# 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 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 .

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 .

Nonmetals : The nonmetals atoms tend to gain electrons during the chemical reaction to form negative ions .

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 chlorine nonmetal tends to gain an electron during chemical reactions to form chloride ion of radius larger than the radius of its atom .

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 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 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 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 .

Application : The first ionization potential of noble gases and alkali metals

The first ionization potential of noble gases is very high , due to the stability of their electronic configuration and it is difficult to remove an electron from a completely filled shell .

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 .

##### Electron affinity

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) + eX(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 .

##### Electronegativity

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 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 increase of the relative values of the electronegativity means the increase in the ability of the element atom to attract the electrons of the chemical bond .

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 .

So , The atoms of nonmetals group 7A ( halogens ) are the greatest in the electronegativity  , while the atoms of the alkali metals group 1A are lowest in the electronegativity .

Fluorine ( F ) is considered to be the most electronegative element , while cesium ( Cs ) is considered to be the lowest electronegative element .

Modern periodic table and classification of Elements

Metallic & nonmetallic property , Acidic & basic property in the periodic table

Graduation of the properties of the elements in the modern periodic table