Download The Periodic Table and physical properties (1)

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The Periodic Table and physical properties (1)
In the Periodic Table elements are placed in order of
increasing atomic number. Elements with the same number of
valence electrons are placed vertically in the same group. The
groups are numbered from 1 to 8 (or 0). Some groups have
their own name:
• Group 1 - alkali metals
• Group 7 - halogens
• Group 8 or 0 - noble gases (sometimes also called rare
gases or inert gases).
THE PERIODIC TABLE
, I
-
Group --------~8(=0)1
r=
,J...,2.
3 4 5
2 Li Be
f- I-
BIC
3NaM
l
5 Rb
1
E Csl
I
01
F Ne
AI Si P S CI Ar
Transitionelements
4 K Ca SclTi V Cr
N
6 7 He
Iw Fe Co Ni Cu Zn
Br
,
Elements with the same outer shell of valence electrons are
placed horizontally in the same period. The transition
elements are located between groups 2 and 3.
I
I
ATOMIC RADIUS
~
0
152
8 CD
112
88
©
77
@
@
CD
70
66
64
G 80008
186
o
231
160
143
117
110
The atomic radius is the distance from the
nucleus to the outermost electron. Since
the position of the outermost electron can
never be known precisely, the atomic
radius is usually defined as half the
distance between the nuclei of two bonded
atoms of the same element.
@
104
99
As a group is descended the outermost
electron is in a higher energy level, which
is further from the nucleus, so the radius
increases.
Across a period electrons are being added
to the same energy level, but the number
of protons in the nucleus increases. This
attracts the energy level closer to the
nucleus and the atomic radius decreases
across a period.
114
atomic radius/10-12 rn
133
244
IONIC RADIUS
It is important to distinguish between positive
ions (cations) and negative ions (anions). Both
cations and anions increase in size down a
group as the outer level gets further from the
nucleus.
Cations contain fewer electrons than protons so
the electrostatic attraction between the nucleus
and the outermost electron is greater and the ion
is smaller than the parent atom. It is also smaller
because the number of electron shells has
decreased by one. Across the period the ions
contain the same number of electrons
(isoelectronic), but an increasing number of
protons, so the ionic radius decreases.
Anions contain more electrons than protons so
are larger than the parent atom. Across a period
the size decreases because the number of
electrons remains the same but the number of
protons increases.
Cations
atom
G
186
2.8.1
11 protons
11electrons
ion
Anions
atom
ONa+
@
98
2.8
11 protons
10 electrons
99
2.8.7
17 protons
17 electrons
ion
(3
Ou+
68
•
•
01'-.!a+
OVlg2+
98
65
e
133
133
~~G G
212
ionic radius/1a'2 m
181
2.8.8
17 protons
18 electrons
190
181
G
196
219
_:'
The Periodic Table and physical properties (2)
PERIODICITY
Elements
in the same group
properties.
across a period.
properties
ionization
pattern
by the different
chemical
and physical
of physical
periods
trends can clearly
energies,
MELTING
and chemical
as periodicity.
is 'known
be seen in atomic
electronegativities
and physical
properties
and melting
radii,
ionic
radii,
points.
POINTS
points
depend
both on the structure
the element
5000,----------------------------------------------------------------.
of
and on the
type of attractive
holding
in chemical
The repeating
shown
These periodic
Melting
tend to have similar
There is a change
C
forces
4000
the atoms
together, Using period
3 as an example:
• At the left of the
period
elements
exhibit
metallic
bonding
Ca
(Na, Mg,
AI), which
increases
in strength
as the
number
1000
K
of valence
electrons
increases.
• Silicon
3
in the middle
of the period
4
6
7
8
10
9
has a
atomic
11
12
13
15
14
16
17
18
19
20
number
macromolecular
covalent structure
• Elements
attraction
with
very strong bonds resulting
5, 6, and 7 (P4' S8' and C12) show simple molecular
between
the molecules.
• The noble gases (Ar) exist as monatomic
Within
groups
molecules
with weak van der Waals'
(single atoms) with extremely
point
decreases
down
the group
decreases.
weak forces of attraction
forces of
between
the atoms.
as the atoms become
larger and the strength
of the metallic
I
I
bond
"
M.ptlK
• In group
point.
structures
there are also clear trends:
1 the melting
• In group
in a very high melting
in groups
li
Na
454
371
7 the van der Waals'
attractive
K
336
Rb
312
forces between
the diatomic
Cs
302
molecules
increase
down
the group
so the melting
points
increase.
M.pt!
K
FIRST IONIZATION
ElECTRONEGATIVITY
Electronegativity
is a relative
measure
that an atom has for a shared pair of electrons
is covalently
bonded
to another
the electronegativity
the value
across a period
increases
a Group.
elements
when
incr-eases, so
and decreases
The three most electronegative
are F, N, and O.
1.0
B
C
N
1.5
2.5
3.0
0
3.5
I
F
4.0
Na
CI
0.9
3.0
K
Br
2.8
0..8
and therefore
e.g. for the group
1 elements,
I
2.5
less energy is required
arrangement
extra electrons
protons
energy
the values
(k] mol+)
increase
are filling
in the nucleus
is further
to remove
it,
li, Na and K.
Element:
First ionization
2.0
energy has been given on page 9.
each group as the outer electron
from the nucleus
Generally
Be
ENERGY
of first ionization
The values decrease down
Electron
H
2.1
Li
it
atom. As the size of
the atom decreases
down
The definition
of the attraction
Li
Na
K
2.1
519
2.8.1
494
2.8.8.1
418
across a period.
the same energy
attract this energy
This is because
the
level and the extra
level closer making
an electron,
Element
Na
Mg
AI
Si
P
5
CI
Ar
11
12
13
14
15
16
17
18
2.8.1
2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.72.8.8
494
736
Number
e.g. for the third
it
harder to remove
period.
of
protons
Electron
arrangement
First ionization
energy(kJ
mol-1)
577
786
1060 1000 12601520
Periodicity 13
~}-
3~2 Physical properties
3.2.1
Define the terms first ionizstion energy and elearonegativity.
First ionization energy
The first ionization energy is the minimum energy required to remove one mole of electrons
from one mole of gaseous atoms (under standard thermodynamic
conditions of 25 "C and I arm).
In general:
X(g) --> X'(g) + e-
For example, the first ionization
H(g)-+H-(g)
energy of hydrogen is given by the tollowing equation:
+ e-
D.H = +1310kJmol-!
The amount energy required to carry out this process for a mole of hydrogen atoms is lJ 10
kilojoules,
Atoms of each element have different values of first ionization energy. The factors [hat control
the values of first ionization energy are discussed in. Chapter 12.
dtew\ls ~
.tV ~ \ B Dl P
Tcdloo +j Huv-WODJ
J
Ib/>LCv
COCiks
@20·/D
71t
PERIODICITY
Electronegativity
The electronegativity of an atom is the ability or power of an atom in a covalent bond to attract
shared pairs of electrons to itself. The greater the electro negativity of an atom, the greater its
ability to attract shared pairs of electrons to itself.
Electronegativity value are based on the Pauling scale. A value of 4.0 is given to fluorine, the
most electronegative atom. The least electronegative elements, caesium and francium, both have
an electronegativiry value of 0.7. The values for all the other elements lie between these two
extremes. Note that electronegativity values are pure' numbers with no units.
History of Chemistry
Linus Pauling (1901-1994) was an American chemist who was awarded two Nobel Prizes: Chemistry (1954)
and Peace (1962). His early work was centred on chemical bonding and intermolecular forces. He developed the
concepts of hybridization and resonance (Chapter 14). He also studied biological molecules and correctly proposed
the a-helix and J3-sheetsas.common secondary structures in proteins. However, he incorrectly predicted a triple
helix structure for DNA. Later in his life he began controversial research into the use of vitamin C as an anticancer compound, both as a preventive measure and to treat it.
Trends in the physical properties of the elements in group' 1
and group 7
.
.~.
3.2.2
Describe
and explain
and melting
the trends in atomic radii, ionic radii, first ionization
energies, electronegativities
points for the alkali metals (Li -7 Cs) and the halogens (F -7 I).
Trends in atomic and ionic radii
At the right of the periodic table the atomic radius is defined as half the distance between the
nuclei of two covalently bonded atoms (Figure 3.18). For example, the bond length for chlorine
atoms in a chlorine molecule (distance between two chlorine nuclei) is O.199nm. Therefore the
atomic radius of chorine is x 198 = 99 pm (1 picometre (pm) = Hrl2 m; 1 nanometre (nrn) =
10-9 m). At the left of the periodic table, the atomic radius is that of the atom in the metallattice
(the metallic radius). For the noble gases the atomic radius is that of an isolated atom (the van
der Waals' radius).
The atomic radius of an atom is determined by the balance between two opposing factors:
+
metallic radius
covalent radius
c
G
van der Waals' radius
(for group 0)
Figure
3.18 Atomic
radius
• the shielding effect by the electrons of the inner shellis) - this makes the atomic radius larger.
The shielding effect is the result of repulsion between the electrons in the inner shell and those"
in the outer or valence shell
• the nuclear charge (due to the protons) - this is an attractive force that pulls all the electrons
closer to the nucleus. With an increase in nuclear charge, the atomic radius becomes smaller.
However, when moving down a group in the periodic table, there is an increase in the atomic
radius as the nuclear charge increases (Tables 3.3 and 3.4 and Figures 3.19 and3.2D). This is the
result of two factors:
• the increase in the number of complete electron shells between the outer (valence) electrons
and the nucleus
• the increase in the shielding effect of the outer electrons by the inner electrons.
Moving down a group, both the nuclear charge and the shielding effect increase. However, the
outer electrons enter new shells. So, although the nucleus gains protons, the electrons are not
only further away, but also more effectively screened by an additional shell of electrons .
• ••
• • • • • __•• _•••••••..•
~•••••••
~ • L._ •••.
..I" ••.••
~
••
Physicaf properties
79
300
Fr
Cs
Rb.
K
E
200Q.
Na
en
OJ
'6
Atom
Atomic
number
Atomic
Li
3
152
Na
11
186
K
19
Rb
37
244
C50
5.5
262
Fr
87
270
------------_._--_.
~
radius/pm
-
U
0
-f"
0
~
100-
231
-----
Table 3.3 The variation
D+-L-~L-~r---~--~----~~~--~----~~~
o
10
20
30
40
50
60
70
80
90
Atomic number
of atomic radii in group
1
Figure
3.19 Bar chart showing
the variation of atomic radii in group 1
200-
E
Atomic
Atom
number
Atomic
9
radius/pm
~
58
~
At
Br
OJ
CI
100
o
CI
99
f"
35
114
~
53
133
85
140
17
Br
-
----
At
------------
10
---
20
30
40
50
60
70
80
90
Atomic number
Table 3.4'The variation of atomic radii in group 7
-Figure 3.20 Bar chart showing the variation
of atomic
radii in group 7
Ionic radii for ions of the same charge also increase down a group for the same reason (Tables 3.5
and 3.6). Ionic radii are the radii for ions in a crystalline ionic compound (Figure 3.22).
Ion
Atomic
number
Li+
-----Na+
3
Ion
Ionic radius/pm
Atomic
number
Ionic radius/pm
68
F-
CI-
17
181
Br
35
196
133
9
11
98
K+
19
133
Rb+
37
148
1-
53
219
cs-
55
167
At-
85
No data
Fr+
87
No data
Table 3.5 The variation of ionic radii in group
Table 3.6 The variation
of ionic radii in group 7
1
atomic radii decrease
Figure 3.21 Summary
of trends in periodicity
...
in atomic radii in the
~
periodic table
-
--
-.-
-.---.'
-
-
.
80
PERIODICITY
Na
K
3.22 The relative
sizes of the atoms and
Figure
ions of group 1 metals
Trends in first ionization energy
On moving down a group, the atomic radius increases as additional electron shells are added. This
causes the shielding effect to increase. The further the outer or valence shell is from the nucleus,
the smaller the attractive force exerted by the protons in the nucleus. Hence.zhe more easily an
outer electron can be removed and the lower the ionization energy. So, within each group, the
first ionization energies decrease down the group. This is shown in Table 3.7 and Figure 3.23.
Atom
Atomic
number
First ionization
energy/kJ
3
Li
mol-1
11
494
K
19
418
------
Rb
E
....,
400-
K
Rb
es
e>
Ql
Ql
300-
c
-
0
37
~
402
N
'c
55
Cs
Na
(5
c
- --
U
500
~
0>-
519
Na
..
200-
.Q
376
"P!
u:: 100
Table 3.7 The variation
of first ionization
energy in group 1
O+-L--T~~r---.---Lo----,-~-.o
10
20
30
40
50
60
Atomic number
Figure
3.23 Bar graph showing the variation of first ionization
energy in group
•
Extension:
Effective nuclear charge
An alternative way to account f-orchanges in icnizacion-energtes-is-to use the concept of
effective nuclear charge (Figure 3.24). This is the nuclear charge experienced by the electrons
after taking into account the shielding effect of electrons. For example, in the atoms of group 2
the effective nuclear charge is + 2, which is calculated by adding the charges of the protons and
shielding electrons. However, moving down group 2 the outer or valence electrons are held less
strongly, being further away from the same effective nuclear charge.
shielding electrons in
inner full shells
/~
Figure
3.24 Shielding in
beryllium,
magnesium
Be
e
Mg
~
e
outer electron
Ca
e
and calcium atoms
••••••••••••••••••
__
••
_ ••••••••••••.
.11 •••
.IIe.ll_ •••
1
Physical properties
Trends in electronegativity
Atom
Atomic number
81
Electronegativity
Electronegativity values generally decrease down a
1.0
li
3
group. Clear decreasing trends in electronegativity
Na
11
0.9
can be found in group 1 (the alkali metals) (Table
3.8) and group 7 (the halogens) (Table 3.9).
K
19
0.8
Electronegativity can be interpreted as a measure
Rb
37
0.8
of non-metallic or metallic character. Decreasing
Cs
55
0.7
electronegativity down a.group indicates a
~-.------~.------~--Fr
87
decrease +n-nen-metaliiccharaceer
and an
increase in metallic character.
Table 3.8 The variation of electronegativity in group 1
The decrease in elecrronegativity down groups
1 and 7 can be explained by the increase in
Atom
Atomic number
Electronegativity
atomic radius. There is therefore an increasing
distance between the nucleus and shared pairs of
9
4.0
electrons, Hence the attractive force is decreased.
CI
17
3.0
Although the nuclear charge increases down
Br
35
2.8
a group, this is counteracted by the increased
2.5
53
shielding due to additional electron shells.
The trends in electronegativity can be used to
At
85
2.2
explain the redox properties of groups 1 and 7.
Reducing power decreases down group 1; oxidizing Table 3.9 The variation of electronegativity in group 7
power increases up group 7 (Chapter 9).
Trends in melting
point
• Group 1
The melting points of the alkali metals decrease down the group (Table 3.10 and Figure 3.25}.
Metals are held together in the solid and liquid states by metallic bonding (Chapter 4). Metals are
composed of a lattice of positive ions surrounded by delocalized electrons which move between
the ions. The delocalized electrons are valence electrons shed by the metal atoms as they enter
the lattice.
The melting points decrease down the group because the strength of the metallic bonding
decreases. This occurs because the attractive forces between the delocalized electrons and the
nucleus decrease owing to the increase in distance. The increase in nuclear charge is counteracted
by the increase in shielding.
I
500
Atom
3
li
Na
-
Melting.pointiK
Atomic number
11
-----
400
Na
K
Rb
Cs
300
Fr
Ol
19
337
Rb
37
312
Cs
55
302
Fr
87
300
__ .. ----
0.
371
K
.
y:
:;:,
c
·0
454
U
c
~
200-
Z
tOO-:
.-....•....
0
0
10
20
30
40
50
60
70
80
90
AtomiC number
Table 3.10 The variation of melting point
in group 1
Figure 3.25 The melting points of the alkali metals
• Group 7
In contrast to the alkali metals, the melting and boiling points of the halogens increase down the
group (Table 3.11 and Figure 3.26). This is because as the molecules become large, the artractive
forces between them increase. These shorter-range attractive forces are known as van der Waals'
forces and increase with the number of electrons in atoms or molecules (Chapter 4) .
I
•
- •••••
- ••••••..•••••••••
' ••
, .•
-r- •••••.•
- •••••••••••
-.
82
PERIODICITY
600
~
:g
500-
-6
Atom
Atomic
number
9
--------------
Melting
~ 400-
pointlK
£
:g
54
-----------------
CI
17
172
oc»
Br
35
266
+J
53
387
85
575
c
~
At
Table 3.11 The variation
300-
200100
O+---~--~._----r_~~----~L-~----._----r_~_,_
a
Atomic number
of melting point
in group 7
Figure
3.26 Melting and boiling points of the halogens
Trends in physical properties of elements across period 3
3.2.3
Describe
and explain
the trends in atomic radii, ionic radii, first ionization
energies and electronegativities
for elements across period 3_
Trends in atomic radii
There is a gradual decrease in atomic radius across period 3 from left to right (Table 3.12 and
Figure 3.27). When moving from group to group across a period, the number of protons and
the number of electrons increases by one. Since the electrons are added to the same shell, there
is only a slight increase in the shielding effect across the period. At the same time additional
protons are added to the nucleus, increasing the nuclear charge. The effect of the increase in
nuclear charge more than outweighs the small increase in shielding and consequently all the
electrons are pulled closer to the nucleus. Hence, atomic radii decrease across period 3. The same
effect is observed in other periods.
Atomic
Atom
186
Na
200 -
radius/pm
Mg
160
AI
143
Na
Mg
E
0.
AI
{iJ
:0
Si
'0
Table 3.12 The atomic
radii in period
3
~
S
CI
16
17
o
5i
117
-E
P
110
<
5
104
CI
99
0
-
a
11
12
13
-
14
15
Atomic number
----No data
Ar
p
100
Figure
3.27 Bar graph of the atomic radii in period 3
Trends in ionic radii
The data in Table 3.13 shows the following trends in ionic radii across period 3.
<b
• The radii of positive ions decrease from the sodium ion, Na" to the aluminium ion, AP+.
• The radii of negative ions decrease from the phosphide ion, P3- to the chloride ion, Cl.
• The ionic radii increase-from the- aluminium ion, Ap· to the-phosph-ide-i0n, P3-.
Element
Sodium
Na+
Ion
------------------------Ionic radius/pm
98
..-.-
Magnesium
Aluminium
Silicon
Mg'·
A13.
(5i4•
65
45
(42and 271)
and 5i4-)
Phosphorus
Sulfur
_
\
Chlorine
p3-
5'-
CI-
212
190
181
Table- 3-. HThe-ionic
-
O.
radii in period 3-
.
Physical properties
83
The data for the silicon ions are calculated values. Silicon does not' form simple ions and its
bonding is covalent.
Isoelectronic
species
Isoelectronic species are atoms and ions that have the same number of electrons. For a specific
number of electrons, the higher the nuclear charge, the greater the forces of attraction between
the nucleus and the electrons. Hence, the smaller the atomic or ionic radius.
Ions of sodium, magnesium and aluminium are isoelectronic species (Table 3.14). The nuclear
charge increases from .the sodium ion to the aluminium ion. The higher nuclear charge pulls. all
the electron shells closer to the nucleus. Hence, the ionic radii decrease.
Similarly, the nuclear charge increases from the phosphide ion to the chloride ion. The higher
nuclear charge causes the electron shells to be pulled closer to the nucleus. Again, the ionic radii
decrease (Table 3.15).
Species
Nuclear
charge
Number
of electrons
Ionic radius/pm
Table 3.14 Atomic
----
Na+
Mg2+
A13+
+11
+12
+13
Nuclear
charge
10
10
10
Number
of electrons
98
65
45
Ionic radius/pm
Species
data for sodium, magnesium and
aluminium ions
p3-
52-
CI-
+15
+16
+17
18
18
18
212
190
181
Table 3.15 Atomic data for phosphide,
sulfide and
chloride tans
The large increase in size from the aluminium ion to the phosphide ion is due to the presence
of an additional electron shell. This causes a large increase in the shielding effect and as a result
the ionic radius increases.
Trends in first ionization
energy
The first ionization energies of the elements in period 3 are listed in Table 3.16. The general
trend is an increase in first ionization energy across the periodic table. When moving across
a period from left to right the nuclear charge increases but the shielding effect only increases
slightly (since electrons enter the same shell). Consequently, the electron shells are pulled
progressively closer to the nucleus and as a result first ionization energies increase.
Element
Sodium
Magnesium
..-- -----------------First ionization
494
736
-
Aluminium
~
energy/kJ
Silicon
577
786
Phosphorus
------1060
Sulfur
Chlorine
------
1000
1260
mol-'
--------Table 3.16 First ionization
energies for the elements in period 3
However, the increasein first .ionizationenergy is not uniform and there are two decreases between magnesium and aluminium and between phosphorus and sulfur. These decreases can
only be explained by reference to sub-shells and orbitals (see Chapter 12).
Comparing electronegativity values
3.2.4
Compare
the.relative electronegativity
values of two or more elements based on their positions in the
periodic table.
The electronegativities of the elements in period 3 are listed in Table 3.17. The general trend is
an increase in first ionization energy across the periodic table. When moving across a period from
left to right the nuclear charge increases but the shielding effect only increases slightly (since
electrons enter the same shell). Consequently, the electron shells are pulled progressively closer
to the nucleus and as 'a result electronegativity values increase.
Element
Electronegativity
Sodium
0.9
Aluminium
Magnesium
Silicon
1.5
1.2
Phosphorus
1.8
2.1
Chlorine
Sulfur
2.5
3.0
------------------Teble 3.17 Electronegativity
••••••.•••••••••••••.•
-.~
••
:-.
•.•••••
values for tile elements in period 3
- •••••
-."WI
••
"WI •
84
PERIODICITY
Generally, the electronegativity
values of chemical elements increase across a period and
decrease down a group (Figure 3.28). This observation can be used to compare the relative
electronegativity
values of two elements in the periodic table. To do this, find the-positions of
the elements in the periodic table. Then simply see which one is further lip and [Q the right; that
is the more electronegative
element (Figure 3.29). The further apart the two elements are in the
periodic table, the larger the difference will be in their electronegativities.
This is important in
determining the type of bonding between the two elements (Chapter 4).
jffi
P
least
electronegative
Figure 3.28 Trends in eJectronegativity for s- and p-block
elements
•
Extension:
Ge-
i
most
electronegative
As-
Figure 3.29 Relative values of electronegativity of elements
in the periodic table
Diagonal relationships
Elecrronegativity
increases across a period and decreases down a group. This
results in what are known as diagonal relationships, where a pair of elements
have similar chemical properties. The most important pairs are lithium
and magnesium, beryllium and aluminium, and boron and silicon.
History of Chemistry
Dimitri Mendeleev was born in 1834 in Tobolsk, Siberia, the youngest of 17 children.
When Dimitri was 13 years old, his father died and his mother's glass-making factory
burnt down. In 1849 the family relocated [Q St Petersburg (formerly Leningrad) and he
later became Professor of Chemistry at the University of Sr Petersburg. In 1862 he married
Feozva Nikitichna Leshcheva. This marriage ended in divorce and in 1882 Mendeleev
married one of his students, Anna Popova. He was dismissed from the University in 1890
for supporting the causes of students against the authorities. In 1893 he was appointed the
Director of the Bureau of Weights and Measures and helped to formulate new standards
for measures such as mass and length. Mendeleev was nominated for the 1906 Nobel Prize
in Chemistry, but narrowly lost to Frenchman Henri Moissan, who had isolated fluorine.
He probably would have been awarded the 1907 Nobel Prize in Chemistry, but died early
in 1907 from influenza. Nobel Prizes cannot be awarded posthumously (after death).
The Periodic Table
Periods and groups
Mendeleyev
is said to ha e made
his discovery
after
a
0
dream.
When he awoke he set out his
chart in virtually
its final form. He
enjoyed
a form of patience
playing
(solitaire) and wrote the properties
of each element
arranged
on cards which
he
into rows and columns.
If you have visited a large supermarket you will appreciate the importance of
a classification system. Similar productsare grouped together to help you find
what you want. In the same way a chemist knows what type of element to find in
different parts of the Periodic Table. The elements are placed in order of increasing
atomic number (Z), which we now know is a fundamental property of the element
- the number of protons in the nucleus of its atoms. As there are no missing
atomic numbers we can be confident that the search for new eIements in nature is
over.
The only way.to extend the Periodic Table is by making elements artificially. Today
there are over 110 elements recognized by the International Union of Pure and
Applied Chemistry (IUPAC). The columns of the table are called groups and the
rows periods.
'j
IUPAC is an international,
governmental
membership
which
0>
non-
body with a
. 1
made up of chemists
has the aim of fostering
worldwide
communication
The position of an element is related to the electron arrangement in its atom. The
element sodium, for example, is in Period 3 as it has three occupied energy levels,
and in Group 1 as there is one electron in the outer shell (Figure 3.1).
in
chemistry.
Figure3.l
The Periodic Table. The
'island' of elements
In the IB Data booklet Periodic Table the main groups are numbered from 1 to 7,
with the last column on the-far right labelled '0'. The gap between Group 2 and
Group 3 is filled by transition elements from the fourth period onwards .
Group
~
2
from Ce to Lu and
3
level.
r~Ci.J -'t-Be•
2
5
6
7
0
.----,-
GJ
from Th to Lr is of little interest at this
4
He
5
B
6
C
7
8
N
0
9
F
10
Ne
.'
f"l'~H "'1'12
3
u
The rows in the Periodic Table
are called periods. The period
number gives the number of
occupied etectron shells.
0
o 4
.~
0..
Na
fFP19
K
Rb
"""55
Sr
56
6
Cs ~Sa
7
Fr
87
faa
I.
IS
AI_ 5i - P
21
-€a Sc
1~.37 -~ 38
5
13
-l'V1g
'-.:t~20
39
Y
57
22
Ti
40
23
V
41
89
42
Zr Nb Mo
72
104
74
73
La Hf Ta
Ra Ac
24
105
Rf Db
25
Cr Mn
W
106
Sg
43
26
Fe
44
27
Co
28
29
45
.6
47
Tc Ru Rh Pd Ag
75
75
Re Os
107
Bh
loa
30
Ni Cu Zn
n
78
79
48
Cd
SO
Ir Pt Au Hg
109
Hs Mt
110
III
112
31
Ga
49
32
Ge
16
S
33
34
As. Se
50
51
52
In Sn Sb~~Te
81
11
In
82
Pb
11.
83
84
Bi _ Po
115
116
17
CI
18
Ar
35
36
Br Kr
53
I
54
Xe
85
117
58
group number gives the number
of electrons in the outer shell.
Visit an interaaive
Periodic Tagle.
Now go to
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insert the express code 4402P and
Ce
90
Th
59
_60
61
Pr Nd Pm
91
Pa
92
U
93
Np
62
Sm
94
63
64
65
Eu Gd Tb
95
96
Pu Am Cm
97
Bk
66
67
Dy Ho
98
Cf
99
68
69
Er Tm
100
101
Es Fm Md
-
70
Yb
102
No
71
Lu
103
Lr
The electronic configuration gives a more fine-tuned description as the block
structure of the Periodic Table is based on the electron sub-levels of the atom.
The position of an element in the Periodic Table is based on the sub-level of the
highest-energy electron in the ground-state atom.
dick on this activity.
• Challenge
:~-'
::,,-:
~ :-,,'-
yourself:
~amefrom
Four elements
a small town
'_0: :'_:0':2 Stockholm.
Try to
118
Os Rg Uub Uut Uuq UUR Vuh Uus Vuo
The columns in the Periodic
Table are called groups. The
86
At Rn
The table below shows the relationship between the period and group of an
element and its electron arrangement. The group number gives the number of
electrons in the outer energy level (valence electrons). The period number gives
the number of occupied energy levels. The electron arrangement of the noble
gases fit this pattern if they are considered to have 0 electrons in their outer shell.
Helium, for example can be considered to have the electron arrangement: 2, O.
Element
I
Period
I
helium
1
lithium
carbon
Electron
arrangement
a
I
I
2
1
J
2, 1
I
2
2
4
I
aluminium
3
3
chlorine
3
7
2,8,7
4
1
4
2
I
I
calcium
I
!
Electron
configuration
I
1S2
I
,
-0
.
152
i -.
1S22s22p2
2,4
2,8,3
potassium
I
I
I
Group
I
2,8,8,.1
2,8,8,2
.
1s22s22p63s23p
1s22s22p63s23ps
1
1s22s22p63s23p64s
The discovery of the elements
an international
illustrated
endeavour.
was
This is
by some of their names.
Some derive from the place where
they were made, some derive from
the origins of their discoverers
1
and
some derive from the geographical
1s22s22p63s23p64s2
origins of the minerals from which
they were first isolated. The Periodic
Table of chemical
The number of electrons in the outer shell of elements with higher atomic
numbers can be deduced from the group number of the element.
L
throughout
the world.
Pi
@
How many electrons are in the outer shell of iodine?
Solution
Find the element in the Periodic Table. It is Group 7 so it has seven electrons in its
outer shell.
1
hangs
classrooms
and in science laboratories
.
Worked example
elements
in front of chemistry
Use the IB Periodic Table to identify
Element
the position
I
of the following
Now go to
:;;-:
www.heinernann.co.uk/hotlinks,
"
insert the express code 4402P and
click on this activity.
elements.
I
Group
Period
See a Chinese Periodic Table.
t'
helium
chlorine
barium
francium
2
Phosphorus
is in Period 3 and Group 5 of the Periodic Table.
(a) Distinguish
between
(b) State the electron
the terms period and group.
arrangement
of phosphorus
and relate it to its position
in the Periodic
Table.
3
How many valence (outer shell) electrons
are present in the atoms of the element
with atomic
numberSl?
Physical properties
The elements in the Periodic Table are arranged to show how the properties of the
elements repeat periodically.
"This periodicity of the elements is reflected in their physical properties. The
. atomic and ionic radii, electro negativity and ionization energy are of particular
. interest as they explain the periodicity of the chemical properties.
'Science is built of facts the way
a house is built of bricks: but an
accumulation
of facts is no more
science than a pile of bricks is a
house' .
(H. Poincare)
Do you agree with this description
The concept of effective nuclear charge is helpful in explaining trends in both
physical and chemical properties.
of science?
-
~~:~~iCiry
.
~~----
"
_
QI~.(-------O
Effective nuclear charge
Ato
The cc
in the
A
The nuclear charge of the atom is given by the atomic number and so increases by
one between successive elements in the table, as a proton is added to the nucleus.
The outer electrons which determine many of the physical and chemical properties
of the atom do not, however, experience the full attraction of this charge as they
are shielded from the nucleus and repelled by the inner electrons. The presence
of the inner electrons reduces the attraction of the n~?leus for the outer electrons
(Figure 3.2). The effective charge 'experienced' by the outer electrons is less than
the full nuclear charge.
C\
attraction
Figure 3.2 An electron in the
hydroqen atom experiences the full
attraction of the nuclear charge, but in
a many-electron atom the attraction
for the nucleus is reduced as the outer
electron is repelled by inner electrons.
descrij
atomi
(Figur
from t
Table l
decrea
the Gr
Consider, for example, a sodium atom as shown in Figure 3.3. The nuclear charge
is given by the atomic number of element. The outer electron in the third energy
level is, however, shielded from these 11 protons by the 10 electrons in the first and
second energy levels.
electrons in the inner
two energy levels shieldthe outer electron
Figure 3.3 The outer electron is
shielded from the nucleus by the inner
electrons.
r,-------A----"
See different Periodic Table formats.
Now go to
www.heinemann.co.uklhotlinks,
The at.
(given.
the Pel
insert the express code 4402P and
cOckon thfs activity.
Consider the
first
four elements in Period
Element
3
as shown in the table below.
Na
Mg
Af
Si
11
12
13
14
2,8,1
2,8,2
2,8,3
2,8,4
.'
Nuclear charge
Electron
arrangement
~.Elem
~
All the
nucleu
a gene:
A chlo:
As a period is crossed from left to right, one proton is added to the nucleus and
one electron is added to the outer electron shell. The effective charge increases
with the nuclear charge as there is no change in the number of inner electrons.
The changes down a group can be illustrated by considering the elements in
Group 1 as shown in the table below.
Element
The effective nUEieafcnarge
experienced by an atom's outer
electrons increases with the
group number ofthe element.
It increases across a period but
remains approximately the
s.a.medcwn a group.
Nuclear charge
4
(a)
(b)
Electron arrangement
Li
3
2, 1
loni'
Na
11
2,8,1
The at,
K
19
2,8,8,1
As 'i¥e descend the group, the increase in the nuclear charge is largely offset by
the increase in the number of inner electrons; both increase by eight between
successive elements. The effective nuclear charge experienced by the outer
electrons remains approximately the same down a group.
Atomic radius
The concept of atomic radius is not as straightforward as you may think. We saw
in the last chapter that electrons occupy atomic orbitals, which give a probability
description of the electrons' locations, but do not have sharp boundaries. The
atomic radius r is measured as half the distance between neighbouring nuclei
(Figure 3.4). For many purposes, however, it can be considered as the distance
from the nucleus to the outermost electrons of the Bohr atom.
Table 8 in the IE Data booklet shows that atomic radii increase down a group and
decrease across a period. To explain the trend down-a group consider, for example,
the Group 1 elements as shown in the table below.
--"
I
Electron
Element
.
-
-
I 152
I 186
0
2,1
2
Na
2,8, 1
3
K
2,8,8,1
4
231
Rb
2,8,8, ..,1
5
1244
0
262
0
2,8,8, ..,..,1
The atomic radius r is
measured as halfthe distance between
neighbouring nuclei.
Radiusl10-2 m
Li
Cs
•••Figure 3.4
..••...
No. of occupied
shells
- arrangement
-2r--
6
0
0
@ The atomic radii of the noble-gases
~-
are not given in Table 8 of the 18
Data booklet. Their inter-nuclei
distances are difficult to measure as
noble gases do not generally bond
to other atoms.
The atomic radii increase down a group, as the number of occupied electron shells
(given by the period number) increases. The trend across a period is illustrated by
the Period 3 elements as shown below.
-I
. Element
radius/10-12
Atomk
m
Na
Mg
AI
Si
P
S
CI
Ar
186
160
143
117
110
104
99
-
AIfthese elements have three occupied energy levels. The enracnon between the
nucleus and the outer electrons increases as the nuclear charge increases so there is
a .general decrease in atomic radii across the period.
A chlorine atom has a radius that is about half that of a sodium atom.
4
(a) Explain what is meant by the atomic radius of an element
(b) The atomic radii of the elements are found in Table 8 of the 18Data book.
(i) Explain why no values for atomic radii are given for the noble gases.
(iil Describe and explain the trend in atomic radii acrossthe Period 3 elements.
Ionic radius
The atomic and ionic radii of the Period 3 elements are shown in the table below.
;_,.:0
Na
Mg
AI
Si
P
S
(I
radltl;St
186
160
143
117
110
104
99
42 (Si4+);
271(Si4+)
212
190
181
(P3-)
(52-)
(CI-)
Element
Atomic
_t-o-12 m_'c:)~~~G:-;_
. IMIc radiu~l';
1O"'lhn
-;:.-.-
98
65
45
(Na")
(Mg2+)
(AI3+)
~\~
I
Five trends can be identified.
• Positive ions are smaller than their parent atoms.
The formation of positive ions involves th~ loss of the outer shell. Na, for
example, is 2, 8,1 whereas Na" is 2, 8.
• Negative ions are larger than their parent atoms.
The formation of negative ions involves the addition of electrons into the outer
shell. Cl for example is 2, 8, 7 and Cl- is 2, 8, 8. 'Fhe increased electron repulsion
between the electrons in the outer shell causes the electrons to move further
apart and so increases the radius of the outer shell.
• The ionic radii decrease from Groups 1 to 4 for the positive ions. The ions Na",
Mg2+,.Al3+ and 5j4+ all have the same electron arrangement 2, 8. The decrease
in ionic radius is due to the increase in nuclear charge with atomic number
across the period. The increased attraction between the nucleus and the
electrons pulls the outer shell closer to the nucleus.
• The ionic radii decrease from Groups 4 to 7 for the negative ions. The ions
5i4-, p3-, 51- and cr have the same electron arrangement 2,8,8. The decrease
in ionic radius is due to the increase in nuclear charge across the period, as
explained above.
The positive ions are smaller than the negative ions, as the former have
.only two occupied electron shells and the latter have three. This explains
the big difference between the ionic radii of the 5i4+ and 5i4- ions and the
discontinuity in the middle of the table.
Two ge
.• 10m
chill
nuc
• Ioni
elec
the
elec
red,
The sn
levels i
elemei
energ:>
energrernov
occur
Then:
trend:
electn
• The ionic radii increase down a group as the number of electron shells increases.
Thee
attrac
meas.
Worked example
Describe and explain the trend in radii of-the following ions:
02-, F-, Ne, Na + and Mg2+.
The following animation illustrates
.atomic and ionic radii.
bond.
Solution
The ions have 10 electrons and the electron arrangement 2, 8. The nuclear charges
increase with atomic number: 0: Z = +8, F: Z = +9,Ne: Z = + 10, Na: Z = + 11
and Mg: Z = + 12. The increase in nuclear charge results in increased attraction
between the nucleus and the outer electrons. The ionic radii decrease as the atomic
number increases.
Now go to
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An e
ongu
the TI
•
El
ill
m
The first ionization energy
First ionization energies are a measure of the attraction between the nucleus and
the outer electrons. They were defined in Chapter 2 (page 55), where they provided
evidence for the electron configuration' of the atoms of different elements (Figure 3.5).
of electrons from one mole of
gaseous atoms.
Figure 3.5 First ionization energies of
the first 20 elements.
•B
Ionization energies
of an element is the energy
required to remove one mole
2500
~
'I
0
E
-.
---
Ne
2000
1500
c
'"
c
0
:p
to
The
least
to th
high
Alth
sam
and
.><:
>en
Qj
t1:
H
Be
1000
N
'c
.2
500
V>
.E
Li
K
0
element in order of atomic number
Cor
thel
~-------------------------------------------------.
Two general trends can be identified from Figure 3.5.
• Ionization energies increase across a period. The increase in effective nuclear
charge causes an increase in the attraction between the outer electrons and the
nucleus and makes the electrons more difficult to remove.
• Ionization energies decrease down a group. The electron removed is from an
electron shell furthest from the nucleus. Although the nuclear charges increase,
the effective nuclear charge is about the same, owing to shielding of the inner
electrons, and so the increased distance between the electron and the nucleus
reduces the attraction between them.
The small departures from these trends provide evidencefor division of energy
levels into sub-levels as discussed in Chapter 2 (page 68). Thus, the Group 3
elements, with the electron configuration ns2 np\ have lower first ionization
energies than Group 2 elements, with the configuration ns2, as p orbitals have higher
energy than sorbitals, The drop between Groups 5 and 6 occurs as the electron
removed from a Group 6 element, unlike a Group 5 element, is taken from a doubly
occupied 2p orbital. This electron is easier to remove as it is repelled by its partner.
The trend in ionization energy is the reverse of the trend in atomic radii. Both
trends are an indication of the attraction between the nucleus for the outer
electrons.
Electronegativity
Theelectronegativity
of an element is a measure of the ability of its atoms to
attract electrons in a covalent bond. It is related to ionization energy as it is also a
measure of the attraction between the nucleus and its outer electrons - in this case
bonding electrons.
An element with a high electronegativity has strong electron pulling power and
an element with a low electronegativity has weak pulling power. The concept was
originally devised by the American chemist Linus Pauling and his values are given in
the IE Data booklet. The general trends are the same as those for ionization energy.
r-:-:<
.~
~£tectronl!gativityiS'the·ability
Wof an atom to attract electrons
in a covalent bond.
E~
'f~--
.•.:.-;.
• Electronegativity increases from left to right across a period owing to the
increase in nuclear charge, resulting in an increased attraction between the
nucleus and the bond electrons.
•
Electronegativity decreases down a group. The bond electrons are furthest from
the nucleus and so there is reduced attraction.
The most electronegative element is on the top right of the Periodic Table and the
least electronegative element on the bottom left. As the concept does not apply
to the Group 0 elements which do not form covalen~ bonds, Pauling assigned the
highest value of 4.0 to fluorine and the lowest value to of 0.7 to caesium.
Although the general trends in ionization energy and electronegativity are the
same, they are distinct properties. Ionizati~t'energiescan
be measured directly
and are a property of gaseous atoms. Electronegativity is a property of an atom in
amolecule and values are derived indirectly from experimental bond energy data.
Melting points
Comparisons between melting points of different elements are more complex as
they depend on both the type of bonding and the structure (Chapter 4). Trends
@
,~
Linus Pauling has the unique
distinction
of winning
two
unshared
Nobel Prizes - one for
chemistry
in 1954 and one for
peace in 1962. His Chemistry
Prize was for improving
understanding
our
of the chemical
bond and his Peace Prize was
for his campaign
weapons
against nuclear
testing.
\\.0
~.
down Groups 1 and 7 can.however, be explained simply, as the elements within
each group bond in similar ways. Trends in melting points down Group 1and
Group 7 are shown in the table below.
-Seecaesium melt.
Now go to
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insert the express code 4402P and
click on this activity.
L
Element
I
Melting pointCK)
I
Element
I
L
Li
I
454
I
Fz
i
54
I
I
266
I
I
i
Na
K
I
Rb
I
Cs
I
371
I
I
337
I
I
I
312
I
I
302
C~i~:'
Brz
12
Atz
Melting·poihi
r
=
.!.
-
9
172
387
I
575
Melting points increase down Group 7. The elements have molecular structures
which are held together by van der Waals' intermolecular forces. These increase
with the number of electrons in the molecule.
Melting points generally rise across a period and reach a maximum at Group 4.
They then fall to reach a minimum at Group G.In Period 3, for example, the
bonding changes from metallic {Na, Mg and Al) to giant covalent (Si) to weak
van der Waals' attraction between simple molecules (P 4' 58' C12) and single atoms
(Ar) (Figure 3.6). All the Period 3 elements are solids at room temperature except
chlorine and argon.
at room temperature and
. atmospheric pressure.
"
(K)
Melting points decrease down Group 1. The elements have metallic structures
which are held together by attractive forces between delocalized outer electrons
and the positively charged ions. This attraction decreases "lith distance.
Only two elements are liquids
8
C~
D
I
-,
-
. ,.
The cr
, c
0;,"
,
arrang
PIOt;:;
alkali :
the ha
prope:
in see:
Gro
aloof:
Figure- 3.6 The melting points show
4500
~
a periodic pattern as the bonding
c
4000
changes from metallic, to giant
covalent, to simple molecular.
3500
::.:::
."·0
c
flo
Ol
Use a database from the Internet
B
c:
1500
500
Now go to
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end'to science?Could we reach a
:point where everything important
iha sCientiflc sense is known?
-
Si
Be -
•
Tt
ionize
Li
N 0
F Ne
0
element in order of atomic number
insert the express code 4402P and
click on this activity,
scheme. Could there ever be an
Tn-
1000
spreadsheet.
go, but it is expected that;;" new
elements will fit into the current
•
•
2500
Qj
E
in physical properties using a
3000
2000
:w
to investigate trends and variations
No one knows how high the
atornrc number of the elements will.
:\lend
T':
5
Explain why sulfur has a higher melting point than phosphorus.
6
Which physical property generally increases down a group but decreases from left to right across
a period?
A melting point
B eleoronegativity
-C ionization energy
A
W
octet.
noble
5 to:their :
mid.:i
D atomic radius
7
The elements in the Periodic Table are arranged in order of increasing:
A
relative atomic mass
B ionic radii
C nuclear charge
D· ionization energy
Grc
All
±
areu:
then
3 ~ Periodicity
. pblock
3A
5E1
7
2p B c
3p AI Si
4p Ga Ge
'5p- In- Sn
i
0
'He
.'
N 0
P S
As Se
Sb :roe
F Ne
(I Ar
Br
Kr
·1· Xe
6p Ti Pb Bi Po At Rn
~;'f J)lock
<..
i'anthanides 4f rr::Tn::TiO:iTn:==r;=:cr;:~'i:T.:;:T,-;--.-"",;--r"":'~
;J0;inideS. 51 ITc'k+i-i'f;;i:::F..'~::t.=::t-;tt~;::+'~;..;.:.;t~~
«
Figure 5 The Periodic Table rearranged into blocks to show the sub-levels (HL only).
The Periodic Table and physical properties
Ionization energies
Periodicity refers to the repeating pattern of physical and chemical
properties that is seen at regular intervals in the Periodic Table. A
graph of the first ionization energies of the elements against atomic
number illustrates periodicity very clearly (Figure 6) .
.tle
~~em\S~'i
Col)rse ~rnrCH1 ron
~ 2000
~
<,
10 DIp 10V\."o. Prv j V(I. t1'I me
Kr
1500
60
~ 1000 H
'"
.C-
o 500
.~
K
'c
.9
~
'"
0
}JeuS$ @ 2001
~~
Rb
i---,-----,------,-------r-----.,..0
10
20
30
40
50
atomic number
Fig.ure 6 First ionization energies plotted against atomic number.
As we saw in Chapter 2, the first ionization energy of an element
refers to the energy required to remove one electron from an atom
of the element in the gaseous state. It is measured in kilojoules per
mole.
.
M(g) ~ M+ (g)
i
•
t
J
+ e:
The elements in Group 1 (the alkali metals) have the lowest values
in each period. As we descend Group 1 from lithium to caesium the
values decrease, because the outer electron is further away from
the nucleus and is therefore already in a higher energy level, so less
energy is required to remove it. As each energy level is successively
filled with electrons an equal number of protons are also being
added to the nucleus. As each electron is added the level is attracted
closer to the nucleus and therefore it becomes lower in energy, so
that ionization energies generally increase across a period .
3
D
Periodicity
The exceptions to the general increase (e.g. boron and oxygen) are
due to the presence of sub-levels within the main energy levels. The
elements with the highest first ionization energy in each period are
the Group 0 elements, the noble gases (He, Ne, Ar, Xe and Rn).
Electronegativities
A covalent bond is formed when one or more pairs of electrons
are shared between two atoms. This is explained fully in Chapter
4 on bonding. When an atom is covalently bonded, its relative
ability to attract a bonding pair of electrons to itself is known as -"electronegativity. Electronegativity is a relative value, not an
absolute value, and so there are different scales of electronegativity
in use. The values used by the IE are attributed to the North
American chemist Linus Pauling (1901-1994) (Figure 7).
He-
H
2.1
Li
1.0
Be
1.5
B
2.0
C
2.5
N
3.0
0
3.5
Na
0.9
Mg
1.2
AI
1.5
Si
1.8
P
2.1
S
2.5
K
0.8-
Ca
Sc
Ti
V
1.0·
1.3
1..5-
i.s
As
Ga Ge
1-.6 1-.8· 2.0
Cr
1.6
Mn Fe
1-.5 L8-
Co
Ni
1-.8 ~.8
Cu
1-.9
Zn
1.6-
~
Ne
4.
CI
3.0
Ar
Se
Br
Kr
2.4
2.8
Xe
Rb
0.8
I
2.5
Cs
0.7
At
2.2
L'$
Rn
Figure 7 Pauling's scale of electronegativities
Electronegativity values also exhibit periodicity. Apart from helium
and the other Group 0 elements, which have no values because
the noble gases either form no compounds or form them only with
difficulty, the electronegativity values decrease down each group. The
values increase across each period. The most electronegative element
is fluorine and the least electronegative element is caesium. Metals
in Groups 1 and 2 are sometimes described as electropositive
elements, as they have relatively low electronegativity values,
whereas non-metals on the right of the Periodic Table in Groups 5, 6
and 7 are often described as electronegative elements. As you will
see in Chapter 4, it is generally the difference in electronegariviey
between two bonding atoms that is important, rather than the
precise value that each atom has. The value of the electronegativity
is related to the size of the atoms. As the atoms become smaller,
the nucleus will tend to attract a pair of electrons more strongly.
However, it cannot be as simple as this: apart from helium, which
has no value, hydrogen is the smallest atom, and on this basis it
might be expected to have the highest value.
Some books explain this by referring to the effective core charge.
This (also known as the effective nuclear charge) is the negative
charge of the Inner electron shells plus the pos-itive charge due to
the protons in the nucleus. Thus the core charge of fluorine is +7,
because there are nine protons in the nucleus, two of which are
"cancelled out" by the inner shell of two electrons. Chlorine, too,
will have a core charge of +7, but because fluorine is smaller it will
have a higher electro negativity. This is not really much better as
an explanation: for example, it does not explain why oxygen, with
3
core charge of +6, has a higher value than chlorine, with a core
Garge of +7, unless size is by far the predominating factor.
covalent radius
I
2
atom X
-,
.. ~
Atomic radii
.
,~'
Logically, the atomic radius should be defined as the distance from
:.!ie centre of the nucleus to the outermost electron. However,
:nis is impossible to measure directly, partly because the precise
position of an electron can never be known at a fixed point in time
(Heisenberg's uncertainty principle: see Chapter 2). What can
be measured, by a technique known as X-ray diffraction, is half
the distance between two nuclei of bonded atoms (see Chapter 4).
If they are bonded covalently it is known as the covalent radius of
an atom, whereas if metallic bonding is involved it IS the metallic
radius. See Figure 8.
The covalent radius is defined as half the distance between the
nuclei of two identical atoms that are covalently bonded together. If
the radius is doubled it is equal to the bond length between the two
atoms.
As a group is descended in the Periodic Table, the atomic radius
increases. For example, the atomic radius of the atoms in Group 1
increases from 0.152 run for lithium to 0.262 nm for caesium. This
increase in atomic radius upon descending a group occurs because
the outer electron is in an energy level that is progressively further
away from the centre of the nucleus: for example, Li 2.1, Na 2.8.1
and K 2.8.8.1. See Figure 9.
•
80 © ® @
r::
y 00000
152
o
112
88
77
70
66
160
143
117
110
104
atomic
radius / x I0-12 m.
®
64
@
99
G
114
231
8
o
244
8
133
262
Figure 9 Atomic radii of the elements.
The atomic radius decreases across a period. As the number of outer
electrons increases upon moving across a period, so the number of
protons in the nucleus also increases. This increase in nuclear charge
increases the attraction to the outer shell, so that the outer energy
level progressively becomes closer to the nucleus. This decrease in
size is quite considerable: a chlorine atom, for example, has a radius
that is only about half that of a sodium atom.
X.
.X
bond length
Figure 8 Covalent radius.
Periodicity
:5.
Periodicity
Ionic radii
Positive ions (cations)
When an atom of a Group 1 element such as sodium loses an
electron, the ion that is formed has a much smaller radius-almost
half the value, in fact. There are two reasons for this. First, there
is now one fewer electrons than there are protons, so the nucleus
attracts the remaining electrons much more strongly. Second, there
is one fewer energy level, because the outer shell has effectively
been removed, and the remaining electrons have the noble gas
electron arrangement of the preceding element. The size of the 'ions
increases as Group 1 is descended, as the outer energy level becomes
progressively further from the nucleus.
It is difficult to compare positive ions going across a period directly,
because, apart from Group 1, unipositive ions tend not to be formed
in compounds, as more than one electron is lost when ions in
Groups 2 and 3 are formed. However, it is easy to see a trend if we
consider isoelectronic ions-that
is, those that contain the same
number of electrons. Sodium ions, Na+, magnesium ions, Mg2+, and
aluminium ions, AP+, all contain 10 electrons and have the electron
configuration of neon (2.8), with the second shel1 completely fu11.
However, sodium has II protons in the nucleus, magnesium has 12,
and aluminium has 13. The 13 protons in the aluminium nucleus
will attract the eight electrons in the outer shell much more strongly
than the 11 protons in the nucleus of the sodium ion. As we move
across the period, the isoelectronic ions will become much smaller:
see Figure 10.
Cations
Anions
atom
ion
atom
ion
8
ONa+
@
@)
98
2.8
11 protons
10 electrons
99
2.8.7
17 protons
17 electrons
186
2.8.1
II protons
11 electrons
ou+
.,-----
®
68
ONa+
98
0 Mg'+
65
OAI>+
45
133
GG0
212
®
radius
133
181
2.8.8
17 protons
18 electrons
I x j.O-"
190
m
181
8
(0
196
219
Figure 10 Radius of ions.
, 4.6.
,.~'{
Negative ions (anions)
When the atoms of elements in Group 7 (the halogens) gain one
electron to form a negative ion, there will be one more electron
in the outer shell and hence more electron-electron repulsion. As
the number of protons in the nucleus is unchanged, each of the
electrons will be attracted less strongly, and the radius of the ion
l
"?
increases to almost twice the radius of the atom. The size of the
negative ions increases down the group as the outer shell is further
away from the nucleus. Across a period we again need to compare
isoelectronic ions. Phosphide ions, p3-, sulfide ions, S2-, and chloride
ions, Cl-, all have the electron configuration of argon (2.8.8).
However, the 18 electrons will be less attracted by the 15 protons in
the nucleus of the phosphide ion than they will be by the 16 protons
in the sulfide nucleus. Similarly, the sulfide ion will be larger than
the chloride ion, where the 18 electrons are more strongly attracted
by the 17 protons in the chloride ion nucleus.
Melting points
There are essentially two factors that determinethe
melting point
of a crystalline substance. When a substance melts, the attractive
forces holding the particles together in the crystal structure of the
soltd are overcome, and the particles are free to move around in the
liquid state. The temperature at which this happens will depend both
on the strength of the attractive forces and on the way in which the
particles are packed in the solid state.
Within each group of the Periodic Table the forces of attraction
tend to be similar. This is certainly true for Group I, in which the
elements all have a metallic structure, and for Group 7, in which
there are only weak forces of attraction between the separate halogen
molecules. It is less clear cut in other groups. For example, in
Group 4, carbon and silicon have strong covalent bonds between the
atoms to form giant covalent molecules, whereas tin and lead-at
the bottom of the group-have
metallic structures. In Group I the
melting points decrease down the group (Table 3). Lithium melts at
181°C, whereas the melting point of caesium is only just above room
temperature. The melting points decrease because, as the atoms get
larger, the forces of attraction between them, which are proportional
to the inverse of the distance squared, decrease.
Table 3 Melting points of Group 1 elements (alkali metals)
-Na
K
Cs--
t]f~~ent:--
li~---.
Rb
-
39
29
~f
Melting pointj"(
181
98
64
Table 4 Melting points of Group 7 elements (halogens)
~
~
Element "
-->'
Me[tingpoint/oC
-
-----
F2
(12
Br2
'2
-220
-101
-7.2
114
In Group 7 the melting points show the opposite trend, and increase
down the group (Table 4). This is because the solid crystals of the
halogens contain non-polar diatomic molecules, which are only
weakly attracted to each other. As you will read in Chapter 4, these
weak forces, which are known as van der Waals' forces, increase as
the mass of the molecules increases.
Across the period there is a large change in the pattern of melting
points as the bonding type changes from metallic (Na, Mg and AI),
to giant covalent (Si), to weak van der Waals' attraction between
simple molecules (P4' S8' C12) and monatomic molecules (Ar). What is
striking is the way the trend shows periodicity, as the pattern repeats
itself with the next period. See Figure 11.
Periodicity
----------------
----
Periodicity
3
0
-'<,<
4000
c
~ 3500
i!!
OJ
~3000
2
2500
2000
Figure 11 Melting points of the first 20 elements.
Thinking about science
How perfect is the Periodic Table?
Mendeleyev's Periodic Table and its subsequent revisions to include
new elements rank as one of the greatest achievements in science. The
Periodic Table is very much the chemist's tool, and chemist is able to
use it to predict the properties of elements (and their compounds) that
he or she is unfamiliar with or which have not yet been synthesized. It
fulfils Popper's criterion for a scientific theory, that it is capable of being
tested by falsification (see Chapter 1). The graph of first ionization
energies against atomic number is a perfect example of periodicity, and
can be explained by the way in which the elements are set out in the
Table. However, sometimes the trends down a group or across a period
are not as perfect as we might expect. Let us look at the formulas for the
highest fluorides formed by the elements in Period 3. Argon is left out,
because it does not form a stable compound with fluorine.
NaF
MgF2
AIF3
SiF4
PFs
SF6
There is obviously a very clear trend here, and we would expect that
a
the highest fluoride of chlorine would have the formula CIFT In fact the
trend breaks down, and the highest fluoride of chlorine has the formula
CIFs' Sometimes, when the Periodic T-able appears to give us the "wrong"
answer, it can lead to an even greater understanding. The clue to the
explanation as to why chlorine does not form CIF7 lies in the fact that
bromine also forms BrFs as its highest fluoride, but iodine does form a
fluoride with the expected formula of 1FT If we look at the covalent radii
of the chlorine, bromine and iodine atoms we can see that the chlorine
and bromine atoms are too small to accommodate seven fluorine atoms
around them, whereas they are able to fit around the larger iodine atom.
However, not all the anomalies are so easy to explain. We have seen
that the values for the electronegativities of the elements increase across
a period and up a group, so that fluorine is the most electronegative
element. Electronegativity is a measure of the relative ability of an atom
of the element to attract a bonding pair of electrons. When a single
electron is added to a gaseous atom of the element we can physically
measure the energy change. This is known as the first electron affinity,
and it is measured in kilojoules per mole. The value is always exothermic
~·•.hen just one electron is added to an atom of any gaseous element,
o
1 Use the graph of first ionization
energies against atomic number
(up to Z = 50) to predict what the
value for the first ionization energy
of caesium (Z = 55) will be. Check
your prediction with the value .given
in the IB Data Booklet
2 The trans-uranium elements with
an atomic number greater than
92 are all radioactive and hence
unstable. Many have been made
artificially, often only in very small
amounts. Element 119, ekafrancium or ununennium, Uue, has
not yet been made. If it follows the.
expected periodic trend it should
have one electron in its outer shell,
which will place it below francium
in Group 1. Use your knowledge
of the Periodic Table to predict the
following about ununennium, Uue:
(a) If it could be obtained in
sufficient quantities, what would
it be expected to look like at
room temperature?
(b) What ion would it be expected
to form when it reacts?
3 Can you suggest a possible reason
wby tbe electron affinity of fluorine
is lower than that of chlorine? Hint:
look at the values for the other
elements in the second period
compared with the corresponding
elements in the third period of the
Periodic Table;
.~3
1,',Periodicity
because the attraction between the nucleus and the added electron is
greater than the electron-electron repulsion.
The values for chlorine, bromine and iodine are
Element
CI
Br
I
Etectronaffinityl kJ mol:"
-364
-342
-314
There is a clear trend, and if we try to explain if we would probably
refer to the size of the atoms and the ions formed, and perhaps to the
number of protons in the nucleus. Almost certainly our explanation would
lead us to predict that the value for fluorine would be even greater,
perhaps in the region of -380 kJ mol='. In fact the value for fluorine
is -348 kJ mol ", The trend has broken down, and there does not at
present seem to be an obvious explanation for this, although several
theories have been proposed.
_.
It is when a hypothesis does not explain all the observed information
that one needs to question and perhaps modify the hypothesis. Scientific
theories continue to evolve, and perhaps the Periodic Table will one day
evolve to become even better and closer to the truth.
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c
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