1.1 INTRODUCTION

Elements are pure substances made up of one type of atoms. They are the basic units of all types of matter. In order to study elements in an organised manner, they need to be classified.

Historical Classifications

• Dobereiner: Grouped the elements in three (triads).
• Newland: Observed that when elements are arranged in increasing order of their atomic mass, every eighth element beginning from any element resembles the first element in its physical and chemical properties (Law of Octaves).
• Dmitri Mendeleev: A Russian chemist gave the first periodic table of elements based on his law which states that "the properties of elements are the periodic functions of their atomic masses". This arrangement enabled him to place 63 elements known at that time in vertical columns (groups) and horizontal rows (periods). He even predicted the existence of elements yet to be discovered.

Modern Periodic Law

These defects were removed when Henry Moseley put forward the modern periodic table. He stated that:

Later on Niels Bohr gave the extended form of the table known as the long form of the modern periodic table.

1.2 SALIENT FEATURES OF THE MODERN PERIODIC TABLE

Groups

The modern periodic table has eighteen vertical columns. They are known as groups, arranged from left to right in the order 1 to 18.

Group No.	Name / Family	Characteristics
Group 1	Alkali metals	(Li, Na, K, Rb, Cs, Fr) They form strong alkalis with water.
Group 2	Alkaline earth metals	(Be, Mg, Ca, Sr, Ba, Ra) They form weaker alkalis as compared to group 1 elements.
Groups 3-12	Transition elements	They have their two outermost shells incomplete. Placed at extreme right of representative metals.
Group 13	Boron family	Boron is the first member of the group.
Group 14	Carbon family	Carbon is the first member.
Group 15	Nitrogen family	Nitrogen is the first member.
Group 16	Oxygen family / Chalcogens	Chalcogen means ore-forming.
Group 17	Halogens	Meaning salt-formers. They form salts.
Group 18	Noble gases / Inert gases	Elements of this group are called the noble gases. They have their outermost orbit complete. Due to stable electronic configuration they hardly react with other elements.

Periods

There are seven horizontal rows in the modern periodic table. They are known as periods.

The number of shells present in an atom determines its period.

• Elements of period one have 1 shell.
• Elements of period two have 2 shells.
• Elements of period three have 3 shells and so on.

Period	Type of period	Number of elements	Atomic no. of elements	Starting & Ending Element
1	Shortest period	2	1 - 2	H to He
2	Short period	8	3 - 10	Li to Ne
3	Short period	8	11 - 18	Na to Ar (Typical Elements)
4	Long period	18	19 - 36	K to Kr
5	Long period	18	37 - 54	Rb to Xe
6	Longest period	32	55 - 86	Cs to Rn
7	Longest period	32	87 - 118	Fr to Uuo

1.3 PERIODICITY

The properties that reappear at regular intervals, or in which there is gradual variation (i.e. increase or decrease) at regular intervals, are called 'periodic properties' and the phenomenon is known as the periodicity of elements.

In a particular group, electrons in the outermost orbit remain the same i.e. electronic configuration is similar. Since chemical properties of elements depend upon the number of electrons in their outermost shell, thus elements of the same group have similar properties.

Example: Periodicity in Halogens (Group 17)

All elements have seven electrons in their respective outermost shells, therefore, they show similar properties, such as:

1. They are coloured non-metals.
2. They form negative ions carrying a single charge. For example chloride ions ($Cl^-$).
3. They are very reactive and are, therefore, found in combined state.
4. They are only slightly soluble in water, but they dissolve much better in organic solvents like carbon disulphide, chloroform, alcohol, etc.
5. Their melting and boiling points increase regularly moving down the group.
6. They are good oxidising agents.

Conclusion: It is thus concluded that periodicity is due to the same number of electron(s) in the outermost orbit of different elements.

1.4 SHELLS (ORBITS) AND VALENCY

Orbits: Electrons revolve around the nucleus in certain definite circular paths called orbits or shells.

(1) Number of shells:

• Down a group (top to bottom): The number of shells increases successively, i.e., one by one, such that the number of shells that an element has, equals the number of the period to which that element belongs.
  Example in Halogens (Group 17): F (9) has 2 shells (2,7); Cl (17) has 3 shells (2,8,7); Br (35) has 4 shells (2,8,18,7); I (53) has 5 shells (2,8,18,18,7); At (85) has 6 shells.
• Across a period (left to right): On moving from left to right in a given period, the number of shells remains the same. For example, in the 2nd period, the number of shells remains two. Similarly in the third period the number of shells remains three and so on.

(2) Valency:

Valency denotes the combining capacity of the atom of an element. It is equal to the number of electrons an atom can donate or accept or share.

• Down a group: On moving down a given group, the number of electrons in the outermost shell, i.e., valence electron, remains the same. Therefore, valency in a group remains the same.
• Across a period: In a given period, the number of electrons in the valence (outermost) shell increases from left to right. But the valency increases only upto Group 14 (where it becomes 4), and then it decreases, i.e., in Group 17 it becomes 1.

Valency Trends in Elements of the 2nd Period

Groups →	1	2	13	14	15	16	17	18
Elements	Li	Be	B	C	N	O	F	Ne
Atomic No.	3	4	5	6	7	8	9	10
Electronic Conf.	2, 1	2, 2	2, 3	2, 4	2, 5	2, 6	2, 7	2, 8
Valency	1	2	3	4	3	2	1	0
Formula of Hydride	LiH	$BeH_2$	$BH_3$	$CH_4$	$NH_3$	$H_2O$	HF	-

1.5 PERIODIC PROPERTIES

The properties of elements which are directly or indirectly related to their electronic configurations and show a regular gradation as we move across a period, from left to right or down the group from top to bottom, are called Periodic Properties.

Important periodic properties are: (i) atomic size (atomic radius), (ii) metallic character, (iii) non-metallic character, (iv) ionisation potential, (v) electron affinity, (vi) electronegativity.

1.5.1 Atomic Size (Atomic Radius)

Atomic size depends upon:

1. Number of shells: An increase in the number of shells increases the size of an atom because the distance between the outermost shell and the nucleus increases.
2. Nuclear charge: It is the positive charge present in the nucleus of an atom, which is equal to the number of protons (atomic number). An increase in nuclear charge decreases the size of the atom because the electrons are then attracted towards the nucleus with a greater force, thereby bringing the outermost shell closer to the nucleus.

Trends in Atomic Size (Atomic Radius)

• (a) Down a group (INCREASES): In a group, the size of an atom increases as one proceeds from top to bottom. This is due to the successive addition of shells (which overweighs the increased nuclear charge).
  Example (Group 1): H (37pm) < Li (152 pm) < Na (186 pm) < K (231 pm) < Rb (244 pm) < Cs (262 pm).
  Example (Group 17): F (64pm) < Cl (99pm) < Br (114pm) < I (133pm) < At (140pm).
• (b) Across a period (DECREASES): In a period, the size of an atom decreases from left to right. This is because the nuclear charge increases while the shells remain the same, bringing the outermost shell closer to the nucleus.
  Example (Period 2): Li (152pm) > Be (112pm) > B (88pm) > C (77pm) > N (70pm) > O (66pm) > F (64pm). Lithium is largest, Fluorine is smallest.
  Example (Period 3): Na (186pm) > Mg (160pm) > Al (143pm) > Si (117pm) > P (110pm) > S (104pm) > Cl (99pm).

1.5.2 Metallic Character

Those elements, which have a tendency to lose their valence electrons and form a positive ion, are considered metals. Metals are good reducing agents (they lose electrons). Greater the tendency to lose electron(s) stronger is the reducing agent.

Reactions: $Na - e^- \rightarrow Na^+$ | $Mg - 2e^- \rightarrow Mg^{2+}$

The metallic character of elements depends on: (i) atomic size (greater size = easier to remove electrons) and (ii) nuclear charge (greater charge = difficult to remove electrons).

Trends in metallic character

• Down a group (INCREASES): The atomic size increases and the nuclear charge also increases. The effect of an increased atomic size is greater. Therefore, metallic nature increases as one moves down a group (can lose electrons easily). In group 1, Lithium is least metallic, Cs is highly metallic.
• Across a period (DECREASES): Nuclear pull increases and atomic size decreases, so atoms cannot lose electrons easily. Metallic nature decreases.
  Example: In 3rd period: Na (Metal) -> Mg (Metal) -> Al (Metal) -> Si (Metalloid) -> P (Non-metal) -> S (Non-metal) -> Cl (Non-metal) -> Ar (Noble gas).

1.5.3 Non-Metallic Character

Those elements, which have a tendency to gain electrons, in order to attain octet in their outermost orbit, are considered as non-metals. They form anions. Non-metals are good oxidising agents.

Reactions: $Cl + e^- \rightarrow Cl^-$ | $O + 2e^- \rightarrow O^{2-}$

Non-metallic character depends on: (i) atomic size (smaller size = greater pull for incoming electrons) and (ii) nuclear charge (greater charge = greater tendency to gain electrons).

Trends in non-metallic character

• Down a group (DECREASES): The atomic size increases over successive periods. The effect of an increasing atomic size is greater than nuclear charge. Therefore, non-metallic nature decreases down the group. (e.g. Group 14 goes from Carbon (non-metal) to Si, Ge (metalloids) to Sn, Pb (metals)).
• Across a period (INCREASES): Tendency to gain electron(s) increases due to an increase in nuclear pull and a decrease in the atomic size. Example: $Na < Mg < Al < Si < P < S < Cl$ (Non-metallic character increases).

1.5.4 Ionisation Potential or Ionisation Energy (I.E.)

I.E. depends on (i) atomic size (greater size = lesser IO energy), and (ii) nuclear charge (greater charge = higher IO energy).

Trends in ionisation energy

• Across a period (INCREASES): Atomic size decreases and nuclear charge increases, requiring more energy to remove the electron.
  Period 2 ($\text{kJ mol}^{-1}$): Li (520), Be (899), B (801), C (1088), N (1402), O (1314), F (1681), Ne (2080).
  Period 3 ($\text{kJ mol}^{-1}$): Na (496), Mg (737), Al (577), Si (786), P (1011), S (999), Cl (1256), Ar (1520).
• Down a group (DECREASES): Increase in atomic size overcomes the effect of increase in nuclear charge. Valence electrons are held less tightly.
  Group 1 ($\text{kJ mol}^{-1}$): H (1312), Li (520), Na (496), K (419), Rb (403), Cs (375).

Note: Helium will have the highest ionisation energy (2372.0 $\text{kJ mol}^{-1}$) while Caesium has the lowest (375.0 $\text{kJ mol}^{-1}$). Metals usually have low I.E. whereas non-metals have high I.E.

1.5.5 Electron Affinity (E.A.) or Electron Gain Enthalpy

Electron affinity depends on (i) atomic size (smaller size = greater EA) and (ii) nuclear charge (greater charge = greater EA).

Trends in electron affinity

• Across a period (INCREASES): Atomic size decreases and nuclear charge increases, so the atom eagerly accepts an electron. E.A. is highest for group 17 (halogens) and least for group 1 (alkali metals).
  Period 2 E.A. Values: Li (-59.8), Be (exception), B (-26.7), C (-122), N (exception), O (-141), F (-328), Ne (0).
• Down a group (DECREASES): Atomic size increases heavily, shielding the nucleus and decreasing the attraction for a new electron.

1.5.6 Electronegativity (E.N.)

It is a dimensionless property. The most widely used scale was devised by Linus Pauling. Fluorine has the highest E.N. (4.0) and Caesium lowest (0.7).

Trends in Electronegativity

• Across a period (INCREASES): Atomic size decreases and nuclear pull increases.
  Period 3 EN values: Na(0.9) -> Mg(1.2) -> Al(1.5) -> Si(1.8) -> P(2.1) -> S(2.5) -> Cl(3.0).
• Down a group (DECREASES): Atomic size increases rapidly overcoming increased nuclear charge.
  Group 17 EN values: F(4.0) -> Cl(3.0) -> Br(2.8) -> I(2.5) -> At(2.2).

1.6 RELATION BETWEEN ATOMIC NUMBER AND MASS NUMBER

Elements which have an even number of protons (e.g. $_6^{12}\text{C}$) generally have their mass numbers twice the atomic numbers ($A = 2Z$). Exceptions exist like $_4^9\text{Be}$ and $_{18}^{40}\text{Ar}$.

Elements which have an odd number of protons (e.g. $_9^{19}\text{F}$, $_{11}^{23}\text{Na}$) have mass numbers twice the atomic numbers + 1 ($A = 2Z + 1$). Exceptions exist like $_7^{14}\text{N}$ and $_1^1\text{H}$.

1.7 COMPARISON OF ALKALI METALS AND HALOGENS

Property	Alkali metals [Group 1]	Halogens [Group 17]
Elements	Li, Na, K, Rb, Cs, Fr	F, Cl, Br, I, At
Occurrence	Combined state (due to reactive nature).	Combined state as salts (due to reactive nature).
Physical State	Shining white solid metals. Soft (cut with knife). Shiny when freshly cut but become dull reacting with air.	Non-metals. Diatomic gas F (yellow), Cl (yellow green), liquid Br (red brown), solid I (dark grey).
Valence Electrons	Possess one valence electron.	Possess seven valence electrons.
Nature & Conduction	Highly electropositive. Good conductors.	Highly electronegative. Non-conductors.
Atomic Size	Largest atomic size in their period (except inert gases). Increases down the group.	Smallest atomic size in their period. Increases down the group.
Ionisation Energy	Lowest I.E. in their period. Decreases down.	High I.E. (lower than noble gases) in their period.
Electron Affinity	Low E.A. values. Decrease down the group.	High E.A. values. Decrease down the group.
Chemical Action	Strong reducing agents (lose electrons). React vigorously with water and acids liberating hydrogen.	Strong oxidising agents (accept electrons). Generally do not react with dil. acids and water.
Compound Formation	Form electrovalent compounds with non-metals (NaCl, KBr).	Form electrovalent (with metals like KCl) and covalent compounds (with non-metals like HCl).

SUMMARY: VARIATIONS IN PERIODIC PROPERTIES

Properties	Across the period (Left to right)	Down the group (Top to bottom)
No. of valence electrons	Increases	Remains same
Atomic size (radius)	Decreases	Increases
Ionisation energy	Increases	Decreases
Metallic character	Decreases	Increases
Non-metallic character	Increases	Decreases
Electron affinity	Increases	Decreases
Electronegativity	Increases	Decreases
Basic nature of oxides	Decreases	Increases
Melting point & Boiling Point	Increases from group I to IV and then decreases	Group I and II decreases. Group V to VII increases
Oxidising nature	Increases	Decreases
Reducing nature	Decreases	Increases

CHAPTER EXERCISE (Exam Prep & Past Board Questions)