PART one: matter
30
2.1
What are elements made of?
The lead compounds in lead paint are made up of lead ions. Metals are always
present as ions. An ion is an atom that has gained a net electrical charge—either
positive or negative. But how does the ion gain or lose its electrical charge?
What makes up an atom to enable this to happen? If we were able to shrink down
to the size of an atom and look inside it, what would we see?
Elements and atoms
A pure substance contains only one
kind of particle. For example, pure
oxygen only contains oxygen particles,
pure water only contains water
particles, and so on. A mixture consists
of two or more pure substances mixed
together. For example, saline solution
is a mixture of water and the salt
commonly known as table salt, which
has the chemical name sodium
chloride. The two different pure
substances in saline solution can be
separated, using a process such as
distillation.
The atoms of an element are the tiny
spherical particles from which it is
made. They are unique; that is, no
other element has the same atoms. It
was once thought that atoms are the
smallest particle from which all
substances are made but we now know
that atoms are built up from even
smaller particles, called protons,
electrons and neutrons. Protons have
a positive charge, electrons have a
negative charge and neutrons have no
charge and are, therefore, electrically
neutral (hence their name). How these
smaller particles are arranged and
how they behave has led to the
development of modern atomic
theory.
practivity 2.1
How can you tell what is inside?
1 Team A places the ball in one of the
small boxes, the wooden block in
another, and the two nails in the last
box. The boxes are then closed.
Elements are the ‘ingredients’ from
which all substances are made. Every
single pure substance is made up of one
or more elements. When it contains two
or more elements, it is called a
compound.
2 Team B has to work out a way of
knowing what is inside each of the
boxes without opening or touching
them.
†† Fig 2.4 Why is pure gold classified as an element?
3 Team B can then touch and examine
the boxes, still without opening them.
•
Was team B more successful at
identifying what was inside the box
when able to touch and examine
the box?
More questions to consider ...
This kind of investigation uses what
scientists call ‘indirect evidence’. It has
been used by many scientists when trying
to work out what is inside the atom.
What you need: 1 ball, 1 soft drink can,
2 nails, 1 wooden block, 3 small boxes
Form two teams (A and B) of three
students to work with each other.
• How might scientists have used
indirect evidence to model what
is inside an atom?
•
Identify at least one other field of
scientific investigation in which
the scientists working in that field
would have to use indirect evidence
to develop their theories.
What do you know about elements and
atoms?
1 Which of the following substances is a
mixture?
a table sugar (sucrose)
b neon
c air
d sodium chloride
2 What is an atom?
3 The English chemist John Dalton said
that the atom is the smallest particle
Looking inside the
atom
In the early 1800s, the English chemist
John Dalton proposed that the atom was
the smallest particle of matter, that
atoms are indivisible and can be neither
created nor destroyed. However, a
century later in the early 1900s, the
physicist Joseph John Thomson (known
as JJ by his colleagues) discovered that
atoms were divisible and were made up
of even smaller particles. His
experiments showed that inside the
atom are far tinier, negatively charged
particles, which we now call electrons.
He also showed that the atom contained
positively charged material, although it
was not yet clear what this material
–
–
–
–
–
–
–
–
–
–
†† Fig 2.5 Thomson’s plum pudding model of the
atom.
that an element is made from. He said
atoms cannot be divided into smaller
parts. This theory fitted the evidence
available at the time. But new evidence
came to hand after this. What do we
now know about the atom?
4Explain the difference between an
element and a compound and give an
example of each.
was. From this evidence, and knowing
that oppositely charged objects attract
each other and move towards each
other, Thomson suggested that the atom
is like a plum pudding, in which the
positively charged material is the ‘cake’
and the electrons are the fruit. This was
called the Thomson plum pudding
model of the atom.
In 1911, Thomson’s former student,
Ernest Rutherford, performed an
experiment to test this theory. He fired
alpha particles (tiny positively charged
particles) at a very thin sheet of gold.
Most travelled straight through the gold
foil but, most unexpectedly, a very
small proportion were reflected back at
a range of angles. Rutherford realised
that these alpha particles must have
been subjected to a very powerful
electrostatic force of repulsion to
behave like that. He argued that the
only way such a strong force could arise
is if all the positive charge in each gold
atom in the foil was concentrated in a
very tiny nucleus instead of being
spread out over its whole volume.
Rutherford concluded that the atom is
mostly empty space. The electrons
travel through this space, revolving
around the tiny nucleus at its centre
like planets around the Sun. Most of the
mass of the atom is due to its nucleus.
This model of the atom is known as the
Rutherford nuclear model.
Not long after this, after examining
further evidence, another scientist,
†† Fig 2.6 The planetary model of the atom.
Neils Bohr, concluded that the electrons
in the atom do not behave quite like the
planets around the Sun but move about
the nucleus in circular orbits that are at
certain distances from the nucleus. The
more energy they have, the further their
orbit is from the nucleus. These sets of
orbits form concentric spheres known
as electron shells. There is a limit to the
number of electrons that can be found
in any of the shells. This is called the
Bohr model of the atom.
Over the next 20 years it was shown
that the nucleus is not a piece of solid,
positively charged material but, in fact,
contains separate particles—protons
and neutrons. Protons are positively
charged, whereas neutrons have no
charge. Rutherford was very involved
in this research.
It was found that the charge on a proton
was the same as that on an electron, but
positive instead of negative. It follows
that when an atom is uncharged (i.e. it
has no net charge), the number of
electrons present must equal the
number of protons present.
Bohr’s proposal about electron shells
has also been modified as new evidence
was gathered by others working in the
field. This led to the development of a
more sophisticated model of the atom,
known as the quantum mechanical
model, or the Schrödinger model, of
the atom.
chapter two: it’s all elemental
31
PART one: matter
32
What’s the theory now?
As more and more data has been
obtained, and scientists have far more
accurate instruments and other
technologies available to assist them in
their research, they have gained a much
deeper understanding of the structure
of the atom. The most recent model of
the atom is known as the quantum
mechanical, or Schrödinger, model.
Some of the features of this model are
listed below.
1 The nucleus of an atom is made up
of protons and neutrons.
2 The mass of the atom is almost
entirely due to the mass of the
nucleus; electrons have very little
mass in comparison.
3 Electrons do not move in precise
circular orbits in concentric shells
as proposed by Bohr, but rather
spend most of their time in regions
of space called atomic orbitals,
which have a variety of shapes.
No more than two electrons can
occupy an atomic orbital.
4 Each shell contains a definite
number of atomic orbitals, which
are collected into sets known as
subshells. Electrons in the same
subshell have the same energy.
5 The further the shell is from the
nucleus, the greater the number
of subshells within that shell and
hence the greater the number of
atomic orbitals in the shell. Each
shell can only hold a certain
maximum number of electrons.
Nucleus
Proton
Neutron
Electron
Electron shell
†† Fig 2.7 A model of an atom of the element lithium.
What do you know about looking inside
the atom?
1 a Describe Thomson’s plum pudding
model of the atom.
b What was Thomson’s reasoning
behind where he thought the
electrons were located?
2 a What is an alpha particle?
b
What was the most important new
understanding of the structure of
the atom that Rutherford deduced
from his experiment with alpha
particles?
4 What is an electron shell? Who
suggested this idea?
5Name and describe the three types
of particles we now know are found
inside the atom.
6 Working with a partner, make a threedimensional model of an atom from
plasticine or other suitable materials.
Make sure you label all parts correctly
and state which model of the atom
you are modelling.
3In his model of the atom, what did
Rutherford say about the electrons?
Atoms and their
mass
All atoms have a mass. But since they
are so small, chemists have devised a
relative mass scale for atomic particles
that is more convenient to work with
than using metric units, such as the
gram. First, however, it is important to
understand that, for most elements, not
all the atoms of the element have the
same mass. This is because they are not
identical. Why is this? What do they
have in common and what is different?
All the atoms of an element have the
same unique number of protons. This is
termed their atomic number, symbol Z,
and is used to identify the element. For
example, all carbon atoms contain 6
protons in their nucleus, so their atomic
number is 6. If you examine the periodic
table of the elements (see Fig 2.13), you
can see that the elements are listed in
order of their atomic number.
However, the number of neutrons in
the atoms of an element can vary. For
example, most carbon atoms have 6
neutrons in their nucleus but some have
7 and some have 8. The different forms
of the atoms of an element that have
different numbers of neutrons are
termed isotopes.
Protons and neutrons have a very
similar mass. They each have a mass
that is almost exactly 1 on the relative
scale of masses used for particles at the
atomic level. The mass of an electron is
negligible in comparison.
As they are both found in the nucleus,
protons and neutrons are both called
nucleons.
The atomic mass of the carbon isotope
that has 6 protons and 6 neutrons (i.e.
12 nucleons) in its nucleus is exactly 12
on this relative mass scale. (Carbon-12
is the only atom that has a relative
mass that is an exact whole number.
This isotope was chosen as the
standard mass for the relative mass
scale, and given a mass of exactly 12
for technical reasons.)
The relative mass of the carbon isotope
that has 6 protons and 7 neutrons
(i.e. 13 nucleons) in its nucleus is very
close to 13. The relative mass of the
carbon isotope that has 6 protons and
8 neutrons (i.e. 14 nucleons) in its
nucleus is very close to 14. Similarly,
the relative mass of any isotope of any
element is very close in value to the
total number of nucleons in its nucleus.
For this reason the total number of
nucleons present is termed the mass
number of the isotope, symbol A.
The conventional name and
representation of an isotope of an
element is shown in Figure 2.8 and the
three isotopes of carbon are shown in
Figure 2.9.
Most elements have more than one
naturally occurring isotope and are
present in nature as a mixture of these
isotopes. In these cases, chemists use
the average mass of the isotopes of an
element for calculations. This is termed
the relative atomic mass of the element.
For example, almost all carbon atoms
exist as the carbon-12 isotope and only
a very small proportion are present as
the two heavier isotopes. Thus, the
Mass number
(total number of protons
and neutrons)
6 protons
6 neutrons
8 neutrons
14
12
carbon-14: 6 C
carbon-12: 6 C
6 protons
7 neutrons
13
A
Z
X
Symbol of
element
Atomic number
(total number of
protons)
†† Fig 2.8 The conventional representation of an
isotope of an element.
carbon-13: 6 C
†† Fig 2.9 The three isotopes of carbon.
mean mass of carbon atoms is very close
to 12 (12.011 correct to 3 decimal places).
The unit for relative atomic mass is the
atomic mass unit (amu). So, the relative
atomic mass of carbon is 12.011 amu.
Maths Lab: Calculating protons,
neutrons and electrons
As stated earlier, for uncharged atoms,
the number of electrons must equal
the number of protons. If we know
the atomic number and mass number
of an isotope, we can calculate the
number of protons, neutrons and
electrons present in an uncharged
atom of that isotope.
EXAMPLE
Aluminium is one of the few elements
with only one isotope. It has an atomic
number of 13 and a mass number
of 27. This means that it has 13
protons, and 14 (27 – 13) neutrons in
the nucleus. Its uncharged atoms will
have 13 electrons.
YOUR TURN
One of the isotopes of iron, iron-56,
has an atomic number of 26 and a
mass number of 56. Calculate the
number of protons, electrons and
neutrons present in its uncharged
atoms.
ANSWER
protons = 26, electrons = 26 and
neutrons = 30
SKILLS LAB
6 protons
The relative atomic masses of the
elements are usually shown in the
periodic table, correct to 1 or 2 decimal
places. Be careful not to mix this up
with their atomic number.
chapter two: it’s all elemental
33
PART one: matter
34
What do you know about atoms and their mass?
1 A student wrote that all the atoms of an element are identical.
Is this correct? Discuss.
3Explain the meaning of the mass number and how this name
arose. Use an example to assist your explanation.
2Explain the meaning of the atomic number and why this
number is used to identify the element. Use an example
to assist your explanation.
4 Using your knowledge of isotopes and a copy of the periodic
table, copy and complete the following table and fill in the
missing gaps.
Isotope symbol
13
6
Isotope name
Atomic number
of element
Number of
protons
Number of
electrons in
uncharged atom
Number of
neutrons
C
oxygen-16
10
30
36
30
How are electrons
arranged in the
atom?
When considering the way electrons are
arranged in an atom, we will consider
their arrangement according to the Bohr
model of the atom. In this model, the
electron shells are named and
numbered from the nucleus outward.
These are shown in Table 2.1, along
with the maximum number of electrons
in each shell. Figure 2.10 is a
diagrammatic representation of how the
electron shells are named.
N
M
L
K
†† Fig 2.10 Each electron shell is named.
29
34
†† Table 2.1 The Bohr model of the atom
Number of shell from
nucleus outwards
1
2
3
4
n
Name of shell
K
L
M
N
–
Maximum number of
electrons in shell
2
8
18
32
2n 2
Table 2.1 shows that the further the
electron shell is from the nucleus, the
larger the number of electrons it can
contain. The maximum number of
electrons a shell can hold is related to
its shell number by a simple formula:
2n2, where n is the number of the shell
from the nucleus.
Bohr also stated that the electrons of an
atom are normally located as close to
the nucleus as possible, as this is a
lower energy state and is more stable.
Therefore, the shells are filled up from
the inside out. However, there is one
restriction. The outermost occupied
shell, that is, the shell containing
electrons that is furthest from the
nucleus, cannot contain any more
than 8 electrons.
The arrangement of electrons in an
atom is termed its electronic
configuration. Let’s consider the
electronic configuration of oxygen:
• Its atomic number is 8, so an
uncharged atom contains 8
electrons.
• The K shell can only hold 2
electrons.
• The L shell holds the other 6
electrons.
• Its electronic configuration is written
as: 2, 6.
Now let’s consider the electronic
configuration of calcium:
• Its atomic number is 20, so an
uncharged atom contains 20
electrons.
• The K shell can only hold 2
electrons.
• There are 18 electrons left to place
in shells. The L shell can only hold 8
electrons.
• There are 10 electrons left to place in
shells. The M shell can only hold 8
electrons.
• The N shell holds the last 2 electrons.
• Its electronic configuration is written
as: 2, 8, 8, 2.
These electronic configurations are often
represented by simple shell diagrams
that show the electron shells as circles.
The electrons are shown in pairs. The
outermost occupied shell of uncharged
atoms is known as the valence shell.
Many substances give off coloured light
when small samples are introduced into
a flame. When this light is seen through
a spectroscope—an instrument that
breaks the light up into its colours—a
pattern of coloured lines is observed.
This pattern is known as an emission
spectrum and is unique for each
element.
Bohr explained this by saying that when
atoms of the elements were given energy
in a flame, the electrons jumped from
their normal shell to one further out
from the nucleus. He described the
electrons as being excited. Because this
higher energy state was unstable, the
electrons then jumped back to their
normal levels in one or more jumps,
almost instantly. For each jump made
by each electron, a certain amount of
energy was given out. This was in the
form of a photon (‘package’) of light of
particular wavelength. Each coloured
2, 6
2, 8, 8, 2
a Oxygen
b Calcium
†† Fig 2.11 The electronic configurations for (a) oxygen and (b) calcium are shown as simple shell diagrams.
line in the spectrum represented one
of these wavelengths of light.
different for each element. In other
words, the energy of the electrons in the
K shell is different for each element,
and so on. Thus, the energy of the light
emitted by each electron as it jumps
back is different for each element.
Each element produces a different
spectrum when excited by heating
because the possible values of energy
for the electrons present are slightly
†† Fig 2.12 The emission
spectrum of hydrogen.
What do you know about how electrons are arranged in the atom?
1 a For the Bohr model of the atom, what is the maximum
number of electrons that the fourth electron shell can contain?
b What letter of the alphabet is used to name this shell?
2 A potassium atom contains 19 protons.
aHow many electrons will be present in an uncharged
potassium atom? Justify your answer.
b What is the electronic configuration of a potassium atom
according to the Bohr model?
c
From the electronic configuration of potassium, it is clear that
electrons do not normally occupy the fifth shell. What could
be done to potassium atoms for electrons to jump into this
shell? Explain.
Element
Atomic
number
Electronic
configuration
Beryllium
9
Magnesium
Neon
2, 8, 3
11
2, 8, 7
Sulfur
Shell
diagram
chapter two: it’s all elemental
35
Big Ideas
2.1
What are elements made of?
Remember
1 Where are each of the following particles found in an
atom and what are their charges?
a protons
b neutrons
c
electrons
2 When an atom is uncharged, what is true of the
number of protons and electrons present?
8 According to the Bohr model of the atom, the
electronic configuration of the uncharged atoms of a
certain element is 2, 8, 8.
a What is the atomic number of the element?
b What element must it be?
c What will be the electronic configuration of
the next element on the periodic table? State
your reasoning.
3 State the symbol for:
a atomic number
b mass number
c the second shell from the nucleus in the Bohr
model of the atom
Analyse
9Element number 52, tellurium, has a relative atomic
mass of 127.6 amu, correct to 1 decimal place. The
mass numbers of its naturally occurring isotopes
are 120, 122, 123, 124, 125, 127, 128 and 130. The
next element, iodine, has a relative atomic mass of
126.9 amu, correct to 1 decimal place. It only has
one naturally occurring isotope, which has a mass
number of 127.
Understand
4 Distinguish between an atom and an element. Give
an example to illustrate your point.
5Explain why the mass numbers of isotopes are exact
whole numbers but the relative masses of all atoms,
except atoms of the carbon-12 isotope, are not exact
whole numbers.
6Element 22 in the periodic table, titanium, has five
naturally occurring isotopes. What will the atoms of
titanium have in common and in what way(s) will they
be different?
Apply
7
235
92
aHow many protons are present in an atom of this
isotope?
bHow many neutrons are present in an atom of this
isotope?
cHow many electrons are present in an uncharged
atom of this isotope?
dOnly 0.7% of the uranium atoms in naturally
occurring uranium exist as this isotope. The
other isotopes present are uranium-234 (0.01%)
and uranium-238 (99.3%). Write the symbols for
these two isotopes.
U is a radioactive isotope of uranium that is used
in nuclear reactors.
a Write the symbol for tellurium-127 and for
iodine-127.
bExplain why the atoms of these two different
elements can have the same mass number.
Evaluate
10 Scientists have had to deduce what it is like inside
the atom from indirect evidence, in the same way
that astronomers have worked out the temperature
and composition of stars. Write a list of the
advantages and disadvantages of using indirect
evidence to develop theories in science.
Create
11Create a poster that shows the models of the atom,
from the original theory that it was a solid sphere,
as proposed by the English chemist John Dalton, to
the Bohr model. Use the Internet to find images of
the scientists involved and place copies onto your
poster. Investigate the year in which each model was
proposed and show a timeline on your chart.
>>CONNECTING IDEAS<<
12Magnesium is element 12 and lead is element 82. Lead ions present in old paint have a net charge of 2+. This means that
they contain 2 more protons than electrons. Likewise, magnesium ions have a net charge of 2+. What is the electronic
configuration of a magnesium ion? Show your reasoning.
chapter two: it’s all elemental
37
PART one: matter
38
2.2
How do scientists classify
elements?
Have you ever wondered about how the periodic table of the
elements is laid out? Columns don’t match and rows start and
end in interesting ways. Why is hydrogen always written first,
but helium, the next element in line, is all the way over on the
other side? With 111 elements officially recognised to date,
and some evidence of more, it is important that we have a
classification system so that we can see the patterns and
relationships between the elements and use the table to make
predictions about properties.
The design of the
periodic table
The periodic table of the elements was
first developed during the 19th century
when chemists realised that if the
elements were listed in order of
increasing atomic weight, repeating
patterns in their properties could be
observed. For example, the English
chemist John Newlands noticed that
every eighth element on the list was
similar—he called this the Law of
Octaves. Dmitri Mendeleev, a Russian
chemist, collected the most data on the
elements and developed a periodic table
that is regarded as the forerunner of the
modern version. He placed elements in
the same vertical column as other
elements with similar properties. Gaps
were left for elements that had not yet
been discovered. From the patterns in
properties he observed, he predicted
the atomic weights and properties of
‘missing’ elements.
When a number of these elements were
discovered, and the predictions were
shown to be correct, chemists knew
that this classification system worked.
But they did not know the reason for
the patterns because they thought the
atom was the smallest particle that
existed.
Since then, many more elements have
been discovered. We now list the
elements in order of increasing atomic
number and know that the reason for
the patterns in their properties is that
there are patterns in their electronic
configurations (Fig 2.13).
Groups and periods
Vertical columns are called groups.
Originally numbered with Roman
numerals, the groups are now
numbered 1 to 18. Elements in the same
group have similar properties.
We now know that this is because,
generally, elements in the same group
have the same number of electrons in
the outermost (valence) shell. For
example, in group 17, the electronic
configuration of fluorine is 2, 7 and of
chlorine is 2, 8, 7. The uncharged atoms
of all other group 17 elements also have
7 electrons in their valence shell.
Hydrogen and helium, however, are
exceptions to this. Hydrogen has
unique properties—no other element is
like it! It was originally placed in group
1, even though all the other elements in
group 1 are metals, simply because its
uncharged atoms have one electron in
their valence shell, like all the other
elements in the group. It has now been
placed alone and is not part of any group.
Helium is placed in group 18. All
members of group 18 except helium have
8 electrons in the valence shell of their
uncharged atoms. Helium atoms only
have 2 electrons. However, this means
that, in all elements in group 18,
including helium, the valence shell of
their uncharged atoms contains as many
electrons as possible. Because helium
also has very similar properties to the
other elements in group 18, it remains
placed there.
Periods are the horizontal rows in the
periodic table. The uncharged atoms of
the elements in the same row are of
similar diameter. We now know that this
is because they have the same number
of occupied electron shells. For example,
the uncharged atoms of the elements in
the third period all have electrons in
the first three shells only—the K, L
and M shells.
1
IA
New designation
Original designation
1
1.01
3
2
3
4
Li
6
Carbon
2
II A
2
Non-metals
13
III A
5
4
Be
B
14
IV A
6
C
He
15
VA
16
VI A
7
N
8
O
17
VII A
9
F
4.00
Helium
10
Ne
6.94
9.01
10.81
12.01
14.01
16.00
19.00
20.18
Beryllium
Boron
Carbon
Nitrogen
Oxygen
Fluorine
Neon
11
12
13
14
15
16
17
Na
Mg
Transition metals
22.99
24.31
Magnesium
3
III B
4
IV B
5
VB
6
VI B
7
VII B
8
Sodium
19
20
21
22
23
24
25
26
K
Ca
Sc
Ti
39.10
40.08
44.95
47.88
Potassium
Calcium
Scandium
Titanium
38
39
Rb
Sr
Y
40
Zr
85.47
87.62
88.91
91.22
Rubidium
Strontium
Yttrium
Zirconium
55
56
Cs
Ba
132.91
137.33
Cesium
Barium
87
7
12.01
Lithium
37
5
C
H
Hydrogen
Atomic number
Chemical symbol
Atomic mass
Name of element
6
1
18
VIII A
Fr
88
Ra
(223)
226.03
Francium
Radium
57
to
71
89
to
103
72
Hf
V
50.94
Cr
52.00
Mn
54.95
Vanadium Chromium Manganese
41
Nb
92.91
42
Mo
95.94
43
Tc
(98)
Fe
9
VII BI
10
27
28
Co
Ta
74
W
75
Re
Si
P
S
Cl
18
Ar
11
IB
12
II B
26.98
28.09
30.97
32.07
35.45
39.95
Aluminium
Silicon
Phosphorus
Sulfur
Chlorine
Argon
29
30
31
32
33
34
35
Cu
Zn
Ga
Ge
As
Se
Br
36
Kr
55.85
58.93
58.70
63.55
65.39
69.72
72.61
74.92
78.96
79.90
83.80
Iron
Cobalt
Nickel
Copper
Zinc
Callium
Germanium
Arsenic
Selenium
Bromine
Krypton
52
53
44
45
Ru
101.07
Rh
102.91
Niobium MolybdenumTechnetium Ruthenium Rhodium
73
Ni
Al
76
46
Pd
Ir
48
Cd
49
In
50
Sn
51
Sb
Te
I
54
Xe
106.4
107.87
112.41
114.82
118.71
121.74
127.60
126.90
131.29
Palladium
Silver
Cadmium
Indium
Tin
Antimony
Tellurium
Iodine
Xenon
78
79
80
81
82
83
84
85
77
Os
47
Ag
Pt
Au
Hg
Ti
Pb
Bi
Po
At
86
Rn
178.49
180.95
183.85
186.21
190.23
192.22
195.08
196.97
200.59
204.38
207.2
208.98
(209)
(210)
(222)
Hafnium
Tantalum
Tungsten
Rhenium
Osmium
Iridium
Platinum
Gold
Mercury
Thallium
Lead
Bismuth
Polonium
Astatine
Radon
104
Unq
(261)
105
Unp
(262)
106
Unh
(263)
107
Uns
(262)
108
Uno
(265)
109
Une
(266)
110
Uun
Mass numbers in parentheses are
from the most stable of common isotopes.
(267)
Metals
57
La
Rare earth elements
Lanthanoid series 138.91
Lanthanum
89
Actinoid series
Ac
58
Ce
59
Pr
60
Nd
61
Pm
140.12
140.91
144.24
(145)
Cerium
Praseodymium
Neodymium
Promethium
91
92
90
Th
Pa
U
93
Np
237.05
62
Sm
150.4
63
Eu
68
Er
69
Tm
70
Yb
71
Lu
168.93
173.04
Erbium
Thulium
Ytterbium Lutertium
(244)
95
227.03
232.04
231.04
238.03
(243)
Actinium
Thorium
Protactinium
Uranium Neptunium Plutonium Americium
96
Cm
(247)
Curium
97
Bk
(247)
162.50
67
Ho
167.26
Am
158.93
66
Dy
164.93
94
157.25
65
Tb
Samarium Europium Gadolimium Terbium Dysprosium Holmium
Pu
151.97
64
Gd
98
Cf
(251)
99
Es
(252)
100
Fm
(257)
Berkelium Californium Einsteinium Fermium
101
174.97
102
103
(258)
(259)
(260)
Mendelevium
Nobelium
Lawrencium
Md
No
Lr
†† Fig 2.13 The periodic table.
What do you know about the design of the periodic table?
1 Briefly outline the contribution Dmitri Mendeleev made to chemistry.
2 Add the missing words:
When the periodic table was first developed, the elements were listed in order of increasing
. They were placed in vertical columns called
had similar
. Now the elements are listed in order of increasing
.
with other elements that
3 Explain why hydrogen was placed in group 1 even though it does not have similar properties to the rest of the group.
4 aName the element that is located in group 17, period 5.
bOutline two facts you can deduce about the electronic configuration of uncharged atoms of this element from its position in
the periodic table.
5 a Write the electronic configuration of the first three elements in group 2 of the periodic table.
bIdentify the pattern in these electronic configurations.
chapter two: it’s all elemental
39
PART one: matter
40
What are the main
types of elements?
The two main types of elements are
metals and non-metals. As well, there
are some elements with properties
in-between those of metals and nonmetals. These are termed metalloids.
Metals
Metals make up nearly three-quarters of
all of the elements. They have many
properties in common. Pure metals are:
• lustrous (shiny)
†† Fig 2.14 Potassium reacting with water produces a
spectacular reaction.
†† Fig 2.15 Magnesium, an alkaline earth metal,
burning.
• able to conduct heat and electricity
• malleable (able to be beaten into a
new shape)
• ductile (able to be drawn into a
wire).
Metals are divided into different groups
according to their position in the
periodic table. We will consider three
of these groups.
Group 1 metals: the alkali
metals
The alkali metals, such as sodium and
potassium, are found in group 1. The
uncharged atoms of all the alkali metals
have just one electron in the valence
shell. The alkali metals have quite a low
melting point, are soft and highly
reactive. They react very strongly—
some violently—with water, producing
hydrogen gas and an alkaline solution.
(An alkali is a soluble base.) The further
you go down the group, the more
violent this reaction.
strongly, producing hydrogen gas and
an alkaline solution. Again, the further
you go down the group, the more
reactive the metal.
Transition metals
The transition metals are found in a
large block of the periodic table that
consists of the ten groups across the
centre (groups 3–12). Many have special
properties that are not shown by group
1 or 2 metals:
The electronic configurations of the
uncharged atoms of the transition
metals across each period are very
similar. For example, the electronic
configurations of the uncharged atoms
of some of the transition metals in the
first row (period 4), known as the first
transition series, are as follows:
• scandium: 2, 8, 9, 2
• titanium: 2, 8, 10, 2
• A small number are magnetic.
• vanadium: 2, 8, 11, 2
• The transition metals gold and
copper are the only metals that are
not silvery in colour.
• zinc: 2, 8, 18, 2
• Many of the transition metals form
coloured compounds.
• Many of the transition metals form
more than one compound with a
Group 2 metals: the alkaline
earth metals
The alkaline earth metals, such as
magnesium and calcium, are found in
group 2. The uncharged atoms of all the
alkaline earth metals have two
electrons in the valence shell. The
alkaline earth metals have quite a low
melting point, are relatively soft and
very reactive, though in general they
are not quite as reactive as group 1
alkali metals. Like the alkali metals,
they also react with water, some
non-metal like chlorine. For
example, iron forms FeCl2 and FeCl3.
†† Fig 2.16 Zinc is a transition metal.
You can see that in the uncharged
atoms of this set of elements, the only
difference in their electronic
configuration is in the second outermost
shell. This shell (the third shell) is
being filled to its maximum capacity.
E XPE RIME NT 2 . 2
Reactivity of metals
Aim
To compare the reactivity of various metals by observing
their reaction with hydrochloric acid.
Remember to wear safety glasses, a lab coat and
latex gloves when you do this experiment.
5 Repeat the above process for the remaining metals.
Results
Metal
Observations
Height of foam
(cm)
Magnesium
Equipment
2 M hydrochloric acid
Detergent
Test tubes and test tube rack
0.5 cm piece magnesium, aluminium, iron, zinc, copper
Steel wool
Ruler
Wooden board
Method
1 Clean the surface of the magnesium with a piece of
steel wool.
Aluminium
Iron
Zinc
Copper
Discussion
• Which metal was the most reactive?
• Which metal was least reactive?
2 Place the magnesium into the test tube.
• Why were the metals cleaned with steel wool first?
3 Add 3 drops of detergent to the test tube.
• Why was detergent added to the test tubes with the
hydrochloric acid?
4 Add 2 cm of hydrochloric acid to the test tube.
Record your observations and the height of the foam
produced in the results table below.
Non-metals
Non-metals, as the name suggests, are
elements that do not show the set of
properties common to all metals. None
are lustrous. None are ductile. A small
number are coloured. Some are brittle.
Also, they have a much larger range of
melting points and boiling points than
the metals. At room temperature, quite
a number of the non-metals are gases
and one is a liquid (bromine), while all
metals except one (mercury) are solids
at room temperature. There are a
number of different groups of nonmetals in the periodic table. We will
consider two of these.
Group 17: the halogens
The halogens, such as fluorine and
chlorine, are found in group 17. The
uncharged atoms of all the halogens
have 7 electrons in their valence shell.
The halogens are mostly known for
their capacity to react with metals to
• What properties would you think the most reactive
metal would also exhibit?
form salts. Indeed, the word ‘halogen’
means salt-forming. Some have
bleaching properties as well.
As you go down the group, their
melting points and boiling points
increase. At room temperature, fluorine
and chlorine are gases, bromine is a
liquid and iodine and astatine are
solids. This is the only group in which
the elements range from gas to liquid to
solid at room temperature. Astatine,
however, is radioactive and very
unstable.
Unlike the metals in groups 1 and 2,
the further you go down this group
of non-metals, the less reactive the
element. Fluorine is the most reactive
non-metal of all and is extremely
dangerous to handle.
Group 18: the noble gases
The noble gases, such as neon and
argon, are found in group 18. The
uncharged atoms of the noble gases
†† Fig 2.17 Fluorine, the most reactive non-metal, is
used to etch glass. It is extremely dangerous to
handle.
chapter two: it’s all elemental
41
PART one: matter
42
have 8 electrons in their valence shell,
except for helium, which has 2. The
noble gases are so-called because they
are all gases at room temperature and
are ‘aloof’ if mixed with other
elements; that is, they are very
unreactive. The first three in the group
do not react with any other element
and have no compounds. It was first
thought that the same was true of
xenon and krypton but, in recent
years, chemists have discovered they
will react with fluorine under certain
conditions. They only have a very
small number of compounds. The
last member of the group, radon, is
very dangerous—not because of any
chemical reactivity but because it is
a radioactive gas.
Metalloids
Metalloids are the small set of
elements along the ‘staircase’ diagonal
boundary between the metals and
non-metals. As might be expected from
this location, they exhibit in-between
properties. Three of them are
semiconductors, which means that
they only conduct electricity in a
certain way under certain conditions.
Examples are silicon and arsenic.
What do you know
about the main
types of elements?
>>FRESH IDEAS<<
Weighing up the costs and benefits
As the saying goes, ‘there are two sides to every coin’. Scientific research is of
benefit to society but that does not mean it is not without its costs. These two
factors must be weighed up against each other to determine the overall cost
to society of the research.
Radiation and medicine
While Marie Curie is arguably the
most famous female scientist, there
is an irony to her death that is really
quite sad. Marie Curie died from
leukaemia after having worked with
the radioactive substances radium and
uranium. Yet a radioactive isotope of
iodine is used to treat thyroid cancer.
Strontium and samarium isotopes are
used to relieve the pain caused by
bone cancers. Radiation is now used
in a number of ways not only to treat
and relieve the symptoms of disease
but also to help diagnose diseases.
Diagnosis covers a wide range of tests,
from fairly routine X-rays to complex
scans and the injection of radioactive
material for nuclear medicine imaging.
In general, radiation therapy involves
delivering a large dose of radiation
to a small area of the body. Radiation
therapy has also been used to treat
coronary artery disease to reduce the
chance of an artery closing.
Marie Curie took a huge risk when she
started investigating the properties
of radioactivity. But where would
modern medicine be without this
information?
1 Quite a number of the elements
have radioactive isotopes. In each
case, it is the nucleus of the atom
that is unstable. Investigate what
kinds of particles and/or rays can be
emitted by radioactive atoms.
2Investigate one radioactive isotope
that is used in medicine. State the
symbol of the isotope and its uses.
Find out how it actually works, in
simple terms. What are the benefits
of using this instead of other
treatments, such as chemotherapy?
What precautions must be taken in
handling this isotope? Does its use
have any side-effects?
3 Where else do we see elements
and chemicals being used to
treat disease? Can you think of
the benefits and risks of such
treatments?
4 What are the side-effects of
radiotherapy and chemotherapy?
Do you think that these ‘costs’
outweigh the potential benefits?
1 What are the main properties of
metals?
2 What is the difference between a
metal and a metalloid?
3Name two properties shown by
some transition metals that are not
shown by group 1 or 2 metals?
4 Predict some of the properties of
the element rubidium and state
your reasoning.
†† Fig 2.18 X-rays use radiation to make images of the bones in the body.
Big Ideas
2.2
How do scientists classify elements?
Remember
Apply
1 What is the name given to the following features of
the periodic table?
a a horizontal row
b a vertical column
c the set of 10 groups from group 3 to group 12
2 State the group number of the following sets of
elements:
a alkaline earth elements
b halogens
c noble gases
d alkali metals
8Only two elements are liquids at room temperature—
bromine and mercury. Bromine is a non-metal and
mercury is a metal. In what ways are these two
liquids likely to appear and behave differently from
each other? Discuss.
Analyse
9 a
3 What is a valence shell?
4 a State the features that elements in the same
group have in common.
b State the features that elements in the same
period have in common.
Understand
5 Suggest why transition metals are much more widely
used than the alkali metals.
6 Give explanations for the following.
aHydrogen was placed in the same group as a set
of metals, even though it is a non-metal.
bHelium was placed in the same group as the
noble gases, even though its uncharged atoms
have a different number of electrons in the
valence shell to those of the other group
members.
7 An inert substance is one that will not react with any
other substance. Originally, group 18 elements were
known as the ‘inert gases’. Suggest why the name
was changed to ‘noble gases’.
Some sodium metal was introduced into a sealed
jar containing chlorine gas. They react to produce
sodium chloride, which is table salt. Would you
expect this reaction to need heat to get it going
or would you expect it to produce heat? Would
you expect it to be a mild reaction or a more
violent one? Justify your answer.
b What two elements would you expect to react
together in the most violent way? Justify your
answer.
Evaluate
10 Before the 1980s, the groups of the periodic
table were numbered with Roman numerals
(see Fig 2.13). Some scientists prefer this version
because the uncharged atoms of the elements
in group III (now 13) have three electrons in their
valence shell, those in group IV (now 14) have four
electrons in their valence shell, and so on. Examine
how the groups of transition metals were numbered
in the old way. Which numbering system do you
think is the most helpful? How can you deduce the
number of electrons in the valence shell from the
new group number?
Create
11 Research one of the elements in group 1, 2, 17 or 18
or a metalloid. Present this information to the class.
>>CONNECTING IDEAS<<
12 The experiment conducted by Ernest Rutherford in which alpha particles were fired at a thin sheet of gold was able to be
performed because gold is a soft metal that can be beaten into very thin sheets, just a few atoms thick, which alpha particles
can pass through. The alpha particles were emitted by a radioactive material that had been discovered just a few years earlier.
Find out more about this experiment. What was the source of the alpha particles? Can alpha particles pass through any
metal? Can any metal stop them? Did Rutherford suffer any long-term effects from his exposure to radioactive materials?
chapter two: it’s all elemental
43