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كيمياء لغات - الباب الثانى | C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
01 Trends And Periodicity Of Properties in the periodic table
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كيمياء لغات - الباب الثانى | C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
C- TYPES OF ELEMENTS IN THE PERIODIC TABLE
1- Noble gases (inert gases)
* Elements in which all their levels are completely filled.
so they are more stable. and chemically inactive. G.R.F.
* They are called zeroth group
* They are elements which have the electronic configuration np6 except Helium which has electronic configuration 1S2 .
2- Representative elements [ main groups elements]
*Elements which all their levels are completely filled except
the highest level , their electronic configuration ends by nS1-2 , nP1-5
* They are elements of S-block and P-block except zeroth group.
*Their highest levels tend to reach the completely electronic configuration
[nS2, nP6 ] by losing , gaining or sharing electrons.
3- Main transition elements
* They are elements of the d-block. (B Groups)
* d-sublevel of these elements are successively filled.
Classification of main transition elements
1st series 2nd series 3rd series
Transition series (3d) Transition series (4d) Transition series (5d)
*It includes the elements in which the sublevel 3d is filled successively.
* Placed in the fourth period.
* Consists of 10 elements. * It includes the elements in which the sublevel 4d is filled successively.
* Placed in the fifth period.
Consists of 10 elements. *It includes the elements in which the sublevel 5d is filled successively.
* Placed in the sixth period.
* Consists of 10 elements.
4- Inner-transition elements nf, (n+1)d and ( n+2 ) S
*They are elements of f-block , f-sublevel filled successively.
* These elements are divided into lanthanides and actinides
have incomplete orbitals 4f,5d and 6S Actinides
have incomplete orbitals 5f,6dand 7S
* Placed in the sixth period.(after 57La)
* Consists of 14 (rare earth’s elements)
* Includes elements in which sublevel 4f is filled successively. * Placed in the seventh period.(after 89Ac)
* Consists of 14( radioactive elements)
* Includes elements in which sublevel 5f is filled successively.
1) Lanthanides are called rare earth’s elements
Because: They are very difficult to be separated
2) lanthanides are similar in behavior [ chemical properties ].
Because: They have the same valence electrons (6S2 )
3) Actinides are called radioactive elements
Because: Their nuclei are unstable
* Write electronic configuration then determine :-
[ block – period – type – sublevels - orbitals ]
or determine [the location of elements in periodic table]
26Fe , 36Kr , 13Al , 20Ca , 2He , 19K , 18Ar
Write the electronic configuration according to nearest inert gas:-
Types and numbers of elements in the periodic table
NUMBER OF PERIOD SUBLEVELS WHICH FILLED SUCC. NUMBER OF ELEMENTS TYPES OF ELEMENTS IN
First 1S2 2 Representative and noble 
Second 2S2 , 2P6 8 Representative and noble 
Third 3S2 and 3P6 filled succ. 8 Representative and noble 
Fourth 4S2 , 3d10 and 4P6 filled succ. 18 Representative , and main transition
[first series] and noble 
Fifth 5S2 , 4d10 and 5P6 filled succ 18 Representative , and main transition [second series]
and noble 
Sixth 6S2, 4f14 , 5d10 and 6P6 filled succ 32 Representative , and main transition [third series ]
and inner transition [lanthanides ] and noble 
D- TRENDS AND PERIODICITY OF PROPERTIES IN THE PERIODIC TABLE
1) The atomic radius
* It’s incorrect to define the atomic radius as the distance from the nucleus to the farthest electron. G.R.F
Because: wave mechanics proved that :-
* It is impossible to determine exactly the location of electrons around
the nucleus. So the atomic radius cannot be defined and cannot be
Atomic radius: * Half the distance between the centers of two similar atoms in a diatomic molecule.
Bond length: *The distance between the nuclei of two bonded atoms.
• Covalent atomic radius = bond length / 2.
Molecule Bond length in Angstrom Covalent atomic radius
H - H 0.60 Ao 0.30A0
F - F 1.28A0 0.64A0
Cl - Cl 1.98A0 0,99A0
Br - Br 2.28A0 1.14A0
I - I 2.66A0 1.33A0
Ionic crystals* Ionic crystals consists of positive ions and negative ions such as sodium chloride crystals and the radius between them called ionic radius.
Ionic radius :- Sum of the two radius of cation and anion
and it depends on the number of electrons losed or gained.
Bond length of = covalent atomic radius of A + covalent atomic radius of B
Ex.1: The bond length in the chlorine molecule (Cl – Cl) equals
1.98 Ao and the bond length between carbon and chlorine atom (C – Cl) = 1.76 Ao calculate the atomic radius of carbon.
Bond length in C - Cl = covalent radius of C + covalent radius of Cl
1.76 = ?? + 1.98/2
1.76 = ?? + 0.99
1.76 - 0.99 = ??
0.77Ao = covalent atomic radius (of C )
Ex.2: The bond length in chlorine ( Cl – Cl ) = 1.98 Ao and the bond length in hydrogen chloride molecule ( H – Cl ) = 1.29 Ao
Calculate the covalent atomic radius of hydrogen
Ex.3: The bond length in nitrogen molecule ( N – N ) = 1.96 Ao and the bond length in hydrogen molecule = 0.62 Ao calculate:
a) Atomic radius of nitrogen. b)Atomic radius of hydrogen
c) Bond length between ( N ) and ( H ) in ammonia molecule NH3
Graduation of atomic radius in the periodic table
* We can observe the following from the periodic table.
In the horizontal periods :-
Atomic radius decreases from left to right across a period. G.R.F
Because :- * number of positive charges inside the nucleus increase and
attraction force of nucleus increases.
So atoms of group 1A have the largest size and group 7A have
the smallest size which called halogens.
In the vertical groups :-
Atomic radius increases as the atomic number increases
from up to down G.R.F
* Number of energy levels increases .
* Screening effect of an extra shell takes place.
* Repulsion force between electrons in the levels increases .
* Attraction force of nucleus decreases.
* Atomic radius increases from up to down across a group.
Obs. Increasing atomic radius when transfer from period to another more
than decreasing atomic radius when transfer from group to another.
Because the effect of increasing energy level more than the effect of
increasing one positive charge across a period.
Changes of cations and anions from their atoms
1.75 Ao 0.95 Ao
A) Cations are smaller than their atoms
Or cationic radius smaller than atoms
Answer: due to increasing the pull of the nuclear charge on the remaining electrons in the cation.
*When number of positive charges of an ion increases , attraction force of the nucleus increases and atomic radius decreases. [ as in Iron ]
1.17 Ao 0.75 Ao 0.60 Ao
Radius of Fe+3 smaller than that of Fe+2 smaller than that of Fe atom
*Because when number of positive charges increase, attraction force of nucleus increases so atomic radius decreases.
b)Radius of anions are larger than their atoms
The radius of chloride ion is larger than the radius of its atom. G.R.F
Answer: due to increasing the number of electrons without increasing
the nuclear charge. [ tug of war ]
*Metals :- Element in which its ionic radius smaller than its atom.
* Non-metal :- Element in which its ionic radius larger than its atom.
2) Ionization energy (Ionization pot.)
Def.: It’s the amount of energy required to remove the most loosely bound electron completely from an isolated gaseous atom.
*and it is determined by a spectral measurements so :-
*It’s possible to remove one, two, three or more electrons from atoms
*So, there are first, second, third ………. etc ionization energy.
* First ionization energy:
- It’s the energy required to convert an atom into ion with one positive charge. M M+1 + e * H = + ve KJ/ mol
* Second ionization energy: M+1 M+2 + 2 e-
- It’s the energy required to convert M+1 into M+2
G.R.F Second ionization energy is higher than first ionization energy
Variation of ionization energy in periods
Ionization energy increases from left to right across period G.R.F
. Because from left to right
*Attraction force of nucleus increases and atomic radius decreases.
Variation of ionization energy in groups
Ionization energy decreases from up to down. G.R.F
Because * No. of levels increases (repulsion force of electrons increase)
* Energy levels will spread out.
* Screening effect takes place
* Attraction force of the nucleus decreases
* Atomic radius increases..
1) Ionization energy of 4Be, 12Mg is high. G.R.F
or elements of group II A have high ionization energy )
Answer: 4Be 1S2, 2S2 12Mg 1S2, 2S2, 2P6, 3S2
* Because the outer most sublevel S is filled with 2 electrons so atom will be more stable so it must need large amount of energy to remove electron.
2) Ionization energy of Nitrogen (7N) , ( 15P ) is high
or elements of group 5A have high ionization energy.
Answer: 7N 1S2, 2S2, 2P3 15P 1S2, 2S2, 2P6, 3S3, 3P3
*Because the outer most sublevel P is half filled, [p3] so atom will be more stable so it must need large amount of energy to remove electron.
3) The first ionization energy of noble gases of zeroth group is very high.
4). The second ionization energy of elements of 1A group is high.
5)- The third ionization energy of elements of 2A group is very high.
*Because it must need large amount of energy to break down the completely filled level due to the stability of atom.
Ionization energies of Magnesium.
(2,8,2)Mg Mg+ + e H +737 k.j / mol
(2,8,1)Mg+ Mg+2 + e H = +1450
(2,8)Mg+2 Mg+3 + e H = +7730
* Ionization energy is inversely proportional to atomic radius.
*Compare between the first ionization energy and the second ionization energy
3) Electron affinity الميل الإلكترونى
Def: It is the amount of energy released (produced) when an extra
electron is added to a neutral gaseous atom .
X + e- X-1 + energy H= - ….. k.j mol-1
Variation in periods: * Electron affinity increases from left to right
*Due to decreasing atomic radius leading to easy to attract new electron
Variation in-groups: - Electron affinity decreases from up to dow
. *Due to increasing atomic radius leading to increasing the distance
between nucleus and valence level
* The magnitude of the electron affinity is high when the added electrons make the orbitals half filled or completely filled due to increasing the stability of the atom.
1) Beryllium has a relatively small electron affinity.
Answer:- Due to the stability of its atom., it has filled orbitals. 4Be 1S2,2S2
2) Nitrogen has a relatively small electron affinity
Answer:- Due to the extra stability of its atom. The outer sublevel is half filled. 7N 1S2 ,2S2 , 2P3
3) Neon has very low E.Aff
*Due to great stability of atoms because all levels are completely filled
4) Electron affinity of fluorine -339.9 kj/mol is less than that of chlorine
-361.7 although electron affinity decrease in vertical groups from up to down.
Because:- Fluorine has small size [ top right of periodic table ]so repulsion force takes place between the new electron and original nine electrons
Def: - It’s the tendency of an atom to attract the electrons of
chemical bond to itself.
*Fluorine is considered to be the most electronegative element
* Diff. in electronegativity between elements plays an important role in determining the nature of chemical bonds [ chapter 3]
Electronegativity increases from left to right Electronegativity decreases from up to down
* Increasing the atomic number.
* Increasing the +ve nuclear charge.
* Increasing the attraction force.
* Decreasing the atomic radius.
* Increasing electron affinity.
* Increasing electronegativity. * Increasing the electron levels.
* Increasing the repulsion among electrons.
* Decreasing attraction force.
* Increasing the atomic radius.
* Decreasing electron affinity.
* Decreasing electronegativity.
1- Electronegativity of Oxygen is higher than nitrogen ( 8O ., 7N )
2- Electronegativity of Fluorine is higher than Bromine ( 9F , 35Br )
• Compare between electronaffinity and electronegativity
* Refers to the atom in its single state *Refers to combined atoms in molecule
*Term of energy *Not term of energy
* It is the amount of energy released when ………………………………. * It is the tendency of an atom to attract …………………………..
* Its unit is Kj / mole * unitless
Compare between ionization energy and electron affinity