03_Periodic Table for engineering students.pdf

The electron configuration of an element describes how electrons are distributed in its
atomic orbitals. Electron configurations of atoms follow a standard notation in which all
electron-containing atomic subshells (with the number of electrons they hold written in
superscript) are placed in a sequence. For example, the electron configuration of sodium
is 1s2
2s2
2p6
3s1
.
The Pauli exclusion principle states that no two electrons in an atom can have the same
four quantum numbers.
Consider an atom has two electron. If one electron has the quantum numbers of n = 1, l =
0, m = 0, and s = +1/2, then, other electron must has n = 1, l = 0, m = 0, and s = -1/2.
Each subshell holds a maximum of twice as many electrons as the number of orbitals in
the subshell. Thus, a 2p subshell, which has three orbitals (with l = -1, 0, and 1), can hold
a maximum of six electrons. The maximum number of electrons in various subshells is
given in the following table.
Electron Configuration
Periodic
Table

Subshell Number of orbitals Maximum number of electron
s (l=0) 1 2
p (l=1) 3 6
d (l=2) 5 10
f (l=3) 7 14

Building up principle’ or Aufbau Principle
Aufbau (German aufbauen, “to build up”) principle tells us that electrons fill the orbitals
in the order of increasing their energy level. In the ground state of an atom, the electrons
tend to occupy the available orbitals in the increasing order of energies, the orbitals of
lower energy being filled first.
The energy of an orbital is determined by the sum of principal quantum number (n) and
the azimuthal quantum number (l). This rule is called (n + l) rule. There are two parts of
this rule :
(a) The orbitals with the lower value of (n + l) has lower energy than the orbitals of
higher (n +l) value. For example, let us compare the (n + l) value for 3d and 4s orbitals.
For 3d orbital n = 3, l = 2 and n + l = 5 and for 4s
Aufbau order of orbitals
orbital n = 4, l = 0 and n + l = 4. Therefore, 4s orbital is filled
before 3d orbital.
(b) When two orbitals have same (n + l) value, the orbital with
lower value of n has lower energy. Similarly, for 4p and 5s orbitals,
the (n + l) values are (4 + 1) and (5 + 0) respectively. In this case 4p
orbital has lesser value of n and hence it has lower energy than 5d
orbital and is filled first.

Building up principle’ or Aufbau Principle

Hund's rule states that:
(1) Every orbital in a sublevel is singly occupied before any orbital is doubly
occupied.
(2) All of the electrons in singly occupied orbitals have the same spin (to
maximize total spin).
Hund's rule
Consider the correct electron configuration of the nitrogen (Z = 7) atom:
The p orbitals are half-filled; there are three electrons and three p orbitals. This is
because the three electrons in the 2p subshell will fill all the empty orbitals first
before pairing with electrons in them
consider oxygen (Z = 8) atom, the element after
nitrogen in the same period; its electron
configuration is: 1s2
2s2
2p4
Oxygen has one more electron than nitrogen; as the
orbitals are all half-filled, the new electron must
pair up.
1s2
2s2
2p3

Dobereiner's Triads:
In 1829, J.W. Dobereiner, a German chemist made groups of three
elements each and called them triads. All three elements of a triad
were similar in their physical and chemical properties. He proposed
a law known as Dobereiner's law of triads.
According to this law, when elements are arranged in order of
increasing atomic mass, the atomic mass of the middle element was
nearly equal to the arithmetic mean of the other two and its
properties were intermediate between those of the other two.
Element Atomic mass
Li 6.9
Na 23
K 39
Periodic Table
A tabular arrangement of elements in rows and columns, highlighting
the regular repetition of properties of the elements, is called a periodic
table.
Johann
Wolfgang
Dobereiner
In the triad of Li, Na and K the atomic mass
of Na (23) is the mean of the atomic masses of
Li and K
6.9 + 39 = 45.9 ÷ 2 = 22.95

Features:
Only a few triads could be identified
System of triads could not continue

John Alexander
Reina Newlands
Newland’s Law of Octaves:
In the year 1864, the British chemist John Newlands attempted the 56 elements known
at that time. He arranged them in an ascending order based on their atomic masses
and observed that every 8th element had similar properties. On the basis of this
observation, Newland’s law of octaves was formulated.
The law of octaves states that every eighth element has similar properties when the
elements are arranged in the increasing order of their atomic masses. An illustration
detailing the elements holding similar properties as per Newland’s law of octaves is
provided below.

Features:
Newlands arranged them in the increasing order of their atomic masses. Every
eighth element had properties similar to the first.
Out of the 56 elements, Newland’s law of octaves held true only for elements up to
calcium. Elements with greater atomic masses could not be accommodated into
octaves.
After Ca every eighth element did not possess properties similar to the first.
Several elements were fit into the same slots in Newland’s periodic classification.
For example, cobalt and nickel were placed in the same slot.
Elements with dissimilar properties were grouped together. For example, the
halogens were grouped with some metals such as cobalt, nickel and platinum.
Inert (noble) gases were not included because they were not discovered. Fifty-six
elements were discovered.

Dmitri Ivanovich
Mendeleev
Mendeleev’s Periodic Table
Dmitri Ivanovich Mendeléev, a Russian chemist, was the most important contributor to
the early development of the periodic table. Many periodic tables were made but the
most important one was the Mendeleev periodic table.
In Mendeleev’s periodic table, elements were arranged on the basis of the fundamental
property, atomic mass, and chemical properties. During Mendeleev’s work, only 63
elements were known. After studying the properties of every element, Mendeleev
found that the properties of elements were related to atomic mass in a periodic way.
He arranged the elements such that elements with similar properties fell into the same
vertical columns of the periodic table.
He believed that atomic mass was the most fundamental property in classifying the
elements and examined the relationship between the atomic masses of elements and
their physical and chemical properties.

Mendeleev’s Periodic Table
Mendeleev’s Periodic
Law
The physical and
chemical properties of
elements are a periodic
function of their atomic
masses.
Features
Periods
Horizontal rows, numbered 1 to 7
Properties of elements in a period show regular gradation from left to right
Groups
Vertical columns, numbered I to VIII. I to VII are further divided into A and B
subgroups

Merits of Mendeleev Periodic Table:
Classification of all known elements: In this periodic table, elements are classified
into groups with similar properties helping to study of the properties of elements.
Position of iodine and tellurium: Iodine (127) is placed just after Tellurium (128)
to give more importance to physical and chemical properties than the atomic mass.
It is to be mentioned here that moving horizontally across the Mendeleev’s periodic
table it is seen that Te is in a group which includes O, S, Se. All these elements have
similar properties. Similarly, I is placed in a group which includes F, Cl, Br. All
these elements resemble to each other chemically.
Prediction of new elements: Mendeleev left some vacant spaces for elements yet
to be discovered in his periodic table. He named some elements such as eka-boron,
eke-aluminium, eka-silicon with respective atomic masses 44, 68 and 72
respectively. Later all these three elements have been discovered known as Sc, Ga
and Ge. The properties of these elements are strikingly similar to those predicted by
Mendeleev.
Noble gases were discovered later and placed in the table without disturbing the
position of other elements.

Demerits of Mendeleev Periodic Table:
Although Mendeleev’s periodic table create tremendous sensation to the chemists at
that time however, it suffers the following defects:
Position of hydrogen: H bears the properties resembles to both the alkali metals
and also halogens, so the position of hydrogen in his table is ambiguous.
Anomalies in the order of atomic mass: Increase in atomic mass was not regular
while moving from one element to another. For examples: Co with higher atomic
mass (58.93) is placed before Ni (58.71).
Placement of isotopes: Isotopes of same elements have different atomic masses.
Each of them should be given a different position. As isotopes are chemically
similar, they were given same position.
Position of some elements: Many dissimilar elements have been placed together such as
Cu, Ag, Au are grouped along with alkali metals. Mn is placed with halogens which totally
differ in the properties. Certain elements which possess similar properties are placed in
different group such as Cu & Hg, Ba & Pb etc.
Position of lanthanides and actinides: Among the lanthanides, only two elements (La &
Ce) were discovered at that time but no proper place are kept for these two elements in the
Mendeleev’s periodic table.

Modern Periodic Table
Atomic number is the most fundamental property of an element and
not its atomic mass – Henry Moseley.
Modern Periodic Law:
The modern periodic law states that the physical and chemical properties of the
elements are the periodic functions of their atomic numbers.
Scientists arranged elements in increasing order of their atomic numbers from left to
right across each row. And discovered that the elements having similar properties
repeat after regular intervals.
Why atomic number and not atomic mass?
Atomic mass is the total mass of the protons and neutrons present in a nucleus of an
atom. Whereas, the atomic number is the number of protons in a nucleus. Also, the
number of protons in the nucleus is equal to the electrons present outside the nucleus.
We know that the nucleus is deep-seated inside an atom. But the electrons outside it,
especially the ones in the outermost shell, are free to move around. Hence they take
part in chemical reactions. For this reason, the properties of an element depend on the
atomic number rather than the atomic mass.

Long Form of Periodic Table
The modern or long form of the periodic table is based on the modern periodic law.
The table is the arrangement of elements in increasing order of their atomic numbers.
The modern periodic table is the present form of the periodic table. And it consists of
18 vertical columns and 7 horizontal rows.
Groups in the Modern Periodic Table:
Groups are the vertical columns in the modern or long form of the periodic
table.
There are 18 groups in the periodic table.
These groups are numbered from 1 to 18.
Each group consists of elements having the same outer shell electronic
configuration.
Periods in the Modern Periodic Table:
Periods are the horizontal rows in the modern or long form of the periodic table.
There are 7 periods in the periodic table.
These are numbered as 1, 2, 3, 4, 5, 6 and 7 from top to bottom.
The 1st period consists of only two elements – Hydrogen and Helium.
While the 2nd and 3rd period consists of 8 elements each.
The 4th and 5th period consists of 18 elements each

On the other hand, the 6th period consists of 32 elements.
The 7th period of the periodic table now has four new elements. They are
113-Nihonium, 115-Moscovium, 117-Tennessine, and 118 –Oganesson. This
addition has completed the 7th period with 32 elements.
Also, the long form of the periodic table consists of a separate panel at the bottom.
It consists of 14 elements of the 6th period called the lanthanoids. And 14 elements
in the 7th period called the actinoids.
Each period represents the number of shells or energy levels present in an atom of
an element.

Period no. Size Elements
1 shortest 2
2 short 8
3 short 8
4 long 18
5 long 18
6 longest 32
7 incomplete see box
*IUPAC announces the verification of the discoveries of four new chemical elements: The 7th period of the
periodic table of elements is complete.
Update 21 Jan 2016: Technical Reports available
http://www.iupac.org/news/news-detail/article/discovery-and-assignment-of-elements-with-atomic-numbers-113
-115-117-and-118.html
*Temporary working
names and symbols
113 ununtrium, Uut
115 ununpentium, Uup
117 ununseptium, Uus
118 ununoctium, Uuo

Groups (number of valence electrons)
Vertical columns, Eighteen, numbered 1-18
Elements in the same group have same number of valence electrons / same outer
electronic configuration, shows same chemical properties.
Group 1, alkali metals
Group 2, alkaline earth metals
Group 17, halogens
Group 18, inert/ noble gases
Metals – left hand side
Non-metals – right hand side
Normal elements – Groups 1, 2 and Groups 13-17.
One outermost shell incomplete
Transition elements – Groups 3-12.
Two outermost shells incomplete
Inert gases – Outermost shell contains 8 electrons

Group
no.
1 2 3-4-5-6-7-8-9-10-
11-12
13-14-15-1
6
17 18
Type Alkali
metals
Alkaline
earth
metals
Transition
elements
Non-metal
s,
metalloids,
metals
Halogens Inert or
noble
gases
Normal elements Normal elements
Inner transition – at the bottom, contain two series, viz. lanthanides, actinides
Lanthanides (Ce – Lu) – 14 elements, atomic numbers 58-71. Placed along with La
(57), Group 3, Period 6. Close resemblance in properties to La
Actinides (Th – Lr) – 14 elements, atomic numbers 90-103. Placed along with Ac
(89), Group 3, Period 7. Close resemblance in properties to Ac
Group 3
Period 6 Lanthanides 14 elements
Period 7 Actinides 14 elements

Elements can be classified into four categories according to their
electron configurations.
s - Block elements:
Elements in which the last electron enters the s- orbital or their respective outermost
shells are called s- block elements. These are present in the left part of the periodic table
in group IA & IIA i.e. 1 & 2 group in modern periodic table. Electronic configuration of
valence shell is ns1-2
(n =1 to 7)
Characteristic of s – block elements:
•They are soft metal (except H & He, non metal, gas) with low melting & boiling points.
•They posses metallic character & reactivity of metal increases down the group.
•They are highly electropositive and having low ionization enthalpies except H & He.
•They have valency +1(in case of alkali metals) & +2 (in case of alkaline earth metals).
•Most of the metals of this block impart characteristic color to the flame.
•Except He, these are strong reducing agents & are good conductor of heat & electricity.

p − block elements:
These are present in right part of the periodic table and constitute the groups IIIA to
VIIA and zero groups except He i.e. group 13 to 18 of the modern periodic table. Most
of these elements are metalloids & non metals but some of them are metals also. The
last electron enters in p − orbital of valency shell and electronic configuration of valency
shell is ns2
np1−6
(n=2 to 7) For example: Silicon: 1s2
2s2
2p6
3s2
3p2
Characteristics of p − block elements:
•It contains both metals & non metals. The metallic character decrease from left to right
along the period and metallic character increases from top to bottom within a group.
•These elements show variable oxidation state.
•Most of these elements are highly electro-negative.
•They mostly form covalent compounds as well as ionic compounds.
•Ionization energy is higher as compared to s – block elements.
•Reducing character increases from top to bottom in a group & oxidizing character
increases left to right in a period.

d - Block elements:
These are present in the middle part of the periodic table (between s & p block element)
and constitute IIIB to VIIB, VIII, IB & IIB i.e. 3 to 12 groups of the modern periodic
table. The outermost electronic configuration is (n-1) d1-10
ns1-2
(n=4 to 7). There are
four series of d-block elements, which are-
3d series − Sc (21) to Zn (30)
4d series – Y (39) to Cd (48)
5d series –La (57), Hf (72) to Hg(80)
6d series- Ac (89), Rf (104) to Cn (112)
Some elements of d-block have special characteristics properties and are called
transition metals. These are metallic elements in which the outermost s sublevel and
nearby d sublevel contains electrons.
The transition elements, called Group B elements, are characterized by addition of
electron to the d orbitals.
Scandium: 1s2
2s2
2p6
3s2
3p6
4s2
3d1
Yttrium: 1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p6
5s2
4d1

Relationships and Differences between d-block elements and transition elements
Given below are some differences between d-block elements and transition elements :
1. d-block elements are chemical elements that have electrons in their d-orbitals. Whereas,
transition elements are chemical elements that have at least one stable cation that has partially
filled d-orbitals.
2. Colourful complexes can be formed by d-block elements or not. Colourful complexes are
formed by transition elements at all times.
3. Many d-block elements are diamagnetic, while some are paramagnetic or ferromagnetic. All
transition elements are paramagnetic or ferromagnetic.
4. Many d-block elements are not solids at room temperature (mercury is a liquid), but others
are, whereas all transition metals are solids at room temperature.
5. Several d-block elements exhibit multiple oxidation states, while others exhibit a single
oxidation state, and transition elements exhibit multiple oxidation states.

f - Block elements:
These are placed separately below the main periodic table and are mainly related to
IIIB i.e. group 3 of the periodic table. There are two series of f–block elements, which
are 4f series – Lanthanides: 14 elements Ce (58) to Lu (71) and
5f series - Actinides: 14 elements Th (90) to Lr (103)
These are known as inner transition metals. These are metallic elements in which the
outermost s sublevel and nearby f sublevel generally contain electrons.
The inner transition metals are characterized by the filling of f orbitals.
Cerium: [Xe]6s2
5d1
4f1
Thorium: [Rn]7s2
6d1
5f1

The noble gases. These are elements in which the outermost s and p sublevels are
filled.
The noble gases belong to Group 0. The elements in this group are sometimes called
the inert gases because they do not participate in many chemical reactions. The
electron configurations for the first four noble-gas elements are listed below. Notice
that these elements have filled outermost s and p sublevels.
Helium 1s2
Neon 1s2
2s2
2p6
Argon 1s2
2s2
2p6
3s2
3p6
Krypton 1s2
2s2
2p6
3s2
3p6
3d10
4s2
4p6
Notice that all of these elements
have filled outermost s and p
sublevels
Characteristics to f – block elements:
•They are heavy metals with high melting & boiling points.
•They show variable oxidation states & their compounds are generally coloured.
•Most of the elements of the actinide series are radioactive.

a) Finding Period of Elements:
Period of the element is equal to highest energy level of electrons or principal quantum
number. Look at following examples for better understanding;
16
S: 1s2
2s2
2p6
3s2
3p4
3 is the highest energy level of electrons or principal quantum
number. Thus period of S is 3.
23
Cr: 1s2
2s2
2p6
3s2
3p6
4s2
3d4
4 is the highest energy level of electrons or principal
quantum number. Thus period of Cr is 4.
b) Finding Group of Elements:
Group of element is equal to number of valence electrons of element or number of
electrons in the highest energy level of elements. Another way of finding group of element
is looking at sub shells. If last sub shell of electron configuration is "s" or "p", then group
becomes A.
19
K: 1s2
2s2
2p6
3s2
3p6
4s1
Since last sub shell is "s" group of K is A.
35
Br: 1s2
2s2
2p6
3s2
3p6
4s2
3d10
4p5
Since last sub shell is "p" group of Br is A. Elements in
group B have electron configuration ns and (n-1)d, total number of electrons in these
orbitals gives us group of element. Look at following examples.
26
Fe: 1s2
2s2
2p6
3s2
3p6
4s2
3d6
6+2=8 B group
Finding Location of Elements in Periodic Table with Examples

Last Orbital Group
ns1
: 1A
ns2
: 2A
ns2
np1
: 3A
ns2
np2
: 4A
ns2
np3
: 5A
ns2
np4
: 6A
ns2
np5
: 7A
ns2
np6
: 8A
Last Orbital Group
ns2(n-1)d1
: 3B
ns2
(n-1)d2
:4B
ns2
(n-1)d3
:5B
ns2
(n-1)d4
or ns1
(n-1)d5
:6B
ns2
(n-1)d5
:7B
Last Orbital Group
ns2
(n-1)d6
:8B
ns2
(n-1)d7
:8B
ns2
(n-1)d8
:8B
ns2
(n-1)d9
or ns1
(n-1)d10
:1B
ns2
(n-1)d10
:2B
Example: Find period and group of 16
X.
16
X: 1s2
2s2
2p6
3s2
3p4
3. period and 2+4=6 A group
Example: Find period and group of 24
X.
24
X:1s2
2s2
2p6
3s2
3p6
4s2
3d4
4. period and 4+2=6 B group
Finding Location of Elements in Periodic Table with Examples
Here are some clues for you to find group number of elements.

Defects of Modern Periodic Table:
The modern periodic table is an essential tool in chemistry that organizes and displays
the elements based on their atomic number, electron configuration, and recurring
chemical properties. However, the periodic table also has a few limitations and defects
that are worth mentioning.
1. Anomalous Position of Hydrogen: Hydrogen is placed separately at the top of Group
1 (IA), which deviates from the pattern followed by other elements. It is because
hydrogen has unique properties and does not fit precisely into any particular group. The
placement of hydrogen in a separate position often causes confusion and makes it
difficult to establish its chemical behavior.
2. Position of Isotopes: Isotopes are atoms of the same element with different numbers
of neutrons. The periodic table does not provide a specific position for isotopes, as it is
primarily based on the atomic number. As a result, isotopes of an element are not
represented individually, leading to a lack of clarity regarding their properties and
characteristics.
3. Lack of Information on Chemical Reactions: The periodic table does not provide
detailed information about the chemical reactions that elements undergo. It only displays
the electron configuration and recurring chemical properties, but not their specific
reactivity. This limitation makes it necessary to refer to other sources or databases to
gather information about the reactions of specific elements.

4. Inadequate Representation of Inner Transition Elements: The inner transition
elements, also known as the f-block elements, are placed below the main body of the
periodic table. The placement of these elements makes the periodic table wider and less
visually appealing. Moreover, the f-block elements are often compressed and not
displayed in their entirety, leading to a lack of clarity.
5. Ambiguity in Electron Configuration: The electron configuration of certain
elements becomes ambiguous when they are placed in different blocks of the periodic
table. For example, the electron configuration of copper (Cu) is represented as [Ar] 3d10
4s1
, even though it is placed in the d-block. This ambiguity can create confusion and
complicate the understanding of electron arrangements.
6. Position of inert gases in periodic table: Except helium, in all other noble gases
there are eight electrons in their outermost shells. According to the principle of modern
periodic table law, the number of outer most shell electron of an atom indicates the
group number. So, helium is supposed to be in IIA and all other inert gases in group
VIII. But He along with other noble gases have been placed in zero group of modern
periodic table.

1. Periodic table has been useful in predicting the existence of new elements.
2. It has been useful in the past in correcting the position of elements in relation to
their properties.
3. Study of elements and their compounds has become systematic and easier to
remember.
4. Position of an element in the periodic table reveals its:
1. atomic number
2. electronic configuration
3. number of valence electrons
4. Properties
5. Nature of chemical bond, formula of compound formed and properties of that
compound can all be predicted from the periodic table.
6. Periodic table helps to predict the types of chemical reactions that a particular
element is likely to participate in.
Applications of Modern Periodic Table:

7. It tells us which elements are more reactive than others.
8. Without periodic table it would be difficult to show the similarities and
dissimilarities between elements.
9. Position of an element in the periodic table reveals:
valency of the element
whether the element is a metal or a non-metal — metals occupy the extreme left
positions of the periodic table while non-metals are at the extreme right of the periodic
table.
10. The reactivity of an element, whether it is likely to conduct electricity or not,
whether it is hard or soft , and many other information.
11. Periodic table has simplified the study of different elements to a large extent.
12. Prediction of undiscovered elements.
13. Creation of enthusiasm to the scientists.

There is only one position for an element in the periodic table.
Discuss the position of hydrogen in the periodic table.
Position of noble (inert gases) in the periodic table.
Diagonal relationship of elements in the periodic table.
Self study……………………

Periodic properties/ trends
Properties repeat after a certain interval of atomic number
Atomic size
•how big the atoms are
Ionization energy
•How much energy to remove an electron
Electronegativity
•The attraction for the electron in a compound
•Electron affinity: How much energy released when an electron is added
to a neutral atom to form an anion.

Radius is influenced by two factors:
Energy Level
Higher energy level is further away
Charge on nucleus
More charge pulls electrons in closer
Atomic radius is the distance between the centre of atom and the outermost
shell
Atomic Radius = half the distance between two nuclei of molecule
}
Atomic radius

H
Li
Na
K
• As we go down in a group
• New shells are added, thereby pushing outermost
electrons farther from the nucleus.
• The atomic radius of atoms generally increases from top
to bottom within a group.
Group trends
Periodic Trends of atomic radius

As you go across a period the radius gets smaller.
❖ Electrons are added to same shell
❖ Experience greater pulls from the nucleus, The atomic radius of atoms
generally decreases from left to right across a period.
Na Mg Al Si P S Cl Ar
Period trends

Periodic trends in atomic radius:

• Electronegativity is defined as an atom’s ability to attract electrons towards
itself in a chemical bond.
• Big electronegativity means it pulls the electron toward it.
• Different elements have different electronegativities based on a number of
factors such as size and number of protons, neutrons, and electrons.
Electronegativity
Electronegativity Trends
From left to right across the period table electronegativity increases. This is
because of the increased number of protons as the atomic number increase.
From top to bottom electronegativity decreases because of the increasing size of
the atoms. As a result, Fluorine is considered the most electronegative element
while cesium is the least electronegative element. Halogens are considered to
have a high electronegativity, while it is low for the alkali metals and alkaline
earth metals.

Electronegativity trends in periodic table:

Ionization Energy
❖ The minimum amount of energy required to remove the outermost electron
from an atom in the gaseous state is called the ionization energy.
❖ First ionization energy (I1): amount of energy required to remove the most
loosely bound electron from an isolated gaseous atom to form a cation.
❖ Second ionization energy (I2): amount of energy required to remove a
second electron from the gaseous monopositive cation to form dipositive
cation.
❖ Ionization energies are usually expressed in electron volts (eV) per atom or in
kilojoules per mol (kJ/mol)
1eV/atom=96.48 kJ/mol
❖ Ionization energies measure hoe tightly electrons are bound to atoms.
❖ Low energies indicate ease of removal of electrons and vice versa

Ionization energy generally increases, moving from left to right across an
element period (row). This is because the atomic radius generally
decreases moving across a period, so there is a greater effective attraction
between the negatively charged electrons and positively-charged nucleus.
Ionization is at its minimum value for the alkali metal on the left side of
the table and at its maximum for the noble gas on the far right side of a
period. The noble gas has a filled valence shell, so it resists electron
removal.
Ionization decreases, moving top to bottom, down an element group
(column). This is because the principal quantum number of the outermost
electron increases moving down a group. There are more protons in atoms
moving down a group (greater positive charge), yet the effect is to pull in
the electron shells, making them smaller and screening outer electrons
from the attractive force of the nucleus. More electron shells are added
moving down a group, so the outermost electron becomes increasingly
distant from the nucleus.

1. Effective nuclear charge
2. Atomic size i.e. atomic radius
3. Principle quantum number
4. Shielding effect
5. Half filled and completely filled orbitals
6. Nature of orbitals
7. The extent of penetration of valence electrons
Factors affecting the magnitude of Ionization Potential:

(1) Effective Nuclear Charge:
Shielding and energy levels.
A, Within an energy level, each electron shields (red arrows) other electrons from the full nuclear charge
(black arrows), so they experience a lower Zeff.
B, Inner electrons shield outer electrons much more effectively than electrons in the same energy level.
The effective nuclear charge is the positive charge that an electron
experiences from the nucleus, equal to the nuclear charge but reduced
by any shielding or screening from any intervening electron
distribution.

Effective Nuclear Charge
Greater the magnitude of effective nuclear charge, higher is the amount of
energy needed to remove the outermost shell electron.
Thus with the increase of the magnitude of effective nuclear charge, the
magnitude of ionization potential also increases.
The effective nuclear charge increases from left to right in a period.
Consider the effective nuclear charge on the 2s electron in the lithium
atom (configuration 1s2
2s1
). The nuclear charge is 3e, but the effect of
this charge on the 2s electron is reduced by the distribution of the two
1s electrons lying between the nucleus and the 2s electron (roughly, each
core electron reduces the nuclear charge by 1e).

(2) Atomic size:
Greater is the atomic size of an atom, more far is the outermost shell electron
from the nucleus and hence lesser will be the force of attraction exerted by the
nucleus on the outermost shell electron. Thus higher the value of atomic
radius of an atom, lower will be the ionization energy.
(3) Principal Quantum Number (n):
Greater is the value of n for the valence shell electron of an atom, further away
this electron will be from the nucleus and hence lesser will be the force of
attraction exerted by the nucleus on it so lesser energy will be required to
remove the valence shell electron. Thus with the increase of the principal
quantum number of the orbital from which the electron is to be removed, the
magnitude of ionization potential decreases.
(4) Shielding Affect:
The magnitude of shielding effect determines the magnitude of the force of
attraction between the nucleus and the valence-shell electron. Greater is the
magnitude of shielding effect working on the valence shell electron. The lesser
the ionization potential.

The screening or shielding effect:
∙ When the number of inner electrons is greater, they shelter the outermost electron from the
nucleus, allowing it to neglect the nuclear pull to some extent.
∙ This is referred to as the shielding or screening effect.
∙ The electrons in the valence shell are pulled to the nucleus in a multielectron atom, and these
electrons are repulsed by the electrons in the inner shells.
∙ As a result of the repulsive forces acting in opposite directions, the actual force of attraction
between the nucleus and the valence electrons is slightly reduced.
∙ The screening effect or shielding effect refers to the decrease in the nucleus's force of attraction
on valence electrons due to the existence of electrons in the inner shell.

(5) Half-filled and completely-filled orbitals:
According to Hund’s rule, half-filled (ns1
, np3
, nd5
) or completely-filled (ns2
.
ns6
, nd10
) orbitals are comparatively more stable and hence more energy is
needed to remove an electron from such orbitals. Thus the ionization
potential of an atom having half-filled or completely-filled orbitals in its
electronic configuration is relatively higher than that expected normally
from its position in the periodic table .
(6) Nature Of Orbitals:
The nature of orbitals of the valence-shell from which the electron is to be removed
also influences the magnitude of ionization potential. The relative order of energy
of
s, p, d and f orbitals of a given nth shell is as:
ns < np < nd < nf
This order clearly shows that to remove an electron from f-orbital will be easiest
while to remove the same from s-orbital will be the most difficult.
(7) The extent of penetration of valence electrons:
The degree of penetration of valence electrons in a given principal energy level
decreases in the order s>p>d>f, since ns electron is more tightly bound than any np

Trends in Ionization Potential:
Ionization energy generally increases from left to right in a period
because of the increase in nuclear charge and decrease in atomic
radius.
Ionization energies generally decrease down a group due to the
shielding effect and increase in atomic size.
Departures from these trends can usually be traced to repulsion
between electrons, particularly electrons occupying the same orbitals.
The minimum amount of energy required to remove the outermost electron
from an atom in the gaseous state is called the first ionization energy.

Ionization energies display a periodic variation when plotted against atomic
number. Within any period, IE values tend to increase with atomic number.
The increase in ionization energy with atomic number in a given
period—can be explained as follows: The outer-shell electrons in the
elements of the same period are arranged in the same shell, hence, the
positive charge on the nucleus increases whereas the distance between the
nucleus and valence electrons decreases. Therefore more energy is required
to remove an electron as we go from left to right in the Period.
Ionization energies tend to decrease going down any column of main-group
elements.
This is because new shells are added, thereby pushing outermost electrons
farther from the nucleus, so the atoms get bigger, and IE is decreases.

Small deviations from this general trend occur.
A IIIA element (ns2
np1
) has smaller ionization energy than the preceding
IIA element (ns2
). Apparently, the np electron of the IIIA element is more
easily removed than one of the ns electrons of the preceding IIA element.
Also note that a VIA element (ns2
np4
) has smaller ionization energy than
the preceding VA element (ns2
np3
). As a result of electron repulsion, it is
easier to remove an electron from the doubly occupied np orbital of the VIA
element than from a singly occupied orbital of the preceding VA element.

Electron Affinity
Electron affinity (Eea
) is the energy change when an electron is added to a
neutral atom in the gas phase. In simple terms, it is a measure of a neutral
atom’s ability to gain an electron. The gas phase atom is used (rather than liquid
or solid) because the atom’s energy levels aren’t influenced by neighboring
atoms. The most common units for electron affinity are (kJ/mol) or (eV).
X (g) + e-
X-
(g)

Electron Affinity Trend on the Periodic Table
Like electronegativity, ionization energy, atomic or ionic radius, and metallic
character, electronegativity displays periodic table trends. Unlike some of these
other properties, there are many exceptions to the trends for electron affinity.
•Electron affinity general increases moving across a row or period of the
periodic table, until you reach group 18 or the noble gases. This is because of the
filling of the valence electron shell moving across a period. For example, a group
17 (halogen) atom becomes more stable by gaining an electron, while a group 1
(alkali metal) must add several electrons to reach a stable valence shell. Further,
the effective nuclear charge increases as you move across a period.
•Noble gases have low electron affinities.
•Generally (with exceptions) nonmetals have a higher or more positive Eea
value
than metals.
•Atoms that form anions that are more stable than the neutral atoms have high
electron affinity values.
•Although usually depicted on a diagram of periodic table trends, electron affinity
does not reliably decrease moving down a column or group.

Variation of Electron Affinity Trend on the Periodic Table

Factors Affecting Electron Affinity:
Atomic size:
The smaller the size of atom smaller will be the distance between the extra
electron and the nucleus. Therefore, electrostatic force of attraction will be
more and the electron affinity will be higher.
Nuclear charge:
More the nuclear charge of the atom more strongly will it attract additional
electron. Therefore, electron affinity increases as the nuclear charge increases.
Electronic Configuration:
Atoms having stable electronic configuration (i.e. those having completely
filled or half filled outer orbitals) do not show much tendency to add extra
electron, so have either zero or very low electron affinities.

Which Element Has the Highest Electron Affinity?
Halogens, in general, readily accept electrons and have high electron affinities.
The element with the highest electron affinity is chlorine, with a value of 349
kJ/mole. Chlorine gains a stable octet when it captures an electron.
The reason why chlorine has a higher electron affinity than fluorine is because
the fluorine atom is smaller. Chlorine has an additional electron shell, so its
atom more easily accommodates the electron. In other words, there is less
electron-electron repulsion in the chlorine electron shell.
Which Element Has the Lowest Electron Affinity?
Most metals have lower electron affinity values. Nobelium is the element with
the lowest electron affinity (-223 kJ/mol). Nobelium atoms have an easy time
losing electrons, but forcing another electron into an atom that’s already huge
isn’t thermodynamically favorable. All of the existing electrons act as a screen
against the positive charge of the atomic nucleus.

63
Summery of Periodic Functions of the elements

03_Periodic Table for engineering students.pdf

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