Name of All Topics
Genesis of PT or Earlier Attempts:
Proust Hypothesis:
According to him all elements are aggregation or made up of hydrogen and its isotopes.
H11 + 2H13 –> Li37
De-Merits:
Atomic mass that can be fraction.
Lavoisier’s Classification (father of chemistry):
He arranged the element into metals and non-metals.
Metals:
Which have tendency to lose e- i.e. electropositive.
Non-Metals:
Which have tendency to gain e- i.e. electronegative.
De-Merits:
It is failed after the discovery of metalloids.
Do Berner’s Triad:
He arranged the elements into set of three elements known as triads in such away that atomic mass of middle element is average of other two or all 3 elements have nearly same atomic mass.Example:
Li37 Na1123 K1939 (7 + 39)/2 = 23, Ca2040 Sr3888 Ba56137 (40 + 137)/2 = 88, Cl1735.5 Br3580 I53127 (35.5 + 127)/2 = 80.De-Merits:
All elements doesn’t follow triads rule. He arranged the elements only in group.Example:
F919 Cl1735.5 Br3580 (19 + 80)/2 = 49.5Newland’s law of octave:
He was the first scientist which corelate properties of elements with atomic mass. According to him if elements are arranged in increasing order of their atomic mass then every 8th element will have similar properties to first element just like musical node. Sa Re Ga Maa Pa Dha Nee Li7 Be9 B11 C12 N14 O16 F19 Na23 Mg24 Al27 Si28 P31 S32 Cl35.5 K39 Ca40Merits:
Coming Soon…De-Merits:
It works well only for lighter elements (Up to Ca). Nothing says about the inert gases.Luther’s Mayer Classification:
He determined atomic Vol. of elements by dividing their atomic mass with density of elements in solid state.Observation:
Elements having similar properties occupies similar position in curve. Alkali metals occupies top position (Crust) of the curve. Alkaline earth metals occupies positions in the position order of the curve. Halogens occupies positions in the ascending position of curve. Heavy metals, metalloids occupies position at trough of the curve.Conclusions:
On the basis of this observation, he concluded that atomic volume (which is a physical properties) is periodic function of atomic mass of elements.Merits:
He was the first scientist who represents the elements on the graph.De-Merits:
Lacks of practically utility. Different curve for different physical properties. For larger elements curve becomes confusing.Mendeleev’s Periodic Table:
MPT is based on Mendeleev’s laws according to which physical and chemical properties of elements are periodic function of their atomic masses. Nearly 63 elements are arranged in their order of increasing atomic masses into horizontal rows and 9 verticals column or group or family. 7 Horizontal rows: 1 2 3 4 5 6 7 9 Verticals columns: 0 I II III IV V VI VII VIII Each group is divide into 2 groups (except 0 and VIII) Group-A (Main or Normal elements i.e. s and p blocks elements) Group-B (Transitive elements i.e. d block elements)Merits:
He has simplified and systemize study of elements. He left some vacant positions in his periodic table for discovery of new elements in future.Note:
eka Aluminum (Al) à Gallium (Ga) eka Boron (B) à Scandium (Sc) eka Silicon (Si) à Germanium (Ge) eka Manganese (Mn) à Tellurium (Te)Mnemonics:
Alina Gaegi B Sc me nhi Sidha Gaegi M Tech me nhi AIIMS me. Corrections of atomic mass.Example:
Be à atomic mass 13.5 to atomic mass 9. At. Wt. = n-factor * equivalent weight atomic mass of Be = 2 * 4.5 atomic mass of Be = 9Note:
U Be In Pt Au U Beimaan Insaan Pt Au churate hainDe-Merits:
Not say anything about the position of hydrogen. (i.e. Group 1 or IA, Group 17 or VIIA). No separate position was given for isotopes as they have different atomic mass. Mendeleev could not explain cause of periodicity. Similar elements placed at different positions while dissimilar elements placed at same position.Mnemonics:
Note:
Ar K Arbaaz Khan Te I Teri Icha Th Po Thanda Peppsi Co Ni Coak NiModern Periodic Table or Long form of Periodic Table or Mosley’s Periodic Table or Bohr Periodic Table:
Based on experiment Mosely proposed modern periodic laws which states that “ Physical and chemical properties of elements are periodic function of their “Atomic Number”. If elements are arranged in increasing order or their atomic number(Z), then after a regular interval elements with similar properties are repeated.Periodicity:
Repetition of elements after a certain interval when the elements are arranged in increasing order of atomic number is known as periodicity.Cause of Periodicity:
Periodic representation of properties of elements after a certain regular interval is due to reoccurrence of similar electronic configuration.Note:
Modern Periodic Table have 7 Horizontals rows is known as Period and 18 columns is known as Group. Contributions of lots of scientist example Bohr, Range, Warner, Barry etc. This modern periodic table is also known as Bohr Periodic Table because it follows the Bohr skim of classification of elements into four types s, p, d, and f. Modern Periodic Table:
Periodic Table Mnemonics:
Coming Soon…
How to remember atomic number of elements of the periodic table:
If we need to remember of atomic number of elements then use magic number:
1.s-block elements:
AM i.e. Group-1:
2, 8, 8, 18, 18, 32
AEM i.e. Group-2:
8, 8, 18, 18, 32
- p-block elements:
Group(13-17):
8, 18, 18, 32
Group(18):
8, 8, 18, 18, 32
- d-block elements:
Group(3 – 12):
18, 32, 32
- f-block elements:
32
Classification of elements:
1.s-Block 2.p-Block 3.d-Block 4.f-block- s-block:
Note:
Be is not considered as AEM but why? Because Oxide of Be is amphoteric in nature.Note:
Be is not considered as AEM but why? Because Oxide of Be is amphoteric in nature.Characteristic:
Elements of G-1 and G-2 constituents s-block. They are present in extreme left on PT. Their General Electronic Configuration: [IG] ns1,2 This block includes metals which shows fixed valency. For AM valency is 1 and for AEM valency is 2. They are good conductor of electricity. They are soft in nature it can be cut with the help of knife it can have shining surface, it is ductile in nature.2.p-block:
The elements in which (n-1)th shell completely filled and differentiating electron i.e. last electron enter into p-subshell of outermost shell i.e. valence shell or nth shell.
Example:
Groups | General configuration | Name of groups | Other name of groups |
13 | ns2np1 | Boron family | |
14 | Ns2np2 | Carbon family | |
15 | Ns2np3 | Nitrogen family | Pnictogen |
16 | Ns2np4 | Oxygen family | Chalcogen |
17 | Ns2np5 | Halogen family | Halogen |
18 | Ns2np6 | Nobel gas | Zero group |
Characteristics:
Elements of G-13 to G-18 constituents p block of PT.
They are present on extreme right of PT
General Electronic configuration:
ns2np1-6
This block includes metal, non-metals and metalloid.
It shows variable valencies.
Solid/liquid/gas are the physical properties of p-block elements.
3.d-block:
The elements in which outer most shell (i.e. nth shell) and penultimate shell (n-1)th shell are incompletely filled and differentiating electron enters
Into d-subshell of outermost shell.
Characteristic:
Elements of G-3 to G-12 constituents d-block of PT.
They are present in middle in PT (in b/w s-block and p-block).
Their properties are in b/w s-block and p-block hence they are also known as “Transition elements”.
Note:
Transition Elements(TE):
Elements which contains at-least unpaired electron in their ground state or any of ion (ES) or in compound are known as “TS”.
Note:
All TE are d-block but all d-block elements are not TE.
(Except Zn, Cd, Hg).
Their general EC is:
[IG] (n-1)d1 to 10 ns1,2
Exception: Pd – 5s0 4d10
Zn (Z=30) à [Ar] 3d10 4s2
Cd (Z=48) à [Kr] 4d10 5s2
Hg (Z=80) à [Xe] 4f14 5d10 6s2
d-block form coordinate compounds due to presence of vacant ‘d’ orbitals.
They act as good catalyst.
It form coloured compound.
Normally show paramagnetic.
Alloys (Homogenous mixture of metals).
Normally solid.
4.f-block:
Elements in which outermost shell i.e. nth shell penultimate shell i.e. (n-1)th shell and anti-penultimate shell are incompletely filled and differentiating electron enters
Into f-subshell of anti-penultimate shell i.e. (n-2)th shell are known as f-block elements.
Characteristic:
4f- Lanthanoids Series: (Ce58 – Lu71)
5f- Actinoids Series: (Th90 – Lr103)
In f-block number of elements in group -3 or group- IIIB is 28 elements.
They are present at the bottom of the PT.
Their General electronic configuration:
(n-2)f1-14 (n-1)d0,1,2 ns2
They all are metals and properties with in each series are quite similar.
Description of Periods:
1st Period have 2 elements (H, He).
This period is also called ‘Very Short Period ’ because of 1s = 2e-.
2nd Period contains 8 elements ( Li- Ne ) called ‘short period’ because of 2s2p = 8e-.
3rd Period contains 8 elements as well. (Na – Ar) also called’ short period’ because 3s3p = 8e-.
4th Period contains 18 elements (K – Kr) called ‘long period’ because of 4s 3d 4p = 18e-.
5th Period contains 18 elements (Rb – Xe) called ‘long period’ because of 5s 4d 5p = 18e-.
6th Period it contains 32 elements (Cs – Rn) called ‘longest’ Period of PT because of 6s 4f 5d 6p = 32e-.
7th Period contains 32 elements (Fr – Og) called also called ‘longest period of PT’ because of 7s 5f 6d 7p = 32e-.
Note:
Maximum elements in a period of PT = 6th and 7th period (32 elements).
Maximum elements in a group of PT = Group-3 or Group- IIIB (32 elements).
Each period start from ns and end to np:
Period | Subshells | Elements |
1 | 1s | 2 |
2 | 2s 2p | 8 |
3 | 3s 3p | 8 |
4 | 4s 3d 4p | 18 |
5 | 5s 4d 5p | 18 |
6 | 6s 4f 5d 6p | 32 |
7 | 7s 5f 6d 7p | 32 |
To determine maximum elements in period:
Maximum elements in a period = mn2
m = number of electrons in an orbital (but in general m = 2 )
How to find the value of n?
Case-I:
If p is odd then n = (p+1)/2
Case-II:
If p is even then n = (p+2)/2
Examples:
IUPAC nomenclature of heavy elements (i.e. Z >= 100)
According to IUPAC elements with atomic number Z >= 100 are represented by 3 latter symbol which are based on 1st latter of number from 0 to 9 (derived from Greek or Latin words). Latin words of various digits of atomic number are written together in order of which they appear in atomic number and suffix “ium” is added at the end.| Digits | Name | Symbol |
| 0 | Nil | n |
| 1 | Un | u |
| 2 | Bi | b |
| 3 | Tri | t |
| 4 | Quad | q |
| 5 | Pent | p |
| 6 | Hex | h |
| 7 | Sept | s |
| 8 | Oct | o |
| 9 | Enn | e |
Examples:
100 à Unnilnilium (Unn) 112 à Ununbium (Uub) 115 à Ununpentium (Uup) 118 à Ununoctium (UuoDetermine period , block, and group of given elements:
Case-I: When electronic configuration is given. Case-II: When atomic number is given. Case-I: When electronic configuration is given. For Period maximum value of ‘n’ in given configuration. For block if np electron is present then block is p-block. and Group = 12 + number of electrons in p-subshells. if ‘np’ electron is absent then we can say that the block is s/d/f. if electronic configuration is like (n-2)f0 (n-1)d0 ns1,2 s-block à Group-1 (ns1) à Group-2 (ns2)If electronic configuration is like this
(n-2)f1-14 (n-1)d0,1 ns2
f-block and group-3
If any other configuration then
d-block and group = 2 + number of electrons in (n-1)d subshell.
Examples:
Note:
Pd (Z=46) à [Kr] 4d10 5s0
5th period , d-block and 10th group
Case-II:
When atomic number is given.
If atomic number :
58 <= Z <= 71 is called “Lanthanoid series”
then we can say that, Period = 6, f-block and Group = 3rd
90 <= Z <= 103 is called “Actinoid series”
then we can say that, Period = 7, f-block and group = 3rd
If any other atomic number is given then follow:
[IG] | He | Ne | Ar | Kr | Xe | Rn | Og |
Atomic number (Z) | 2 | 10 | 18 | 36 | 54 | 86 | 118 |
Period (p) | 1 | 2 | 3 | 4 | 5 | 6 | 7 |
Note:
Period = Period of next inert gas from the given table
Group = 18 + (Atomic number of given element) – (Atomic number of next inert gas)
Block = Decide block from group
Examples:
Exception electronic configuration:
Note:
No any exception in electronic configuration for s and p-block.
Exception are present in d and f block
D-block:
- Group-3
- Group-4
- Group-5:
V(Z=23) 3d34s2
Nb(Z=41) 4d45s1
- Group-6:
Cr(Z=24) 3d54s1
Mo(Z=42) 4d55s1
W(Z=74) 5d46s2
- Group-7
- Group-8:
Fe(Z=26) 3d64s2
Ru(Z=44) 4d75s1
- Group-9:
Co(Z=27) 3d74s2
Rh(Z=43) 4d55s1
- Group-10:
Ni(Z=28) 3d84s2
Pd(Z=46) 4d105s0
Pt(Z=78) 5d96s1
- Group-11:
Cu(Z=29) 3d10rs1
Ag(Z=47) 6d105s1
Au(Z=79( 6d107s1
- Group-12
Note:
These given below d-subshells are not possible in d-block:
3d series : d4 and d9
4d series: d3, d6 and d9
5d series: d8
General Terminology
Normal/Main/representative elements:
Elements in which only one shell i.e. valence shell is incomplete.
‘s’ and ‘p’ block elements except inert gases.
Typical elements:
Elements of 3rd period which represents properties of their respective group are known as typical elements (except inert gas).
Why not 2nd period elements are typical elements?
Second period elements:
Li Be B C N O F
The main reasons are:
Due to small size of second period elements
Absence of d-orbital(they can’t expand their octoate)
High ionisation energy
High electronegativity
Bridge elements:
Elements which can show similarity in properties between two group. (Mendeleev’s PT).
Na (IA)
K (IA) Cu (IB)
And
Mg (IIA)
Ca (IIA) Zn (IIB)
Diagonal relationship:
There are some elements which shows relationship with element present diagonally opposite to them and it is known as diagonal relationship.
period-2: Li Be B
period-3 Na Mg Al Si
Reason: same ionic potential (phi)
Phi = charge / size
Transition elements:
2 outermost shell incomplete (Zn, Cd, Hg) is not transition it is only d-block elements. Because transition elements is incomplete d-subshell any ground state or exited state.
Inner-transition elements:
3 outermost shell are incomplete.
(n-2), (n-1), and nth f-block
Trans-uranic elements:
Elements after uranium (Z>90) all elements are synthetic and radioactive.
Trans-fermium:
Elements after Z>100 known as TFE.
Periodic Properties:
Atomic properties:
which can be explain by a single atom.
examples: Zeff, IE, EA and EN
Molecular properties:
which can’t be explain by a single atom.
examples: Radius, MP, BP, and density etc.
Effective Nuclear charge (Zeff) :
Effective Nuclear charge (Zeff) :
It is for multi electron species i.e. more than 1 electrons like He, Li, Be, Na, K, …etc.
Shielding Effect or screening effect:
Repulsion created by inner shell electrons or all other electron on testing electron is known as ‘SE’.
Sigma:
Shielding created by other electrons.
i.e. number of protons which utilize in over coming the repulsion created by other element.
Actual nuclear charge experienced by differentiating electron.
Zeff = Total nuclear charge – Screening effect created by other electron
i.e. Zeff = Z – Sigma
where,
Z= Atomic number
Sigma= Slater’s constant
Note:
s-orbital:
it is less diffused in nature and attraction is minimum therefore, shielding power is maximum.
p-orbital:
it is less-more diffused in nature and attraction is more than previous therefore, shielding power is less than s- orbital.
d-orbital:
it is more diffused in nature than p-orbital and attraction is more than p-orbital therefore, shielding power is less than s-orbital.
f-orbital:
it is more diffused in nature than d-orbital and attraction is more than d-orbital therefore, shielding power is least .
Penetration power:
it is the order of shielding power.
Therefore, we can see that s> p> d> f
Rules for determination of sigma:
Write electronic configuration of given elements.
(1s) (2s,2p) (3s,3p) (3d) (4s, 4p)(4d) …
Rule:
All electron in principal shell higher than the electron in given question contribute Zero to Sigma.
i.For (ns, np) electrons:
All electron in same shell = .35
All electron in (n-1)th shell = .85
All electron in (n-1)th, (n-3)th, etc. = 1
(Except 1s e- , Sigma = .3 for He)
ii.For (nd) e-
All e- in same shell i.e. nth shell = .35
All e- in inner shells i.e. (n-1) (n-2) etc. = 1
Note: Along the period on moving from left to right
Ziff will be increasing therefore, size will be decreasing
Note: On moving down the group remains same Zeff but size increasing due to the increasing of number of shells.
Atomic Radius:
Atomic Radius:
Conceptually atomic radius is defined as distance between the nucleus and outermost electron but it is not possible to calculate exact value of atomic radius.
Because,
Atom is too small.
Atom doesn’t have well defined boundaries. (because of no fixed number of shells.)
Wave nature of electron.
Therefore, we will calculate atomic radius of any element in bonded state so it is simply half of the inter nuclear distance.
Nature of radius depends on nature of bonding.
Types of atomic radius:
1.Covalent radius (CR)
2.Vander Waal radius (VR)
3.Metallic radius or crystal radius (MR)
4.Ionic radius (IR)
Covalent Radius
It is defined as half of the internuclear distance between two atom bounded by single covalent bonds.
It of two types:
Homoatomic and Heteroatomic
Case-I:
Homoatomic: like H2, Cl2, F2, Br2, N2 etc.
When electronegative of both atom is same.
dA-B = rA + rB
Case-II:
Heteroatomic: like HCl, HBr, etc.
dA-B = rA + rB – 0.09 (XA – XB) is equation is called Stevenson and shoemaker equation.
Where, rA = covalent radius of atom A
rB = covalent radius of atom B
XA = Electronegativity of atom A
XB = Electronegativity of atom B
Note:
Bond length of pure single bond (BO = 1)= 1.54 Ao or 154 pm
therefore, CR = 154/2 = 77pm
Bond length of pure double bond (BO = 2) = 1.34 Ao or 134 pm
Therefore, CR = 134/2 = 68pm
Bond length of pure triple bond (BO = 3) = 1.2 Ao or 12 pm
therefore, CR = 120/2 = 60pm
Note:
BO inversely proportional to the Bond length
Vander Waal Radius (VR)
Vander Waal Radius (VR):
It is the half of distance between adjacent atom in 2 nearest neighboring molecular in solid state.
VR is not applicable for metal.
Its magnitude depends Upon packing of atoms/molecules in solid state.
Vender weal force attraction is a weak force of attraction hence, VR is smaller then CR.
For inert gases only VR is defined.
Metallic Radius (MR)
MR:
It is defined as half of the distance between two nearest atoms in closely packed metallic lattice.
Order (for some elements)
CR < MR< VR
Notes:
Trends in AR:
Along period:
L à R
Number of shells same and Zeff increase therefore size is decrease .
Along Group:
T à B
Zeff same and number of shells therefore size is increasing.
Exception in Atomic Size
Exception in atomic size:
3d Series:
Sc Ti V Cr Mn Fe Co Ni Cu Zn
For 3d series initially size decrease then for some elements size remains constant and finally at increasing.
Sc> Ti> Zn> V> Cr= Mn= Fe= Cu= > Co> Ni
Mnemonics:
Scoo Ti me Zn Vaise Braber Coi Ni
for Group-3
3d < 4d < 5d
But for Group-4 to Group-12
3d < 4d = 5d
Lanthanoid Contraction:
In f-block elements with increase in atomic number proton increase by 1 unit and that extra proton goes into the nucleus which increase nuclear charge but extra electron
e- goes into for subshell which has very poor shielding due which this increase nuclear charge attracts outermost e- towards itself.
Note:
Lanthanoid contraction (LC) is observed in lanthanoids as well as all elements appearing after lanthanoids.
Ce Pr Nd Pm Sm En Gd Tb Dy Ho Er Tm Yb Lu
Size decreasing continuously Z increasing by 1 unit but e- goes into f-subshell
Zeff increase, size decreasing
B-family:
B < Ga < Al < In < Tl
Note:
d-orbital has very less shielding due to their diffused nature hance Z-eff increasing implies that size decrease known as ‘Transition Contraction’.
Ionic Radius
the distance from the nucleus of an ion up to which it has an influence on its electron cloud. Ions are formed when an atom loses or gains electrons. When an atom loses an electron it forms a cation and when it gains an electron it becomes an anion.
Cation (Loss of electron) | Anion (Gain of electron) |
It contains +ve charge | It contains –ve charge |
During formation of cation Zeff increase so size is decrease i.e. cation is always smaller than parent atom. | Formation of anion Zeff decrease and size increase anion is always bigger than parent atom. |
rcation < rneutral < ranion
1.Size of anion increase on moving down the group.
Examples:
Li+ < Na+ < K+ < Rb+ < Cs+
Be+2 < Mg+2 < Ca+2 < Sr+2 < Ba+2
F- < Cl- < Br- < I-
2.For same elements, higher is the charge
+ve à size decrease
-ve à size increase
Due to Zeff.
Examples:
Mn+2 > Mn+4 > Mn+6 > Mn+7
O- < O-2
3.Isoelectronic species:
That species which have same number of electrons.
Examples:
O-2 F- N-3
e- 10 10 10
Z 8 9 7
Order of size:
O-2 < F- < N-3
Note:
for Isoelectronic species higher is the value of Z smaller will be the size.
Exceptions:
F- H- Cl- Br-
Correct order of ionic radius:
F- < Cl- < Br- < H-
H- O-2 F-
Correct order of ionic radius:
F- < O-2 < H-
Li+ Be+2 Mg+2 Na+ Al+3
Correct order of ionic radius:
Be+2 < Mg+2 < Li+ < Na+ < Al+3
Ionization Energy (IE)
Ionization Energy (IE):
Energy required to removed the electron of outermost orbit to infinity from isolated gaseous atom is known as ionisation energy.
Note
ΔH is Heat of reaction
It is of two types:
1.Heat absorbed is called endothermic. i.e. ΔH > 0.
2.Heat released is called exothermic i.e. ΔH < 0.
Note:
Unit of IE:
eV/atom = KJ/mole = kcal/mole
eV à electron volt
1eV = 1.6 * 10-19 J
= 96 KJ/mole
Successive Ionisation energy:
M(g) à M+(g) + e- first IE = IE1
M+(g) à M+2(g) + e- 2nd IE = IE2
M+2(g) à M+3(g) + e- 3rd IE = IE3
Note:
IE3 > IE2 > IE1 for same elements.
After formation of cation it is difficult to removed e- because of increment in Zeff so, successive IE are higher.
Factor affecting:
1.Size of atom/ion:
IE is inversely proportional to the size of atom/ion.
2.Effective Nuclear Charge:
IE is directly proportional to the Zeff.
3.Penetration Power of subshell:
IE s > p > d > f
4.Electronic configuration:
Half and full filled orbitals are relatively more stable so, it is difficult to remove e- so, their IE will be high.
Half filled: p3, d5, f7
Full filled: p6, d10, f14
Trends ( periodicity in IE):
1.Period:
L à R
Zeff increases therefore size will be decreases so, IE will also be increases.
2.Group:
T à B
Zeff remains same but size will be increases due to the number of shells increases top to bottom so, IE will be decreases.
Note:
For same elements electronic configuration is not considered whether it is half or full filled only charge is considered.
Note:
For cations higher is the charge more will be the ionisation energy.
(IE)cation ∝ charge on ion
Note:
For Anions higher is the charge smaller will be the ionisation energy.
(IE)Anion ∝ 1/ Charge on ion
Exception in IE:
1.Boron family:
IE: B > Tl > Ga > Al > In
2.Carbon family:
IE: C > Si > Ge > Pb > Sn
3.d-block:
G-3: 5d < 4d < 3d
G-4,5,6,10 : 3d < 4d < 5d
G-7,8,9,11,12: 4d < 3d < 5d
Application of IE:
1.Basic Nature/ Metallic character/ Oxidising tendency of elements:
BN/ MC/ OT ∝ 1/ IE o f element
Along the period L to R.
IE increases, MC decreases, OT decreases and BN decreases
Metallic Compound: Basics nature and electron losing tendency.
Non-metallic Compound: Acidic nature and electron gaining tendency.
2.Basic nature of oxide:
Basic strength ∝ 1/ IE of element
along the period i.e. L to R:
IE increases, Basic nature of oxides decreases, and acidic nature increases.
along the group i.e. T to B:
IE decreases, Basics nature increases
Questions:
- Arrange the following compounds in increasing order of its basic strength.
Na2O, K2O, Cs2O
- Arrange the following compounds in increasing order of its acidic strength.
Li2O, CO2, N2O5, BeO, B2O3.
Electron Affinity (EA) or Electron gain enthalpy
Energy released when an e- is added to outer most orbit of an isolated gaseous atom is known as EA and enthalpy change in the process is known as e- gain enthalpy.
Note:
- EA1 is normally +ve because for addition of e- energy is released but EA2 is always –ve because energy should be given to over come e- – e- repulsion.
2.In the formation of polyvalent anion energy is always observed.
3e- + N(g) à N-3(g)
2e- + S(g) à S-2(g)
Factor Affecting of EA:
1.Size of atom/ion:
EA ∝ 1/Size of atom/ion
2.Zeff :
EA ∝ Zeff
3.Electronic configuration of element:
elements having half filled and fully filled configuration will be reluctant to take up extra e- because they will that’s why their EA value is very low (0 or -ve).
G-2 G-12 G-15 G-18
ns2 (n-1)d10ns2 ns2np3 ns2np6
Trends in EA:
1.Period:
L à R
Zeff increases, size decreases therefore, EA increases
2.Group:
T à B
Zeff remains same, Size increases therefore, EA decreases
Question:
Which is more stable?
Li- or Be-
Question:
Removing of electron is easy from Li- or Be-?
Exception in EA:
1.Noble gas (Inert gas):
for noble gas value of EA is –ve or very less because extra e- goes into next higher shell which becomes unstable due to its higher energy.
For G-14, 15, 16 and 17 value of EA for 2nd period element is smaller than 3rd period elements because in P-2 element e- goes into 2nd shell which is relatively small so e-
Will significantly repulsion but in P-3 element electron goes into 3rd shell which is larger in size so extra e- is easily accommodated.
Note:
Minimum IE à most reactive metal
à Cs (Cation)
Maximum EA à Most reactive non-metal
à Cl (anion)
Maximum IE and Minimum EA à Most unreactive Non-metal
à Inert gases
Note:
Order of EA:
Cl > F > Br > I
Order of EA:
S > Se > Te > O
G-17: lowest EA i.e. Iodine(I) > G-16: Highest EA i.e. Sulphur(S)
Electronegativity(EN)
Tendency of any atom to attract shared pair of e- towards itself in a covalently bonded molecules is known as ‘EN’.
Value of EN is depends upon value of EA and IE of the elements and higher is the value of EA and IE.
Difference between EA and EN:
EA | EN |
It is property of an isolated gaseous atom. | It is property of an atom bonded by covalent bond. |
It is an absolute value. | It is an relative value |
Unit: eV/atom or KJ/mole or Kcal/mole | It is unitless |
Factor affecting EN of element:
1.Size of atom:
EN ∝ 1/size of atom
2.Zeff:
EN ∝ Zeff
3.Charge on cation:
Higher is the charge on cation more will be Zeff i.e. EN increases.
EN ∝ charge on cation
4.Charge on anion:
Higher is the charge on anion, smaller will be Zeff i.e. EN decreases.
EN ∝ 1/Charge on anion
5.Hybridization of elements:
Higher is the %age s-character of element in the given compound more closes will be the orbital to the nucleus and more will be the attraction for the nucleus on the
for the nucleus on the e- Hence higher will be EN.
Note:
% s-character | sp | sp2 | Sp3 |
½*100 = 50% | 1/3*100 = 33.33% | ¼*100 = 25% |
Trends in EN (Periodicity):
along the period i.e. L à R, Zeff increases, size decreases, EN increases
along the group i.e. T à B, Zeff remains same, size will be increases due to number shells increases, EN decreases
Some Important EN
1 | 2 | 3 | 4 | 5 | 6 | 7 | 8 | 9 | 10 | 11 | 12 | 13 | 14 | 15 | 16 | 17 | 18 | |
1 | H(2.1) | C(2.5) | N(3) | O(3.5) | F(4) | |||||||||||||
2 | P(2.1) | S(2.5) | Cl(3) | |||||||||||||||
3 | Br(2.8) | |||||||||||||||||
4 | I(2.5) | |||||||||||||||||
5 | ||||||||||||||||||
6 | ||||||||||||||||||
7 |
Note:
F> O > N = Cl > Br > I > = S = C > P = H
F O N a Cl Biyar I S C PaHal Waan
EN Scale:
There is no directed method to determine EN but by following we can calculate it easily:
1.Pauling Scale:
Linus Pauling developed a method to determine relative values of EN using following formula:
xA – xB = .208 √(EA-B – √(EA-A∗EB-B)) unit of Bond energy is Kcal/mol
Where,
𝑥𝐴 à EN of atom A
𝑥B à EN of atom B
𝐸𝐴−𝐵 à Bond energy of A-B
𝐸𝐴−A à Bond energy of A-A
𝐸B−𝐵 à Bond energy of B-B
Note:
1 Cal = 4.2 Joule
xA – xB = .1017 √(EA-B – √(EA-A∗EB-B)) unit of Bond energy is KJ/mol
2.Mullikan Scale:
According to Mullikan EN is average of IE and EA.
i.e.
EN = (IE + EA)/2
remember EA and IE is in eV/atom
Application of EN:
1.Nomenclature of compound:
Covalent bond is formed by non-metal and compound containing two non-mental are known as Binary Compound. Non-metal haing more EN is written with
suffix ‘ide’ and name of less EN atom is written first.
Examples:
ICl à Iodine chloride
HCl à Hydrogen chloride
etc.
2.%age of Ionic Character:
Neither any bond is 100% ionic nor 100% covalent each bond has some %age of ionic as well as covalent character.
%age of ionic character = 16|Δx| + 3.5|Δx|2
Where,
Δx à | xA -xB|
3.Strength of Bond:
Higher is the EN difference more will be the strength of bond.
Examples:
HF > HCl > HBr > HI
EN difference: 1.9 .9 .7 .4
4.Metallic and Non-metallic character:
EN à low (Metallic (Basic))
EN à high (Non-Metallic (Acidic))
5.Nature of AOH type of Compound:
AOH à A+ + OH-
Base
AOH à AO- + H+
àOH with metal or polyatomic cation è Base
Examples:
NaOH, KOH, Ca(OH)2, NH4OH etc.
àOH with non-metal è Acid
Examples:
HNO2 or NO(OH) Nitrous acid
H2SO4 or SO2 (OH)2 Sulphuric acid
H3BO3 or B(OH)3 Boric acid
Nature of Oxide
Nature of Oxide:
Oxide which reacts with bases and this on dissolving in water forms oxyacid.
1.Oxy-acid: Non-metal + Oxygen + Hydrogen
2.Hydra –acid: Non-metal + Hydrogen
1.Oxy-acid:
H2SO4, HCl etc.
Non-metallic Oxide are acidic in nature.
Examples:
CO2, SO2, P4O6, N2O4 etc.
Note:
CO2 + 2NaOH à Na2CO3 + H2O
acid Base
OS= +4 OS= +4
of C-atom of C-atom
And
CO2 + H2O à H2CO3
OS = +4 OS = +4
Basic Oxide:
Oxides which reacts with acid and on dissolving in water it form hydroxide.
Examples:
Na2O, K2O, Fe2O3 etc.
Metallic Oxides are basic in nature.
Note:
Na2O + H2SO4 à Na2SO4 + H2O
Base Acid
And
Na2O + H2O à 2NaOH
Higher is the %age of Oxygen in the oxide more will be its acidic strength.
Question:
Arrange in increasing order of acidic strength?
i.N2O3, N2O5, N2O4
ii.SO2, SO3
iii.MnO, MnO2, MnO3, Mn2O7
iv.Cr2O3, CrO3
v.Na2O, CO2, CuO, N2O5, Li2O
Note:
MnO3 + H2O à H2MnO4 (Manganic acid)
Mn2O7 + H2O à 2HMnO4 (Permanganic acid)
On moving left to right in periodic table acidic strength of oxide increases and on moving top to bottom basic strength increases.
Neutral Oxide:
Oxides which neither reacts with acid nor reacts with base is known as ‘Neutral oxide’.
Examples:
N2O, NO, CO and H2O
Amphoteric Oxides:
Oxides which reacts with acid as well as base is known as ‘Amphoteric oxides’.
Examples:
Be Sn Ga Al Zn Pb
Be Sn Ga Al Znbaan Pb
Note:
ZnO + 2HCl à ZnCl2 + H2O
ZnO + 2NaOH à Na2ZnO2 + H2O
Sodium Zincate