First, Before we talk about anything else , we need to clear this one thing up ...What is the Periodic Table?
Periodic Table may be defined as the arrangement of all the known elements according to their properties in such a way that the element of similar properties are grouped together in a tabular form.
Its a periodic representation of all the elements, in a table ( But, you will be surprised by the amount of information the scientist were able to pack in one chart! )
*The Periodic Table is divided into
1) 18 vertical columns , also called Groups
2) 7 horizontal rows , also called Periods
3) Its divided into 4 blocks (namely s,p,d and f) -- This tells us about the electronic configuration of the element , the subshells or orbitals present in the orbits
If you are wondering what "s, p ,d and f" orbital is:-
Within the shells, electrons are further grouped into subshells of four different types, identified as s, p, d, and f in order of increasing energy. We can think of the orbitals as small shells within the bigger shells, which determine its holding capacity. The orbitals are arranged in the following order 👇
Hence from the above diagram 👆 (not exactly as it is in an atom , made only for easier visualisation) , we understand now, how the 2n^2 rule works for each energy shell
*K shell - 2 (1 ^2) = 2
-- S orbital within the K shell holds 2 electrons
*L shell - 2 (2 ^2) = 8
-- S orbital and P orbital of the L shell respectively hold 2 and 6 electrons respectiely
-- S orbital + P orbital = 2 + 6 electrons = 8 electrons and so on
By studying the periodic table we find out about these subshells too, as each block represents and different subshell
* Basically they represent the orbitals in which an element's valence electrons resides.
Example as 'Li' is in the S block :-
This means that the valence electron of the Lithium atom is present in the S subshell ( under L shell.)
(Again ,Not exactly how it looks in an Lithium atom, diagram given just for easier visualisation)
Now that we know what a
1. Period
2. Group and
3. Block
let's get really started and go through this fashion show of the latest trends in the periodic table!
1- VALENCY
Valency is defined as the combining capacity of the atom of an element with the atoms of another element in order to aquire 8 electrons (octet rule ) or 2 electrons in exceptional cases (Looking at you Hydrogen, Lithium and Helium) in the outermost shell to obtain stability
Points to remember
1) Noble gases' valency = 0 ( as they have already achieved octet configuration)
2) Elements with valence electrons = 1,2 or 3 tend to loose electrons easily hence show valency of 1,2 or 3 respectively and obtain positive charge after combining ( these elements are also known as Metals)
3) Elements with higher number of electrons in valence shell like 5,6 or 7 tend to gain electron more easily and hence show valency of (8-5 =) 3, (8-6=) 2 , (8-7=) 1 valency respectively and obtain negative charge after combining ( also known as Non - Metals )
* In Groups :-
All elements of the same group (remember - They are vertical columns) have the same valency (same number of valence electrons in outermost shell )
* In a Period :-
Valency first increases from 1 to 4 then decreases from 4 to 0
for example -- 
The diagram shows how the 17th grp and 18th grp elements are bonded ( again not exactly how they appear, just for easier visualisation ( yep, thats the third time )
Okay, few of you might have been thinking "Hey, something is off in this example" ( If you didn't go, back and check it once again)
(What it roughly looks like)
Here, we can see that Be has a fully filled last orbital ( S orbital)
Nitrogen has a half filled orbital ( The maximum occupancy of P orbital is 6)
In Group :- Down the group, the electronegativity decreases as the size increases ( decreasing the effective force the nucleus has on the outermost shells)
2. ATOMIC RADIUS
What is Atomic Radius?
It is the distance between outermost sell and nucleus , ( basically , the radius of the atom = Atomic radius 😎)
Usually this is calculated as the half of the nuclear distance between two atoms( as few elements are not stable on their own hence can only exist as pairs (combined form ) in the nature.)
*In Groups :-
On moving down the group of the periodic table, the size of the atom increases , as the number of shells increases as we go from top to bottom.
* In Periods:-
On moving from left to right in a period, the size of the atom goes on decreasing as the nuclear charge increases while the number of shells remain same. This results in more pull of the outermost shell by the nucleus as we go from left to right.
Egs:- Li > Be > B > C > N > O > F < Ne
( Exception :- The inert gases at the end of the period always have a greater atomic radius than the 17th group. This is because of the Wanderwaal's Bonding they posses)
The diagram shows how the 17th grp and 18th grp elements are bonded ( again not exactly how they appear, just for easier visualisation ( yep, thats the third time )
Hence, we see how the Covalent bond has a reduced internuclear distance than in the Vanderwaal's bond ( cause its weaker), so the Vanderwaal's bonded atoms have a greater atomic radius than Covalent bonding.
3. Ionic Radius
Cationic Radius :
Cations are when the neutral atom loses an electron, hence aquiring a positive charge.
Atomic radius > Cationic Radius
Egs:- In lithium, when the atom loses an electron and becomes Li+ , It also loses a shell ( the L shell) as there are no electron in that shell anymore, and the number of protons in the nucleus remains same though the number of electrons reduce. This causes the electrons to get pulled towards the nucleus and causing the atomic radius to reduce.
Anionic Radius :
Anions are formed when the neutral atoms gains electrons, hence aquiring a negative charge
Anionic Radius > Atomic radius
In an Anion, as the electrons now are in majority than the protons, the inter electronic repulsion increases. This causes the distance between the electrons and nucleus to increase and ultimately the anion sizes increases.
4 .Metallic and Non Metallic Characters
Metals -- They are generally electro-positive as the have a strong tendency to loose electrons ( because of their valency)
Non-Metals -- They are electro-negative as they have a tendency to gain electron to form negative ions.
*In Groups :-
As we go down, the metallic character or electro-positive character increases.
When we move down the group, the atomic radius increases, hence the tendency of the atoms to loose an electron increases .The force of attraction between the nucleus and the outermost electrons decreases making it easier to loose electrons
Similarily, the non-metallic character goes on decreasing as we move down the group. As the atomic radius and effective nuclear attraction on the outermost shell decreases , the chance of the atom to attract and hold electrons from other atoms reduces. Hence the non-metallic character decreases.
*In Period :-
As we move across the period, The metallic character decreases and Non-metallic character increases. This is because of the trend shown by the periodic table in relation to the valency as we go from left to right. ( scroll up to recap about the valency)
5. Ionization energy
Ionization energy / Ionization energy / Enthalpy is the minimum energy required to remove the valence electrons from neutral 'gaseous' atom. (isn't that a gassy definition?)
Hence if E1 is the Ionization energy for an element M then:-
M + E1 = M⁺ + e⁻
In Group:-
As we go down the group, the ionization energy required to remove the valence electrons decreases .As the atomic size increases, the pull of the nucleus also has lesser effect hence making it easier to remove the valence electrons
In Period:-
Across the Period, as we move left to right,the ionization energy required increases as the atomic size decreases.
Across the period, the non-metallic character increases with the decrease in metallic character, this means that the atom is more likely to gain charges than loose it. ( Its like asking a person to run a race, she/he is less likely to quit as she/he sees the finish line approaching, now imagine the valence electrons as their persistence, and the race they run is the race for stability in electronic configuration) . The closer the atom is to the next octet configuration or filling their valence shells, the more energy it will take to remove their valence electrons.
for example :-
Once you have gone through the catchy humming of the Periodic table song (If you haven't listened to the song , check out this link The Periodic Table Song! ), You might notice that the places of Berylium , Boron and Nitrogen , Oxygen have been interchanged. To understand this we need to do a tiny flash back.
Remember the segment where we spoke about Orbitals and Blocks ? Well, they play a very important role here.
One way of writing the electronic configuration of an element is through mentioning the orbitals, for example:-
For now, we don't need to worry about the entire order, we just need to concentrate on the last term of each sequence because they are the ones which contain the electron which usually react with other elements, loosing and gaining of electrons occur here (If you are curious , basically the numerical coefficient talks about which energy level / shell that particular orbital occupies , and the exponent on the letter specifies the number of electron
the particular orbital ,
For example:- When we say Hydrogen -- 1s¹ , it means in the K shell ( 1st energy level ), the S orbital is occupied by 1 electron
Similarily, Li -- 1s² , 2s¹ tells us that in the K shell of lithium , the S orbital is occupied by two electron and in the L shell ( 2nd energy level) the S orbital has 1 electron.
(What it roughly looks like)As per the SYMMETRICAL DISTRUBTION of Electrons,
Symmetry leads to stability. The orbitals in which the sub-shell / orbital is exactly half-filled or completely filled are more stable.
Basically, when the orbitals have exactly half number of electrons than their maximum capacity or have it filled are considered relatively stable. ( but, Fully filled ones are more stable than half filled)
Nitrogen has a half filled orbital ( The maximum occupancy of P orbital is 6)
And Neon has its last orbital ( P sub shell) completely filled.
As we have seen in the symmetrical distribution, a completely filled orbital is relatively stable than the partially filled ones, hence Berylium is considered more stable than Boron. When an atom become stable, it requires more energy to remove the electron from it compared to the partially filled atoms, hence Beryllium has a higher Ionization potential than Boron.
Similarily, Nitrogen having a half filled sub shell, is more stable than Oxygen. Hence, It will require more energy to remove an electron from Oxygen than from Nitrogen.
6. Electron negativity
It is the tendency of an atom to attract the shared pair of electron towards itself.
When a bond is formed, there is a tranfer or sharing of electron between two atoms, one atom tends to attract the electron shared or tranferred towards itself than the other partner it forms the compound with, Hence the first one is called to be more electro'negative' than the latter ( No one likes hoghead now, do they ?)
In Period:- From left to right, the Electro Negativity increases as the size decreases.
Phew, The Periodic Table fashion show has finally ended, I hope after reading this showstopping array of trends , you will never look at the periodic table the same way again.😝
It's just one of the mysterious clothing styles and trends chemistry has to offer, so keep your eyes peeled for more of chemistry blogs!
BRING ON THE EXAMS!
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