Home CHEMISTRY TOPIC 6: PERIODIC CLASSIFICATION | CHEMISTRY FORM 2

TOPIC 6: PERIODIC CLASSIFICATION | CHEMISTRY FORM 2

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TOPIC 6: PERIODIC CLASSIFICATION | CHEMISTRY FORM 2

TOPIC 6: PERIODIC CLASSIFICATION | CHEMISTRY FORM 2

Constructing the modern periodic table has been a major scientific achievement. The first steps towards working out this table were taken long before anyone had any idea about the structure of atoms.

The number of elements discovered increased steadily during the 19th century. Chemists began to find out patterns in their properties.

The Law of Triads
In 1817, the German scientist Johann Dobereiner noticed that calcium,
strontium and barium had similar properties, and that the atomic weight
of strontium was halfway between the other two. He found the same
pattern with chlorine, bromine and iodine and also with lithium, sodium
and potassium.
So, he put forward the law of Triads: “If elements are
arranged in groups of three in order of increasing atomic weights,
having similar properties, then the atomic weight of the middle element
is the arithmetic mean of the atomic weights of the other two elements”,
E.g.
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The
following are examples of Dobereiner’s triads:(Lithium, Sodium and
Potassium)(Calcium, Strontium and Barium)(Chlorine, Bromine and Iodine)
and(Iron, Cobalt and Nickel)
The Law of Octaves
In
1863 John Newlands, an English chemist noted that there were many pairs
of similar elements. In each pair, the atomic weights differed by a
multiple of 8. So, he produced a table with the elements in order of
increasing atomic weights, and put forward the Law of Octaves: “If
elements are arranged in order of their increasing atomic weights, the
properties of the 8th element, starting from a given one, are a kind of
repetition of the first element”.
This finding was comparable to the 8th note of music, hence the use of the word “octave”.
This
was the first table to show a periodic or repeating pattern of
properties. But it was not widely accepted because there were too many
inconsistencies. For example, he put copper and sodium in the same
group, even though have very different properties. Also iron was placed
in the same group as oxygen and sulphur.
The Periodic Law
Dmitri
Mendeleev was born in Siberia, Russia, in 1834. By the time he was 32,
he was a professor of Chemistry. In 1869 Mendeleev advanced the work
done by Newlands and contributed very useful new ideas. He began by
listing all the known elements in order of increasing atomic mass. He
spotted that elements with similar properties appear at regular
intervals or periods down the list. His findings were the basis for the
Periodic Law: “The properties of elements are a periodic function of
their atomic masses”.
Mendeleev
placed similar elements into groups. He realized that not all elements
had been discovered. So he left gaps for new ones in the correct places
in his table. He also swapped the order of some elements to make them
fit better. He predicted the properties of the missing elements from the
properties of the elements above and below them in the table. He also
listed separately some elements which did not appear to fit into any
group i.e. iron, cobalt, nickel, etc.
Table 6.1: Mendeleev’s short form of the Periodic Table
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The
table had 9 vertical columns which he called Groups. The groups were
numbered from 0 to 8. The elements in group 0 were not known by then,
but were discovered later on. Groups 1 to 7 were subdivided into A and B
subgroups. Group 0 included the transition elements. Noble gases were
later placed in group 0.
There
were 7 horizontal rows which he called periods. All vacant positions in
the table stood for new elements yet to be discovered.
Usefulness of Mendeleev’s classification

The table summarized a large amount of information about the elements based on their chemical properties.

The
table was very useful in predicting the existence and properties of
undiscovered elements, for which gaps had been left in the table.

The table was also used in checking relative atomic masses of elements.

Limitations of Mendeleev’s classification

In
three cases, pairs of elements had to be included in one group based on
inverse order of their atomic weights so as to fit into groups of
elements having similar properties. These pairs were argon (39.9) and
potassium (39.1), cobalt (58.9) and nickel (58.9); plus tellurium
(127.5) and iodine (126.9). This difficulty was resolved when the basis
of classification was based on the atomic number instead of the atomic
mass.

The elements that were placed in group VIII formed an incompatible mixture.

The placing of two different families in one group e.g. K and Cu; Ca and Zn, etc.

The
periodic table is the chemists map. It helps you understand the
patterns in chemistry. Today we take it for granted. But it took
hundreds of years, and work of hundreds of chemists, to develop.
The
Modern Periodic Table is similar to that of Mendeleev, but contains
several improvements. Elements are arranged in order of atomic number
instead of atomic mass. This means that elements no longer have to swap
places to fit correctly. Many new elements have been discovered and
slotted into the spaces left by Mendeleev. Also metals and non-metals
are clearly separated. The Modern Periodic Table is shown in Figure 6.1.
Figure 6.1: The Modern Periodic Table
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The
long form of the periodic table is the commonly used form of the
periodic table. The elements in the table are arranged based on their
atomic weights, starting from hydrogen (1), helium (2), lithium (3),
beryllium (4) and so on. The elements appear in vertical columns and
horizontal rows.
The
vertical columns in the table are called Groups, numbered I, II, III,
IV, V, VI, VII and 0, which is also known as group VIII. Group I
contains the elements lithium (L), sodium (Na), rubidium (Rb), caesium
(Cs) and francium (Fr). Group II consists of elements starting from
sodium (Na) down to radium (Ra). Some of the groups have special names.
  • Group I is often called the alkali metals.
  • Group II the alkaline earth metals.
  • Group VII the halogens.
  • Group 0 the noble gases.
The
transition metals (or elements) form a separate block in the middle of
the periodic table between group II and III. The atoms of these elements
have more complicated electron arrangements. Note that the group
contains many common metals such as iron (Fe), Nickel (Ni), copper (Cu),
and Zinc (Zn). One of the interesting properties of these elements is
that they form coloured compounds.

Main features of the Modern Periodic Table

The elements in the table are placed in order of their atomic numbers instead of their atomic masses.

There are a total of 18 groups and 7 periods.

There are 5 blocks of similar elements in the periodic table as shown in figure 6.2.

The
normal (non-transition) elements (groups 1-7) have their outermost
shells incomplete, meaning that they can allow additional electrons to
enter into their outermost orbital (valency shell). But each of their
inner shells is complete.

The transition metals have their outermost as well as their penultimate (second last) shells incomplete.

Elements
of group 0 (noble gases) have their shells complete. These elements
show little reactivity. That is why they wereonce called „inert‟ gases
because they are very unreactive; or „rare gases‟ because they were
rarely found.

Gaps left by Mendeleev for undiscovered elements
(now occupied by the transition elements and the noble gases) have been
filled by the respective elements following their discovery. Man-made
elements have also found a place in the periodic table.

Metals
have been clearly separated from non-metals. Metalloids or semi metals
(poor metals) have also been included. Metalloids are elements whose
properties are intermediate between metals and non-metals. They include
boron (B), silicon (Si), germanium (Ge), arsenic (As), antimony (Sb) and
tellurium (Te). In some publications, germanium and antimony are
usually classed as poor metals and the rest as non-metals.

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Periodicity

The Concept of Periodicity
Explain the concept of periodicity

Consider the electronic configuration of the first twenty elements of the periodic table shown in the table below.

Table 6.3: Electronic configurations of the first 20 elements
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You
will notice that elements in the same vertical columns (groups) have
the same number of electrons in the outermost shells of their atoms.
Because the outer electrons determine the chemical properties of an
element, then the elements in each period tend to resemble each other
closely in chemical behaviour. For instance, the noble gases, He, Ne and
Ar show a chemical inertness which is characterised by the stable outer
electron octet or duplet. Due to this reason, the compounds of the
noble gases with other elements have not been found.

Attempts
to classify elements by arranging them in order of increasing atomic
weights shows that the properties of elements were periodic. This means
elements with similar or comparable properties appear after a certain
specific interval in a given arrangement. The occurrence of successive
groups of elements showing strong chemical similarity in this way is
called periodicity.
Therefore,
periodicity is the repetition of similar chemical properties of
elements after a certain specific interval in a given arrangement. The
repetition in properties is due to repetition of similar electronic
configuration of outermost shells of elements after certain intervals.

General Trends

This
refers to change in some properties of elements across the periods and
down the groups in the periodic table. These trends become more obvious
if we leave aside the noble gases in Group 0. In this case, we shall
concentrate our efforts on variations in the most important properties
of the elements only. The following is a summary of the change in some
properties of elements down the groups and across the periods.
The Change in Properties of Elements Across the Periods
Explain the change in properties of elements across the periods
Atomic and ionic size
The
sizes of atoms and ions may be given in terms of atomic radius and
ionic radius units respectively. The number of shells an atom or ion
posses and the nuclear charge determines the size of an atom or ion.

This is how the two properties vary along the period and down the group:

Atomic size

Along the period:

Considering the normal elements only, the size of the atoms decrease
from left to right across the period. This is because as atomic number
increases across the period, the nuclear charge (due to increasing
protons) increases and electrons in shells are pulled closer to the
nucleus.

Ionic size

Positive ions (cations):Across the period; The ionic size does not change, i.e. remains the same, as you move across the period from either direction.

Negative ions (anions):A
negative ion is larger compared to the corresponding neutral atom
because on forming an ion, one or more electrons are added to the atom.

The added electron(s) is/are repelled by the electron(s) already present
in the outermost shell, hence leading to an increase in the size of an
atom, even though no new shell is formed.Down the group and along the period: Ionic size increases down the group, and along the period, i.e. from left to right.

Atomic radii (singular: radius)

Along the period: In the period, atomic radii decrease from left to right with increase in the atomic number.

Electronegativity

Electronegativity

is the tendency of an atom to attract the shared pair of electrons
towards itself in a molecule. The electronegativity values of elements
in group 0 (inert gases) is zero.

Along the period: Electronegativity increases while moving across the period from left to right in the periodic table.

Metallic character (or electropositivity)

Electropositivity
is the tendency of an element to lose the valency electron(s) and
donate the same to other elements (usually non-metallic elements). This
process occurs during the formation of new substances e.g. molecules and
compounds. Literally, such reactions occur between metals and
non-metals whereby metals donate electrons and non-metals receive these
electrons. So, metals are electropositive elements while non- metals are
electronegative elements.

Along the period:
Generally, metallic character decreases along the period from left to
right.The gradation in metallic properties across the period is as
follows: Metals → poor metals → metalloids → non-metals → noble gases
Chemical reactivity
Reactivity is the tendency of an element to lose or gain electrons in a chemical reaction.
Along the period: For metals, the reactivity decreases from left to right in a period while it increases for non-metals.
Ionization Energy or Ionization Potential (I.E or I.P)
This
refers to the minimum amount of energy required to remove the most
loosely bound electron from an isolated atom or ion in its gaseous
state. The smaller the value of ionization energy, the easier it is to
remove the electron from the atom.M(g) →M+(g) + e-
Along the period: It increases along the period from left to right with the increase in atomic number.
Electron affinity (Ea):
This
is just opposite to I.E. It is defined as the amount of energy released
when an extra electron is added to an isolated neutral atom in its
gaseous state.
Along the period: The value increases along the period from left to right.
Density and melting point
The
density of a substance is the ratio of its mass to its volume, while
the melting point is the temperature at which a solid substance turns
into liquid at standard atmospheric pressure.
  1. Density-Across the period: Densities decrease across the period from left to right.
  2. Meting point-Across the period: Melting points of elements decrease across the period from left to right.
The Change in Properties of Elements Down the Groups
Explain the change in properties of elements down the group
Atomic and ionic size
  1. Atomic size-Down the group: Atomic size increases as you move down the group.
  2. Ionic size- Positive ions (cations)-Down the group:
    On descending the group, the nuclear charge increases and the number of
    shells increase by one at each step so, the ionic size also increases. A
    positive ion is smaller than the corresponding neutral atom because on
    forming the ion, the metal atom loses both the valency electron(s) and
    the outermost shell. Valency electron(s) refer(s) to the electron(s) in
    the outer-most shell of an atom. Any further removal of electron(s) from
    the ion will decrease the ionic size further.Negative ions (anions)-Down the group and along the period: Ionic size increases down the group, and along the period, i.e. from left to right.
Atomic radii (singular: radius)
Atomic radius is the distance from the centre of the nucleus to the outermost shell (valency shell). Down the group: Atomic radii of elements increase down the group with increase in atomic size.
Electronegativity
Down the group: Electronegativity decreases while moving downwards in a group.
Metallic character (or electropositivity)
Down the group: Metallic character (electropositivity) increases down the group
Ionization Energy or Ionization Potential (I.E or I.P)
Down the group: It decreases gradually down the group.
Why
is there a decrease in I.E as you go down the group? This is because
electrons are held in their shells by their attraction to the positive
nucleus, and as you go down the group, the size of the atom increases
(increasing atomic radius). So, the outermost electron(s) of an atom
gets further and further away from the attraction or pull of the
positive nucleus, hence requiring little energy to remove from the atom.
Electron affinity (Ea)
Down the group: The value of electron affinity decreases down the group.
Density and melting point
  1. Density-Down the group: Densities of elements increase down the group.
  2. Meting point-Down the group: Melting points of elements decrease down the group as the elements become less metallic in nature.
Electronic Configuration to Locate the Positions of Elements in Periodic Table
Use electronic configuration to locate the positions of elements in periodic table
The
modern periodic table is based on electronic configurations of the
elements. Look at table 6.3 and study the electronic configurations of
the first twenty elements and where they are placed in the periodic
table.
Beryllium, magnesium and calcium have two electrons in the outer shell. These elements are in Group 2.
This
pattern continues to Group 3, Group 4 and so on. The group number in
the periodic table is the same as the number of electrons in the
outermost shell. The halogens are the elements in Group 7. Bromine is
one of the halogens. How many electrons does each bromine atom have in
its outer shell?
As
we move down each group, the number of shells increases by one at each
step. Each atom of an element has one complete shell than the one above
it.
As
we move across each period, the outer shell is being filled by one
electron at each step. Certain electronic configurations are found to be
more stable than others are. The noble gases at the end of each period
have full outer shells. They have stable duplet (2 electrons) or octet
(8 electrons) in their outermost shells. This makes them more difficult
to break up, and this fits well with the fact that they are so
unreactive.

The
outer electrons of an atom are mainly responsible for the chemical
properties of an element. Therefore, elements in the same group will
have similar chemical properties.

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