= 0.77% + 0.08% = 0.85%
% uncertainty in volume of NaOH = 0.16%
% uncertainty in volume of HCl = 0.43%
overall uncertainty = 1.44%
This converts to an absolute uncertainty of:
0.1394 x 0.0144 = 0.002
therefore:
concentration of HCl = 0.1394 ± 0.002mol dm-3
Significant figures
When you have calculated a value from
measured data you should always consider how
many significant figures (sf) to include in your
final answer. Don’t just write down all the figures
shown on your calculator display. You should not
quote your final answer to more significant
figures than the least precise value used in the
calculation. In this calculation, the least precise
measurement was the mass of sodium hydroxide
(measured to 3sf). Therefore the final answer for
the concentration of the hydrochloric acid should
be quoted to three significant figures, i.e.
concentration of HCl = 0.139 ± 0.002mol dm-3
“And don’t forget to include operator error in your
analysis.”
One way to reduce this uncertainty would be to
use a 3 decimal place balance to weigh out the
sodium hydroxide. If an identical mass of sodium
hydroxide were used, the percentage error in the
measurement would drop to 0.077%.
A value for the concentration of the hydrochloric
acid has been obtained but it may not be
accurate (i.e. close to the true value). Using
sodium hydroxide to make a standard solution
could introduce a systematic error into the
procedure. This is because sodium hydroxide
absorbs carbon dioxide from the atmosphere.
When weighing out sodium hydroxide, we
cannot be certain about the amount (in moles) of
chemical used. An experienced chemist would
probably choose a different substance (such as
anhydrous sodium carbonate, which is much
more stable than sodium hydroxide) when
preparing the standard solution. A good way to
check the accuracy of the result from the above
experiment would be to compare the value
obtained with one obtained using a standard
solution of sodium carbonate.
However, irrespective of which chemical is used
for the standard solution, random errors cannot
be avoided. Judgements about whether the
bottom of the meniscus is touching the
calibration line on the pipette still have to be
made. Other sources of random errors include
temperature variations in glassware and
solutions.
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The accuracy of the final value also depends on the experimental procedures used. Taking care
to transfer all the sodium hydroxide from the weighing boat to the beaker and transferring the
rinsings from the beaker into the volumetric flask are two examples of good experimental
technique; remembering steps such as these can have a significant effect on the overall
accuracy in the final value.
This article has introduced you to some of the important ideas in error analysis. It is as well to
remember that however carefully you may have carried out your experiment, there will always
be some uncertainty associated with your final answer. We don’t live in a perfect world where
all measurements are exact!
Alasdair Thorpe is Head of Chemistry at Ampleforth College and a member of the CHEMISTRY
REVIEW Editorial Board.
Key concepts
Experimental uncertainty
Combining uncertainties
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Chemical Ideas 6.1
Light and electrons
Lyman series
When electrons absorb energy they can be promoted from a lower energy level to a higher
energy level. When the excited electrons drop back to a lower energy level, energy is emitted as
electromagnetic radiation such as visible light, ultraviolet light or infrared light. The frequency
of the radiation emitted depends upon the difference between the two energy levels. This
equation shows the relationship:
ΔE = hv
where ΔE is the difference in energy level
h is the Planck constant, 6.63 x 10-34 J Hz-1
v is the frequency of the radiation
Electrons in an atom can only exist at certain energy levels so each transition from one energy
level to another produces radiation of a characteristic frequency. The Lyman series is the
spectrum of ultraviolet light obtained when electrons fall from higher energy levels to the
ground state, the lowest available energy level.
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Spectra
Emission spectra are obtained by splitting up the light emitted from hot atoms. In the visible
region, they comprise coloured lines on a black background.
This is called an emission spectrum because it arises from light emitted from excited atoms.
Absorption spectra are obtained by passing white light through a vapour of atoms. In the
visible region they comprise black lines in the full spectrum or red, orange, yellow, green, blue,
indigo and violet light.
This is called an absorption spectrum because it arises from the light remaining after
absorption by atoms.
Each element produces a characteristic spectrum. It is possible to identify an element in a
mixture using spectra. By studying the intensity of each line in the spectrum the elements
abundance in the mixture can be determined. For each element, the frequencies absorbed are
the same as those emitted so its emission and absorption spectra can be exactly superimposed
to produce a full spectrum.
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Chemical Ideas 2.2
Nuclear reactions
Half-life
The time taken for half of the nuclei of a particular isotope to decay is called the half-life.
Different isotopes have different half-lives. For example, uranium-238 has a half-life of 4.5
billion years; strontium-90 has a half-life of 28 years and polonium-214 has a half-life of just 150
microseconds.
The half-life of an isotope cannot be changed. It stays the same whatever the pressure or
temperature and it does not change if the isotope is part of a compound.
If we know the half-life of an isotope we can work out:
the mass of the isotope left or its activity if you also know how long it has decayed;
the time needed for the mass of the isotope or its activity to fall to a particular amount;
Plotting a graph of activity or mass against time produces an exponential decay curve.
Whichever part of the curve we study, the time taken for the activity or mass to halve is always
the same.
10
Radioactive decay of bismuth-210
9
8
7
6
Mass/g
5
4
3
2
1
X
0
0
5
10
15
20
25
30
Time in days
From the graph it can be seen clearly that the time taken for the mass of bismuth to drop from
8g to 4g is 5 days. The time taken for the mass of bismuth to drop from 4g to 2g is again 5 days.
The half-life is constant.
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Carbon dating
The relatively short half-life (compared to other radioisotopes) of just less than 6000 years
means that C-14 has been used extensively to back up other dating methods used for
archaeological artefacts. The age of archaeological specimens can be calculated by looking at
the amount of carbon-14 in a sample. The method is a form of radio-dating called carbon
dating. Radio-dating can also be used to date rocks.
How is carbon-14 formed?
The isotope carbon-14 is created at a constant rate in the upper atmosphere by cosmic rays
acting on nitrogen. The carbon-14 which is formed is radioactive and decays to produce
nitrogen again. There is therefore a fixed amount of carbon-14 in the environment which is a
balance between the rate at which it is formed in the atmosphere and the rate at which it
decays back to nitrogen.
How does carbon dating work?
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All living things take in carbon from the environment. Plants take in carbon during
photosynthesis. Animals take in carbon when they eat food because food contains carbon. All
living things therefore have carbon-14 in them at the same amount which is present in the
environment. This amount is small. Only one in 850 billion carbon atoms are the isotope
carbon-14. The others are not radioactive. They are carbon-12 (about 99%) and carbon-13
(about 1%).
When a living thing dies, it stops taking in carbon from its environment. The amount of carbon14 in it will start to decrease as the carbon-14 slowly decays. The further back in time that
something died, the less carbon-14 will be present in it today.
The half-life of carbon-14 is 5730 years. Measuring the amount of carbon-14 in a sample today
can tel you how long ago the thing died and therefore the age of the sample. Carbon dating is
very useful but also has its limitations.
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1. Some rocks contain radioactive isotopes that can be used to date the
rocks.
Dating of rocks requires being able to accurately measure the amount of
both original (parent) radioisotope and finishing (daughter) stable isotope.
Suggest two assumptions that must be made if a radioisotope is to be used
for dating a rock.
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
__________________________________________________________ [2]
2. Much of our knowledge of outer space is based on spectroscopic data.
Absorption spectra give information about the elements present in stars.
(i)
Describe the main features in the appearance of an atomic
absorption spectrum.
________________________________________________________
________________________________________________________
________________________________________________________
_____________________________________________________ [3]
(ii)
How does an atomic emission spectrum differ in appearance from an
absorption spectrum?
________________________________________________________
________________________________________________________
_____________________________________________________ [2]
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3. The Shroud of Turin is a linen cloth imprinted with a faint image of a man.
Some people believe the image to be that of Jesus Christ.
In order to measure the Shroud’s age, small samples were analysed using
radiocarbon dating.
(a) Radiocarbon dating involves measuring the amount of a radioactive
isotope 14C in a sample of the cloth. The amount of 14C is compared to
the amount of 12C in the cloth.
The half-life for the decay of 14C is 5730 years.
Explain what is meant by the term half-life.
___________________________________________________________
________________________________________________________ [2]
(b) The linen in the Shroud was made from plants. The amount of 14C in
plants remains constant until that plant dies. The amount of 14C then
falls steadily. The amount of 12C does not change.
(i) Complete the table below for the decay of 14C.
(ii) Use the figures in the completed table to plot a decay curve for 14C on
the axes below (next page):
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time after death of plant/years
[2]
(iii)
The sample of linen cloth from the Shroud gave a 14C : 12C ratio of
0.920 x 10-12.
Use your graph to estimate the age of the cloth.
age of cloth = _______________ years [1]
(c) Other radioactive isotopes are used as medical tracers.
Suggest why it is not advisable to use an isotope with either a very short
or a very long half-life as a medical tracer.
very short half-life ___________________________________________
___________________________________________________________
very long half-life ___________________________________________
________________________________________________________ [2]
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4. Another use of radioisotopes is for the dating of geological materials such
as ancient lava flows from extinct volcanoes.
One method involves measuring the amounts of potassium-40 (40K) found
in the lava compared with the amounts of its decay product argon-40 (40Ar).
The half-life of potassium-40 is approximately 12,000 million years.
Suggest and explain why this long half-life makes potassium-argon dating
unreliable for determining the age of lavas that are only thousands of years
old.
_____________________________________________________________
__________________________________________________________ [1]
5. Argon was identified by its atomic emission spectrum.
Explain the occurrence of an atomic emission spectrum in terms of changes
in electronic energy levels and explain why such spectra are unique for
individual elements.
In your answer, you should use appropriate technical terms, spelled correctly
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
__________________________________________________________ [4]
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6. Stable lead-207 atoms are the final product in a series of steps starting with
the unstable uranium-235 isotope which is found in some rocks.
(i) What term is used to describe unstable isotopes of elements?
_______________________________________________________ [1]
(ii) Complete the following nuclear equation for the first step of the
breakdown of uranium-235 by alpha decay.
[3]
(iii) The ratio of uranium-235 atoms to lead-207 atoms can be used to
determine an approximate age for some rocks.
The dating relies on the fact that the half-life of any given unstable
isotope is fixed.
Explain the term half-life.
___________________________________________________________
________________________________________________________ [2]
7. A Russian agent died in London in November 2006, possibly as a result of
drinking tea to which the radioisotope polonium-210 had been deliberately
added. Polonium-210 has a half-life of 138 days.
(a) (i) Polonium-210 decays by emitting α-radiation.
Use the Data Sheet to write a nuclear equation for this decay.
[2]
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(ii) It is thought that less than a microgram (1.0 x 10-6g) of polonium-210
could be fatal.
Calculate the number of moles of polonium in 1.0 x 10-6g of
polonium-210.
Give your answer to two significant figures.
number of moles = __________ [2]
(b) Tiny amounts of polonium-210 were later found around London.
Suggest one reason why contact with this polonium-210 was unlikely to
pose a health risk.
___________________________________________________________
________________________________________________________ [1]
(c) A student measured the count rate of a different radioisotope over a
period of one week. The measurements were taken at the same time
each day. The results for this experiment are tabulated below:
(i) Plot the results of the student’s experiment on the graph paper on
the next page. Draw a suitable line showing how the count rate
reduces over a period of 10 days.
[3]
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(ii) Explain the term half-life.
Use your graph to measure the value of the half-life in days.
Show your working on the graph.
_______________________________________________________
_______________________________________________________
half-life = __________ days [2]
(iii) Suggest whether the half-life of this particular radioisotope would
be suitable for its use as a ‘tracer’ in the body. Give your reasoning.
_______________________________________________________
_______________________________________________________
____________________________________________________ [2]
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8. The structure of the atom is an example of a model used in chemistry which
has gradually become more sophisticated as new experimental evidence
has become available.
What feature of modern atomic structure does the occurrence of emission
spectra support?
_____________________________________________________________
__________________________________________________________ [1]
9. The electrons in the gaseous atoms of mercury in a low energy bulb are
excited when the bulb is switched on. Energy is then emitted as UV and
visible light.
Analysis of the UV radiation shows it to be an atomic emission spectrum.
(i) Describe the main features of an atomic emission spectrum.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [3]
(ii) Passing UV and visible light through a cool sample of mercury vapour
produces an atomic absorption spectrum.
Describe one difference between an atomic absorption spectrum and
an atomic emission spectrum.
__________________________________________________________
_______________________________________________________ [1]
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10. The half-life of the 18F isotope is around 100 minutes.
Explain the term half-life.
Suggest why an isotope with a half-life much longer or much shorter than
100 minutes would be unsuitable for use as a radiotracer.
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
__________________________________________________________ [3]
11. The presence of helium in the Sun was detected when its atomic
absorption spectrum was recorded during a solar eclipse.
(i) Describe the appearance of an atomic absorption spectrum.
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
(ii) Explain in terms of electronic energy levels, why the atomic absorption
spectrum of a particular element is unique.
In your answer, you should include the relationship between energy
and the radiation absorbed.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [3]
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12. The isotope carbon-14 can be used to estimate the time since an organism
died.
There is a roughly constant amount of carbon-14 in living material. When
an organism dies, the amount of carbon-14 within it gradually decays and is
not replaced.
Carbon-14 undergoes β-decay with a half-life of approximately 6000 years.
(i) Write an equation for the β-decay of carbon-14.
[2]
(ii) Fossil remains from an organism that died during the last ice age can be
dated by determining the count rate of the remaining carbon-14.
A fossil has a count rate that is 12.5% of the value in living material.
Calculate the number of years since the organism died.
number of years = __________ [2]
(iii) Suggest two assumptions that have to be made if the above
determination has to be valid.
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [2]
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EL2: The molecules of life
Learning objectives:
 Draw and interpret simple electron ‘dot-and-cross’ diagrams to show how atoms bond
through ionic, covalent and dative covalent bonds and be able to describe a simple
model of metallic bonding;
 Describe some limitations of these models;
 Recall the typical properties (melting point, solubility in water, ability to conduct
electricity), characteristic of giant lattice (metallic, ionic, covalent network) and simple
molecular structure types;
 Use the electron pair repulsion principle to predict and explain the shapes of simple
molecules (such as CH4, NH3, H2O and SF6) and ions (such as NH4+) with up to six outer
pairs of electrons (any combination of bonding pairs and lone pairs);
Key definitions:
Compile a glossary by writing your own definitions for the following key terms related to the
learning objectives above.
Key term
cation
anion
electrostatic bond
dot-and-cross diagram
ionic bond
covalent bond
single bond
bonding pairs
lone pairs
Definition
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Key term
double bond
dative covalent bond
lattice
delocalised
metallic bonds
tetrahedral
linear
planar triangular
trigonal bipyramidal
octahedral
Definition
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Chemical Ideas 3.1
Chemical bonding
The covalent bond
A covalent bond is a shared pair of electrons between two non-metal atoms. The electrons
involved are found in the outer shells of the atoms. The shared electrons count as part of the
outer shell of both bonded atoms.
If one pair of electrons is shared, a single bond forms
If two pairs of electrons are shared, a double bond forms
If three pairs of electrons are shared, a triple bond forms
Electron pairs that form bonds are called bonding pairs. Electron pairs that are not involved in
bonding are called lone pairs.
If you know an element’s group number you can work out how many bonds it should form. The
number of bonds is eight minus the group number (there are some exceptions however).
Dot-and-cross diagrams are used to show covalent bonding. The dots and crosses represent the
outer electrons of the atoms in the molecule.
In a dative covalent bond, both bonding electrons come from one of the atoms in the
molecule. Once it has formed, a dative covalent bond is the same as an ordinary covalent bond.
A single covalent bond is shown as a single line in chemical structures. A dative covalent bond is
shown as an arrow with the arrow pointing away from the atom that contributes both
electrons.
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Ionic bonding
An ionic bond is an electrostatic force of attraction between ions with opposite charges. Ionic
compounds have a giant ionic lattice structure in which ions are regularly arranged.
Metal atoms lose electrons from their outer shell to form positive ions while non-metal atoms
gain electrons to form negative ions (see table below).
Metals
Non-metals
Group
1
2
3
5
6
7
Charge on ion
+1
+2
+3
-3
-2
-1
It is unusual for Group 4 elements to form ions – they usually form covalent bonds instead. The
transition metals form positive ions as do the metals in Groups 1, 2 and 3. However, transition
metals can form more than one ion. For example, copper can form Cu + and Cu2+ ions.
Ionic bonds are very strong and it takes a lot of energy to overcome them. This means that ionic
compounds are solids at room temperature and have high melting points and boiling points.
The strength of ionic bonds increases as:
the charge on the ions increases;
the size of the ions decreases;
Ionic compounds do not conduct electricity when they are in the solid state. However, they do
conduct electricity when they are molten or dissolved in water because their ions are free to
move.
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Chemical Ideas 3.2
The shapes of molecules
The basic shape of a molecule or compound is determined by the number of pairs of electrons
around the central atom although the presence of lone pairs of electrons around the central
atom may distort it.
Two pairs of electrons
Where there are two pairs of electrons (or two groups of electrons) around the central atom
the molecule is linear. The bond angle is 180°.
180°
Cl
Be
Cl
linear molecule
Three pairs of electrons
Where there are three pairs of electrons (or three groups of electrons) around a central atom
the molecule is trigonal planar. The bond angles are all 120°.
F
F
B
120°
F
boron trifluoride is a triangular molecule
Four pairs of electrons
H
C
H
109°
H
H
methane is a tetrahedral molecule
Where there are four pairs of electrons
around the central atom the molecule is
tetrahedral. The bond angles are about
109°.
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Six pairs of electrons
Where there are six pairs of electrons the molecule is octahedral.
F
F
F
S
F
F
F
Summary
Pairs of electrons around
central atom
Shape
Examples
2
linear
BeCl2, CO2
3
planar triangular
BF3
4
tetrahedral
CH4, NH4+, NH2-
6
octahedral
SF6
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1. Radiowaves can provide information about the molecules found in some
regions of space.
One molecule found in the coldest regions of outer space has the formula
H2CO. This molecule can be represented as:
Draw the ‘dot-and-cross’ diagram for this molecule.
[2]
2. Use your diagram from Q1 to help you describe and explain the shape of
H2CO, giving the bond shape.
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
_____________________________________________________________
__________________________________________________________ [5]
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3. The strength of metals and metal alloys is due to their strong metallic
bonding.
Draw a labelled diagram to show a simple model of metallic bonding.
[3]
4. The arrangement of water molecules in ice is very regular. The shape of
individual water molecules is important in building up the regular structure.
(i) Draw a ‘dot-and-cross’ diagram for a water molecule.
[2]
(ii) Use the electron pair repulsion principle to describe and explain the
shape of a water molecule and suggest the bond angle.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [4]
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5. Bartlett’s work led to several other noble gas compounds being prepared,
including the gas xenon hexafluoride, XeF6.
The diagram below is a representation of the shape of this molecule.
Explain the significance of the dotted lines and wedges in this diagram.
_____________________________________________________________
__________________________________________________________ [2]
6. The Beijing torch was made of an alloy of aluminium and magnesium.
The strength of metals is due to the nature of metallic bonding.
Draw a labelled diagram to represent a simple model of metallic bonding.
[3]
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7. Sulphur is a vital element in living organisms and a key industrial chemical.
However, many covalent compounds of sulphur smell particularly bad.
Methanethiol, CH3SH, one of the molecules causing bad breath, is
particularly smelly.
Complete a ‘dot-and-cross’ diagram below for methanethiol.
[2]
8. The strong smell of cut onions is the result of volatile sulphur compounds
getting into the atmosphere. One of these compounds also makes you cry.
Its structure is given in Fig. 3.1 with two bond angles indicated by ‘a’ and
‘b’.
(i) What is the molecular formula of the compound shown in Fig. 3.1?
_______________________________________________________ [1]
(ii) The bond angle indicated by ‘a’ in Fig. 3.1 is 109° whereas ‘b’ is 120°.
Explain these bond angles.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
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__________________________________________________________
__________________________________________________________
_______________________________________________________ [4]
9. A particular type of gunpowder explodes to produce a mixture of nitrogen
gas, carbon dioxide gas and solid potassium sulphide, K2S.
(i) Complete and balance equation 2.1 below to show this reaction.
KNO3(s) +C(s) + S(s)  K2S(s) + _____(g) + _____(g)
equation 2.1
[2]
(ii) Potassium sulphide is an ionic compound.
Draw a ‘dot-and-cross’ diagram for K2S.
Show the outermost electron shells only.
[3]
10. Sulphur and carbon are both covalently bonded elements but sulphur has a
simple molecular structure whereas carbon has a giant covalent network.
Some of the physical properties of carbon and sulphur are very different
because of this difference in structure type.
(i) Name one physical property, apart from electrical conductivity, that will
be very different and state how it will differ for the two elements.
__________________________________________________________
_______________________________________________________ [1]
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(ii) In solid sulphur, the molecules are in the form of ‘puckered’ S8 rings as
shown below.
Suggest a value for the
Give your reasoning.
bond angle in the S8 molecule.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [4]
11. The high melting point of tungsten is a result of very strong metallic
bonding.
The diagram below illustrates a model of metallic bonding.
Write appropriate labels in the two boxes which help explain this model of
metallic bonding.
[3]
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12. Fluorine-containing compounds are used in health-related applications. For
example, sodium fluoride is present in some toothpastes.
(i) Draw a ‘dot-and-cross’ diagram to show the ionic bonding in sodium
fluoride.
Show outer electron shells only.
[2]
(ii) Sulphur hexafluoride is used in some types of eye surgery.
Use the electron pair repulsion principle to predict and explain the
shape of the sulphur hexafluoride, SF6, molecule and suggest the F-S-F
bond angle.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [4]
13. The physical properties of ethylene glycol, sodium chloride and paraffin
wax are important to the working of a lava lamp.
The physical properties of these substances are a result of their bonding
and their structure type.
Complete the table on the next page, to show the relationship between the
bonding and the structure of these substances and a characteristic physical
property of each one.
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[3]
14. Flint is a form of silicon dioxide.
(i) Silicon dioxide has a melting point above 1600°C, does not conduct
electricity and does not dissolve in water.
Suggest the structure and bonding of silicon dioxide.
__________________________________________________________
_______________________________________________________ [2]
(ii) Silicon is in the same group as carbon.
Draw a ‘dot-and-cross’ diagram for carbon dioxide in the space below.
Show outer electrons only.
[1]
(iii) A ‘dot-and-cross’ diagram for silicon dioxide is different from the ‘dotand-cross’ diagram for carbon dioxide. In what way is it different?
__________________________________________________________
__________________________________________________________
_______________________________________________________ [1]
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15. Titanium has a giant metallic structure.
Draw a labelled diagram to show a simple model of metallic bonding.
[3]
16. The table below shows the structure and properties of carbon and silicon
and some of their compounds.
Complete the table by ticking the appropriate boxes to show the link
between the type of structure and melting point.
[1]
17. Carbon forms two oxides, CO2 and CO.
Representation of the bonding in these oxides is shown below.
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(i) Use the electron pair repulsion principle to state and explain the shape
of a carbon dioxide molecule.
__________________________________________________________
__________________________________________________________
__________________________________________________________
__________________________________________________________
_______________________________________________________ [3]
(ii) State the type of bond represented by the arrow in the carbon
monoxide structure.
_______________________________________________________ [3]
(iii) Draw a ‘dot-and-cross’ diagram for the carbon monoxide molecule.
Show outer electrons only.
[2]
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EL3: What are we made of?
Learning objectives:
 Explain and use the terms: atomic number, mass number, isotope, Avogadro constant,
relative isotopic mass, relative atomic mass, relative formula mass and relative
molecular mass;
 Use the concept of amount of substance to perform calculations involving: masses of
substances, empirical and molecular formulae, percentage composition;
 Write and interpret balanced chemical equations including state symbols;
Key definitions:
Compile a glossary by writing your own definitions for the following key terms related to the
learning objectives above.
Key term
atomic number
mass number
isotope
Avogadro constant
relative isotopic mass
relative atomic mass
relative formula mass
relative molecular mass
empirical formula
molecular formula
Definition
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Chemical Ideas 1.1
Amount of substance
Relative atomic mass
Each type of atom has a different mass. The link between the mass of an element and the
number of atoms it contains is the relative atomic mass (Ar) of the element.
This link allows chemists to work out chemical formulae.
Definition
RELATIVE ATOMIC MASS is the mass of an atom of a substance compared to
the mass of the carbon-12 isotope (12C).
In chemistry, approximate relative atomic masses of elements are used most of the time and
they DO NOT have units because they are relative values. Relative atomic masses are there only
for atoms.
A hydrogen atom is 12 times lighter than a 12C atom so its Ar = 1; magnesium has Ar = 24
(magnesium is 2 times heavier than 12C.
 Moles of atoms
The mole (abbreviation – mol) is the unit that measures amount of substance.
Chemical amounts are defined so that the mass of one mole is equal to the relative atomic
mass (Ar) in grams.
MASS OF 1 MOLE OF ELEMENT = Ar = MOLAR MASS
mass
mol of atoms =
g
{ mol =
molar mass
}
g mol-1
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Relative formula mass and relative molecular mass
Molecules are different from the single atoms. They consist of a few atoms chemically joined
together unless they are diatomic elements (H2, N2, O2, F2 and the rest of the halogens) which
only have one type of atoms chemically joined together.
Chemists use relative formula mass to compare ionic substances and they have NO UNITS.
For substances where molecules are formed by covalent bonding chemists use relative
molecular mass to compare them and they DO NOT HAVE UNITS.
Both relative formula mass and relative molecular mass are given the symbol Mr.
Definition
RELATIVE FORMULA MASS is the mass of 1 molecule of a compound
relative to the mass of a 12C atom.
To find the value of Mr you need to add the individual Ar of the elements together.
Formula units
Formula units can be single atoms (all metals as elements), molecules (diatomic elements;
covalent compounds) or groups of ions (in ionic compounds).
 Moles of formula units
The relative formula mass (Mr) in grams is equal to the molar mass of a compound:
methane Mr = 16  molar mass is 16g
mass (g)
= amount of moles of formula units
molar mass
The Avogadro constant
The number of formula units (atoms, molecules, ions, electrons) in one mole of a substance is a
constant. It is called the Avogadro constant (symbol L) and its value is 6.02 x 1023 formula units
mol-1.
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Chemical formulae
We use moles to find out the formula of a compound. Chemists use two types of the formulae:
empirical and molecular.
The empirical formula of a substance is the simplest formula. It tells you the ratio of the
numbers of different types of atom in the substance.
The molecular formula tells you the actual numbers of different types of atom.
What does the chemical formula tell you?
(NH4)3PO4
4 moles of oxygen
1 mole of phosphorus
3 moles of nitrogen
12 moles of hydrogen
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1. The characteristic noise produced when Christmas crackers are pulled apart
is caused by a small amount of gunpowder.
Gunpowder typically contains 10 percent sulphur, 15 percent carbon and
75 percent potassium nitrate, KNO3, by mass.
(a) Calculate the number of moles of KNO3 in 100g of gunpowder.
number of moles = _______________________________
[1]
(b) The mole ratio of carbon atoms to sulphur atoms in gunpowder is about
4 to 1.
 State, in terms of the Avogadro constant, what is meant by ‘a mole’.
 Explain why the mole ratio is greater than the mass ratio.
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [3]
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EL4: Looking for patterns in elements
Learning objectives:
 Write and interpret balanced chemical equations including state symbols;
 Recall that the Periodic Table lists elements in order of atomic (proton) number and
groups elements together according to their common properties;
 Use given information to describe trends in a group of the Periodic Table and to make
predictions concerning the properties of an element in this group;
 Describe periodic trends in the properties of elements, in terms of melting point and
boiling point;
 Recall that the position of an element in the Periodic Table is related to its electron
structure (main energy levels or electron shells) and vice versa;
 Describe and compare the following properties of the elements and compounds of Mg,
Ca, Sr, and Ba in Group 2: reactions of the elements with water, acid-base character of
oxides and hydroxides, thermal stability of the carbonates, solubilities of hydroxides and
carbonates;
 Understand how Mendeleev developed the Periodic Table by leaving gaps and
rearranging some elements from their atomic mass order and how subsequent research
validated this knowledge;
 Given relevant information, discuss other examples of how scientific research can be
used to assess the validity of a discovery;
 Describe and explain the main stages in the operation of a time-of-flight mass
spectrometer;
 Use data from a mass spectrometer to: (i) calculate relative atomic mass and the
relative abundance of isotopes; (ii) work out the relative molecular mass of molecules
and understand that other peaks are caused by fragments of the molecule (no detail
required at this stage);
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Key definitions:
Compile a glossary by writing your own definitions for the following key terms related to the
learning objectives above.
Key term
ionisation source
‘time of flight’ mass
spectrometry
drift region
mass spectrum
molecular ion
fragment
base peak
Definition
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Chemical Ideas 1.2
Balancing equations
Write down the reactants and products in a reaction and the relative amounts involved.
Equation is balanced so that there are equal numbers of each type of atom on both sides.
CH4 (g) + 2O2(g)  CO2(g) + 2H2O(g)
Steps for balancing equations
The reaction between calcium and water:
Step 1
Decide what the reactants and products are.
calcium + water  calcium hydroxide + hydrogen
Step 2
Write formulae for the reactants and products including state symbols.
Ca(s) + H2O(l)  Ca(OH)2(aq) + H2(g)
Step 3
Balance the equation so there is the same number of each type of atom on both sides of the equation.
Ca(s) + 2H2O(l)  Ca(OH)2(aq) + H2(g)
Equations can only be balanced by putting numbers in front of the formulae! You cannot balance them
by altering the formulae themselves because you would create different substances.
State symbols
State symbols are included in chemical equations to show the physical state of
the reactants and products:
(g) gas; (I) liquid; (s) solid; (aq) aqueous solution.
The Law of Conservation of Mass
In 1774, French chemist Antoine Lavoisier found that if nothing is
allowed to enter or leave a reaction vessel, the total mass is the same
after the reaction as it was before.
Atoms are not created or destroyed in chemical equations, they are
simply rearranged . . . so equations must balance!
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Chemical Ideas 6.5
Mass spectrometry
Overview
The mass spectrometer is an instrument that can:
measure the mass and relative abundance of different isotopes in an element;
measure the atomic or molecular mass of different particles (atoms or molecules) in a
sample;
at a more complex level to investigate the structure of molecular compounds;
The mass spectrometer can be divided into three fundamental parts – the ionisation source,
the analyser and the detector. These are all maintained under high vacuum to give the ions a
chance of travelling from one end of the instrument to the other without any hindrance by
collisions with air molecules.
Acceleration area
Ionisation area
Sample inlet
Flight path
Light ions
Ion-detector
Heavy ions
Drift region
Vacuum chamber
Time measurement
Ionisation
An electron gun produces a stream of high-energy electrons from a heated metal filament.
These high-energy electrons bombard the sample of the element. When a high-energy electron
hits an atom in the sample it knocks an electron from the outside of the atom forming an ion
with a single positive charge:
X(g) + e- → X+(g) + 2 eAnalyser
A versatile, sensitive and rapid method of analysis called ‘time of flight’ mass spectrometry. The
ions are produced as separate pulses, not continually. A time of flight mass analyser identifies
charged sample atoms or molecules by measuring their flight time.
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An electric field accelerates all the ions to the same kinetic energy in the acceleration area.
Time of flight mass spectrometers identify ions by measuring the time that sample ions, all
starting with the same kinetic energy, take to fly a known distance (the flight path) in a constant
electric field – this area is called the drift region. If all the ions of different masses have the
same kinetic energy then heavier ions will move more slowly than lighter ions and will arrive at
the detector later, having a longer time of flight.
Detection
When an ion hits the detector, a tiny electrical current is released. This is amplified and the
signal is fed to a computer that analyses the signal. A graph of relative abundance against mass
is produced called a mass spectrum. The more of a particular ion that reaches the detector, the
bigger the signal, and the higher its bar in the mass spectrum.
Mass spectrometry: fragmentation and interpreting spectra
In a mass spectrum, the peak with the highest mass will be due to the molecular ion. The
molecular ion can fragment or break up into smaller ions and radicals. The ions will be detected
by the mass spectrometer but the free radicals will not.
The more stable the ion, the more intense the peak it produces. The most stable ion produces
the base peak and the heights of the peaks produced by the other ions are expressed relative
to this. Some ions are more stable than others and are commonly represented by peaks in mass
spectra.
Some ions are so unstable, they are barely represented at all. The table shows some common
fragments lost because of fragmentation or seen in mass spectra:
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Chemical Ideas 11.1
Periodicity
Atomic radius
The atomic radius decreases as you go across Period 3.
Going across the period:
the number of protons in the nucleus and the nuclear charge increase;
there are more electrons but the increase in shielding is negligible because each extra
electron enters the same shell;
the force of attraction between the nucleus and the electrons increases so the atomic
radius decreases;
First ionisation enthalpy
In general, the first ionisation enthalpy increases as you go across Period 3.
Going across the period:
there are more protons in each nucleus so the nuclear charge increases;
the force of attraction between the nucleus and outer electron increases;
there is a negligible increase in shielding because each successive electron enters the
same shell;
so more energy is needed to remove the outer electron;
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Density
The density of an element is its mass per unit volume. It is a measure of how closely its particles
are packed. Solids and liquids have closely packed particles so they tend to have a high density.
The particles in gases are far apart so gases have a low density.
The density of the metals increases across Period 3. The solids silicon, phosphorus and sulphur
are denser than sodium and magnesium. Chlorine and argon are both gases at room
temperature so their densities are very low.
Melting and boiling points
Change of state
When a substance melts, some of the attractive forces holding the particles together are
broken or loosened so that the particles can move freely around each other but are still close
together. When a substance boils, most of the remaining attractive forces are broken so the
particles can move freely and far apart. The stronger the attractive forces are the more energy
is needed to overcome them and the higher the melting and boiling points.
Sodium, magnesium and aluminium
Sodium, magnesium and aluminium are all metals. They have metallic bonding in which metal
cations are attracted to delocalised electrons. Going from sodium to aluminium:
the charge on the metal ions increases from +1 to +3;
the number of delocalised electrons increases;
so the strength of the metallic bonding increases;
and so the melting points and boiling points increase;
Silicon
Silicon has a very high melting point and boiling point because its atoms are held together by
strong covalent bonds in a giant covalent structure.
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Phosphorus, sulphur, chlorine and argon
Phosphorus, sulphur and chlorine exist as simple molecules with strong covalent bonds
between their atoms. Argon exists as single atoms. These elements have low melting and
boiling points because only weak instantaneous dipole-induced dipole attractions need to be
broken rather than strong metallic or covalent bonds.
Phosphorus exits as P4 molecules, sulphur as S8 molecules, chlorine as Cl2 molecules and argon
as individual atoms. As the strength of instantaneous dipole-induced dipole attractions
decreases as the size of the molecule decreases, the melting and boiling points decrease in
order S8 > P4 > Cl2 > Ar.
Electrical conductivity
For an element to conduct electricity, it must contain electrons that are free to move. In
general, metals are good conductors of electricity and non-metals are poor conductors of
electricity.
Sodium, magnesium and aluminium
Sodium, magnesium and aluminium are all metals. They have metallic bonding in which metal
cations are attracted to delocalised electrons. The delocalised electrons are free to move and
carry charge.
Going from sodium to aluminium, the number of delocalised electrons increases so there are
more electrons that can move and carry charge. As a result, the electrical conductivity
increases.
Silicon
Silicon is a metalloid, an element with some of the properties of metals and some of the
properties of non-metals. It has a giant covalent structure similar to that of diamond. Each
silicon atom is covalently bonded to four other silicon atoms in a tetrahedral arrangement.
Silicon is a semiconductor because some electrons can be promoted to higher energy levels
when it is warmed up and these delocalised electrons can move and carry charge.
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Non-metals
The remaining elements in Period 3 do not conduct electricity. In phosphorus, sulphur and
chlorine the outer electrons are not free to move and carry charge because they are held
strongly in covalent bonds. In argon which exists as single atoms, the outer electrons are not
free to move.
Relative electrical conductivities of the elements in
Period 3
1.0
0.9
0.8
0.7
Electrical conductivity
compared with
aluminium
0.6
0.5
0.4
0.3
0.2
0.1
0
Na
Mg
Al
Si
P
S
Cl
Ar
Phosphorus, sulphur and chlorine are simple molecules with no free electrons.
Argon exists as single atoms.
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Chemical Ideas 11.2
Physical trends in Group 2
Atomic radius
The atomic radius increases as you go down Group 2.
Going down the group:
there are more filled energy levels between the nucleus and the outer electrons;
this means that the outer electrons are more shielded from the attraction of the
nucleus;
so the electrons in the outer energy levels are further from the nucleus and the atomic
radius increases;
Atomic radii of the elements in Group 2
Ba
0.20
Sr
0.19
0.18
Ca
0.17
Mg
0.16
Atomic radius/nm
0.15
0.14
0.13
Be
0.12
0.11
0.10
0
10
20
30
40
50
Proton number
There are more filled energy levels between the nucleus and the outer
electrons which are more shielded from its attraction.
60
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First ionisation enthalpy
The first ionisation enthalpy decreases as you go down Group 2.
Going down the group:
there are more filled energy levels between the nucleus and outer electron which shield
it from the attraction of the nucleus;
the radius of the atom increases so the distance between the nucleus and the outer
electron increases;
so the force of attraction between the nucleus and the outer electron is reduced and
less energy is needed to remove it;
Atomic radii of the elements in Group 2
Ba
0.20
Sr
0.19
0.18
Ca
0.17
Mg
0.16
Atomic radius/nm
0.15
0.14
0.13
Be
0.12
0.11
0.10
0
10
20
30
40
50
Proton number
There are more filled energy levels between the nucleus and the outer
electrons which are more shielded from its attraction.
60
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Electronegativity
Electronegativity decreases as you go down Group 2.
Going down the group:
the atomic radius increases;
the outer electrons are more shielded from the attraction of the nucleus;
so bonding electrons are less strongly attracted to the nucleus;
Melting point
In general, melting point decreases as you go down Group 2. The elements in the group are all
metals with metallic bonding where metal cations in a metal lattice are attracted to delocalised
electrons.
Going down the group:
the number of delocalised electrons remains the same and the charge on each metal
cation stays the same at 2+;
but the ionic radius increases;
so the attraction between the delocalised electrons and the metal cations decreases;
Melting points of the elements in Group 2
1600
Be
1500
1400
Melting point/K
1300
1200
Ca
1100
Sr
Ba
1000
Mg
900
800
0
10
20
30
40
50
Proton number
The strength of the metallic bonding decreases because the radius of the
metal ions increases
60
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Chemical Ideas 11.2
Chemical properties of Group 2
The elements
The elements in Group 2 become more reactive as you go down the group. If they react with
water, they produce a metal hydroxide and hydrogen: M(s) + 2 H2O(l)  M(OH)2(aq) + H2(g)
Thermal decomposition of Group 2 carbonates
The thermal stability of Group 2 carbonates increases going down the group. Magnesium
carbonate readily decomposes when it is heated but calcium carbonate needs strong,
prolonged heating. Group 2 carbonates decompose to produce a metal oxide and carbon
dioxide: MCO3(s)  MO(s) + CO2(g)
Solubility of Group 2 compounds
Going down the group:
the carbonates become less soluble;
the hydroxides become more soluble (and the base strength of the metal hydroxides
increases);
Trends in solubility of carbonates of Group 2
0.060
0.06
0.05
0.04
g per 100g water
0.03
0.02
0.01
0.0013
0.001
CaCO3
SrCO3
0.002
0.00
MgCO3
BaCO3
Calcium carbonate, strontium carbonate and barium carbonate
have similar solubilities, however.
Trends in solubility of hydroxides of Group 2
0.04
g per 100g water
3.7
0.03
0.02
1.0
0.01
0.0012
0.12
0.00
Mg(OH)2
Ca(OH)2
Sr(OH)2
Ba(OH)2
This means that their strength as bases also increases as you
go down the group.
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1. A time-of flight mass spectrometer can be used to measure the amount of
parent and daughter isotopes.
Write the following labels on the diagram of a time-of-flight mass
spectrometer below:
[4]
2. In the early 1800s, chemists had to learn the individual properties of all
elements. Fortunately they then developed ways of grouping the elements
so that patterns of chemical and physical behaviour emerged. This led to
our modern Periodic Table.
(a) The Russian chemist Mendeleev organised the known elements into a
pattern widely regarded as the first Periodic Table.
A version of Mendeleev’s Periodic Table is shown below:
SACKVILLE SCIENCE DEPARTMENT
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(i) Why did Mendeleev leave gaps in his Periodic Table?
___________________________________________________________
________________________________________________________ [1]
(ii) How did later evidence from the scientific community support
Mendeleev’s decision to leave gaps in his Periodic Table?
___________________________________________________________
___________________________________________________________
________________________________________________________ [2]
(iii) Which group in the modern Periodic Table is missing from
Mendeleev’s table?
________________________________________________________ [1]
(b) Before Mendeleev, John Newlands, another chemist, had been thinking
on similar lines. He also identified patterns in the behaviour of elements.
The table below shows how Newlands grouped the elements.
(i) Newlands and Mendeleev both put the elements in order of atomic
mass. What property is used to order the elements in the modern
Periodic Table?
________________________________________________________ [1]
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(ii) Suggest why Newlands’ arrangement was less useful than
Mendeleev’s arrangement.
________________________________________________________ [1]
(c) One physical property chemists examined was the melting point of the
elements. This is tabulated below for the Period 2 elements.
(i) Describe the pattern in melting point as you go across the period.
_______________________________________________________ [1]
(ii) The change in melting point across the period can be explained in
terms of the structure and bonding of the elements.
Describe the changes, both in type of bonding and in structure, as
the period is crossed from left to right.
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [4]
(d) The electron structure of an element is 2.8.8.2.
In which group and period of the modern Periodic Table is this element
found?
Group _____
Period _____
[1]
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3. Beryllium is a Group 2 metal and has several modern day uses. It also
played an important role in the development of idea about the structure of
the atom.
(a) In 1932, James Chadwick fired α-particles at beryllium metal and found
that particles were emitted that were not deflected by electric fields.
Chadwick had discovered the neutron. Give the mass number and
atomic number of this particle.
mass number = _____
atomic number = _____
[2]
(b) Alloys of beryllium and copper are used for aircraft parts because of
their high strength and resistance to corrosion.
A typical copper-beryllium alloy contains 1.75% by mass of beryllium.
Assume all the rest of the alloy is copper.
(i) Calculate the number of moles of beryllium and copper in 100g of the
alloy.
moles Be = __________ moles Cu = __________
[2]
(ii) Calculate the percentage of atoms of beryllium in the alloy.
Give your answer to two significant figures.
% Be atoms = ____________________
[3]
(c) Many beryllium compounds are covalent in character. An example is
beryllium chloride. The left-hand box below shows the bonding in a
beryllium chloride molecule in the vapour state. Complete the ‘dot-andcross’ diagram for this molecule in the right-hand box.
[2]
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(d) The chlorides of the other Group 2 metals are all ionic.
Name two physical properties of ionic chlorides which you would expect
to be different from those of simple covalent chlorides. State how the
properties would differ.
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [3]
4. A small amount of the element lead that originates from natural sources is
always found in the atmosphere. Ice cores analysed using a ‘time-of-flight’
mass spectrometer show changes in the concentration of this atmospheric
lead over time.
In the ‘time-of-flight’ mass spectrometer:
(i) What causes lead atoms to lose electrons and become cations in the
ionisation area?
___________________________________________________________
________________________________________________________ [1]
(ii) What causes the acceleration of lead ions in the acceleration area?
___________________________________________________________
________________________________________________________ [3]
(iii) What property of different isotopes causes their ions to take different
times to cross the drift region?
___________________________________________________________
________________________________________________________ [3]
Data from ice core measurements show that the concentration of the lead208 isotope rose sharply between 1930 and 1960 but has now fallen back
again.
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(iv) How would the appearance of a mass spectrum of a sample taken from
a 1950 core differ from the mass spectrum of a sample from a 1930
core?
___________________________________________________________
___________________________________________________________
________________________________________________________ [1]
5. Many compounds of the Group 2 element barium are poisonous. For
example, solid barium carbonate is used as a rat poison. It reacts with the
hydrochloric acid in the stomach to produce soluble barium chloride which
is poisonous.
(a) Write an equation for the reaction of solid barium carbonate with
hydrochloric acid. Carbon dioxide is produced in the reaction. Show
state symbols.
[3]
(b) Barium carbonate will also produce carbon dioxide gas when heated
strongly. There is a trend in the thermal stabilities of the Group 2
carbonates.
 Describe a simple experimental method you could use in the
laboratory to determine this trend. Your method should involve the
use of lime water.
 Describe how you would make this a fair test.
 Describe what you would expect to observe.
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
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___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [5]
(c) Another compound of barium, barium sulphate, BaSO4, is used when
medical X-rays of the digestive system are taken. Barium sulphate can
be swallowed by patients and causes no harm because it is insoluble.
A data book gives the solubility of barium sulphate as 2.20 x 10-4g per
100g of water.
Calculate the solubility of barium sulphate in mol per 100g of water.
Give your answer to three significant figures.
solubility of barium sulphate = ____________ mol per 100g water
[3]
(d) The position of barium in the Periodic Table is related to its electronic
structure.
Explain how the electronic structure of barium is related to the group
and period of the Periodic Table in which it is found.
___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [2]
6. Forest fires can be devastating, as large bush fires in Australia in recent
years have shown.
The ash that remains consists mainly of potassium carbonate that does not
decompose in that fire.
(a) Potassium forms 1+ ions in its compounds.
Write down the electronic configuration for the K+ ion.
________________________________________________________ [1]
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(b) Limestone, impure calcium carbonate, CaCO3, would have decomposed
in the extreme heat of the fire.
Write down the equation for the decomposition of calcium carbonate.
Show state symbols.
[2]
7. One concern linked to global climate change is that the sea level may rise,
leading to flooding. Oxygen isotope ratios, determined from geological
material, have varied over time and can be used to interpret past sea level
changes.
The two oxygen isotopes used are 18O and 16O.
(a) Explain how a time-of-flight mass spectrometer works. You should
include the following terms in your answer.
detector
drift region
ions
kinetic energy
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
___________________________________________________________
________________________________________________________ [5]
(b) The mass spectrum of a sample of material showed the abundance of
the two isotopes as 16O, 99.64% and 18O, 0.3600%.
Calculate a value for the relative atomic mass of oxygen based on these
figures.
Give your answer to four significant figures.
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relative atomic mass = _______________ [3]
(c) The oxygen isotope ratio of some fossilised shells has been determined.
Such shells are mainly calcium carbonate but occasionally magnesium
replaces some of the calcium in carbonate shells. Use your knowledge of
the Periodic Table to suggest and explain why magnesium can replace
calcium in carbonate shells.
___________________________________________________________
___________________________________________________________
________________________________________________________ [2]
8. Use the positions of carbon and sulphur in the Periodic Table and your
knowledge of periodic trends to suggest how the melting point of sulphur
would differ from that of carbon.
State why in terms of structure type.
_____________________________________________________________
_____________________________________________________________
__________________________________________________________ [2]
9. ‘Old fashioned’ light bulbs use tungsten metal as a filament which glows
white hot when an electric current is passed through it.
(i) Tungsten is used because it has a very high melting point.
The graph below shows the melting points of some of the elements on
either side of tungsten in Period 6 of the Periodic Table.
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A student attempts to use the graph to estimate a value for the melting
point of tungsten by using the two lines of best fit shown.
Estimate a value for the melting point of tungsten.
Clearly show on the graph how you arrived at your answer.
melting point of tungsten = _________________ K
[2]
(ii) There is a pattern (or trend) shown in the melting points of the
elements as this part of Period 6 is crossed. A similar pattern is shown in
Period 5.
Describe this pattern in the melting points of the elements.
__________________________________________________________
_______________________________________________________ [1]
10.Mendeleev was one of the first scientists to arrange the known elements
into groups according to their properties.
(i) By what property did Mendeleev order the elements?
_______________________________________________________ [1]
(ii) Mendeleev realised that, when he arranged the elements, some
elements’ properties did not fit with those above and below them in the
table.
Give one change that Mendeleev made to solve this problem.
_______________________________________________________ [1]
(iii) The elements in a modern Periodic Table are arranged by atomic
number.
Explain the meaning of the term atomic number.
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SACKVILLE SCIENCE DEPARTMENT
SALTERS AS CHEMISTRY
11.The molecular formula of FDG is C6H1118FO5. Traces of FDG in a patient’s
urine can be detected by mass spectrometry for several hours after a scan.
The mass spectrum of FDG shows several peaks.
(a) There is a peak in the mass spectrum of FDG at m/z 181.
Explain why you would expect to see a peak at this value.
Include a calculation in your answer.
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(b) What has happened to FDG to produce other peaks at lower m/z values
in the mass spectrum?
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12.In 1875, a French chemist noticed two lines in the atomic emission
spectrum of a sample of zinc ore. These lines did not match the known
spectrum of zinc. They were lines from an undiscovered element, later
named as gallium, Ga.
(a) Explain how the lines in an atomic emission spectrum are formed and
why the frequencies of these lines are unique for a given element.
In your answer, you should use appropriate technical terms, spelled correctly
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SACKVILLE SCIENCE DEPARTMENT
SALTERS AS CHEMISTRY
(b) The French chemist was aware of Mendeleev’s prediction that there
were elements still to be discovered.
In what way did Mendeleev’s Periodic Table allow for the discovery of
new elements, such as gallium?
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13.The isotopic composition of a sample of helium can be determined by using
a time-of-flight mass spectrometer.
Describe the main stages in the operation of a time-of-flight mass
spectrometer and explain why it is able to separate different isotopes of
helium.
In your answer, you should use appropriate technical terms, spelled correctly
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14.Mendeleev, in his early Periodic Table, arranged the known elements in
order of their ‘atomic weights’. However in several places, including
between calcium and titanium, he left gaps.
(a) Explain why he left gaps and how subsequent research confirmed this
decision.
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SACKVILLE SCIENCE DEPARTMENT
SALTERS AS CHEMISTRY
(b) State the property used in a modern Periodic Table to arrange the
elements.
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(c) What feature of an element’s atomic structure is related to the
element’s position in the Periodic Table?
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15.Strontium ions can replace calcium ions in tooth enamel because they both
have the same charge.
(a) State the charge on a strontium ion.
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(b) Calcium reacts with water.
Describe two observations that you would make when calcium reacts
with water.
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(c) The isotope variation in strontium across Britain is closely related to the
type of underlying rocks.
Some of these rocks contain strontium carbonate.
(i) The carbonates in Group 2 show a trend in thermal stability.
Write and equation to show the decomposition of strontium
carbonate on heating.
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SACKVILLE SCIENCE DEPARTMENT
SALTERS AS CHEMISTRY
(ii) Samples of strontium carbonate and calcium carbonate are taken,
each having the same particle size and number of moles. They are
heated in separate test-tubes under the same conditions.
Describe how the difference in thermal stability could be shown,
giving the results you would expect.
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16.Oxygen has three stable isotopes, with the O-16 isotope being by far the
most abundant.
The oxygen combined in tooth enamel is derived from drinking water.
A mass spectrum of water shows a large peak at an m/z value of 18.
A second, much smaller, peak is found at an m/z value of 20.
(a) What information is given by the height of a peak in a mass spectrum?
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(b) Give the formula of the ion responsible for the peak at an m/z value of
18.
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(c) Suggest why there is a peak at an m/z value of 20.
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