Download The Modern Periodic Table

The Modern Periodic Table
In the modern Periodic table the elements are arranged according to electron configuration of the atoms of
the elements. The elements are placed in the increasing order of their atomic numbers.
The first periodic table was arranged by the Russian scientist Dimitri Mendeleeve based on his periodic
law: "The physical and chemical properties of elements are the periodic functions of their atomic masses".
After "Mosley" discovered the atomic # he arranged the elements in the increasing order of their atomic #s
based on his periodic law: "The physical and chemical properties of elements are the periodic functions of
their atomic numbers".
The modern Periodic table is the modification of Mosley’s periodic table by ‘Glenn Seaborg’ an American
Scientist.
Modern Periodic law: “The physical & chemical properties of elements are the periodic functions of the
electron configuration of their atoms.”
The modern Periodic table consists of 7 Horizontal rows called periods and 18 vertical columns called groups
or families.
• The first period is the smallest it has only two elements H and He.
• The second and third periods have 8 elements each.
• The fourth and fifth periods have 18 elements each.
• The sixth period has 32 elements. This is the largest period. 14 elements (Atomic # 58-71) of this period
following the element Lanthanum (La) are placed separately from the body of the periodic table. These
elements are called Lanthanides. Lanthanides are a part of sixth period.
• The seventh period is incomplete. Some of the elements are not discovered yet. 14 elements (Atomic # 90103) of this period following the element Actinium (Ac) are placed separately from the body of the
periodic table. These elements are called Actinides. Actinides are a part of seventh period.
o The 18 vertical columns are further grouped as A and B.
o Elements occurring in the same column or group or family have similar chemical properties. They
have same number of valence electrons.
CLASSIFICATION OF THE ELEMENTS IN TO BLOCKS:
(Relationship between electron configuration, atomic structure and the periodic table.)
The periodic table is also classified in to 4 blocks. s,p,d and f.
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The s block consists of first two columns (Group 1A and 2A) of the periodic table.
These elements are included in the s block their last electrons enter in the ns orbital.
(Where n= principle quantum number or valence energy level which is equal to the # of
period) General Electron Configuration (ends on): Group1A = ns1; Group 2A= ns2
All the s block elements are active metals.
The p block consists of last 6 columns at the left of the periodic table.
These elements are included in the p block because their last electrons enter in the np
orbital. (Where n= principle quantum number or valence energy level which is equal to the
# of period) General Electron Configuration (ends on): ns2 np1-6
P block consists of metals non-metals and metalloids.
The s and p block elements are called representative elements.
d block: The elements from 3rd to 12th column (groups IIIB to IIB) are called d block or
transition metals.
These elements are included in the d block because their last electrons enter in the (n-1)d
orbital. (Where n= principle quantum number or valence energy level which is equal to the
# of period)
General Electron Configuration (ends on): ns1-2 (n-1)d1-10
f blocks The Lanthanides and actinides are called f blocks or Inner transition elements.
These elements are included in the f block because their last electrons enter in the (n-2) f
orbital.
(Where n= principle quantum number or valence energy level which is equal to the # of
period)
General Electron Configuration (ends on): ns1-2 (n-1)d1 (n-1)f1-14
Locating the element in the periodic table with the help of electron
configuration
Example:1 name the element that has electron configuration ending on 3p 5
The element is in 3rd period and in the 5th
column of p block that is group 7A.
Therefore it is Chlorine [Cl]
Example: 2: name the element that has electron configuration ending on 4s 2
The element is in 4th period and in the 2nd
column of s block that is group 2A.
Therefore the element is Calcium [Ca]
Example: 3: name the element that has electron configuration ending on 4d 2
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The element is in 5th period and in the 2nd
column of d block that is group 4B.
Therefore the element is Zirconium [Zr]
Example 4: name the element that has electron configuration ending on 4f 2
The element is in 6th period and in the 2nd
column of ‘f’ block. Therefore the element is
Praseodymium [Pr]
Valence Electrons
Electrons present in the outermost energy level are called Valence Electrons.
All the elements occurring in a group or family will have same # of valence electrons. This
is equal to the group # for representative elements (s and p block elements).
Valency: # of hydrogen atoms (or double the #of oxygen atoms) with which the atom of an
element can make bonds is called valency.
Valency = group # (for groups 1-4) OR Valency = 8-group # (for groups 5-8)
Electron Dot Structure: represents the valence electrons (as dots) around the symbol of
an element. (See page 183)
Ion formation and Ionic charges
The noble gases (group 8A) have 8 electrons in their valence shell. (They have s and p
orbitals completely filled.) That’s why they are stable and extremely un-reactive. This is
called octet.
Atoms of all other elements are reactive because they don’t have 8 electrons in their
valence shell. Therefore, they try to get octet by:
• Losing one or more electrons
• Gaining one or more electrons
• Sharing one or more electrons
# of electrons lost = group #
Metals lose electrons and become cations, positively charged ions.
# of electrons gained = 8- group #
Non metals gain electrons and become anions, negatively charged ions
Periodic Properties of the elements
Properties of the elements that have a clear trend across a period or in a group are called
periodic properties. Some important periodic properties of the elements are:
Atomic radius
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Ionic radius
Ionization energy
Electro negativity
Atomic radius:
Average distance between the centers of the nuclei of two adjacent atoms bonded together.
Trend down the group: Atomic radius increases. Its because at every step a new energy
level is added to the atom.
Trend across a period: Atomic radius decreases. Its because the elements in a period have
the same outer most shell. When we move from left to right across a period, # of Electrons
increases in the same shell. At the same time # of protons increases in the nucleus. This
causes greater attraction between the positively charged nucleus and the negatively
charged valence electrons. As a result the valence energy level shrinks, therefore the
atomic radius decreases.
Element
#Protons
# Valence e
Na
+11
-1
Mg
+12
-2
Al
Si
P
+13
+14
+15
-3
-4
-5
Atomic Radius decreases
S
+16
-6
Cl
+17
-7
Ar
+18
-8
The above elements belong to 3rd period; it means they all have three energy levels.
Sodium has 11 positive and negative charges to attract each other. Where as Argon has 18
positive and negative charges, therefore, the attraction between the nucleus (holding 18
positively charged protons) and the valence electrons will be much higher than that of
sodium. That’s why the Ar atom is much smaller than the sodium (Na) atom.
Ionic
Radius:
Ionic
Radius:
When
electron
an
is
removed from the outermost energy level, there are more positively charged
protons in the nucleus to attract fewer valence electrons. Therefore the size
drastically decreases.
Cations have smaller ionic radius than their respective atomic radius.
Similarly higher the positive charge on the cation, smaller will be the ionic radius.
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When an electron is added to the outermost energy level of an atom, there are
more electrons (negative charges) to be attracted by protons (positive charges) in
the nucleus. Therefore, the attraction between the nucleus and electrons decreases
by increase in the # of electrons or negative charge on the anion.
In an iso-electronic series (series of ions having same # of electrons) greater the
negative charge, greater will be the ionic radius and greater the positive charge
smaller will be ionic radius.
Element
N3O2F1Na1+
Mg2+
Al3+
# Electrons -10
-10
-10
-10
-10
-10
#Protons
+7
+8
+9
+11
+12
+13
Ionic Radius decreases
Ionization Energy:
Energy required to remove the most loosely held electron from the valence shell of
a free neutral gaseous atom is called Ionization Energy. Its measured in
Joules/mole.
Ionization energy depends on the size of the atom. If the atom is smaller, the
valence electrons are closer to the nucleus. As the nuclear attraction is more,
removal of the electron will be difficult. Therefore more ionization energy will be
required.
If the atom is bigger, valence electrons are far from the nucleus, therefore less
nuclear attraction, hence low ionization energy.
The ionization energy decreases down the group as the atomic size increases.
The Ionization energy increases across a period as the atomic size decreases.
Electron configuration also affects the ionization energy.
Orbitals that are completely filled are more stable than half filled. Half filled
orbitals are more stable than the orbitals that are neither half filled nor completely
filled.
Ex: Li (3):
1s2 2s1 Half filled 2s orbital. Less stable than Be.
Be (4):
1s2 2s2 Completely filled 2s orbital. More stable than Li
Ex: C (6):
1s2 2s2 2p2 Neither completely filled nor half filled.
N (7):
1s2 2s2 2p3 Half filled 2p orbital. More stable than C and O.
O (8):
1s2 2s2 2p4 Neither completely filled nor half filled.
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Elements with stable electron configuration will have higher 1 st ionization energy
than their neighboring elements therefore, will have higher Ionization energy. Look
at the graph.
Ionization energy required to remove first electron is called the 1st IE. Energy required to
remove second electron is called 2nd IE. 2nd IE is always greater than 1st IE.
Electronegativity:
Relative ability of the atoms of an element to attract the shared pair of electrons in a
covalent bond is called electro negativity.
Its measured in “Pauling”.
As the noble gases don’t generally involve in chemical bonding, there is no
Electronegativity value for them.
Trends:
Increases across a period from left to right as the atomic size decreases.
Decreases down the group as the atomic size increases.
Fluorine has highest Electronegativity (3.98) and Francium has least.
Application:
1. Electronegativity is helpful in writing the molecular formulas of the elements. In a
molecular formula, element with least Electronegativity is written first.
Example: Water is H2O not OH2. Because the Electronegativity of H is 2.2 and that
of O is 3.44.Therfroe, the least electronegative element H is written first.
2. It’s also used in determining % ionic character of a polar covalent bond. (Learn
Chemical bonding)
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