Classification of Elements
Elements
- Science has come a long way since Aristotle’s theory of Air, Water, Fire, and Earth.
- Scientists have identified 90 naturally occurring elements, and created about 28 others.
- As more elements were discovered in the 19th century chemists started to note similarities in their properties.
- Early attempts to order the elements in a regular fashion were hampered by various difficulties.
Mendeleef’s Periodic Law’ and Periodic Table:
Attempts were made to classify the elements in a number of ways. In 1869, a Russian scientist, Dmitri Mendeleef made
the most significant contribution towards
the classification of elements. Mendeleef observed that when all the 65 elements (known at that time) were arranged
according to increasing atomic weights,
similarities and differences in their properties ‘would be apparent. This was enunciated in the form of a PERIODIC
LAW which was stated as:
the most significant contribution towards
the classification of elements. Mendeleef observed that when all the 65 elements (known at that time) were arranged
according to increasing atomic weights,
similarities and differences in their properties ‘would be apparent. This was enunciated in the form of a PERIODIC
LAW which was stated as:
The physical and chemical properties of elements are a periodic function of their atomic weights, i.e., if the
elements are arranged in the increasing order of their atomic weights, the properties of the elements (i.e., similar elements) are repeated after definite regular intervals.
elements are arranged in the increasing order of their atomic weights, the properties of the elements (i.e., similar elements) are repeated after definite regular intervals.
Mendeleef gave a detailed comparison of the physical and chemical properties of the elements and set up a periodic
system in which the elements were arranged in horizontal ROWS (series) and vertical COLUMNS (groups) according
to increasing atomic weights. Mendeleef arranged the elements in the form of a table which is known as Mendeleeff’s
Periodic Table after his name.
system in which the elements were arranged in horizontal ROWS (series) and vertical COLUMNS (groups) according
to increasing atomic weights. Mendeleef arranged the elements in the form of a table which is known as Mendeleeff’s
Periodic Table after his name.
Mosley’s Modern Periodic Law:
With the advancement or the knowledge about atomic structure and discovery of new elements, in 1913, Mosley, a British
physicist, predicted that most of the defects of Mendeleef’s periodic table disappear, if the basis of classification of elements is
changed to atomic number in place of atomic weight.
physicist, predicted that most of the defects of Mendeleef’s periodic table disappear, if the basis of classification of elements is
changed to atomic number in place of atomic weight.
Accordingly Mosley put forward Modern Periodic Law which is stated as follows:
The physical and chemical properties of the elements are periodic function of their atomic numbers, i.e., if the elements
are arranged in the increasing order of their atomic numbers, the properties of the elements (i.e. similar elements) are
repeated after regular definite intervals.
are arranged in the increasing order of their atomic numbers, the properties of the elements (i.e. similar elements) are
repeated after regular definite intervals.
The Modern Periodic Table (Extended or Long form of Periodic Table):
The original Periodic Table suggested by Mendeleef has undergone many modifications to remove the defects of
Mendeleef’s periodic table, although the basic features have been maintained in all the modified forms. Out of the
various tables the Extended Long Form of Periodic Table which is based on the Aufbau principle (building up of the
atomic electronic configurations) is the most simple and is widely accepted.
Mendeleef’s periodic table, although the basic features have been maintained in all the modified forms. Out of the
various tables the Extended Long Form of Periodic Table which is based on the Aufbau principle (building up of the
atomic electronic configurations) is the most simple and is widely accepted.
Characteristics of Modern Periodic table:
In the modern periodic table 110 elements have been arranged according to the increasing atomic number.
The main two parts of the periodic table are
- Groups
- Period
GroupsThe vertical columns shown in the periodic table are called groups or families or simply columns.
(a) There are nine groups in all including VIII group consisting of’ three triads
(Fe, Co, Ni; Ru, Rh, Pd; Os, Ir, Pt) and zero groups of inert gases. Groups I to VII are sub-divided into sub-groups A and B.
(Fe, Co, Ni; Ru, Rh, Pd; Os, Ir, Pt) and zero groups of inert gases. Groups I to VII are sub-divided into sub-groups A and B.
Thus there are 18 vertical columns which are: IA, IIA, IIIA, IVA, VA, VIA, VIIA, Zero, IB, IIB, IIIB, IVB, VB, VIB, VIIB and three
columns of Group VIII.
columns of Group VIII.
(b) Elements of groups IA, IIA, IIIA, IVA, VA, VIA and VIIA have their outermost shells incomplete while each of their inner
shell is complete. These elements are called normal or representative elements. These elements consist of’ some metals,
all non-metals and metalloids.
shell is complete. These elements are called normal or representative elements. These elements consist of’ some metals,
all non-metals and metalloids.
(c) Elements of groups IB, IIB, IIIB (only Sc, Y, La and Ac), IVB, VB, VIB, VIIB and VIII have their two outermost shells
incomplete. These are called transition elements. These elements are placed in the middle of the table. All these elements are
metals.
incomplete. These are called transition elements. These elements are placed in the middle of the table. All these elements are
metals.
(d) Elements of group zero have satisfied octet in their outermost shell. These elements are called noble gases. These are
placed at the extreme right of the table.
placed at the extreme right of the table.
(e) Two groups of 14 elements lying in group IIIB [Ce (Z=58) to Lu (Z=71) and
Th (Z=90) to Lr (Z=103)] have their three outermost shells incomplete. These are called Lanthanides and Actinides
respectively and have been placed at the bottom of the table.
Th (Z=90) to Lr (Z=103)] have their three outermost shells incomplete. These are called Lanthanides and Actinides
respectively and have been placed at the bottom of the table.
Periods
The horizontal rows shown in the periodic table are called periods or simply rows. There are seven periods in the table.
(a) 1st period consists of 2 elements which are H (Z= 1) and He (Z=2).
(b) 2nd and 3rd periods have 8 elements each.
2nd period → Li (Z=3) to Ne (Z=10)
3rd period → Na (Z=11) to Ar (Z=18)
Both these periods are called short periods.
(c) 4th and 5th periods have 18 elements each while 6th period has 32 elements as shown below:
4th period → K (Z=19) to Kr (Z=36)
5th period → Rb (Z=37) to Xe (Z=54)
6th period → Cs (Z=55) to Rn (Z=86)
All these three periods are called long periods. 6th period also includes 14 rare earths or
lanthanides [Ce (Z=58) to Lu (Z=71)].
lanthanides [Ce (Z=58) to Lu (Z=71)].
(d) 7th period is an incomplete period and at present it consists of 24 elements which are Fr (Z=87) to Ds (Z=110).
All these elements of this period are radioactive. This period also includes 14 actinides [Th (Z=90) to Lr (Z=103)].
All these elements of this period are radioactive. This period also includes 14 actinides [Th (Z=90) to Lr (Z=103)].
General Characteristics of Groups:
1. Number of valency electrons: On moving down a given group the number of valence electrons does not change,
i.e. remains the same.
i.e. remains the same.
Element
|
Atomic No.
|
Electronic configuration
|
Valence shell electronic configuration
|
Li
|
3
|
1s1 2s1
|
2s1
|
Na
|
11
|
1s2 2s2 2p6 3s1
|
3s1
|
K
|
19
|
1s2 2s2 2p6 3s2 3p6 3d10 4s1
|
4s1
|
Rb
|
37
|
1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 5s1
|
5s1
|
2. Valency✯1: The valencies of all the elements of the same group are the same.
3. Properties of elements: All the elements of a given group possess very similar physical and chemical properties.
There is a regular gradation in their properties when we move from top to bottom in a group.
There is a regular gradation in their properties when we move from top to bottom in a group.
For example:
(a) The alkali metals (group IA) resemble each other and their base-forming tendency increases from Li to Cs.
(b) The reactivity of halogens (group VIIA) decreases as we pass from F to I.
4. Size of atoms: Size of atoms increases on descending a group. For example in group IA, atomic size increases from Li to Cs.
Thus, Li<Na<K<Rb<Cs
Thus, Li<Na<K<Rb<Cs
5. Metallic character: The metallic character of the elements increases in moving from top to bottom in a group.
This is particularly apparent in groups IVA, VA and VI A, which begin with non-metals (namely C, N and O respectively),
and end with metals (namely Pb, Bi and Po respectively). For example, in group VA, N and P are non- metals, As and Sb are
metalloids and Bi is a typical metal.
This is particularly apparent in groups IVA, VA and VI A, which begin with non-metals (namely C, N and O respectively),
and end with metals (namely Pb, Bi and Po respectively). For example, in group VA, N and P are non- metals, As and Sb are
metalloids and Bi is a typical metal.
Thus the metallic character of these elements increases from N to Bi as shown below:
It is because of a gradual increase of the metallic character of the elements from top to bottom that the oxides of the elements
become more and more basic in the same direction. For example:
become more and more basic in the same direction. For example:
Oxides of the elements of group VA-
6. Number of electron shells: In going down a group the number of electron shells increases by one at each step and
ultimately becomes equal to the number of the period to which the element belongs as shown below for the elements of
Group IA.
ultimately becomes equal to the number of the period to which the element belongs as shown below for the elements of
Group IA.
Elements
|
Electronic configuration
|
No. of shells
|
Li (3)
|
2, 1
|
2
|
Na (11)
|
2, 8, 1
|
3
|
K (19)
|
2, 8, 8, 1
|
4
|
Rb (37)
|
2, 8, 18, 8, 1
|
5
|
Cs (55)
|
2, 8, 18, 18, 8, 1
|
6
|
Fr (87)
|
2, 8, 18, 32, 18, 8, 1
|
7
|
Li (3) → 1s2, 2s1
Na (11) → 1s2, 2s2, 2p6, 3s1
K (19) → 1s2, 2s2, 2p6, 3s2, 3p6, 4s1
Rb (37) → 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 5s1
Cs (55) → 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10,5s2, 5p6, 6s1
Fr (87) → 1s2, 2s2, 2p6, 3s2, 3p6, 3d10, 4s2, 4p6, 4d10, 4f14, 5s2, 5p6, 5d10, 6s2, 6p6, 7s1
General characteristics of periods:
1. Number of valency electrons: Number of valency electrons increases from 1 to 8 when we proceed from left to right in a
period.
period.
2. Valency: The valency of the elements with respect to (w.r.t) hydrogen in each short period increases from 1 to 4 then falls to
one while the same with respect to oxygen increases from 1 to 7 as shown below for the elements of 2nd and 3rd period:
one while the same with respect to oxygen increases from 1 to 7 as shown below for the elements of 2nd and 3rd period:
Elements of 2nd period
|
Li
|
Be
|
B
|
C
|
N
|
O
|
F
| ||||||
Hydrides of the elements
|
LiH
|
BeH2
|
BH3
|
CH4
|
NH3
|
H2O
|
HF
| ||||||
Valency of the elements w.r.t. Hydrogen
|
1
|
2
|
3
|
4
|
3
|
2
|
1
| ||||||
Elements of 3rd period
|
Na
|
Mg
|
Al
|
Si
|
P
|
S
|
Cl
| ||||||
Oxides of the elements
|
Na2O
|
MgO
|
Al2O3
|
SiO2
|
P2O5
|
SO3
|
Cl2O7
| ||||||
Valency of the elements w.r.t. Oxygen
|
1
|
2
|
3
|
4
|
5
|
6
|
7
| ||||||
3. Size of atoms: Size of atoms decreases from left to right in a period.
Thus alkali metals have the largest size while the halogens have the smallest size.
Thus alkali metals have the largest size while the halogens have the smallest size.
4. Properties of elements: The properties of the elements of a given period differ considerably but the elements in the two
adjacent periods show marked similarity between them. For example, when we consider the elements of 2nd and 3rd periods,
we find that Na resembles Li, Mg resembles Be, Al resembles B, Si resembles C,
P resembles N, S resembles O, Cl resembles F and Ar resembles Ne.
adjacent periods show marked similarity between them. For example, when we consider the elements of 2nd and 3rd periods,
we find that Na resembles Li, Mg resembles Be, Al resembles B, Si resembles C,
P resembles N, S resembles O, Cl resembles F and Ar resembles Ne.
5. Metallic character: On moving from left to right in a period the metallic character of the elements decreases.
For example in 3rd period, Na, Mg and Al are metals while Si, P, S and Cl are non-metals as shown below:
For example in 3rd period, Na, Mg and Al are metals while Si, P, S and Cl are non-metals as shown below:
6. Acidity and Alkalinity: The gradual decrease of the metallic character from left to right shows that, the oxides of the
elements become less and less basic in the same direction. For example:
elements become less and less basic in the same direction. For example:
Oxides of the elements of 3rd period
|
Na2O
|
MgO
|
Al2O3
|
SiO2
|
P2O5
|
SO3
|
Cl2O7
|
Strongly basic
|
Basic
|
Amphoteric
|
Freely acidic
|
Acidic
|
More acidic
|
Most acidic
|
7. Number of shells: In going from left to right in a period the number of electron shells remains the same and the
number of a period corresponds to the number of the shells found in the elements of that period, e.g. all the elements of
2nd period have the electrons only in first two shells as shown below:
number of a period corresponds to the number of the shells found in the elements of that period, e.g. all the elements of
2nd period have the electrons only in first two shells as shown below:
Elements of 2nd period
|
Li
|
Be
|
B
|
C
|
N
|
O
|
F
|
Ne
|
Atomic number
|
3
|
4
|
5
|
6
|
7
|
8
|
9
|
10
|
Electronic configuration
|
2, 1
|
2, 2
|
2, 3
|
2, 4
|
2, 5
|
2, 6
|
2, 7
|
2, 8
|
No. of shells
|
2
|
2
|
2
|
2
|
2
|
2
|
2
|
2
|
8. Diagonal relationship: Sometimes, an element in the periodic table shows similarity of properties with another
element of the next group and the next period diagonally. Such types of relationship between two elements are known as
diagonal relationship. The following are the important examples of diagonal relationship found in the periodic table.
i. Li→Mg diagonal relationship
ii. Be→Al relationship
iii. B→Si relationship
Diagonal relationship is the resemblance of the properties of the elements of 2nd period with their diagonally opposite
members lying in 3rd period.
members lying in 3rd period.
Classification of elements of the periodic Table:
The elements displayed on the periodic table are classified as:
- Metalloids
- Non-metals
- Alkali metals
- Halogens
- Alkaline earth metals
- Noble gases
- Transition metals
- Rare earth elements
- Other metals.
✿Metalloids: The 5 elements classified as “metalloids” are located in groups 13 (IIIA), 14 (IVA), 15 (VA) and 16 (VIA)
of the periodic table.
of the periodic table.
These elements have properties of metals and non-metals. Some are semi-conductors and can carry an electrical charge
making them useful in calculators and computers.
making them useful in calculators and computers.
The metalloids are:
- Boron (B-5) [group 13]
- Silicon (Si-14) [group 14]
- Arsenic (As-33) [group 15)
- Selenium (Se-34) [group 16]
- Tellurium (Te-52) [group 16]
✿ Alkali metals: The 6 elements classified as “alkali metals” are located in group 1 (IA) of the periodic table.
The alkali metals are: Lithium (Li-3) ; Sodium (Na-11) ; Potassium (K-19) ;
Rubidium (Rb-37) ; Cesium (Cs-55) ; Francium (Fr-87)
Rubidium (Rb-37) ; Cesium (Cs-55) ; Francium (Fr-87)
These elements are collectively called alkali metals, since they form strongly alkaline oxides and hydroxides.
Alkali metals are very reactive metals that do not occur freely in nature. They are malleable, ductile and good conductors
of heat and electricity. Fr is a radioactive element.
of heat and electricity. Fr is a radioactive element.
✿ Alkaline Earth Metals: The 6 elements classified as “Alkaline Earth Metals” are located in Group 2 of the Periodic Table.
The Alkaline Earth Metals are: Beryllium (Be-4); Magnesium (Mg-12); Calcium (Ca-20)
Strontium (Sr-38); Barium (Ba-56); Radium (Ra-88).
Strontium (Sr-38); Barium (Ba-56); Radium (Ra-88).
The oxides of the three metals viz., Ca, Sr and Ba were known much earlier than the metals themselves and were called
alkaline earths, since they were alkaline in character and occurred in nature as earths [lime (CaO), strontia (SrO) and baryta
(BaO)]. Later, when Ca, Sr and Ba were discovered, they were named alkaline earth metals. Now this term is used to include
all the elements of Group II A.
alkaline earths, since they were alkaline in character and occurred in nature as earths [lime (CaO), strontia (SrO) and baryta
(BaO)]. Later, when Ca, Sr and Ba were discovered, they were named alkaline earth metals. Now this term is used to include
all the elements of Group II A.
Alkaline Earth Metals are all found in the Earth’s crust, but not in the elemental form as they are so reactive. Instead, they are
widely distributed in rock structures. Although Ra has similar properties as alkaline earth metals, it is a radioactive element.
widely distributed in rock structures. Although Ra has similar properties as alkaline earth metals, it is a radioactive element.
✿ Transition Metals: The elements classified as “Transition Metals” are located in Groups 3- l2 (group VIII and subgroup B)
of the Periodic Table. The d-block elements are called transition elements because they exhibit transitional behaviour
between highly reactive ionic compound forming s-block elements (electropositive elements) on one side and mainly
covalent compound forming
p-block elements (electronegative elements) on the other side.
of the Periodic Table. The d-block elements are called transition elements because they exhibit transitional behaviour
between highly reactive ionic compound forming s-block elements (electropositive elements) on one side and mainly
covalent compound forming
p-block elements (electronegative elements) on the other side.
Transition Metals are ductile, malleable, and conduct electricity and heat.
The Transition metals are:
The Transition metals are:
Scandium (Sc-21)
|
Yttrium (Y-39)
|
Titanium (Ti-22)
|
Zirconium (Zr-40)
|
Vanadium (V-23)
|
Niobium (Nb-41)
|
Chromium (Cr-24)
|
Molybdenum (Mo-42)
|
Manganese (Mn-25)
|
Technetium (Tc-43)
|
Iron (Fe-26)
|
Ruthenium (Ru-44)
|
Cobalt (Co-27)
|
Rhodium (Rh-45)
|
Nickel (Ni-28)
|
Palladium (Pd-46)
|
Copper (Cu-29)
|
Silver (Ag-47)
|
Zinc (Zn-30)
|
Cadmium (Cd-48)
|
Lanthanum (La-57)
|
Actinium (Ac-89)
|
Hafnium (Hf-72)
|
Rutherfordium (Rf-104)
|
Tantalum (Ta-73)
|
Dubnium (Db-105)
|
Tungsten (W-74)
|
Seaborgium (Sg-106)
|
Rhenium (Re-75)
|
Bohrium (Bh-107)
|
Osmium (Os-76)
|
Hassium (Hs-108)
|
Iridium (Ir-77)
|
Meitnerium (Mt-109)
|
Platinum (Pt-78)
|
Darmstadtium (Ds-110)
|
Gold (Au-79)
| |
Mercury (80)
|
✿ Rare earth elements: The elements classified as “Rare earth elements” are located in group 3 of the Periodic Table and
in the 6th and 7th periods. The Rare Earth Elements are of the Lanthanide and Actinide series.
in the 6th and 7th periods. The Rare Earth Elements are of the Lanthanide and Actinide series.
They are hardly being found in earth. Most of the elements in the Actinide series are synthetic or man-made. The Lanthanide
and Actinide series of Rare Earth Elements are:
and Actinide series of Rare Earth Elements are:
Lanthanide elements
|
Actinide elements
|
Lanthanum (La-57)
|
Actinium (Ac-89)
|
Cerium (Ce-58)
|
Thorium (Th-90)
|
Praseodymium (Pr-59)
|
Protactinium (Pa-91)
|
Neodymium (Nd-60)
|
Uranium (U-92)
|
Promethium (Pm-61)
|
Neptunium (Np-93)
|
Samarium (Sm-62)
|
Plutonium (Pu-94)
|
Europium (Eu-63)
|
Americium (Am-95)
|
Gadolinium (Gd-64)
|
Curium (Cm-96)
|
Terbium (Tb-65)
|
Berkelium (Bk-97)
|
Dysprosium (Dy-66)
|
Californium (Cf-98)
|
Holmium (Ho-67)
|
Einsteinium (Es-99)
|
Erbium (Er-68)
|
Fermium (Fm-100)
|
Thulium (Tm-69)
|
Mendelevium (Md-101)
|
Ytterbium (Yb-70)
|
Nobelium (No-102)
|
Lutetium (Lu-71)
|
Lawrencium (Lr-103)
|
✿ Other metals: The 10 elements classified as “Other metals” are located in Groups 13 (IIIA), 14 (IVA), 15 (VA) and 16(VIA)
of the Periodic Table.
of the Periodic Table.
All of these elements are solid, have a relatively high density and are opaque. The “Other Metals” are:
Group 13 (IIIA)
|
Group 14 (IVA)
|
Group 15 (VA)
|
Group 16 (VIA)
|
Aluminum (Al-13)
|
Germanium (Ge-32)
|
Antimony (Sb-51)
|
Polonium (Po-84)
|
Gallium (Ga-31)
|
Tin (Sn-50)
|
Bismuth (Bi-83)
| |
Indium (In-49)
|
Lead (Pb-82)
| ||
Thallium (Ti-81)
|
✿ Non-Metals: The 7 elements classified as “Non Metals” are located in Groups 1 (IA), 14 (IVA), 15 (VA) and 16 (VIA) of the
Periodic Table.
Periodic Table.
Group 1 (IA)
|
Group 14 (IVA)
|
Group 15 (VA)
|
Group 16 (VIA)
|
Hydrogen (H-1)
|
Carbon (C-6)
|
Nitrogen (N-7)
|
Oxygen (O-8)
|
Phosphorus (P-15)
|
Sulfur (S-16)
|
Non-metals are not easily able to conduct electricity or heat and do not reflect light. Non-metallic elements are very brittle,
and cannot be rolled into wires or pounded into sheets. They exist at room temperature, in two of the three states of matter:
gases (such as oxygen) and solids (such as carbon).
and cannot be rolled into wires or pounded into sheets. They exist at room temperature, in two of the three states of matter:
gases (such as oxygen) and solids (such as carbon).
✿ Halogens: The 5 elements classified as “Halogens” are located in Group 17 (VIIA) of the Periodic Table.
The term “Halogen” means “salt-former” and compounds containing halogens are called “salts”. The halogens exist, at room
temperature, in all three states of matter— gases such as fluorine & chlorine, solids such as iodine and astatine and liquid as
bromine. The Halogens are:
temperature, in all three states of matter— gases such as fluorine & chlorine, solids such as iodine and astatine and liquid as
bromine. The Halogens are:
|
Gas
|
|
Liquid
|
|
solid
|
Astatine was discovered in 1940 and is an unstable element of radioactive origin. The other four elements are stable and
resemble each other in physical and chemical properties.
resemble each other in physical and chemical properties.
✿ Noble gases or inert Gases: The 6 elements classified as “Noble Gases” are located in Group 18 (0) of the Periodic Table.
The outermost orbit in all these elements is completely filled. Therefore, these elements are chemically inert. The Noble
Gases
on the periodic Table are:
Gases
on the periodic Table are:
- Helium (He-2)
|
- Argon (Ar-18)
|
- Xenon (Xe-54)
|
- Neon (Ne-10)
|
- Krypton (Kr-36)
|
- Radon (Rn-86)
|
Classification of elements on the basis of electronic configuration
According to the electronic configurations, the elements may be divided into four types such as:
- The Inert Gases (Elements of 0 group).
- The Representative Elements (s and p block elements).
- The Transition Elements (d block elements).
- The Inner Transition Elements (f block elements).
The Inert Gases:
- The noble or inert gases (zero group elements) have been placed at the end of each period in the periodic Table.
It appears that all these elements have satisfied octet in their outermost orbitals. - Helium has 2s2 stable arrangement and all other inert gases have s2p6 outer configurations:
He (2) = 1s2 Ne (10) = 1s2 2s2 2p6 Ar (18) = 1s2 2s2 2p6 3s2 3p6
- The configuration shows duplet and octet in the outermost energy levels. Therefore, they are chemically inert.
As a result, their valencies are zero. - Therefore, the position of the noble gases should be in zero groups.
- It may be noted that no atom has a complete energy level except helium and neon.
- These elements are colorless gases.
The Representative Elements (s and p block elements):
- These elements generally belong to A sub-group of the Periodic Table.
- s-block elements: The elements in which the last electron(s) enters the s-orbital of their outermost energy layer are
called s-block elements. - Thus the alkali metals (Group IA), alkaline earth metals (Group IIA) are s block or s orbital elements.
Example: Na (11) – 1s2 2s2 2p6 3s1
- p-block elements: The elements in which the last electron(s) enters to the p-orbital of their outermost energy layer are
called p-block elements. - The valence electrons of all the elements from boron to halogens (groups IIIA to VIIA vertically) occupy p orbitals.
Hence these elements are called p block or p orbital elements.
Example:
Al (13) : 1s2 2s2 2p6 3s2 3p1
Cl (17) : 1s2 2s2 2p6 3s2 3p5
The Transition Elements (d block elements):
- The elements in which the last electron(s) enters to the d-orbital which is inner to the outer-most shell are called
d-block elements. - The elements of group VIII and sub-group B are generally the d-block elements.
- These elements contain two incomplete energy levels because of the building up of the inner d electrons.
Example: Sc (21) : 1s2 2s2 2p6 3s2 3p6 3d1 4s2
Fe (26) : 1s2 2s2 2p6 3s2 3p6 3d6 4s2
- Elements which have normally the same number of electrons in the outermost level but have a progressively greater
number of electrons in an inner level (such as d level) are called “Transition Elements”. - In the Periodic Table we come across four such transition series in which the additional electrons enter the 3d, 4d, 5d
and 6d orbitals.
Fig. List of Transition Elements (d block elements)
- The first transition series of elements involving the completion of 3d level start from Sc (21) to Zn (30).
- The second series of transition elements start from Y (39) up to Cd (48) involving 4d energy level.
- The third group of transition metals starts from La (57) but with a break from Ce (58) to Lu (71) which are classified as
inner transition metals and proceed up to Hg (80) involving 5d energy level.
Properties of Transition Elements:
- All the elements are of high melting points, electropositive and heavy metals.
- These metals have almost the same atomic and ionic sizes. There is only slight increase in the ionization energy of the
formation of M+2 ions. - All these elements show positive oxidation states of +2 and +3 generally and form mostly ionic compounds. Higher
oxidation states are also exhibited in some cases. - As a general rule, the transition elements form colored compounds.
- These elements are also effective catalytic agents.
- All these form quite a large number of complex compounds.
These properties are due to the influence of the incomplete inner d orbitals in the transition elements. The properties are
similar in the case of inner transition elements where f orbitals are being completed.
similar in the case of inner transition elements where f orbitals are being completed.
The inner transition elements (f block elements):
- The elements in which the last electron(s) enters into the (n-2)f-orbital are called f-block elements.
- These elements are located in group IIIB and have three incomplete outer levels.
- Since (n-2)f orbital lies comparatively deep within the kernel (being inner to the penultimate shell), these elements
are also called inner-transition elements. - The f-block elements consist of two series of elements which are placed in two rows at the bottom of the periodic table.
- The first series of 14 elements (atomic numbers 58 to 71) in which 4f level is being build up follows lanthanum (57) and
are called Lanthanides. Example:
Ce (58) → 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f1 5d1
→ 1s2, 2s2p6, 3s2p6d10, 4s2p6d10f1, 5s2p6d1, 6s2
→ 2, 8, 18, 19, 9, 2
- Another series of 14 elements (atomic numbers 90 to 103) in which 5f level is being filled follows actinium and is known
as Actinides. Example:
Th (90) → 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f1 6d1
→ 1s2, 2s2p6, 3s2p6d10, 4s2p6d10f14, 5s2p6d10f1, 6s2p6d1, 7s2
→ 2, 8, 18, 32, 19, 9, 2
The inner transition elements (lanthanides and actinides) are all metals and show variable oxidation states. Their compounds
are highly colored.
are highly colored.
Defects of Mendeleef’s Periodic Table:
Although the Mendeleef’s Periodic Table was the first successful attempts for the classification of elements but it suffers from
the following defects:
the following defects:
(i) Position of hydrogen: Hydrogen resembles both the alkali metals (group IA and halogen (group VIIA) in properties.
Therefore, its position in the periodic table is anomalous.
Therefore, its position in the periodic table is anomalous.
(ii) Position of lanthanides and actinides: A group of 15 elements (At. No. 57 to 71) which is called rare earths or lanthanides
does not find its proper place in the table and has been placed at one place in group III and period 6. Similarly, another group
of 15 elements (At. No. 89 to 103) called actinides does not find its proper place and has been put at one place in group III and
period 6.
does not find its proper place in the table and has been placed at one place in group III and period 6. Similarly, another group
of 15 elements (At. No. 89 to 103) called actinides does not find its proper place and has been put at one place in group III and
period 6.
These two groups do not find its proper place in the table and have been placed all together separately.
(iii) Existence of four anomalous pairs of elements: The order of increasing atomic weight has been ignored in case of four
pairs of elements in order to place them in a position justified by their properties.
pairs of elements in order to place them in a position justified by their properties.
Thus elements of higher atomic weights precede those of lower atomic weight at four places as shown below
|
K (Z=19, at. wt. =39.0)
|
|
Ni (Z=28, at. wt. =58.6)
|
|
I (Z=53, at. wt. =126.9)
|
|
Pa (Z=91, at. wt. =231)
|
(iv) Similar elements are separated and dissimilar elements are placed in the same group: Elements with similar
properties like Cu and Hg, Ag and Th, Ba and Pb are separated while dissimilar elements like Cu, Ag, and Au are grouped
along with the alkali metals. Mn is grouped with the halogens.
properties like Cu and Hg, Ag and Th, Ba and Pb are separated while dissimilar elements like Cu, Ag, and Au are grouped
along with the alkali metals. Mn is grouped with the halogens.
(v) Position of isotopes: If the elements are arranged in the order of their increasing atomic weights, it is not possible to
accommodate large number of isotopes in the periodic table.
accommodate large number of isotopes in the periodic table.
(vi) Group does not represent valency: Excepting osmium, elements placed in group eight do not show a valency of 8. Also
the elements lying in the middle of long periods show two or more valencies .e.g. Cr, Mn etc.
the elements lying in the middle of long periods show two or more valencies .e.g. Cr, Mn etc.
Mosley's Modern Periodic Law (Characteristics)
With the replacement of atomic weight by atomic number as the basis of classification of elements, many of the irregularities in
the Mendeleef's table disappear as shown below:
the Mendeleef's table disappear as shown below:
1. Position of hydrogen. The dual role of hydrogen is explained by the fact that it has one electron in its outer orbit. It has
equal tendency of gaining or losing one electron for assuming a stable configuration. When it loses one electron to give H+,
it resembles alkali metals (which give Li+, Na+, K+, etc. ions) while when it gains one electron to give H–, it resembles halogens
(which give Cl–, Br– etc.).
equal tendency of gaining or losing one electron for assuming a stable configuration. When it loses one electron to give H+,
it resembles alkali metals (which give Li+, Na+, K+, etc. ions) while when it gains one electron to give H–, it resembles halogens
(which give Cl–, Br– etc.).
2. Anomalous pairs of elements. This anomaly disappears altogether and the pairs Ar—K, Co—Ni, Te—I and Th—Pa are
found arranged in the table in the order in increasing atomic numbers as shown below:
found arranged in the table in the order in increasing atomic numbers as shown below:
Pairs of elements→
|
Ar
|
K
|
Co
|
Ni
|
Th
|
I
|
Ta
|
Pa
|
Atomic numbers→
|
18
|
19
|
27
|
28
|
52
|
53
|
90
|
91
|
Atomic weights→
|
40
|
39
|
59.9
|
58.6
|
127.6
|
126.9
|
232.12
|
231
|
3. Position of rare earths. The arrangement of extranuclear electrons in all the rare earth elements can be represented as
2, 8, 18 (18 + x), 9, 2, where x varies from 0 (for La) to 14 (for Lu). With this general arrangement of electrons, all of them
possess the same valency and similar chemical properties. This justifies their grouping at the same place.
2, 8, 18 (18 + x), 9, 2, where x varies from 0 (for La) to 14 (for Lu). With this general arrangement of electrons, all of them
possess the same valency and similar chemical properties. This justifies their grouping at the same place.
4. Position of isotopes. Since isotopes of the same element possess the same atomic number, all of them should occupy one
and the same place in the periodic table.
and the same place in the periodic table.
5. Justification for dissimilar elements being placed together.
The length of the periods is determined by arrangement of electrons in different orbits. The end of every period results from
the completion of the last orbit (last number is always an inert gas). Different periods carry 2, 8, 18 and 32 elements.
the completion of the last orbit (last number is always an inert gas). Different periods carry 2, 8, 18 and 32 elements.
When 18 elements are to be distributed among 8 groups; groups 1-7 get two elements each while group 0 gets only one.
The three elements which cannot be arranged elsewhere are placed in a special group VIII.
The three elements which cannot be arranged elsewhere are placed in a special group VIII.
This lack of space is enough justification for group VIII.
Out of the two elements which every long period adds to a group, one resembles the typical element, the other does not.
This gives rise to the formation of subgroups.
This gives rise to the formation of subgroups.
This explains why dissimilar elements have been grouped together.
Merits of Long Form of Periodic Table over Mendeleef's Periodic Table
The long form of periodic table has a number of merits over the Mendeleef's periodic table in the following respects
(1) The classification of the elements is based on a more fundamental property
viz., atomic number.
viz., atomic number.
(2) It relates the position of an element to its electronic configuration. Thus each group contains elements with similar
electronic configuration and hence similar properties. For example, all the alkali metals have similar valence-shell electronic
configuration
viz. ns 1 configuration and hence have similar properties.
electronic configuration and hence similar properties. For example, all the alkali metals have similar valence-shell electronic
configuration
viz. ns 1 configuration and hence have similar properties.
Alkali Metals
|
Atomic No.
|
Complete Electronic configuration
|
Valence shell electronic configuration
|
Li
|
3
|
2, 1
|
2s1
|
Na
|
11
|
2, 8, 1
|
3s1
|
K
|
19
|
2, 8, 8, 1
|
4s1
|
Rb
|
37
|
2, 8, 18, 8 1
|
5s1
|
(3) It explains the similarities and variations in the properties of the elements in terms of their electronic configurations and
brings out clearly the trends in chemical properties across the long periods.
brings out clearly the trends in chemical properties across the long periods.
(4) The inert gases having completely filled electron shells have been placed at the end of each period. Such a location of the
inert gases represents a logical completion of each period.
inert gases represents a logical completion of each period.
(5) In this form of the periodic table, the elements of the two sub-groups have been placed separately and thus dissimilar
elements do not fall together.
elements do not fall together.
(6) It provides a clear demarcation of different types of the elements like active metals, transition metals, non-metals,
metalloids, inert gases, lanthanides and actinides. The elements of group IA and IIA are active metals and are located at the
extreme left of the table.
metalloids, inert gases, lanthanides and actinides. The elements of group IA and IIA are active metals and are located at the
extreme left of the table.
The transition metals are found in the middle of the table. The elements lying to the right of a dark line shown in the long form
of periodic table in the form of ladder (i.e. steps) are noble gases, metalloids and non-metals while those lying to the left of the
line are metals (active metals and transition metals).
of periodic table in the form of ladder (i.e. steps) are noble gases, metalloids and non-metals while those lying to the left of the
line are metals (active metals and transition metals).
(7) It is easier to remember, understand and reproduce.
Defects of the modern periodic table
Some defects of the periodic table are as follows
⇨The problem of placing hydrogen remains unsolved.
⇨It fails to accommodate the lanthanides and actinides in the main body of the table.
⇨The arrangement is unable to reflect the electronic configuration of many elements.
Position of hydrogen in the periodic table:
Placing hydrogen in group IA:
Electronic configuration: Hydrogen has the same electronic configuration as the alkali metals (group IA).
Element
|
Atomic No.
|
Complete Electronic configuration
|
Valence shell electronic configuration
|
Li
|
3
|
2, 1
|
2s1
|
Na
|
11
|
2, 8, 1
|
3s1
|
K
|
19
|
2, 8, 8, 1
|
4s1
|
Rb
|
37
|
2, 8, 18, 8 1
|
5s1
|
2. Valency: Hydrogen has the valence of “1” like the alkali metals.
3. Electro positivity: Hydrogen is an electropositive element like the alkali metals.
4. Reducing property: Like the alkali metals it also acts as a reducing agent.
5. Formation of stable compound: Hydrogen forms stable compounds with oxygen and halogen like H2O, HX similar to the
compounds Na2O and NaX.
compounds Na2O and NaX.
6. Formation of M+ ions: Like alkali metals Hydrogen is a strong electropositive element and has the tendency to lose its
electron to form a unipositive cation.
electron to form a unipositive cation.
H–e–→H+
L–e–→Li+
7. Affinity for non-metals: Both hydrogen and alkali metals have a strong affinity for non-metals and little affinity for mortals.
Placing hydrogen in group VIIA:
1. Hydrogen should be placed just before Helium. Therefore it should be placed in group VIIA.
2. Valency: Hydrogen has the valence of similar to the halogens.
3. Non Metal: Hydrogen is a non-metal like halogens.
4. Atomic state: Like the halogens hydrogen is a diatomic gas.
5. Combination with non-metals: Hydrogen forms compounds like CH4 and SiH4 with non-metals similar to Chlorine forming
CCl4 and SiCl4.
CCl4 and SiCl4.
6. Formation of Negative ions: Like halogens, Hydrogen also gains one electron to form negative ion.
H+e–→H–
F+e–→F–
7. Formation of hydride: Hydrogen forms hydride with some metals NaH, CaH2, like halogens form the halides e.g. NaCl, KBr.
Thus, there is a debate whether hydrogen should place group IA or VIIA. But since it is an s-block element it is placed with the
alkali metals.
alkali metals.
Usefulness of periodic table
The followings are the important aspects of periodic table.
- Classification of elements
The classification of elements of similar properties into groups, simplified their study.
For example Na a member of alkali metals , reacts with water vigorously giving hydrogen gas and forming NaOH, which is a
strong base. The other alkali metals also react with water in a similar manner.
strong base. The other alkali metals also react with water in a similar manner.
- Prediction of undiscovered elements
At present all the elements from atomic number 1 to 109 have been discovered and their properties are more or less known.
But a very remarkable use of the periodic table was made by the Mendeleev in predicting a number of undiscovered elements,
which were shown by a number of gaps in the periodic table.
But a very remarkable use of the periodic table was made by the Mendeleev in predicting a number of undiscovered elements,
which were shown by a number of gaps in the periodic table.
Mendeleev’s table contained only 65 elements with a large number of vacant places. Mendeleev predicted the existence and
properties of 6 elements corresponding to the gaps. These elements have since been discovered and are Sc, Gallium,
Germanium, Technium, Rhenium, and Polonium.
- Correction of atomic weight
Atomic weight of some of the elements at the time of Mendeleev gave a wrong position in the periodic table. The properties of
these elements required their placement somewhere else.
these elements required their placement somewhere else.
For instance the element indium was placed in a vacant place in the periodic table between Cd (112.4) and Sn (118.7) and
indium with weight of about 114 fitted very well in between Cd and Sn.
indium with weight of about 114 fitted very well in between Cd and Sn.
- Periodic table in industrial research
The periodic table has been found to be quite useful in industrial researches. Several of the light metals and their alloys used
in modern mechanical equipment’s, jet engines and air crafts were first studied in detail because of their position in the
periodic table.
in modern mechanical equipment’s, jet engines and air crafts were first studied in detail because of their position in the
periodic table.
✯1: Definition of Valence
The term valence (or valency) is often used to state the potential or capacity of an element to combine with other elements.
At one time, it was useful to define valence of an element as: the number of hydrogen atoms or twice the number of
oxygen atoms with which that element could combine in a binary compound (containing two different elements only).
At one time, it was useful to define valence of an element as: the number of hydrogen atoms or twice the number of
oxygen atoms with which that element could combine in a binary compound (containing two different elements only).
In hydrogen chloride (HCl), one atom of chlorine is combined with one atom of hydrogen and the valence of chlorine is 1.
In magnesium oxide (MgO), since one atom of magnesium holds one atom of oxygen, the valence of magnesium is 2.
In magnesium oxide (MgO), since one atom of magnesium holds one atom of oxygen, the valence of magnesium is 2.
As already stated, there are three different types of bonds that are known to join atoms in molecules.
Although no precise definition of valence is possible, we can say that: Valence is the number of bonds formed by an atom in
a molecule.
a molecule.
Periodicity of Properties and Magic Number
The repetition of the elements with similar properties at certain regular intervals of atomic number in the periodic table is
termed periodicity of properties. In order to understand the concept of periodicity of properties we may consider the properties
of the elements of groups I A (Alkali metals), zero (Inert gases) and VII A (Halogens) given below
termed periodicity of properties. In order to understand the concept of periodicity of properties we may consider the properties
of the elements of groups I A (Alkali metals), zero (Inert gases) and VII A (Halogens) given below
The examination of the properties of these elements will show that the elements belonging to the same group have similar
properties, In other words we can say that the atomic number intervals at which the elements with similar properties
reappear are 2, 8, 8, 18, 18, 32 and 32, i.e. we have to pass 2, 8, 8, 18, 18, 32, and 32 elements before we come across an element
with similar properties. The numbers 2, 8, 8 and 32 are called magic numbers.
properties, In other words we can say that the atomic number intervals at which the elements with similar properties
reappear are 2, 8, 8, 18, 18, 32 and 32, i.e. we have to pass 2, 8, 8, 18, 18, 32, and 32 elements before we come across an element
with similar properties. The numbers 2, 8, 8 and 32 are called magic numbers.
Shielding or Screening Effect of inner-shell Electrons on the Valence shell Electron.
The decrease in the attractive force exerted by the nucleus on the valence shell electron, which is obviously due to the
presence of the electrons lying between the nucleus and valence-shell electrons, (called intervening electrons) is called
shielding effect or screening effect. In other words, the intervening electrons screen or shield the
valence-shell electrons from the nucleus.
presence of the electrons lying between the nucleus and valence-shell electrons, (called intervening electrons) is called
shielding effect or screening effect. In other words, the intervening electrons screen or shield the
valence-shell electrons from the nucleus.
Factors Affecting the Magnitude of Shielding Effect
Following are the important factors on which the magnitude of shielding effect caused by the inner-shell electrons on the
valence-shell electron depends.
valence-shell electron depends.
(i) No. of inner-shell electrons or inner shells. Greater is the number of inner-shell electrons or inner shells, greater is the
magnitude of shielding effect caused by the inner electrons on the valence- shell electron. Thus as we move down a group,
the number of inner-shells or inner shell electrons increases and hence the shielding effect also increases. For example in the
elements of group IA.
magnitude of shielding effect caused by the inner electrons on the valence- shell electron. Thus as we move down a group,
the number of inner-shells or inner shell electrons increases and hence the shielding effect also increases. For example in the
elements of group IA.
Effective nuclear charge
Effective nuclear charge, Z is defined as the actual nuclear charge, Z minus the screening effect caused by the electrons
intervening between the nucleus and the outer electrons. It is due to the shielding effect of the inner electrons on the
outer-electrons that the valence electron experiences less attractive pull from the nucleus. The decrease in the attractive
force reduces the nuclear charge, Z represented by the atomic number of the element.
intervening between the nucleus and the outer electrons. It is due to the shielding effect of the inner electrons on the
outer-electrons that the valence electron experiences less attractive pull from the nucleus. The decrease in the attractive
force reduces the nuclear charge, Z represented by the atomic number of the element.
This decreased nuclear charge is called effective nuclear charge and is represented by Zeff. It is given by the relation:
••Why the size of atom aka atomic radius increases from top to bottom while decreases from left to
right?
right?
The atomic size decreases across the period because effective nuclear charge on the valance shall electrons increases.
In a period the valence shell remains same for a elements of a period. Let us consider example of second period elements
the valance shell for these elements is 2s and 2p orbital as electrons are being filled in these orbitals.
In a period the valence shell remains same for a elements of a period. Let us consider example of second period elements
the valance shell for these elements is 2s and 2p orbital as electrons are being filled in these orbitals.
The electrons are filled in firstly 2s orbital and then a2p orbital is being filled. As distance from nucleus and valance shell
remains constant as shown in figure.
remains constant as shown in figure.
The concentration of positive charge in the nucleus increases with increase in atomic number. The nucleus of the atom gains
protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the
nucleus upon the electrons. Although the electrons are also added as the elements move from left to right across a period,
but these electrons reside in the same energy shell and do not offer increased shielding. Thus, moving from left to right across
a period, the atomic radius decreases.
protons moving from left to right, increasing the positive charge of the nucleus and increasing the attractive force of the
nucleus upon the electrons. Although the electrons are also added as the elements move from left to right across a period,
but these electrons reside in the same energy shell and do not offer increased shielding. Thus, moving from left to right across
a period, the atomic radius decreases.
So in fluorine there are nine protons that are attracting the electrons but in Lithium there are only 3 protons that are attracting
valance shell. So size of fluorine would be less as valance shell would contract due to strong nuclear attraction. Thus the new
size of the fluorine would be less than expected.
valance shell. So size of fluorine would be less as valance shell would contract due to strong nuclear attraction. Thus the new
size of the fluorine would be less than expected.
Atomic size increases down a group
The atomic radius increases moving down a group.
Each level (shell) only certain number of electron can occupy. If the number of electron further increased then, levels are
increased. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount
of the protons' positive charge.
Each level (shell) only certain number of electron can occupy. If the number of electron further increased then, levels are
increased. The new energy shells provide shielding, allowing the valence electrons to experience only a minimal amount
of the protons' positive charge.
It happens because each succeeding element has an additional level or shell of electrons. Each new level is shielded from
the pull of the nucleus by the layers below it, so it is further out from the nucleus, making the atom bigger.
the pull of the nucleus by the layers below it, so it is further out from the nucleus, making the atom bigger.
In other words, the number of shells increases as we go down the group. The outermost electrons are repelled by the inner
shell electrons (Screening effect) and hence, the atomic size increases.
shell electrons (Screening effect) and hence, the atomic size increases.
Atomic volume
Atomic volume is defined as the volume in c.c, occupied by one gram atom of the element in the solid state and hence is
commonly called gram atomic volume. It is obtained by dividing the atomic weight of the element by its density; i.e.
commonly called gram atomic volume. It is obtained by dividing the atomic weight of the element by its density; i.e.
•Why the atomic volume increases from top to bottom while first decreases then increases from
left to right?
left to right?
Variation of Atomic Volume in a Period and a Group.
(a) In a group. Atomic volume increases more or less regularly in going down a group (See). The increase in atomic
volume in going down a group is due to the increase in the number of shells. The larger the number of shells, the bigger
is the volume of the atom.
volume in going down a group is due to the increase in the number of shells. The larger the number of shells, the bigger
is the volume of the atom.
(b) In a period.
In going from left to right in a period, it varies cyclically, i.e., it decreases at first for some elements, becomes minimum
in the middle and then increases (See).The variation of atomic volume in going from left to right in a period is influenced by
the following two factors
in the middle and then increases (See).The variation of atomic volume in going from left to right in a period is influenced by
the following two factors
(i) Nuclear charge. We know that the nuclear charge (i.e. atomic number) increases by one, as we move from left to right in
a period. The increased nuclear charge attracts each electron more strongly towards the nucleus, resulting in a decrease in
the volume of the atom.
a period. The increased nuclear charge attracts each electron more strongly towards the nucleus, resulting in a decrease in
the volume of the atom.
(ii) No. of valence electrons.
Towards the close of a period, due to an increase in the number of valence-electrons (i.e. electrons in the valence shell) the
volume of the atom increases, so that it may accommodate all the electrons. These two factors, one causing an increase and
the other causing a decrease, combine to result that in a period atomic volume decreases at first for some elements, becomes
minimum in the middle and then increase.
Electronegativity
In a molecule A – B the electrons forming the covalent bond are attracted by atom A as well as by B. This attraction is
measured in terms of what we call electronegativity, EN.
measured in terms of what we call electronegativity, EN.
It may be defined as: The attraction exerted by an atom on the electron pair bonding it to another atom by a covalent
bond.
bond.
Trend in electronegativity
The variations in electronegativities of elements in the Periodic table are similar to those of ionisation energies and electron
affinities.
affinities.
(1) Increase across a Period
The values of electronegativities increase as we pass from left to right in a Period.
Thus for Period 2 we have
This is so because the attraction of bonding electrons by an atom increases with increase of nuclear charge (At. No.) and
decrease of atomic radius. Both these factors operate as we move to the right in a Period.
decrease of atomic radius. Both these factors operate as we move to the right in a Period.
(2) Decrease down a Group
The electronegativities of elements decrease from top to bottom in a Group.
Thus for Group VII we have
The decrease trend is explained by more shielding electrons and larger atomic radius as we travel down a Group.
Importance of Electronegativity
The electronegativities of elements are widely used throughout the study of Chemistry. Their usefulness will be discussed at
appropriate places. The important applications of electronegativities are listed below.
appropriate places. The important applications of electronegativities are listed below.
(1) In predicting the polarity of a particular bond. The polarity of a bond, in turn, shows the way how the bond would break
when attacked by an organic reagent.
when attacked by an organic reagent.
(2) In predicting the degree of ionic character of a covalent bond.
(3) In predicting of inductive effects in organic chemistry.
(4) In understanding the shapes of molecules.
Electron Affinity
A neutral atom can accept an electron to form negative ion. In this process, in general, energy is released.
Electron affinity (EA) of an element is the amount of energy released when an electron is added to a gaseous atom to
form an anion.
form an anion.
Ionization Energy
Amount of energy required to remove an electron from the ground state of a gaseous atom or ion.
- First ionization energy is that energy required to remove first electron.
- Second ionization energy is that energy required to remove second electron, etc.
Group Trend
- As you go down a column, ionization energy decreases.
- As you go down, atomic size is increasing (less attraction), so easier to remove an e-.
Periodic Trend
- As you go across a period (L to R), ionization energy increases.
- As you go L to R, atomic size is decreasing (more attraction), so more difficult to remove an e- (also, metals want to lose
e-, but nonmetals do not).
- It requires more energy to remove each successive electron.
- When all valence electrons have been removed, the ionization energy takes a quantum leap.
Trends in First Ionization Energies
As one goes down a column, less energy is required to remove the first electron.
For atoms in the same group, Zeff is essentially the same, but the valence electrons are farther from the nucleus.
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