Download The Periodical Table and chemical properties

PERIODIC TABLE
Periodic table
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The periodic table is a tabular arrangement of the
chemical elements, organized on the basis of their
atomic numbers, electron configurations (electron
shell model), and recurring chemical properties.
Elements are presented in order of increasing atomic
number (the number of protons in the nucleus).
The standard form of the table consists of a grid of
elements laid out in 18 columns and 7 rows, with a
double row of elements below that.
The table can also be deconstructed into four
rectangular blocks: the s-block to the left, the p-block
to the right, the d-block in the middle, and the f-block
below that.
Periodic table which is used in
Europe
The history of the periodic
table reflects over a
century of growth in the
understanding of
chemical properties.
The most important
event was the
publication of the first
periodic table by Dmitri
Mendeleev in 1869.
The table is a visual representation of the periodic
law which states that certain properties of elements
repeat periodically when arranged by atomic
number. The table arranges elements into vertical
columns (groups) and horizontal rows (periods) to
display these commonalities.
 All versions of the periodic table include only
chemical elements, not mixtures, compounds, or
subatomic particles. Each chemical element has a
unique atomic number representing the number of
protons in its nucleus.
 In the standard periodic table, the elements are listed
in order of increasing atomic number (the number of
protons in the nucleus of an atom).
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A new row (period) is started when a new
electron shell has its first electron.
Columns (groups) are determined by the
electron configuration of the atom. Elements
with the same number of electrons in a
particular subshell fall into the same columns
(e.g. oxygen and selenium are in the same
column because they both have four electrons
in the outermost p-subshell
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Groups
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A group or family is a vertical column in
the periodic table. Groups usually have
more significant periodic trends than
periods and blocks, explained below.
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Elements in the same group tend to show
patterns in atomic radius, ionization
energy, and electronegativity.
Groups
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From top to bottom in a group, the atomic
radii of the elements increase.
Since there are more filled energy levels,
valence electrons are found farther from
the nucleus. From the top, each
successive element has a lower ionization.
Similarly, a group has a top to bottom
decrease in electronegativity due to an
increasing distance between valence
electrons and the nucleus.
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Periods
A period is a horizontal row in the periodic table.
 Elements in the same period show trends in
atomic radius, ionization energy, electron affinity,
and electronegativity. Moving left to right across
a period, atomic radius usually decreases. This
decrease in atomic radius also causes the
ionization energy to increase when moving from
left to right across a period. Electronegativity
increases in the same manner as ionization
energy. Electron affinity also shows a slight
trend across a period. Metals (left side of a
period) generally have a lower electron affinity
than nonmetals (right side of a period), with the
exception of the noble gases.
The Periodic Table and physical
properties
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
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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.
Atomic radius
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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.
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As group is descended the outermost electron is
in a higher energy level, which is further from the
nucleus, so the radius increases.
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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.
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
electrostatic attraction between the nucleus and
the outermost electron is greater and the ion is
smaller because the number of electron shells
has decreased by one. Across the period the
ions contain the same number of electrons, but
an increasing number of protons, so the ionic
radius decreases.
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Anions contain more electrons than protons, so
more electrons than in the parent atom. Across a
period the size decreases because the number
of electrons remains the same but the number of
protons increases.
Periodicity
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Elements in the same group tend to have similar
chemical and physical properties. There is a
change in chemical and physical properties
across a period. The repeating pattern of
chemical and physical properties shown by the
different periods is known as periodicity.
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These periodic trends can clearly be seen in
atomic radii, ionic radii, ionization energies,
electronegativities and melting points.
The Periodical Table and
chemical properties
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Chemical properties of elements in the
same group
Group 1- the alkali metals
 Lithium, sodium, and potassium all contain
one electron in their outer. They are all
reactive metals and are stored under liquid
paraffin to prevent them reacting with air.
They react by losing their outer electron to
form the metal ion. Because they can readily
lose an electron they are good reducing
agents. The reactivity increases down the
group as the outer electron is in successively
higher energy levels and less energy is
required to remove it.
They are called alkali metal because they all
react with water to from an alkali solution of the
metal hydroxide and hydrogen gas. Lithium
floats and reacts quietly, sodium melts into a ball
which darts around on the surface, and the heat
generated from the reaction with potassium
ignites the hydrogen.
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2Li(s) + 2H20(I) 2Li+(aq) + 20H-(aq) + H2(g)
2Na(s) + 2H20(I )
2K(s) + 2H20(I)
2Na+(aq) + 20H-(aq) + H2(g)
2K+(aq) + 20H-(aq) + H2(g)
Group 7 - the halogens
The halogens react by gaining
one more electron to form halide
ions. They are good oxidizing
agents. The reactivity decreases
down the group as the outer shell
is increasingly at higher energy
levels and further from the
nucleus. This, together with the
fact that there are more electrons
between the nucleus and the
outer shell, decreases the
attraction for an extra electron.
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Chlorine is a stronger oxidizing agent than
bromine, so can remove the electron from
bromide ions in solution to form chloride
ions and bromine. Similarly both chlorine
and bromine can oxidize iodide ions to
form iodine.
Cl2(aq) + 2Br-(aq)
2CI-(aq) + Br2(aq)
Cl2(aq) + 2I-(aq)
2CI-(aq) + I2(aq)
Br2(aq) + 2I-(aq)
2Br-(aq) + I2(aq)
Test for halide ions
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The presence of halide ions in solution can
be detected by adding silver nitrate
solution. The silver ions react with the
halide ions to form a precipitate of the
silver halide. The silver halides can be
distinguished by their colour. These silver
halides react with light to form silver metal.
This is the basis of photography.
Metalloid
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A metalloid is a chemical element that has properties
that are in between those of metals and nonmetals.
The six elements commonly recognised as metalloids
are boron, aluminium, germanium, arsenic,
antimony and tellurium. On a standard periodic
table all of these elements can be found in or near a
diagonal region of the p-block, having its main axis
anchored by boron at one end and astatine at the
other. Some periodic tables include a dividing line
between metals and nonmetals and it is generally the
elements adjacent to this line or, less often, one or
more of the elements adjacent to those elements,
which are identified as metalloids.
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Physically,
metalloids usually
have a metallic
appearance but they
are brittle and only
fair conductors of
electricity;
chemically, they
mostly behave as
(weak) nonmetals.
They can form
alloys with metals.
Here are the general physical
properties exhibited by metalloids.
Metalloids are solids at room
temperature.
They are lustrous i.e. they have
a shiny surface like metals.
Some of them are ductile in nature and
can be drawn into pipes and wires.
silicon
They are fair conductors of heat and electricity,
but not as good as metals. Hence, they are also
known as semi-conductors. Metalloids like boron,
germanium, and arsenic are used as dopants in
glasses for use in semiconductor chips.
They are usually brittle in nature.
These elements mostly exist in several allotropic
forms.
The density of metalloids is lower than that of poor
metals, but higher than that of non-metals.
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germanium
The temperature coefficient of resistance can
be positive (arsenic and antimony) or negative
(boron, silicon, germanium, tellurium) for
metalloids.
 They have an open crystal structure, as
opposed to the closed crystal structure of
metals.
 They have abnormally high values of enthalpy
of fusion.
 They exhibit electrical conductivity even in
the liquid form.
 In metalloids, fewer valence electrons are
available as "free electrons".
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Chemical Properties
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Metalloids tend to have an
intermediate property between metals
and non-metals. Given below are
some general chemical properties of
metalloids.
In a standard periodic table layout, you
will observe that the metalloids that are
placed on the upper right side of the
diagonal line through the p-block,
display increasing non-metallic
behavior. However, those placed to the
lower left of the line are more metallic
in character. This diagonal line is
called the 'stair-step' or 'staircase'.
Arsenic
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The general chemical behavior of metalloids is
similar to that of non-metals.
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Many metalloids have multiple of oxidation states or
valences. However, in most chemical reactions, they
may behave either as metals or non-metals.
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The metalloids usually form amphoteric oxides. The
oxides formed by metals are basic oxides, while nonmetals generally form acidic oxides.
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The ionization energy of metalloids is higher than
that of metals, but lower than that of non-metals.
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Many metalloids have different allotropes. For a
given metalloid, one of its allotrope may react as a
metal and the other allotrope may behave as a nonmetal. For example, carbon in its diamond allotrope
acts like a true non-metal and is a bad conductor of
electricity, but its graphite allotrope is a fairly good
electrical conductor.
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Allotropes of tin, phosphorous and bismuth exhibit
borderline behavior.
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Metalloids form anions (negatively charged ions) in
water.
These elements can form ionic as
well as covalent bonds.
 The oxidation number for
metalloids can be positive or
negative.
 Metalloids can react with metals to
form alloys.
 The oxides of metalloids are
weakly acidic. They are amorphous
and when heated beyond a certain
temperature, which is known as the
glass transition temperature), form
glass.
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tellurium
Transition Metal
The transition metals, also called the d-block elements,
are found in groups 3-12 of the periodic table. These
elements make the transition between the
representative metals in groups 1 and 2 and the
metalloids, representative metals, and nonmetals in
groups 13-18. Moreover, it is in this block of elements
in the periodic table that the d-orbitals are being filled
with electrons. For example, elements in the first
group of the d-block (group 3 of the periodic table:
scandium, yttrium, lanthanum, and actinium) each
have one d-electron. Likewise, elements in the eighth
group of the d-block (group 10 of the periodic table:
nickel, palladium, and platinum) each have eight delectrons when in the +2 oxidation state.
The transition metals have several features
in common: unlike representative metals,
most transition metals have variable
valence, meaning that they have more
than one possible oxidation—or valence—
state. For example, platinum exists most
commonly in the +2 and +4 oxidation
states, but it can also be found in the +5
and +6 oxidation states.
Properties
Because they possess the properties of metals, the
transition elements are also known as the transition
metals. These elements are very hard, with high
melting points and boiling points. Moving from left
to right across the periodic table, the five d orbitals
become more filled. The d electrons are loosely
bound, which contributes to the high electrical
conductivity and malleability of the transition
elements. The transition elements have low
ionization energies. They exhibit a wide range of
oxidation states or positively charged forms.
Properties
The positive oxidation states allow transition
elements to form many different ionic and
partially ionic compounds. The formation of
complexes causes the d orbitals to split into two
energy sublevels, which enables many of the
complexes to absorb specific frequencies of light.
Thus, the complexes form characteristic colored
solutions and compounds. Complication reactions
sometimes enhance the relatively low solubility
of some compounds.
The formation of coloured compounds
Some common examples
The diagrams show aproximate colours for some
common transition metal complex ions.
 When white light passes through a solution of one of
these ions, or is reflected off it, some colours in the light
are absorbed. The colour you see is how your eye
perceives what is left.
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CHANGE FROM METALLIC TO NON-METALLIC
NATURE OF THE ELEMENTS ACROSS PERIOD 3
Metals tend to be shiny and are good conductors
of heat and electricity. Sodium, magnesium, and
aluminium all conduct electricity well. Silicon is a
semi-conductor and is called a metalloid as it
possesses some of the properties of a metal and
some of a non-metal. Phosphorus, sulfur,
chlorine, and argon are non-metals and do not
conduct electricity. Metals can also be
distinguished from non-metals by their chemical
properties. Metal oxides tend to be basic,
whereas non-metal oxides tend to be acidic.
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Sodium oxide and magnesium oxide are
both basic and react with water to form
hydroxides, e.g.
Na20(s) + H20(l)
MgO(s) + H20(l)
2NaOH(aq)
Mg(OH)2
Aluminium is a metal but its oxide is amphoteric, that
is, it can be either basic or acidic depending on
whether it is reacting with an acid or a base.
The remaining elements in period 3 have acidic
oxides. For example, sulfur trioxide reacts with
water to form sulfuric acid, and phosphorus
pentoxide reacts with water to form phosphoric (V)
acid.
S03(g) + H20(l)
H2S04(aq)
P4O10(s) +6H20(I) 4H3P04(aq)
Chlorine itself reacts with water to some extent to
form an acidic solution.
Cl2(aq) + H20(l)
HCI(aq) + HCIO(aq)
Thank you for your attention !