Atomic radii decrease across a period as nuclear charge increases. Cations are smaller than their parent atoms. Among isoelectronic species, the one with the larger positive nuclear charge will have the smallest radius. Ionization energy generally increases across a period as it is more difficult to remove electrons, and decreases down a group as shielding increases. Electronegativity follows similar trends as ionization energy.

2. Atom is a basic unit of matter, that consists of a dense
central dense nucleus surrounded by a cloud of negatively
charged electrons.

3. Element is a substance in which all the atoms have
same mass and properties
For ex: H is an element in all atoms have properties
such as atomic number, atomic mass, electronic
configuration, etc….

4. ‘Sum of proton and neutron number in the nucleus of atom’

5. Total number protons in the nucleus of an atom

6. periodic table is a tabular display of chemical
elements, organized on the basis of their atomic
numbers, electronic configuration and properties
Period: elements repeat their chemical and physical
properties after a periodic intervals
Table: arrange in tabular form

7. In 1789 antoine lavoiser published a list of 33 elements, group them
into gases, metal and non metals
Dobereiner’s traids
In 1829 dobereiner classified elements into groups of three elements with
similar properties on the basis of atomic weight. Atomic weight of middle
element is roughly average of first and third elements, is known as law of
traids
Eg. Li(7), Na(23) & K(39)
Means of atomic masses:(7+39)/2=23
Similarly Cl(35.5), Br(80), I(127) & Ca(40), Sr(88), Ba(137)
Limitations: he could not arrange all the elements only such three traids
that have been mentioned.

8. Newland’s octaves(1864)
He arranged the elements in order of increasing atomic mass, the
properties of eighth element similar to the first as like in musical
note.
Sa re ga ma pa dha ni
Li Be B C N O F
Na Mg Al Si P S Cl
Limitations: this classification was successful up to the element
calcium

9. Mendeleef’s periodic table
Mendeleefs periodic table based on Mendeleefs periodic law which
state that ‘the physical and chemical properties of the elements are
a periodic function of their atomic masses.
At the time of Mendeleef, only 63 elements were known.
He classified the elements into seven horizontal rows called periods
and eighteen vertical columns called groups.

11. Limitations:
position of hydrogen atom: hydrogen atom has been placed in group
1A(alkali metal) but it resembles to halogens of group V11A. thus its
position in the Mendeleefs periodic table is controversial.
Position of isotopes: as Mendeleefs classification based on atomic
weight. Isotopes placed in different position due to their different
atomic weight.
Eg. There is three isotopes hydrogen which occupy the different
position
Anomalous position of some elements:
Eg. Co( atomic weight=58.9) is placed after Ni(atomic weight=58.7) in
case we consider the atomic weight

12. Modern periodic table(1913)
In 1913 Moseley modified Mendeleefs periodic law. He stated that
‘physical and chemical properties are the periodic function of their
atomic number’. That is when arrange the elements in increasing order of
atomic numbers, the properties of elements repeated after certain
regular intervals 2, 8, 8, 18, 18, 32. These number called magic numbers
and cause periodicity in properties due to repetition of similar electronic
configuration .
Periodicity: the repetition of similar properties after intervals is called
periodicity
Causes of periodicity: causes of periodicity in properties due to
repetition similar of electronic configuration
The periodic table divided into four main blocks s, p, d, & f

14. *Elements in which last electron enter the outermost s orbital.
*The elements of group1(alkali metal) and group2 (alkaline earth
metal) which has outermost electronic configuration is ns1-2.

15. Properties:
*They are all reactive metal so they never found in pure form
with low ionization enthalpy
*Metallic character and reactive increases from top to bottom
*They are predominantly form ionic bond
*They are soft metal with low melting and boiling point
*Most of the metal import specific color to flame( Na-orange, Li-
red etc…)
*They are strong reducing agent

16. *Elements in which last electron enter outermost p orbital
*p-block elements belongs to group 13-18
*The outermost electronic configuration is ns2 np1-6

17. Properties:
*It has metal, non metals and metalloids.
*Ionization energy higher than s-block
*Most of them form covalent bond
*Some of p-block elements shows more than one oxidation
states(common oxidation states are +3, +5, -3)
*Elements of oxygen and halogen family have high electron
affinity
*Metallic character decreases from left to right and increases
from top to bottom

18. *The elements in which last electron enters d orbital
*The elements belongs to group 3-12
*The outermost electronic configuration is ns1-2 (n-1)d1-10
*They have inner incomplete shell so known as transition elements

19. Properties:
*They are all metal, hard and ductile with high
melting and boiling point
*They shows variable oxidation states
*They form both ionic and covalent bond
*They are generally colored and paramagnetic

20. *The last electron added to each element is filled in f- orbital
*The two rows of elements at the bottom of the Periodic Table, called
the Lanthanides, Ce(Z = 58) – Lu(Z = 71) and Actinoids, Th(Z = 90) –
Lr (Z = 103) are characterized by the outer electronic configuration
(n-2)f1-14 (n-1)d0–1 ns2.

21. Properties:
*They are heavy metals with high melting and boiling point
*They shows variable oxidation states
*Compounds of these elements are generally colored
*They have high tendency to form complex
*Actinoid elements are radioactive, many of them are man made.
The elements after uranium(92) are called transuranium elements

22. On moving from left to right and top to bottom we
see the some variation
*Atomic radius
*Ionization Enthalpy
*Electron Gain Enthalpy
*Electronegativity

23. *The atomic radius of a chemical element is a measure of the size of its
atoms, it is distance from the center of the nucleus to the boundary of the
surrounding cloud of electrons.
* Calculating atomic radius is difficult because the size of an atom (~ 1.2 Å
i.e., 1.2 × 10–10 m in radius) is very small, also electron cloud surrounding
the atom does not have a sharp boundary.
*Atomic Radius is refer to covalent , ionic, metallic radius or van der
Waals radius.
*Atomic radii can be measured by X-ray or other spectroscopic methods.

24. Covalent radius:
The covalent radius refers to the size of atom, it is the half
of the distance between two atoms when they are bound
together by a single bond in a covalent molecule.
Ex. The bond distance in the chlorine molecule (Cl2) is 198
pm and half this distance (99 pm), is taken as the atomic
radius of chlorine.

25. Metallic radius:
Metallic radius is defined as half of the inter nuclear
distance between the two adjacent metal atoms in the
metallic crystal.
Ex. The distance between two adjacent copper atoms in solid
copper is 256 pm; hence the metallic radius of copper is
assigned a value of 128 pm.

26. Ionic radius:
*The ionic radius is the distance between the nucleus and the
electron in the outermost shell of an ion.
*The removal of an electron from an atom results in the formation
of a cation, whereas gain of an electron leads to an anion.
*A cation is smaller than its parent atom because it has fewer
electrons while its nuclear charge remains the same.
For example, The atomic radius of sodium is 186 pm compared
to the ionic radius of 95 pm for Na+.

27. *The size of an anion will be larger than that of the parent
atom because the addition of one or more electrons would
result in increased repulsion among the electrons and a
decrease in effective nuclear charge.
For example, the ionic radius of fluoride ion (F– ) is 136 pm
whereas the atomic radius of fluorine is only 64 pm.

28. *Atoms and ions which contain the same number of electrons, are
called isoelectronic species.
For example, O2–, F–, Ne, Na+ and Mg2+ have the same number
of electrons (10). Their radii would be different because of their
different nuclear charges.

30. *The atomic size generally decreases across a period, It is
because within the period the outer electrons are in the
same valence shell and the effective nuclear charge
increases as the atomic number increases resulting in the
increased attraction of electrons to the nucleus.
Ex. Li(1s2,2s1), Be(1s2,2s2), B(1s2,2s2 2p1), C(1s2,2s2 2p2), N(1s2,2s2
2p3), O(1s2,2s2 2p4), F(1s2,2s2,2p5), Ne(1s2,2s2,2p6) i.e ns2,np1-6
*Within a vertical column of the periodic table, the atomic
radius increases regularly with atomic number. It is because
the principal quantum number (n) increases and the
valence electrons are farther from the nucleus. This
happens because the inner energy levels are filled with
electrons, which serve to shield the outer electrons from
the pull of the nucleus. Consequently the size of the atom
increases.

32. Solution:
Atomic radii decrease across a period. Cations
are smaller than their parent atoms. Among
isoelectronic species, the one with the larger
positive nuclear charge will have a smaller
radius. Hence the largest species is Mg; the
smallest one is Al3+.

33. The energy required to remove an electron from an isolated gaseous atom (X)
in its ground state is called ionization enthalpy.
the ionization enthalpy is expressed in units of kj mol-1
the first ionization enthalpy for an element X is the enthalpy change for the
reaction depicted in equation
X(g) → X+
(g) + e–
Mg(g) → Mg+
(g) + e–
We can define the second ionization enthalpy as the energy required to remove
the second most loosely bound electron
X+
(g) → X2+
(g) + e–
Mg+
(g) → Mg2+
(g) + e–

34. *Energy is always required to remove electrons from an atom and
hence ionization enthalpies are always positive.
*The second ionization enthalpy will be higher than the first
ionization enthalpy because it is more difficult to remove an
electron from a positively charged ion than from a neutral atom.
*ionization enthalpy and atomic radius are closely related
properties.

36. *The first ionization enthalpy generally increases as we go across a
period ( from Li-F) because successive electrons are added to
orbitals in the same principal quantum level.
* The shielding effect of inner shell electrons does not increase very
much, Consequently, ionization energy increases, so that the
outermost electrons are held more and more tightly towards
nucleus.
*Ionization enthalpy generally decreases from top to bottom
because successive electrons are added to orbitals in the different
principal quantum level.
*The shielding effect of inner shell electrons increase very much,
Consequently, ionization enthalpy decreases, so that the outermost
electrons are held less tightly towards nucleus.

39. *When an electron is added to a neutral gaseous atom (X) to convert
it into a negative ion, the enthalpy change occur in this process is
defined as the Electron Gain Enthalpy.
X(g) + e– → X–(g)
*Electron gain enthalpy becomes more negative with increase
in the atomic number across period. Electron gain enthalpy
becomes less negative from top to bottom. It depends on
nuclear charge and atomic radius.

40. *The ability of an atom in a chemical compound to
attract shared electrons to itself is called
electronegativity
*it is not a measureable quantity.
*Linus Pauling, assigned arbitrarily a value of 4.0 to
fluorine, the element considered to have the greatest
ability to attract electrons

42. *Electronegativity generally increases across a period
from left to right and decrease down a group in the
periodic table.
*Electronegativity is proportional to nuclear charge
and inversely proportional to atomic size
*The trend is similar to that of ionization energy.

44. *Chemical combination between atoms known as chemical bonding. In
which atoms can combine either by transfer of valence electrons
from one atom to another (gaining or losing) or by sharing of valence
electrons in order to have an octet in their valence shells.
Types of chemical bond:
*Covalent bond
*Ionic bond
*Metallic bond
*Coordinate bond
*Hydrogen bond

45. Covalent bond
*Bond formed by the sharing of electrons between nonmetallic
elements of similar electronegativity.
Examples; O2, CO2, C2H6, H2O, SiC
Nonpolar covalent bond:
when electrons are shared equally
Ex: H2, cl2
Polar covalent bond:
when electrons are shared but shared unequally
Ex:H2O, HCl, HF

46. Ionic bond
*Bond formed between two ions by the transfer of electrons from one
atom to another.
*Ionic compounds result when metals react with nonmetals
Ex: NaCl, MgCl2, KCl, K2O

48. Coordinate bond
*A co-ordinate bond (also called a dative covalent bond) is a covalent
bond (a shared pair of electrons) in which both electrons come from
the same atom.

49. Hydrogen bond
*Hydrogen bond is a weak bond formed between hydrogen atom and
an electronegative atom such as O, F, N
Intra molecular hydrogen bonding
Inter molecular hydrogen bonding