2. a table of the chemical elements
arranged in order of atomic number,
usually in rows, so that elements with
similar atomic structure (and hence
similar chemical properties) appear in
vertical columns..

3. Historical Background
 Al-Razi’s classification was based on physical
and chemical properties of substances.
 Dobereiner, a German Chemist in 1829,
arranged then known elements in groups called
Triads, each contained three elements with
similar properties.
 Newland, English Chemist in 1864 classified 62
elements, known at that time in increasing order
of their atomic masses. Every eighth element had
some properties in common with the first one.
This classification was based on principle, known
as “Law of Octaves”

4. Mendleev’s Periodic Table
 In 1871, a Russian Chemist, Dmitri Mendleev gave
more Comprehensive classification of elements.
 Elements are arranged in the periodic table in the
increasing order of their relative atomic masses.
 Mendeleev divided his periodic table in
eight groups and seven periods.
 Elements having similar chemical properties appeared
at regular intervals.
 He left some gaps in his periodic table for
undiscovered elements and predicted their properties.
 For example, he left the place of Germanium and
predicted its properties which was undiscovered at
that time, a few years later Ge was indeed discovered.

6. Improvements in Mendleev’s
Periodic Table
 Mendleev’s work had some drawbacks which were
eliminated by Modern Periodic Law.
 In 1911, Mosely discovered the Atomic Numbers.
 It was noticed that elements could be classified more
satisfactorily by using their Atomic numbers.
 So, the periodic table was improved by arranging the
elements on the basis of their increasing atomic
numbers instead of atomic masses.

7. Improvements in Mendleev’s
Periodic Table
 Modern Periodic Law states “if the elements are
arranged in ascending order of their atomic numbers,
their chemical properties repeat in a periodic manner”
 Another group was added in periodic table at extreme
right for Noble gases which were missing in
Mendleev’s periodic table.
 The elements like Be, Mg, Ca, Sr, Ba and Zn, Cd, Hg
were placed in single group while according to their
properties they belonged to two different categories.
 In modern periodic table, these elements have been
placed in two different groups A and B.

8. The modern periodic table
1.groups and periods
 Modern periodic table is divided in to groups and periods.
 Groups 08
 Groups are further divided into sub group A and B.
 Group A representative or normal elements
 Group B transition elements
 Periods 07
 Period 1 02 elements
 Period 2 & 3 08 elements
 Period 4 & 5 18 elements
 Period 6 & 7 long periods (contain lanthanides and
actinides)

10. 2.Some families in periodic table
 Group 1A Alkali metals because they
form alkalis with water.
 Group 2A Alkaline earth metals due to their
alkali nature and abundance in earth crust.
 Group 7A halogens due to their salt
forming properties.
 Group 8A noble gases
 These families are useful for quick recognition of
elements in periodic table.

11. 3.Blocks in periodic table
This
classificatio
n is based
on valence
orbital that
is involved
in bonding

12. 4.metals, non metals and
metalloids

13. Periodic trends in physical
properties
Some important points
 Group no tells about the number of electrons in outer
most shells.
 Period no tells about the total no of shells
So
Atomic size increases down the group due to increase
in number of shells

14. 2 shells
3 shells
4 shells
5 shells
6 shells

15. Atomic size decreases along
the period because no of
electrons and protons
increases but no of shells
remain same.

16. 1.Atomic size
‘the distance between the nucleus and outermost shell of an
atom is called atomic radius’
 Size of atom is so small that it can not be directly measured
so a technique is developed which can measure the distance
between the nucleus of two similar atoms. Half of this is
taken as atomic radius.
 TREND (as discussed earlier)
 Lanthanide contraction due to poor shielding effect
of d and f subshell, reduction in atomic radius is very
prominent in lanthanides. This is called lanthanide
reduction.

18. b) Ionic radius
Why size of cation is smaller than neutral atom?
 Firstly the removal of one or two electrons result in
loss of outer most shell so radius is reduced.
 Secondly electron proton ratio os imbalanced so due to
greater attraction of nuclear charge, size is reduced.
 E.g radius of Na is 157pm while radius of Na+ is 95pm

19. Why size of anion is greater than neutral atom?
 Addition of one or more electron in shell of neutral
atom increases replulsion between the electrons
causing expansion of shells.
 Electron proton ration is imbalanced.
 e.g size of F atom is 72pm while size of F- is 136pm.

20.  Graph
 TREND same as that of atomic radius

21. 2. Ionization energy
 “minimum amount of energy which is required to remove
an electron from the outer most shell of its isolated
gaseous atom in its ground state”
 E.g
Na Na+ + e_ I.E=496kj/mole
 First ionization energy is the energy required to remove
first electron.
Mg Mg + + e_ I.E(1)=738kj/mole
Mg + Mg + + + e_ I.E(2)=1451kj/mole

22. Why 2nd ionization energy is always higher than 1st
ionization energy?
 Because by removal of one electron nucleus attracts
the remaining electron more powerfully towards its
self
 M0re energy is required to remove electron from a
positive ion rather than a neutral atom.

23. Variation within group
 Factors effecting ionization energy
1) Magnitude of nuclear charge
2) Size of atom
3) Shielding effect (it is actually the repulsion due to
electron that are present between nucleus and
outermost shell)

24. As we move down the group
 Nuclear charge decreases
 No of shells increase hence size as well as shielding
effect increases
 So ionization energy decreases down the group.
 That is why it is easier to remove electron form cesium
atom than lithium.

25. Along the period
 Generally, smaller the atom with greater nuclear
charge, more strongly the electrons are bound to the
nucleus and hence higher the ionization energy of
atom.
 While moving along the period,
1. Nuclear charge increases
2. No of shells remain same
3. Size reduces
So ionization energy increases along the period and that
is why inert gases have highest ionization energies.

26. Electron Affinity
 The electron affinity is the amount of energy which is
released when an electron is added to a gaseous atom
to forma negative ion(anion).
F + e– → F– – ∆H = Affinity = 328 kJ/mol

27. Why addition of 2nd electron requires energy?
 Energy is released when first electron is added into an
atom.
O + e– → O– – ∆H = Affinity = –141 kJ/mol
 Energy is absorbed when second electron is added.
O + e– → O – – ∆H =Affinity= +780 kJ/mol
 It is due to the fact that when second electron is added
to uni-negative ion, it is repelled by previously added
electron so more energy is required to add second
electron

28. Trends in Periodic Table
 Trends- same as that of Ionization Energy

29. Metallic and Non-Metallic Character
Metals
 All the elements which have tendency to form positive
ions by loosing electrons are considered as metals.
 All metals are good conductor of heat and electricity.
 Metals form basic oxides, which gives bases when
dissolved in water.
Na2O + H2O → 2NaOH

30. Trends in Periodic Table
 Metallic character increase from top to bottom in a
group.
 Because size of atom increases down the group, so it
becomes easy to remove electron from atom.
 Metallic character decreases from left to right in a
period because of small size of atoms.
 The Halogens have least metallic character.

31. Non-Metals
 The elements which gain electrons and form negative
ions are called Non-Metals.
 All gases are non metals.
 Poor conductor of heat and electricity.
 Non-Metals form acidic oxides, when dissolved in
water oxides yield Acids.
SO3 + H2O → 2H2SO4

32. Trends
 Non-metallic character decreases with increase in
atomic size.
 Non-Metallic character decreases from top to bottom.
 Fluorine, Oxygen, Nitrogen are the most non metallic
elements of their respective groups.

33. Melting and Boiling Points
 Melting and Boiling points tell us how strong the
atoms are molecules in them are bound together.
Variation in a Period
 Across the short periods, the M & B.P increases with
increase on valence electrons upto group IVA and then
decrease upto noble gases.
 M.P of IIA is higher than IA because it provides two
electrons to form bond.
 Graph

34. Graph of Variation

35.  Carbon has the maximum number of binding
electrons, so it have very high m.p in diamond.
 The element with giant covalent structures have very
high m.p.
 As we move to VA, VIA,VIIA the lighter elements have
small, covalent molecules and have very weak
Intermolecular forces, so they have low M & B.P.

36. Variation in Groups
 The melting and boiling points of IA and IIA decrease
from top to bottom due to increase in atomic size.
 Because the binding forces become weak in large
atoms.
 For group VIIA, the M& B.P increases down the group,
because large molecules exerts stronger force of
attraction.
 Reference, Factors affecting London dispersion forces,
1st year chemistry.
 Graph

37. Graph of Variation

38. Oxidation State
“Oxidation state of an atom in a compound is
defined as the charge(with the sign), which it
would carry in the compound”.
 In ionic compounds, O.S is usually the number of
electrons gained or lost by the atom.
 In NaCl, Na has +1 O.S while Cl has -1 O.S.
 In covalent compounds, oxidation state depends on
the relative electronegativities.
 In SnCl4, Sn has +4 O.S, while each Cl has -1 O.S.
 The oxidation state of an element is zero in its free
state.

39.  The O.S also relates with group number of elements.
 The elements of group IA to IVA have the same O.S as
their group numbers are.
 For VA and VIA, O.S depends on valence electrons or
the number of vacancies available in these shells.
 For example, N, P, As and Sb shows -3 as well as +5 O.S
 In H2SO4, S shows +6 O.S, in H2S, S shows -2 O.S
 In VIIA, The elements mostly show -1 O.S
 Noble gases show zero Oxidation state,

40.  Transition elements, shown as B sub-groups also show
the O.S equal to their group number as Cu(I), Zn(II),
Cr(VI), Mn(VII).
 Due to greater number of valence electrons, Transition
elements usually show more than one oxidation state.

41. Electrical Conductance
 Property of metals
 Due to the presence of loose electrons in valence shells
and ease of their movement in solid lattices.
 The electrical conductance generally increases down
the group in IA and IIA.
 Trend is not same for all groups, there are variations.
 Metals of IB, Knows as Coinage metals(Gold, Silver,
Copper) have very high values of E.C.
 The elements of VIA and VIIA have very low E.C as
they can be considered as non-conductors.

42.  There’s no general trend of E.C in Transition metals.
 Carbon in the form of Diamond is non-conductor as
all four available electrons are tightly bonded.
 Carbon in the form of Graphite is conductor as one of
the four valence electrons is free to move.
 The lower elements of IVA(Tin, Lead) are good
conductors.

43. Hydration Energy
“The hydration energy is the heat evolved or absorbed
when one mole of gaseous ions dissolve in water to
give an infinitely dilute solution.”
 Hydration energy highly depend on charge to size ratio
of ions.
 In IA, charge to size ratio decreases down the group so
does the hydration energy.

44.  Hydration energy increases from left to right as the
charge to size ratio increases.

45. Halides
“binary compounds which halogen form with other
compounds are called halides”
 Physical properties of halides are determined by nature of
bonding.
 (a)ionic halides strongly electropositive
elementshaving greater electronegativity difference with
halogens form ionic halides.
 E.g halides of group 1 A (NaCl,NaF).
 They have 3 D lattice,high M.P and B.P.
 Among them, florides have highest lattice energy due to its
small size.so flourides have highest M and B.P which
decrease in order
Flourides>chlorides>bromides>iodides

46.  (b)covalent halides SiCl4(highly
polar),PCl3,S2Cl2(less polar)
 (c)polymeric halides when a halogen atom
acts as a bridge between two other atoms.
Less electropositive elements like Be, Ga and Al form
these halides. They have partly ionic bonding with
layers.

47. On moving from left to right, electronegativity
difference decreases and thr trend shifts towards
covalent halides.
As cavalent halides have weak van der Waal’s forces so
they are mostly gases,liquids or low melting solids.
Physical properties of covalent halides are influenced
by the size and polarizabilty of halogen atoms so
iodides which have largest size and most polarizable
have strong van der waal forces hence have highest M
and B.P.

48.  Variation in bonding character is also available in
halogen group.
Flouride >chloride >bromide >iodide
AlF3 is purely ionic. Have 1290 degree centigrade
m.p,good conductor
AlI3 is mainly covalent, m.p 198 degree centigrade
,electrically non conductor.

49.  In case of an element forming more than one halides,
metal halide in lower o.s tends to be ionic while in
higher o.s tends to be covalent.
 e.g PbCl2 and PbCl4
+2 +4 Pb4+ has high
polarizing power.

50. Hydrides
 Binary compounds of Hydrogen with other
elements.
 Hydrides can be classified into three types:
a) Ionic
b) Covalent
c) Intermediate

51. Ionic Hydrides
 The elements of IA and heavier members of IIA(Ca,
Sr, Ba) forms ionic hydrides.
 Ionic hydrides are crystalline solids with High M.P &
B.P
 Can conduct electricity in molten(solution) form.
 The tendency toward covalent hydrides increases by
moving from left to right.
 e,g. LiH, NaH, CaH2

52. Intermediate Hydrides
Hydrides of Be and Mg are considered as
Intermediate hydrides.
The properties of intermediates are in between ionic
and covalent hydrides.
They have polymeric structures and covalent in nature.
e,g. BeH2, MgH2, CdH2

53. Covalent Hydrides
 Covalent hydrides are usually gases or volatile
liquids.
Non-conductors
Dissolve in organic solvents. (Like dissolve like)
Bond energies depends upon the size of atom or
electronegativity difference.
Stability of covalent hydrides increases from left to
right and decreases from top to bottom.
Fluorine forms most stable hydrides while Lead,
Bismuth forms least stable.

54.  These hydrides are formed by elements with
electronegative values greater than 1.8(Pauling scale).
 E.N of H is 2.1, so most of the hydrides are polar
covalent bonds as H carries slightly positive charge.
 Left to right E.N difference increases between H and
other bonding atom so bond becomes more polar.
 Due to high polarity, H2O and HF are capable of
forming H-Bonding.
 Down the group, B.P Increases down the group due to
high polarizibility.
 NH3, H2O, HF shows exceptional behavior due to H-
Bond.

55. Melting and Boiling Points of
Hydrides of IVA and VIA

56. Oxides
 “Compounds of Oxygen are called Oxides”
 Many of these have unusual properties
 Oxides are classified into two ways:
a) Based upon the type of bonding
b) Acidic or basic nature
 Metallic oxides are basic in nature while non-
metallic are acidic.
 Metallic and non-metallic oxides reacts with each
other to give salts. For example,
 Third type of oxides are amphoteric.

57. Classification of Oxides
 Oxides of IA ad IIA are basic except Be containing O2-
 O2- can’t exist independently in Water, it takes
proton immediately and forms OH-
 Oxides of non-metallic elements are acidic in
nature and dissolve in water to give acidic solution
 Oxides of less electropositive elements(Al2O3,
BeO, ZnO) are amphoteric in nature.
 They acts as acids with base and as base with acids.

58.  In a period, Oxides progress from:
 Strongly basic weakly basic Amphoteric Weakly
acidic Strongly acidic
 e,g. Na2O, MgO, Al2O3, P4O10
 Basicity of oxides Increases down the group
 BeO<MgO<CaO<SrO<BaO
 This trend is reverse in transition metal oxides
 Oxidation state also affects the nature of oxide
 The acidity increases with increase in O.S
 MnO<Mn2O3<MnO2<Mn2O7

59. Classification of oxides

60. Position of Hydrogen
 Hydrogen is not a metal, but it is still placed at the top
of IA
Similarities oh Hydrogen with Alkali Metals
i. Like alkali metals, H has one electron in its s-
subshell, which it can loose to form H+
ii. Both have strong tendency to combine with
electronegative elements such as halogens
iii. Both form ionic compounds, which dissociates in
H2O

61. Dissimilarities between Hydrogen
and Alkali metals:
i. Hydrogen is non metal
ii. H does not loose its electron as easily as IA metals do
iii. H exists in open environment, IA metals doesn’t
Similarities of Hydrogen with Halogens
i. H is gas like most of the halogens
ii. H is stable in diatomic form such as F2, Cl2, Br2
iii. Such as halogens H needs only one electron to
complete its valence shell and form H-
iv. Hydrogen and halogens both forms stable ionic
compounds with alkali metals

62. Dissimilarities between Hydrogen and
Halogens
a) Hydrogen can form positive ion by loosing its
electron but Halogens do not form positive ions
b) Hydrogen can form stable oxide by combining with
Oxygen while halogens lack this property
Similarities of Hydrogen with IVA Elements
a) Both have half filled valence shell
b) Both can combine with other elements through
covalent bond
c) Both have reducing properties

63. Dissimilarities of Hydrogen with
IVA Elements
 Member of IVA, C and Si can form long chains by
catenation while H doesn’t have this property.
 C can form multiple bonds at the same time while H
can form only one bond at the same time
 H is unique, its properties do not match exactly with
any of the group in periodic table
 But due to partial resemblance with alkali metals and
monovalent nature, Hydrogen is placed at the top of
Alkali metals(IA).