Radius property , Ionization potential , Electron affinity & Electronegativity
The chemical properties & some of the physical properties of the elements depend on their electronic structure and especially on the valence electrons ( the electrons of the outermost level ) , The concept of bond length differs in the covelent compounds from that in the ionic compounds , we can calculate atomic radius and ionic radius .
The radius property
The atomic radius can not be calculated from the distance between the nucleus and the farthest electron , Because it is impossible to determine the precise location of an electron around the nucleus ( as the wave mechanics revealed ) .
But the atomic radius can be calculated by knowing the covalent bond length which is measured by Angestrom unit ( A° ) , Covalent bond length ( 2r ) is the distance between the nuclei of two bonded atoms , Atomic radius ( r ) is half the distance between the centers of two similar atoms in a diatomic molecule .
Covalent bond length = sum of the two atomic radii of the molecule
The atomic radius ( r ) = ½ × bond length in a diatomic element molecule ( 2r )
Ionic radius : The ionic compounds as sodium chloride are found in a crystalline from and consist of positive ions ( cations ) and negative ions ( anions ) , Ionic bond length is the distance between the centers of the nuclei of two bonded ions , The ionic radius depends on the number of electrons lost or gained to form ions .
The ionic bond length = the sum of two ionic radii of the formula unit
The effective nuclear charge concept ( Zeff )
The valence electrons are not affected by the complete nuclear charge ( the charge of the nucleus protons ) , This is because the inner electrons ( of the inner energy levels ) screen a part of this charge from the valence electrons , so , the actual charge affecting an electron is called the effective nuclear charge Zeff .
Effective nuclear charge Zeff is the actual nuclear charge ( positive charge ) which effects on an electron in an atom .
The effective nuclear charge Zeff is always less than the nuclear charge ( the total number of protons present in a nucleus ) , due to the screening effect of the electrons of the inner energy levels on a part of the nuclear charge affecting the electrons under study .
Graduation of atomic radius property in the periodic table
In the horizontal period : The atomic radius decreases as we go from the left to the right across a period by increasing the atomic number from 1A to zero group .
This is due to the gradual increase in the effective nuclear charge ( Zeff ) which increases the nuclear attraction force on the valence electrons leading to the reduction of the atomic radius .
In the vertical group : The atomic radius increases as we go down the group by increasing the atomic number from the first period to seventh .
This is due to the increase in the number of the energy levels in each new period , The increase in the number of the filled energy levels that having a screening effect on the pull of the nuclear charge on the outer electrons , The increasing of the repulsive forces between electrons .
The atoms of the first group elements ( alkalis ) are the biggest atoms , while the atoms of the seventh group elements ( halogens ) are the smallest atoms , The biggest atom in size is cesium ( Cs ) .
Example : 7N , 9F → ( N ) atom is bigger than ( F ) atom , because the atomic size decreases in the same period as we go from left to right by increasing the atomic number .
56Ba , 4Be → ( Ba ) atom is bigger than ( Be ) atom , because the atomic size increases in the same group as we go down the group by increasing the atomic number .
The relation between the radii of atoms and their ions
The radii of atoms differ from the radii of their ions , The ionic radius decreases as the effective nuclear ( positive ) charge of the ion increases .
Metals : The metal atom tend to lose their valence electrons during the chemical reaction to form positive ions .
The positive ion radius is smaller than its atomic radius , because the number of positive protons in the cation ( positive ion ) is larger than the number of negative electrons , So , the pull of the effective nuclear charge on remaining electrons increases leading to decrease the size .
Application : The sodium metal tends to lose its valence electron during chemical reactions to form sodium ion of radius smaller than the radius of its atom .
The negative ion radius is larger than its atomic radius , because the number of negative electrons in the anion ( negative ion ) is larger than the number of positive protons , so , the repulsive forces between electrons increase due to increasing the number of electrons without any increasing in the nuclear charge leading to increase the size .
Application : The atomic radius of iron atom ( Fe ) > the ionic radius of iron ( II ) ion Fe2+ > the ionic radius of iron ( III ) ion Fe3+ , because the atomic radii of metals are bigger than the radii of their ions , As ionic radius of positive ion decreases , its charge increases .
Ionization potential ( Ionization energy )
If an energy is supplied to an atom , electrons may be excited and transferred to higher energy levels , but if a sufficient energy is supplied , the most loosely bound electron may be completely removed , giving a positive ion , The minimum amount of this energy is called ionization potential .
Ionization potential ( Ionization energy ) is the amount of energy required to remove the most loosely bound electron completely from an isolated gaseous atom .
ΔH of the ionization process has a positive sign , because the ionization energy is an absorbed energy .
Na(g) + Energy → Na+(g) + e– , ΔH = + 496 kJ/mol
The atom of the same element has more than ionization energy :
First ionization potential : It is amount of energy required to convert an isolated gaseous atom to an ion carries one positive charge .
M(g) + Energy → M+(g) + e– , ΔH = ( + )
Second ionization potential : It is the amount of energy required to remove an electron from a positive ion carries one positive charge .
M+(g) + Energy → M2+(g) + e– , ΔH = ( + )
Third ionization potential : It is the amount of energy required to remove an electron from a positive ion carries two positive charges .
M2+(g) + Energy → M3+(g) + e– , ΔH = ( + )
The first ionization potential < the second ionization potential < the third ionization potential .
Example : 19Ne : [ 2He ] , 2s² , 2p6 , 18Ar : [ 10Ne ] , 3s² , 3p6
The first ionization energy of energy alkali metals is lower than that of all elements , due to the easily loss of the valence electron .
Example : 11Na : [ 10Ne ] , 3s¹ , 19K : [ 18Ar ] , 4s¹
The ionization potentials of magnesium : The second ionization energy of magnesium is greater than the first one , due to the increasing of the effective nuclear charge ( Zeff ) , The third ionization potential of magnesium is much greater than that of its first and second ones , because it results in the breaking up of a completely filled energy level .
Mg(g) → Mg+(g) + e– , ΔH1 = ( + 737 kJ/mol )
Mg+(g) → Mg2+(g) + e– , ΔH2 = ( + 1450 kJ/mol )
Mg2+(g) → Mg3+(g) + e– , ΔH3 = ( + 7730 kJ/mol )
The first ionization potential of potassium 19K is lower than that of calcium 20Ca , while the second ionization potential of potassium is much greater than that of calcium .
19K : [ 18Ar ] , 4s¹ , 20Ca : [ 18Ar ] , 4s²
The first ionization potential of potassium is lower than that of calcium due to the easily loss of valence electron , while the second ionization potential of potassium is much greater than that of calcium because it results in the breaking up of a completely filled shell .
The graduation of ionization potential in the periodic table :
In the same period : The first ionization potential increases as we move from left to right , This due to the increase of the effective nuclear charge and the decrease of the atomic radius , which would lead to increase the attraction force of the nucleus on the valence electrons , which need higher energy to separate them from the atom .
In the same group : The first ionization energy decreases as we go down the group , This due to the extra shells of electrons are added which increase the atomic radius , The decrease of attraction force of the nucleus on the valence electrons , so , the energy required to remove the valence electrons decreases .
So , the ionization potential is inversely proportional to atomic radius .
The ionization potential of oxygen 8O is lower than that of nitrogen 7N , although oxygen comes next nitrogen through the same period , because the atom becomes more stable when the 2p sub-level is half-filled electrons as in nitrogen atom and removing an electron from it will decrease its stability .
Example : Mention which atom is higher in ionization potential in each of the following pairs of atoms :
13Al , 16S : The ionization potential of 16S is greater than that of 13Al , because the ionization potential increases in the same period as we move left to right by increasing the atomic number .
3Li , 55Cs : The ionization potential of 3Li is greater than that of 55Cs , because the ionization potential decreases in the same group as we move down the group by increasing the atomic number .
The removal of an electron from the atom will convert it into a cation , which requires an amount of energy named by the first ionization energy , on the other hand , if the atom gained an extra electron , it will be converted into a negative ion , this is associated with releasing an amount of energy named by electron affinity .
Electron affinity is the amount of energy released when an extra electron is added to a neutral gaseous atom , The magnitude of the electron affinity is high when the added electron makes the sub-level , half-filled or completely filled , as in both cases it helps in the stability of the atom .
X(g) + e– → X−(g) + Energy , ΔH = ( − )
The graduation of electron affinity in the periodic table
In the same period : The electron affinity increases as we move from the left to right , this due to the increase of the atomic number leading to decrease the atomic radius ( atomic size ) , which makes it easier for the nucleus to attract a new electron .
In the same group : The electron affinity decreases as we go down the group , This is due to the increase of the atomic number leading to increase the atomic radius ( atomic size ) , so , the ability of the nucleus to attract the new electron decreases .
The electron affinity values for beryllium , nitrogen and neon are close to zero , 4Be : 1s² , 2s2 , 7N : 1s² , 2s2 , 2p3 , 10Ne : 1s² , 2s2 , 2p6
Because the atom will be more stable when the sublevel : 2s is full-filled as in case of beryllium atom 4Be , 2p is half-filled as in case of nitrogen atom 7N , 2p is full-filled as in case of neon atom 10Ne and the addition of an electron to any atom of them will decrease its stability .
The electron affinity of chlorine ( − 348.6 kJ/mol ) is greater than the electron affinity of flourine ( − 328 kJ/mol ) , although chlorine comes next fluorine through the same group .
Because fluorine atom is smaller in size as it has smaller radius than chlorine atom , so , the entering of an electron will suffer a strong repulsive force with the nine electrons already existing around the fluorine nucleus which decreases the released energy due to consuming a part of it to overcome this repulsive force .
When two atoms of two different elements combine together , the ability of one atom of them to attract the electrons of the chemical bond towards itself differs from that of the other atom , this attraction force is named by electronegativity , Electronegativity is the tendency of an atom to attract the electrons of the chemical bond to itself .
The electron affinity differs from the electronegativity , where the electron affinity is an energy term which refers to an atom in its single state , while the electronegativity of the elements is represented by relative values and it refers to a combined atom .
The difference in electronegativity between elements plays a very important role in determining the nature of the bond formed between them .
The graduation of electronegativily in the periodic table :
In the same period : The electronegativity increases as we move from left to right , this due to the increase of the atomic number leading to decrease atomic radius , so , the ability of atom to attract the electrons of the bond towards itself increases .
In the same group , The electronegativity decreases as we go down the group , this due to the increase of the atomic number leading to increase the atomic radius , so , the ability of atom to attract electrons of the bond towards itself decreases .
Fluorine ( F ) is considered to be the most electronegative element , while cesium ( Cs ) is considered to be the lowest electronegative element .