Third Row Elements

Reasons for The Study of The Third Row Elements:

  • The elements are arranged in the order of increasing atomic numbers, in such a way that elements with similar properties fall in the same vertical column of the periodic table. There are eighteen vertical columns known as groups and seven horizontal rows are known as periods.
  • All these elements have vacant 3d orbitals. These elements are capable of transferring their valence electrons to these vacant d orbitals. Thus they are capable of expanding their octet which is not possible in the second-row elements.

Third Row Elements:

  • A group of eight elements namely, sodium (Na), magnesium (Mg), aluminium (Al), silicon(Si), phosphorous (P), sulphur (S), chlorine(Cl) and argon (Ar) belongs to the third period of the periodic table are called as third row elements.
  • The properties of each element are character­istics of the group to which they belong.  Each one of these elements can, therefore, be considered as the representative of the whole group to which it belongs.  Hence the third-row elements are known as typical or representative elements.
  • The properties of these elements gradually change from metallic or basic to non-metallic or acidic character across the third period. Sodium, magnesium and aluminium are metals.  They have a metallic lustre, can conduct heat and electricity.  They are malleable and ductile.  Sodium is very soft metal. Silicon is hard solid and is metalloid. Phosphorous is yellow waxy solid.  It is non-metal.  Sulphur is yellow coloured solid.  It is a non-metal.  Chlorine and argon are gases at room temperature.  They are non-metals.

Position Of The Third Row Elements In The Periodic table:

Third Row Elements 01

  • In the modern periodic table, third row elements have been placed in 1, 2, 13, 14, 15, 16, 17 and 18 groups. Except for Sodium and Argon, the third-row elements are the second member of their group. Sodium and Argon  are the third members of their groups 1 and 18 respectively
  • For sodium and magnesium, the last electron to be configured enter into ‘s’ orbitals. Hence sodium and magnesium are s block elements. For all the other third row elements the last electron to be configured enter into ‘ ‘p’ orbitals hence they are p block elements.
  • There are three orbits and valence electrons are present in third main energy level.


Electronic Configuration of Third Row Elements:

  • Electronic configuration of third row elements of the third row is as fol­lows

Third Row Elements 02

  • It can be seen that Sodium ( Z = 11)  has a single electron in its 3 s valency orbital. With the increase in atomic number, in magnesium, aluminium, silicon, phosphorous and chlorine, the electrons successively occupy 3s and 3p valence orbitals until another closed shell configu­ration 1s22s2 2p63s23p6 is reached at argon (Z = 18).

Periodic Trends in Third Row Elements:

Ionisation Enthalpy:

  • The minimum energy required to remove the most loosely attached electron from the outermost shell of a neutral gaseous, isolated atom of an element in its ground state to produce gaseous cation is known as ionisation enthalpy of that element.
  •  Ionisation potential is expressed in terms of kJ per mole

M(g)       +    I. P.       →     M(g)+     +       e

Where M is third row element



  • Factors Affecting the ionisation Enthalpy:

    • The size (atomic radius) of an atom i.e. the distance of the outermost electron from the nucleus.
    • The charge on the nucleus or the nuclear charge i.e. protons present in the nucleus.
    • The screening effect.
    • The type or the geometry of the subshell in which the electron is present.
    • For the third row elements, the screening effect for all the elements is almost the same.
  • Trends in the ionisation Enthalpy Across the Period:

    • With some minor exceptions, the trend in third row elements is that as we move from left to the right i.e. from Sodium to Argon along period, ionisation enthalpy of these elements goes on increasing steadily.
    • The values of ionization enthalpy of third row elements are given below.
Element Na Mg Al Si P S CI Ar
Atomic Number (Z) 11 12 13 14 15 16 17 18
First Ionization Enthalpy in kJ/mole 496 737 577 786 1012 999 1252 1520
  • The irregularity in the trend can be observed from the following graph.

Third Row Elements 03

  • The ionisation enthalpy of magnesium is more than aluminium. The ionisation enthalpy of phosphorous is more than sulphur. The last element argon has the highest ionisation enthalpy. Magnesium, phosphorous and argon have more ionisation enthalpy than expected due to the extra stability of their exactly half filled and completely filled orbitals.
  • Ionisation enthalpy of third row elements goes on increasing steadily along the period. 

    • Ionisation enthalpy of an element depends upon i) atomic size ii) nuclear charge iii)  screening effect.
    • If the element has smaller atomic size, greater nuclear charge and less dense inter electronic cloud, then it has greater ionisation enthalpy. If the element has bigger atomic size, low nuclear charge and denser electronic cloud then it has low ionisation potential.
    • Along third period i.e. from Na to Ar, as atomic number increases, the nuclear charge goes on increasing, atomic size goes on decreasing, the number of valence electron goes on increasing. The attractive force on the valence electron of an atom increases. Hence Ionisation enthalpy of third row elements goes on increasing steadily along the period.


  • The ionisation enthalpy of magnesium is more than aluminium.

    • The atomic number of magnesium is 12. Its electronic configuration is 2, 8, 2. Its detailed configuration is 1s22s2 2p63s2. The box diagram for final orbit configuration for magnesium is as below. We can see that Mg has completely filled ‘3s’ orbitals.

    • The atomic number of aluminium is 13. Its electronic configuration is 2, 8, 3. Its detailed configuration is 1s22s2 p63s23p1. The box diagram for final orbit configuration for aluminium is as below. We can see that aluminium has partially filled ‘3p’ orbitals.

Third Row Elements 05

    • It has been observed that the extra stability is associated with vacant, half filled and completely filled orbitals. magnesium has thus extra stable electronic state due to completely filled ‘3s’ orbital. It is difficult to remove an electron from pair due to extra stability. While aluminium has no such extra stable state.  There is one un­paired electron in the 3p subshell of aluminium.
    • s orbitals are closer to the nucleus than p orbitals. So energy required to remove the electron from s orbital is always more than to remove it from p orbital. In case of magnesium, the electron is to be removed from s orbital while in case of aluminium it is to be removed from p orbitals. Thus More energy will be required to remove valence electron from magnesium atom than that of aluminium. So ionisation enthalpy of magnesium is greater than that of the aluminium.
  • The ionisation enthalpy of phosphorous is more than sulphur.

    • The atomic number of Phosphorous is 15. Its electronic configuration is 2, 8, 5. Its detail configuration is 1s2, 2s2 2p6, 3s2 3p3. The box diagram for final orbit configuration for phosphorous is as below. We can see that the p orbitals are exactly half filled.

Third Row Elements 06



    • The atomic number of Sulphur is 16. Its electronic configuration is 2, 8, 6.Its detail configuration is 1s2, 2s2 2p6, 3s2 3p4. The box diagram for final orbit configuration for sulphur is as below. We can see that sulphur has partially filled ‘3p’ orbitals.

Third Row Elements 07

    • It has been observed that the extra stability is associated with vacant, half filled and completely filled orbitals. Phosphorous has thus extra stable electronic state due to exactly half filled ‘3p’ orbital. Sulphur has no such extra stable state as there are two un­paired electrons in 3p subshell. In sulphur, removing one electron from 3p orbital, it becomes half filled and attains the stable state. Thus sulphur tries to lose electron fast. Thus more energy will be required to remove valence electron of phosphorous than that of sulphur.  So ionisation potential of phosphorous is greater than that of the sulphur.
  • Argon has the highest ionisation potential.

    • The atomic number of argon is 18. Its electronic configuration is 2, 8, 8. Its detail configuration is 1s2, 2s2 2p6, 3s2 3p6. The box diagram for final orbit configuration for argon is as below. We can see that argon has completely filled ‘3s’ and ‘3p’ orbitals.

Third Row Elements 08

    • It has been observed that the extra stability is associated with empty, half filled and completely filled orbitals. Argon the as most stable electronic configuration. It has a complete octet of electrons in the outer most shell. All electrons are paired. It is very difficult to remove valence electrons. Hence much more energy is required to remove an electron from such a stable configuration. So Argon has the highest ionisation potential.
  • Concept of Higher Ionisation Potentials:

Third Row Elements 09

  • Aluminium rarely forms the tri-positive cation. 

    • Aluminium has atomic number 13. Its electronic configuration is 1s22s2 p63s23p1. To form a tri-positive ion of aluminium all the three electrons of the outermost shell are to be removed.
    • Formation of mono-positive ion i.e. cation involves the removal of an electron from a neutral atom. Hence the first ionisation potential is always low. Formation of di-positive ions involves removal of an electron from mono-positive cation while the formation of tri-positive ion involves removal of an electron from di-positive cation. Due to the positive charge on the mono-positive and dipositive cations, the outgoing electron has to overcome larger attractive force.
    • Hence more energy is required to remove the second electron and still more energy is required to remove the third electron. Hence Aluminium rarely forms tri-positive ions.

Metallic and Non Metallic Characters or Electro Positive and Electro negative Characters:

  • Metallic Character: The tendency of an atom to lose electrons to form positively charged ion is called its metallic character or electropositive character.
  • Non-Metallic Character: The tendency of an atom to gain electrons to form negatively charged ion is called its non-metallic character or electronegative character.
  • Factors affecting metallic and non-metallic characters :
    • Size of an atom
    • ionisation enthalpy
    • Nuclear charge
  • Trends in metallic and non-metallic characters:

    • As we move across the periodic table from left to the right i.e. from sodium to argon in the third period metallic character decreases and non-metallic character increases. Sodium, magnesium, aluminium are typical metals. Silicon is metalloid. (weakly non-metallic). Phosphorous, Sulphur, Chlorine are non-metals. Argon is an inert gas.
    • Sodium is the most metallic element while chlorine is the most non-metallic element. Both the extreme elements sodium and chlorine are extremely reactive.
    • Argon is neither electropositive nor electronegative.
  • As we move across the periodic table from left to the right i.e. from sodium to Argon in the third period metallic character decreases and non-metallic character increases.

    • As we move from left to right In the third period the atomic number goes on increasing. Thus as we move from sodium to agon the nuclear charge increases and from left to right additional electron is added to the same i.e. third orbit. Due to increase in the nuclear charge the attractive force on electrons in outermost orbit increases and thus the size of atom decreases.
    • As we move from left to right In the third period the ionisation enthalpy increases. From left to right In the third-period electronegativity increases i.e. the tendency of losing electrons decreases and tendency of gaining electrons increases. As we move from left to right In the third period number of valence increases. Hence as we move across the periodic table from left to the right i.e. from sodium to Argon in the third period metallic character decreases and non-metallic character increases.


  • Sodium, Magnesium, Aluminium are typical metals.

    • Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionisation potential. Sodium, magnesium, aluminium have atomic numbers 11, 12, 13 respectively. They contain 1, 2, 3 valence electrons in their outermost shell.
    • Compared to other elements in the third row their atomic sizes are larger. Similarly, the nuclear charges are also less. Due to this attractive force on valence electrons is less. Compared to other elements in the third row their ionisation enthalpies are lower. Thus the tendency of these elements is to lose their valence electrons. Hence sodium, magnesium, aluminium are typical metals.
  • Phosphorous, Sulphur, Chlorine are non-metals.

    • Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionisation potential. Phosphorous, sulphur, chlorine have atomic numbers 15, 16, 17 respectively. They contain 5, 6, 7 valence electrons in their outermost shell.
    • Compared to other elements in the third row their atomic sizes are smaller. Similarly, the nuclear charges are also more. Due to this attractive force on valence electrons is more. Compared to other elements in the third row their ionisation enthalpies are higher. Thus the tendency of these elements is not to lose their valence electrons but to gain electrons. Hence phosphorous, sulphur, chlorine are non-metals.
  • Argon is neither electropositive nor electronegative element.

    • Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionisation potential. Argon has atomic number 18. It contains 8 valence electrons in their outermost shell. Argon has completed octet and completely filled s and p orbitals thus it has the most stable electronic configuration in third row elements which it will not disturb by accepting or losing the electron.
    • Compared to other elements in the third row its atomic sizes are larger but the effect of filled s and p orbitals is dominating. Compared to other elements in the third row its ionisation enthalpy is the highest. Thus the tendency of Argon to lose or gain valence electrons is almost absent. Hence Argon is neither electropositive nor electronegative element.
  • Sodium Is most metallic element.

    • Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionisation potential. Sodium has atomic numbers 11. It contains 1 valence electron in its outermost shell.
    • Compared to other elements in the third row its atomic size is larger. Similarly, the nuclear charge is also less. Due to this attractive force on valence electron is less. Compared to other elements in the third row its ionisation enthalpy is the lowest. Thus the tendency of Sodium is to lose its valence electron. Hence Sodium is the most metallic element.
  • Chlorine is the most non-metallic element. 

    • Metallic and non-metallic characters of the elements depend on their atomic size, nuclear charge and ionisation potential. Chlorine has atomic number 17. It contains 7 valence electrons in its outermost shell.
    • Compared to other elements in the third row its atomic size is the smallest. Similarly, the nuclear charge is also more. Due to this attractive force on valence electrons is more. Compared to other elements in the third row its ionisation enthalpy is the highest. Thus the tendency of Chlorine is not to lose its valence electrons but to gain electrons. Hence Chlorine is the most non-metallic element.

Oxidising and reducing property:

  • Oxidation: The process in which an atom, a molecule or an ion loses one or more electrons is called oxidation. It is also known as de-electronation.

e.g.            Na      →    Na+   +   e

In this case, the oxidation of sodium is taking place.



  • Reduction: The process in which an atom, a molecule or an ion gains one or more electrons is called reduction. It is also known as electronation.

e.g.             CI      +    e    →     Cl

In this case, the reduction of chlorine is taking place.

  • Reducing agent: A substance ( an atom, a molecule or an ion) which forces another substance to accept electrons and it itself undergoes oxidation by losing electrons is called reducing agent. Reducing agent is electron donor. e.g.  Na, Al, Mg etc.
  • Oxidising agent: A substance ( an atom, a molecule or an ion) which forces another substance to lose electrons and it itself undergoes reduction by accepting electrons is called oxidising agent. The oxidising agent is an electron acceptor. e.g. Cl, F, Br, O etc.
  • Reducing property: The tendency of an element to lose electrons is called its reducing property. By virtue of this property, the substance itself undergoes oxidation.
  • Oxidising property: The tendency of an element to gain electrons is called its oxidising property. By virtue of this property, the substance itself undergoes reduction.
  • Factors affecting the oxidising and reducing property :

    • Ionisation potential
    • electropositivity or electronegativity
    • Atomic size
    • Metallic and non-metallic character
    • Number of valence electrons
  •  Trend in oxidising and reducing property :

    • As we move from left to right i.e. from sodium to chlorine along the third row, the oxidising property goes on increasing while reducing property goes on decreasing.
    • Sodium, magnesium, aluminium are good reducing agents. Silicon, phosphorous and sulphur are weak reducing agents.
    • Sodium is the strongest reducing agent. Chlorine is the strongest oxidising agent. Argon is neither oxidising agent nor reducing agent.
  • As we move from left to right i.e. from Sodium to Chlorine along the third row, the oxidising property goes on increasing while reducing property goes on decreasing.

    • Oxidising and reducing strength of elements depends upon the atomic size, ionisation enthalpy, electropositive and electronegative character and the number of valence electrons.
    • Those elements who have bigger atomic size, lower ionisation enthalpy and few valence electrons tend to donate electrons and hence are reducing agents. Hence sodium, magnesium, aluminium are reducing agents.
    • Those elements who have greater ionisation enthalpy, smaller atomic size and more valency electrons tend to accept electrons and hence are the oxidising agent. Hence Chlorine is an oxidising agent.
    • It is observed that, as we move from left to right along the third period atomic size gradually decreases, ionisation enthalpy increases and the number of valence electrons increases. Hence electron donating tendency of elements goes on decreasing and that of electron gaining tendency of elements goes on increasing. Hence as we move from left to right i.e. from sodium to chlorine along the third row, the oxidising property goes on increasing while reducing property goes on decreasing.


  • Sodium, magnesium and aluminium are good reducing agents.

    • Oxidising and reducing strength of elements depends upon the atomic size, ionisation enthalpy, electropositive and electronegative character and number of valence electrons.
    • Atomic numbers of sodium, magnesium and aluminium are 11, 12 and 13 respectively. They have 1, 2, and 3, valence electrons respectively. Thus removal of 1, 2, 3 electrons from sodium, magnesium, aluminium respectively would give the stable inert gas configuration of neon.
    • Compared to other third row elements these elements have bigger atomic sizes and lower ionisation potentials.  Hence they readily lose their valency electrons and are thus strong reducing agents. Reducing strength decrease in the order.  Na > Mg > Al.
  • Sodium is strongest reducing agent. 

    • Oxidising and reducing strength of elements depends upon atomic size, ionisation enthalpy, electropositive and electronegative character and number of valence electrons.
    • Atomic number of sodium is 11. It consist of single unpaired electron (3s1). Thus removal of 1 valence electron would give sodium a stable inert gas configuration of neon.
    • Sodium has the largest atomic size among third low elements and the lowest ionisation potential among third row elements. Hence Sodium readily lose its  valence electron and is thus strongest reducing agents.

Na      →    Na+   +   e

  • Silicon , phosphorous and sulphur are weak reducing agents.

    • Oxidising and reducing strength of elements depends upon atomic size, ionisation enthalpy, electropositive and electronegative character and number of valence electrons.
    • Atomic numbers of Silicon , phosphorous and sulphur are 14, 15 and 16 respectively. They have 4, 5, and 6, valency electrons respectively.
    • They have comparatively smaller atomic size and higher ionisation potential. Hence they show less tendency to lose their valency electrons. They are less electropositive. Hence they are weak reducing  agents.
    •  They act as weak reducing agents when treated with strong oxidising agents like fluorine. They also act as weak oxidising agent with strong reducing agents like sodium.
  • Chlorine is strongest oxidising agent in third row elements.

    • Oxidising and reducing strength of elements depends upon atomic size, ionisation enthalpy, electropositive and electronegative character and number of valence electrons.
    • Chlorine has high electronegativity. Atomic number of chlorine is 17. It consist of seven electrons in valence shell. It requires only one electron to complete its octet and attain stable electronic configuration of argon.
    • Chlorine has the smallest atomic size among third low elements.
    • Chlorine has very high ionisation potential and very high electron affinity.
    • Hence it has strong tendency to gain an electron and it is highly electronegative. Hence Chlorine is the strong oxidising agent.

 CI      +    e    →     Cl

  • Argon is neither oxidising agent nor reducing agent.

    • Oxidising and reducing strength of elements depends upon atomic size, ionisation enthalpy, electropositive and electronegative character and number of valence electrons.
    • argon is neither electropositive nor electronegative element Atomic number of Argon is 18. It consist of eight electrons in valence shell. Thus it has completed octet. s orbital and p orbitals are completely filled. Hence it has stable electronic configuration.
    • Argon has very high ionisation enthalpy. Hence it has strong tendency not to gain or lose electrons. Hence Argon is neither oxidising agent nor reducing agent.

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