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INTRODUCTION OF
INORGANIC CHMISTRY
Rana Atif Javed Kalander
MAY 30, 2021
PUNJAB COLLEGE, SAMBRIAL
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Introduction of Inorganic Chemistry
Inorganic Chemistry
Inorganic Chemistry deals with the study of inorganic compounds, which are formed by the combination of
two or more chemical elements. Although majority of the compounds formed by carbon with different
elements are termed as ‘Organic Compounds’ and are studied under ‘Organic Chemistry’, but the compounds
of carbon which do not contain a Carbon-Hydrogen bond including some carbonates and some cyanides are
termed as ‘inorganic’ in nature and are studied under ‘Inorganic Chemistry’. Further, inorganic compounds
may be formed by any type of chemical bonding such as covalent bonding, coordinate bonding, electrovalent
bonding and metallic bonding.
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Prof
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Where is Inorganic Chemistry used?
‘Inorganic Chemistry’ is an extremely significant branch of Chemistry. It is due to the fact that different
inorganic compounds find numerous applications in our daily life as well as in exceptional research works.
There are a number of places where inorganic chemistry is used, few of them are mentioned below for your
knowledge.
Sustains life
Inorganic Chemistry plays an important role in the sustenance of life. This is because ‘water’ which is essential
for our survival, is itself an inorganic compound. Not only this, table salt which forms a crucial part of our
diet, is chemically Sodium Chloride, which is again an inorganic compound. This way, inorganic chemistry
plays a vital role in maintaining the life cycle.
Synthesis of Ammonia
Inorganic Chemistry is used in the synthesis of Ammonia. It is an extremely useful inorganic compound which
finds its usage in both households as well as the industrial sector. The production of fertilizers, manufacture
of household cleaning products and refrigeration are some examples of prominent applications of Ammonia.
Medicinal applications
Inorganic compounds find a number of applications in the healthcare industry. Satraplatin, Budotitane and
sodium aurothiomalate are some examples of inorganic compounds which are used for the treatment of dreary
diseases like arthritis and cancer. Thus, inorganic chemistry is widely used in medicinal applications.
Catalytic Applications
Catalytic Applications are one of the most important applications of inorganic chemistry. Chemical reactions
require some amount of energy to take place. This energy is known as ‘Activation Energy’. At times, when
the activation energy required for chemical reactions to take place is extremely high and it becomes difficult
or challenging to supply it externally, the role of catalysts comes into play. Catalysts lower down the energy
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of activation and thus, promote the rate of the reaction. Vanadium oxide and Titanium(III) chloride are two
examples of extremely important inorganic catalysts.
Coating Applications
Coatings of inorganic materials serve an important purpose of corrosion protection. Anodizing, porcelain
enamel and conversion coatings are some examples of useful inorganic coatings. Thus, inorganic Chemistry
finds important application in the coating industry.
Careers in Inorganic Chemistry
Inorganic Chemistry finds an abundance of applications in the industrial sector, thus, being an inorganic
chemist, you get to enjoy a promising career in different professional fields. Below mentioned are some of
them.
Ceramic Chemists
Ceramic materials are materials having advanced properties such as high strength and thermal resistance.
Ceramic Chemists put all their knowledge into developing such advanced ceramic materials which find usage
in important industrial sectors such as Aerospace industry, cutting tool manufacturing industry and insulation
industry. Further, it is because of the wide range of applications which ceramics find in different industrial
sectors, such as the ones mentioned above, that ceramic chemists always remain high in demand. Also, ceramic
chemists enjoy a good pay scale and thus, lead a professional life of fulfillment.
Food Scientists
Inorganic Chemists have a good scope of employment in the food safety and inspection industry. They are
involved in conducting different kinds of research work to enhance the nutritional value of the food items and
also ensuring the safety of packaged food items. Not only this, food scientists are also involved in ensuring
proper distribution of food items across different channels of food delivery. Furthermore, food scientists are
employed at a high salary package in both the private as well as the government sector. It is only because of
the meticulous efforts of food scientists that we can purchase safe packaged food which is fit for consumption.
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Environmental Chemists
Inorganic Chemists utilize their vast knowledge to deal with different types of environmental issues and
problems. It is with the rise in the levels of environmental pollution that environmental problems are also
witnessing a splurge. The water we drink, the air we breathe and the soil we use to raise crops, all are being
polluted by pollutants such as industrial effluents and unburnt fuel particles from automobiles. All this is
resulting in a severe damage to the planet and its inhabitants be it animals, plants or human beings. Thus,
numerous government and private industries are working towards the development of advanced techniques
and technologies which can reduce different types of environmental pollution. For this, they conduct extensive
recruitments for environmental Chemists. Moreover, both these sectors offer handsome salary packages to
environmental chemists.
Mining Analytical Chemists
Inorganic Chemists get employed in the mining industry as mining analytical chemists. The job role involves
analyzing the mined ores to determine their chemical composition and then using them for various
applications. Along with this, many important elements exist in combined form as compounds beneath the
surface of the Earth. Analytical Mining Chemists also shoulder the responsibility of synthesizing these
elements from their compounds in pure state. Thus, inorganic chemists have a great role to play in the mining
industry and get employed at handsome pay packages.
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The Periodic Table and Periodic Trends Notes
Definition:
The periodic table is a tabular array of the chemical elements organized by atomic number, from
the element with the lowest atomic number, hydrogen, to the element with the highest atomic number. The
atomic number of an element is the number of protons in the nucleus of an atom of that element.
History and development:
Periodic table, in full periodic table of the elements, in chemistry, the organized array of all
the chemical elements in order of increasing atomic number—i.e., the total number of protons in the
atomic nucleus. When the chemical elements are thus arranged, there is a recurring pattern called the
“periodic law” in their properties, in which elements in the same column (group) have similar
properties. The initial discovery, which was made by Dmitry I. Mendeleyev in the mid-19th century,
has been of inestimable value in the development of chemistry.
It was not actually recognized until the second decade of the 20th century that the order of elements in the
periodic system is that of their atomic numbers, the integers of which are equal to the positive electrical charges
of the atomic nuclei expressed in electronic units. In subsequent years great progress was made in explaining
the periodic law in terms of the electronic structure of atoms and molecules. This clarification has increased
the value of the law, which is used as much today as it was at the beginning of the 20th century, when it
expressed the only known relationship among the elements. The early years of the 19th century witnessed a
rapid development in analytical chemistry—the art of distinguishing different chemical substances—and the
consequent building up of a vast body of knowledge of the chemical and physical properties of both elements
and compounds. This rapid expansion of chemical knowledge soon necessitated classification, for on the
classification of chemical knowledge are based not only the systematized literature of chemistry but also the
laboratory arts by which chemistry is passed on as a living science from one generation of chemists to another.
Relationships were discerned more readily among the compounds than among the elements; it thus occurred
that the classification of elements lagged many years behind that of compounds. In fact, no general agreement
had been reached among chemists as to the classification of elements for nearly half a century after the systems
of classification of compounds had become established in general use. J.W. Döbereiner in 1817 showed that
the combining weight, meaning atomic weight, of strontium lies midway between those
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of calcium and barium, and some years later he showed that other such “triads” exist (chlorine, bromine,
and iodine [halogens] and lithium, sodium, and potassium [alkali metals]). J.-B.-A. Dumas, L. Gmelin, E.
Lenssen, Max von Pettenkofer, and J.P. Cooke expanded Döbereiner’s suggestions between 1827 and 1858
by showing that similar relationships extended further than the triads of elements, fluorine being added to the
halogens and magnesium to the alkaline-earth metals, while oxygen, sulfur, selenium, and tellurium were
classed as one family and nitrogen, phosphorus, arsenic, antimony, and bismuth as another family of
elements. Attempts were later made to show that the atomic weights of the elements could be expressed by
an arithmetic function, and in 1862 A.-E.-B. de Chancourtois proposed a classification of the elements based
on the new values of atomic weights given by Stanislao Cannizzaro’s system of 1858. De Chancourtois plotted
the atomic weights on the surface of a cylinder with a circumference of 16 units, corresponding to the
approximate atomic weight of oxygen. The resulting helical curve brought closely related elements onto
corresponding points above or below one another on the cylinder, and he suggested in consequence that “the
properties of the elements are the properties of numbers,” a remarkable prediction in the light of modern
knowledge.
In 1864, J.A.R. Newlands proposed classifying the elements in the order of increasing atomic weights, the
elements being assigned ordinal numbers from unity upward and divided into seven groups having properties
closely related to the first seven of the elements then known: hydrogen, lithium, beryllium, boron, carbon,
nitrogen, and oxygen. This relationship was termed the law of octaves, by analogy with the seven intervals of
the musical scale.
Then in 1869, as a result of an extensive correlation of the properties and the atomic weights of the elements,
with special attention to valency (that is, the number of single bonds the element can form), Mendeleyev
proposed the periodic law, by which “the elements arranged according to the magnitude of atomic weights
show a periodic change of properties.” Lothar Meyer had independently reached a similar conclusion,
published after the appearance of Mendeleyev’s paper.
The first periodic table
Mendeleyev’s periodic table of 1869 contained 17 columns, with two nearly complete periods (sequences) of
elements, from potassium to bromine and rubidium to iodine, preceded by two partial periods of seven
elements each (lithium to fluorine and sodium to chlorine), and followed by three incomplete periods. In an
1871 paper Mendeleyev presented a revision of the 17-group table, the principal improvement being the
correct repositioning of 17 elements. He, as well as Lothar Meyer, also proposed a table with eight columns
obtained by splitting each of the long periods into a period of seven, an eighth group containing the three
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central elements (such as iron, cobalt, nickel; Mendeleyev also included copper, instead of placing it in Group
I), and a second period of seven. The first and second periods of seven were later distinguished by use of the
letters “a” and “b” attached to the group symbols, which were the Roman numerals.
Periodic system of elements with periods demarcated by noble gases
Periodic system of elements with periods demarcated by noble gases.
long-period form of periodic system of elements
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Long-period form of periodic system of elements.
With the discovery of the noble gases helium, neon, argon, krypton, radon, and xenon by Lord
Rayleigh (John William Strutt) and Sir William Ramsay in 1894 and the following years, Mendeleyev and
others proposed that a new “zero” group to accommodate them be added to the periodic table. The “shortperiod” form of the periodic table, with Groups 0, I, II,…VIII, became popular and remained in general use
until about 1930.
Short-period form of periodic system of elements, 1930
Short-period form of periodic system of elements, listing the elements known by 1930. At that time it was not
clear that thorium (90), protactinium (91), and uranium (92) were part of the actinide series, and they were
often placed in groups IVa, Va, and VIa, respectively, because they showed some similarities to hafnium (72),
tantalum (73), and tungsten (74).
Based on an earlier (1882) model of T. Bayley, J. Thomsen in 1895 devised a new table. This was
interpreted in terms of the electronic structure of atoms by Niels Bohr in 1922. In this table there are
periods of increasing length between the noble gases; the table thus contains a period of 2 elements,
two of 8 elements, two of 18 elements, one of 32 elements, and an incomplete period. The elements
in each period may be connected by tie lines with one or more elements in the following period. The
principal disadvantage of this table is the large space required by the period of 32 elements and the
difficulty of tracing a sequence of closely similar elements. A useful compromise is to compress the
period of 32 elements into 18 spaces by listing the 14 lanthanoids (also called lanthanides) and the
14 actinoids (also called actinides) in a special double row below the other periods.
CLASSIFICATION
Definition:
The arrangement of elements having same properties in the same group and separating them
from elements with different properties is called classification of elements.
Significance of Classification:
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With the increase in the number of elements discovered, an attempt was made to arrange them on
the basis of the similarities in their physical and chemical properties.
Such grouping of elements was useful because
It made their study easy.
It enabled the scientists to understand the reasons for such similarities in their properties.
1. DOBEREINER’S CLASSIFICATION
(LAW OR RULE OF TRIADS)
In 1829 Dobereiner classified the chemically similar elements in group of three. He noticed that the
atomic mass of the middle element is almost the arithmetic mean of the other two. The group of
three elements was called the law or rule of triads.
Central atom of each set of triad had an atomic mass almost equal to the arithmetical mean of the
atomic masses of the other two elements.
The atomic masses of Li and K are 7 and 39 respectively. The average of these two numbers is
23 which is the atomic mass of Na.
The atomic masses of Cl and I are 35.5 and 126.5 respectively. The average of these two
numbers is 81 which is the atomic mass of Br.
The atomic masses of Ca and Be are 40 and 137 respectively. The average of these two
numbers is 80 which is the atomic mass of sr.
Drawbacks:
This law or rule cannot be extended to the classification of all the elements, because it is true
only in the cases of very few elements.
2. NEWLAND’S LAW OF OCTAVE
In 1864 and English chemist, Newland’s proposed that if the elements are arranged in ascending
order of their atomic masses, every eighth element will have properties similar to the first
Statement:
If elements are arranged in the order of increasing atomic masses, the eighth element starting
from a given one, has similar properties as first one i.e. its properties are kind of repetition of
the first , like the eighth note in an octave of music.
Example:
Lithium (Li) and Sodium (Na) resemble with each other.
Beryllium (Be) resembles magnesium (Mg)
Fluorine (F) resembles Chlorine (Cl).
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Element
Li
Be
B
C
N
O
F
Atomic Mass
7
9
11
12
14
16
19
Element
Na
Mg Al
Si
P
S
Cl
Atomic Mass
23
24
27
28
31
32
35.5
Drawbacks:
This law failed because it held good for the first sixteen elements but did not work after
seventeenth element. Moreover hydrogen was not included in this sequence.
3. LOTHER MEYER CLASSIFICATION
Group:
In december 1869 Julius Lother Meyer, a German scientist published a periodic table in which he
arranged 56 elements on the basis of their atomic masses in nine vertical coulmns or groups I to IX.
He laid down emphasis on the physical properties of elements.
Atomic Volume:
Lother Meyer calculated the atomic voulmes of elements. The atomic volume of an element is the
volume which would be occupied by 1 gram Atomic weight (1 mole) of atoms of elements of it
were a solid.
Formula:
Atomic mass=
Gram atomic weight
Density
Graph:
He plotted a graph between atomic volume against increasing atomic masses of the elements. He
included about 50 elements. The curve obtained consists of sharp peaks and broad minima.
FACTS SEEN BY THE STUDY OF GRAPH
Similar Physical Properties:
The elements with physical properties like boiling points occupy similar positions on the
curve.
Position Of Alkali Metals:
Alkali metals like lithium (Li), Sodium (Na), potassium (K), rubidium (Rb), cesium (Cs)
occupy the peaks of the curves showing that these elements have largest atomic volume.
PERIODICITY OF ELEMENTS
Periodicity:
Similar elements were located at the similar positions of the curves. The regular spacing
of the highest positions confirms the idea of periodicity.
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MENDELEEV’S PERIODIC LAW
Statement:
The physical and chemical properties of elements are a periodic function of their atomic weight.
Mendeleev’s Periodic Table:
In Mendeleev’s periodic table, the elements were arranged in the increasing order of their
atomic masses in such a way that similar elements were repeated after regular intervals
and were placed one above the other.
SALIENT FEATURES OF MENDELEEV’S PERIODIC TABLE:
Periods and Groups:
In the table 12 horizontal rows from left to right are called periods. The 8 vertical
rows are called groups.
Similar Properties:
Elements in each vertical coulms have similar properties.
Vacant Spaces For Undiscovered Elements:
Mendeleev’s left certain vacant spaces in his table to place similar undiscovered
elements in the same group. He proposed their names as eka-boron, eka-aluminum
and eka-silicon.
Vacant Number:
The group number indicates the highest valence number that can be attained by
elements of that group.
Group
Group II Group III Group IV V
Group
VI
Group
VII
Be=9
Mg=24
Ca=40
B=11
Al=27.3
___=44
C=12
Si=28
Ti=48
N=14
P=32
B=51
O=16
S=32
Cr=52
F=19
Cl=35.5
Mn=55
Cu=63
Rb=85
Nz=65
Sr=87
___=68
Yt=88
___=72
Zr=90
As=75
Nh=94
Se=78
Mo=96
Br=80
___=100
Ag=108
Cs=133
-----
Cd=112
Be=137
-----
In=113
Di=138
--Er=178
Sn=118
Ce=140
--La=180
Sb=122
----Ta=182
Te=125
----W=184
I=127
-------
Row
1
2
3
4
Group I
H=1
Li=7
Na=23
K=39
5
6
7
8
9
10
11
Au=199 Hg=200 Tl=204
Pb=207
Bi=208
--12
------Th=231
--U=240
ADVANTAGES OF MENDELEEV’S PERIODIC TABLE:
Periodicity:
-----
Group VIII
Fe=56,Co=59
Ni=59,Cu=63
Ru=104,Rh=104
Pb=106,Ag=108
Os=195,Ir=197
Pt=198,Au=199
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It helped in systematic study of elements for example the study of sodium helps means
toa large extant in predicting the properties of other alkali metals as potassium, rubidium
and cesium. It proved the concept of periodicity.
Prediction Of New Elements:
Prediction of new elements was made possible example the physical and chemical
properties of eka-boron, eka-aluminum and eka silicon were predicted by Mendeleev.
This helped in their discovery. These have been named as scandium, gallium and germanium.
Their properties are the same as were predicted by Mendeleev.
Correction Of Atomic Masses:
Mendeleev’s periodic table helped in correcting many doubtful atomic masses.
Gradual Change In Physical Properties:
In Mendeleev’s table elements of any group resembled chemically with each other, but there was a
gradual change in physical properties going down the group.
Example:
The melting point of alkali metals in group I decrease gradually from top to the bottom.
DEFECTS IN MENDELEEV’S PERIODIC TABLE:
Arrangements Of Six Pair Of Elements:
There is three pair of elements i.e. elements of higher atomic masses placed before elements of
lower atomic masses i.e.
Argon (40) placed before Potassium (39)
Cobalt (59.9) placed before Nickel (58.6)
Tellurium (127.6) placed before Iodine (126.9)
No Place For Isotopes:
Placement Of Dissimilar Elements In The Same Group:
Dissimilar elements placed in the same group i.e. Alkali metals (Li, Na, K, Rb Cs, Fr) were
placed with the coinage metals (Ag, Cu Au)
Placements Of Similar Elements In Dissimilar Group:
Similar elements placed in different groups for example Barium (Ba) and Lead (Pb)
resembles in many properties but they are placed in separate groups.
No Idea Of Atomic Structure:
It filled to give the idea of atomic structure.
MODERN PERIODIC TABLE:
Modern periodic table is the result of the discovery of the atomic number by Moseley in 1914.
The physical and chemical properties of all elements are periodic function of their atomic
number.
Bohr’s Long Form of Periodic Table:
Modern periodic table is also known as Bohr’s long form of periodic table in which the elements are
arranged in order of their increasing atomic number. The elements having similar properties are
repeated after regular intervals. The modern periodic table contains seven horizontal rows called
periods and sixteen vertical columns called groups.
PERIODS:
The modern periodic table contains seven horizontal rows called periods.
The elements with a period have dissimilar properties from left to right across any period.
The physical and chemical properties of elements change from metallic to nonmetallic along a
period.
All periods except the firs starts with an alkali metal with one electron in their valence shell
end up with zero group elements with valence shell having eight electron except helium He which
has only two electrons.
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The First Period (Shortest Period)
It contains only two elements i.e. H and He.
This period signifies the completion of K shell or first orbit.
It is the shortest period with two elements.
Each of these contains eight elements.
They signify the filling of L shell and M shell respectively.
The second period starts with Lithium Li and end up with Ne whereas the third period starts
with Na and ends at Ar.
Each of these contains eighteen elements.
In these periods electrons fill M and N shells.
Fourth periods start from K and ends at Kr
Fifth period start from Rb and ends at Xe.
The Sixth Period (Longest Period)
It contains thirty two elements.
It starts from Cs and ends with Rn.
Besides fourteen elements called lanthanides are placed at the bottom of the periodic table.
The Seventh Period (Incomplete Period)
It starts with Francium Fr. This period is incomplete as to date about 109 elements have been
discovered.
This period also includes a group of fourteen elements starting from actinium. These
elements are called actinides. They are also placed at the bottom of the table.
Their valence shell contains one electron only and on reaction they lose this electron and
form univalent positive ions (M+1).
They are highly reactive metals with low melting points.
Fr is radioactive.
Their atomic radii, atomic volumes, ionic radii increases from Li to Cs due to the addition of
extra shell to each elements and due to same reason, the melting and the boiling points decreases
downward.
They are called alkali metals because they form water soluble base such as NaOH and KOH.
Elements:
Li (Lithium)
Na (Sodium)
K (Potassium)
Rb (Rubidium)
Cs (Caesium)
Fr (Francium
Group IIA (The Alkaline Earth Metals) (Beryllium Family)
Their valence shell contains two electrons.
On reaction they lose these two electrons and form divalent positive ions (M+2).
Ra is radioactive.
These elements are bit harder, having higher melting and boiling points than the alkali
metals, but they have similar atomic, ionic radii and atomic volume.
Down the group they do not show a regular trend in melting and boiling points and densities.
Elements:
Be (Beryllium)
Mg (Magnesium
Ca (Calcium)
Sr (Strontium
Ba (Barium)
Ra (Radium)
Group IIIA (The Boron Family)
Their valence shell contains three electrons.
They exhibit a valence of 3and form M+3 ions.
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Except boron they are highly electropositive elements i.e. having metallic character which
increases down the group due to increase in atomic volume.
Boron is metalloid. A metalloid is an element which has some properties of metals and some
properties of nonmetals.
Their valence shell contains four electron, C, Si and Se form covalent compounds whereas Sn
and Pb exhibit a variable valence of 2 and 4.
Of these elements C is nonmetal, Si and Ge are metalloids, Sn and Pb are metals.
Down the group atomic radii increases due to addition of a new shell and for the same reason
metallic character increases down the group.
C and Sn exist in different allotropic forms.
Elements:
C (Carbon)
Si (Silicon)
Ge (Germanium)
Sn (Tin)
Pb (Lead)
Group V (Nitrogen Family)
Of these elements N and P are nonmetals, As and Sb are metalloids and Bi is a metal.
Their valence shell contains five electrons.
There is a large variation of properties as we go down the group.
Nitrogen exists as diatomic molecules (N2) and forms a number of oxides as NO and N2O
Due to small atomic size and large ionization potential, nitrogen has a tendency to accept
three electrons to form nitride ion.
Phosphorous exists as P4 molecule.
Except nitrogen all exists in more than one allotropic form.
Elements:
N (Nitrogen)
P (Phosphorus)
As (Arsenic)
Sb (Antimony)
Bi(Bismuth)
Group VIA (Oxygen Family)
Of these oxygen and sulphur are nonmetals, selenium, tellurium are metalloids and polonium is
metal.
All the elements exhibit the property of allotropy. For example allotropic forms of oxygen
(O2) and ozone (O3).
Oxygen and sulphur form divalent negative ions O¯2 and S¯2. Their valence shell contains six
electrons.
Except astatine (which is metalloid) all others are nonmetals and exist as diatomic molecules.
At room temperature F2 and Cl2 are gases, bromine is a liquid and iodine is a solid.
Their valence shell contains seven electrons.
They have, high ionization energies and large negative electrons affinities hence they easily
accept an electron to form halide ion (X¯1) i.e. (F¯1, Cl¯1, Br¯1, I¯1)
Elements:
Their valence shell contains eight electrons, except helium which has two electrons.
With the exception of krypton and Xenon (which have large atomic volume so slightly
reactive under drastic conditions) the rest of these elements are totally inert chemically. The
reason is that these have completely filled outer shells, a condition that represents greater
stability.
Elements:
He (Helium)
Ne (Neon)
Kr (Krypton)
Xe (Xenon)
Rn (Radon)
Group IB To VIII B (Transition Elements)
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Characteristics:
These are metals.
In these elements, besides the valence shell penultimate shell is also incomplete.
They show more than one valence in chemical reaction.
These elements in compounds have characteristic colors.
METALS
Characteristics:
They are electropositive elements i.e. they lose electrons to form cat ion.
They form basic oxides.
All of them have luster and are malleable (i.e. can be spread out into sheet) and ductile (i.e.
can be drawn into wire), are good conductors of heat and electricity.
Examples:
In the periodic table elements of group IA, IIA and all transition elements are metals. Some of the
elements of group IIA, IVA, VA, and VIA are also metals. Sodium, Calcium, Iron, Gold, Silver
etc.
NON METALS
Characteristics:
They are electronegative elements i.e. they gain electrons to form an ions.
They form acidic oxides.
They are bad conductor of heat and electricity.
Most of them are gases.
Examples:
In the periodic table, majority of elements of p-block i.e. group IIIA, IVA, VA, VIA, VIIA and
VIIIA are nonmetals.
Chlorine, Oxygen, Bromine, Sulphur, Carbon etc.
METALLOIDS
Characteristics:
These are the elements which exhibit dual characters. That is they show the properties of both
metals as well as nonmetals. Their oxides are amphoteric i.e. have basic as well as acidic nature.
Examples:
Boron (B) of group IIIA.
Silicon (Si) and Germanium (Ge) of group IVA.
Arsenic (As) and Antimony (Sb) of group VA.
Tellurium (Te) and Polonium (Po) of group VIA.
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Astatine (At) of group VIIA.
DIFFERENCE BETWEEN GROUP IA AND VIIA
GROUP IA
GROUP VIIA
Group IA contains Lithium (Li) Sodium (Na), Group VIIA contains Fluorine (F), Chlorine
Potassium (K), Rubidium (Rb), Cesium (Cs)
(Cl), Bromine (Br), Iodine (I) and Astatine
and Francium (Fr).
(At).
Elements of this group are called alkali metals. Elements of this group are called halogens.
They contain one electron in their outer most
They contain seven electrons in their outer
shell.
most shell.
They form only ionic bonds.
They form ionic as well as covalent bonds.
They exist in monoatomic form e.g. Li, Na, K. They exist in diatomic form e.g. Cl2, Br2, I2
They have tendency to lose their single
They have tendency to gain electron.
electron.
They are strongly electropositive.
They are highly electronegative.
Their oxides give strong alkali when dissolve
Their oxides are acidic in nature.
in water
Na2O + H2O
2HOCl
Cl2O + H2O
2NaOH
Periodic Trends
The physical properties of elements that exhibit periodicity in the periodic table i.e. the physical
properties vary from element to element with the change in atomic number from left to right in
periods or from top to bottom in group are called Periodic Properties.
Periodic Trends
Horizontal & vertical trends can be seen in the elements for:
1. Atomic radius
2. Ionization energy
3. Electron affinity
4. Electronegativity
ATOMIC RADII
The atomic radii may be defined as half the distance between two adjacent nuclei of two similar
atoms in touch with each other.
Unit:
It is measured in Angstrom unit (Aº or A.U)
1Aº = 10 –8 cm = 10 –10 m
Dependence:
The atomic radii depend upon the number of shells and nuclear charge in an atom.
Trend In Groups:
In the periodic table the atomic radii increases down the group due to addition of new shell in
each atom.
Trend In Period:
In a period the atomic radii decreases from left to right due to increase in number of protons i.e.
increase in nuclear charge, which results in stronger pull on orbiting electrons by the nucleus.
or
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ATOMIC RADIUS
To find atomic radius, atoms are assumed to be spheres. The electron cloud size determines the
atomic radius for an atom. The radius values are only estimates. These values are measured by finding
the distance between 2 nuclei and dividing the distance by 2.
GROUP TREND: Atomic radius increases as you move from top to bottom in a family. This is
because major energy levels (1-7) are being filled with more & more electrons. The electrons get
farther & farther from the nucleus.
PERIOD TREND: Atomic radius generally decreases from left to right as atomic number
increases. This is because extra electrons are entering the same level while the nucleus gets larger
& more positive. This draws the electron cloud in towards the nucleus.
[Atomic Size of Elements]
[Atomic Size of Ions]
ATOMIC RADIUS OF IONS: When an atom loses an electron it has a positive charge. The
radius of the atom decreases because there’s a smaller electron cloud. When an atom gains an
electron it has a negative charge. The radius of the atom increases because the electron cloud is
larger.
IONIZATION ENERGY
It is defined as the minimum energy required to removes an electron from a gaseous atom in its
ground state.
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Unit:
It is measured in KJ/mole or electron volt (e.v) per atom.
Dependence:
Ionization energy depends upon atomic size and nuclear charge. The higher the ionization energy
the more difficult is to remove an electron. The ionization energy of hydrogen is 1312 KJ/mole
i.e.
H (g) + energy
H+ + e– I.E
=
+1312 KJ/mol
GROUP TREND: In vertical groups, ionization energy decreases from top to bottom. This is
because electrons are farther from the nucleus & filled levels cause a shielding effect.
SHIELDING EFFECT: Inner electrons shield outer electrons from the positive nucleus. This means
outer electrons are not held as tightly.
PERIOD TREND: Ionization energy tends to increase as you move from left to right toward the
noble gases. This is because metals tend to lose electrons & nonmetals tend to gain electrons. All of
them want to as stable as the noble gases.
ELECTRON AFFINITY
Electron affinity is the ability of an atom to attract and hold an extra electron.
GROUP TREND: Electron affinity decreases from top to bottom of a group. This is because it’s
easier for small atoms with the nucleus closer to the outer electrons to gain another electron.
PERIOD TREND: Electron affinity in a horizontal period increases from left to right. This is
because the desire to gain an electron increases the closer you get to fill the energy level. What do
you think the electron affinity of the noble gases is? Zero, they are happy like they are.
[Diagram of Ionization Energy]
[Diagram of Electron Affinity]
ELECTRONEGATIVITY
Definition:
Electronegativity is defined as the relative tendency of an atom in a molecule to attract shared
pair of electron to itself. It is denoted by a number.
Unit:
It has no unit.
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Trend In Group:
Down a group the electro negativity decreases. Due to addition of new shell, the power of a
nucleus to attract electron decreases.
Trend In A Period:
In a period from left to right it increases in nuclear charge.
LANTHANIDE
In the 6th period, after Lanthanum (57La) there are fourteen elements with atomic numbers 58 to
71. These elements have six electronic shells. The electron is gradually added in the 4f orbitals of
the 4th shell of these elements. The first member of the series is Lanthanum so all the 14 elements
are called Lanthanides. All lanthanides resemble with each other.
ACTINIDES
In the 7th period after actinium (89Ac) there are 14 elements with atomic number 90 to 103. These
elements have seven electronic shells. Electrons are gradually added in the 5f orbitals of the 5th
shell. Lanthanides and actinides are f-block elements.
Redox potential
Redox potential is a measure of the propensity of a chemical or biological species to either acquire
or lose electrons through ionization.
Oxidation potential and reduction potential are two types of electrode potential values for chemical
species given in Volts at standard conditions. Therefore, we name them standard oxidation potential
and standard reduction potential. The value of these potentials determines the ability of a particular
chemical species to undergo oxidation/reduction.
The key difference between oxidation potential and reduction potential is that oxidation potential
indicates the tendency of a chemical element to be oxidized. In contrast, reduction potential indicates
the tendency of a chemical element to be reduced.
What is Oxidation Potential?
The oxidation potential is a value that indicates the tendency of a chemical species to be oxidized. In
other words, it is the ability of an electrode to lose electrons (to get oxidized). Usually, this value is
given at standard conditions; hence, we should name it as standard oxidation potential. The
denotation for this term is SOP. It is measured in Volts. And, this is very similar to the standard
reduction potential, but they are different in the sign of the value, i.e. the value of the standard
oxidation potential is the negative value of the standard reduction potential. We can write the
oxidation potential as a half-reaction. The general formula for an oxidation reaction and the oxidation
potential for copper is given below:
Half reaction of copper oxidation: Cu(s) ⟶ Cu2+ + 2e–
The value for standard oxidation potential for the above reaction (oxidation of copper) is -0.34 V.
What is Reduction Potential
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Reduction potential is the tendency of a particular chemical species to undergo reduction. That
means; this particular chemical species is willing to accept electrons from outside (to get reduced).
It is measured in Volts and usually measured under standard conditions. Therefore, we can name it
as standard reduction potential. The denotation for this term is SRP. We can write it in the form of a
reduction half-reaction. The general formula and copper as an example are given below:
Half reaction of copper reduction: Cu2+ + 2e– ⟶ Cu(s)
The value for standard reduction potential for the above reaction (reduction of copper) is 0.34 V,
which is the exact value, but the opposite sign from that of the oxidation potential of the same
chemical species, copper. Therefore, we can develop a relationship between the standard oxidation
and reduction potentials as follows:
E00(SRP) = -E00(SOP)
What is the Difference between Oxidation Potential and Reduction Potential?
Oxidation potential and reduction potential are two types of electrode potential values for chemical
species given in Volts at standard conditions. The key difference between oxidation potential and
reduction potential is that oxidation potential indicates the tendency of a chemical element to be
oxidized, whereas the reduction potential indicates the tendency of a chemical element to be reduced.
Since these potential values are measured at standard conditions, we should name them as standard
oxidation potential and standard reduction potential. Moreover, we denote them as SOP and SRP.
Furthermore, there is a relationship between these two terms; the standard oxidation potential is the
exact same value but with a different sign from that of the standard reduction potential.
Measurement
Three-electrode setup for measurement of electrode potential
The measurement is generally conducted using a three-electrode setup (see the drawing):
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1. Working electrode
2. Counter electrode
3. SHE (standard hydrogen electrode or an equivalent).
In case of non-zero net current on the electrode, it is essential to minimize the ohmic IR-drop in the
electrolyte, e.g., by positioning the reference electrode near the surface of the working electrode (e.g.,
see Luggin capillary), or by using a supporting electrolyte of sufficiently high conductivity. The
potential measurements are performed with the positive terminal of the electrometer connected to the
working electrode and the negative terminal to the reference electrode.
Sign conventions
Historically, two conventions for sign for the electrode potential have formed:[2]
1. convention "Nernst–Lewis–Latimer" (sometimes referred to as "American"),
2. convention "Gibbs–Ostwald–Stockholm" (sometimes referred to as "European").
In 1953 in Stockholm IUPAC recognized that either of the conventions is permissible; however, it
unanimously recommended that only the magnitude expressed according to the convention (2) be
called "the electrode potential". To avoid possible ambiguities, the electrode potential thus defined
can also be referred to as Gibbs–Stockholm electrode potential. In both conventions, the standard
hydrogen electrode is defined to have a potential of 0 V. Both conventions also agree on the sign
of E for a half-cell reaction when it is written as a reduction.
The main difference between the two conventions is that upon reversing the direction of a half-cell
reaction as written, according to the convention (1) the sign of E also switches, whereas in the
convention (2) it does not. The logic behind switching the sign of E is to maintain the correct sign
relationship with the Gibbs free energy change, given by Δ G = -nFE where n is the number of
electrons involved and F is the Faraday constant. It is assumed that the half-reaction is balanced by
the appropriate SHE half-reaction. Since Δ G switches sign when a reaction is written in reverse, so
too, proponents of the convention (1) argue, should the sign of E. Proponents of the convention (2)
argue that all reported electrode potentials should be consistent with the electrostatic sign of the
relative potential difference.
Potential difference of a cell assembled of two electrodes
Potential of a cell assembled of two electrodes can be determined from the two individual electrode
potentials using
ΔVcell = Ered,cathode − Ered,anode
or, equivalently,
ΔVcell = Ered,cathode + Eoxy,anode.
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This follows from the IUPAC definition of the electric potential difference of a galvanic cell,
according to which the electric potential difference of a cell is the difference of the potentials of the
electrodes on the right and the left of the galvanic cell. When ΔVcell is positive, then positive electrical
charge flows through the cell from anode to cathode.