The Periodic Table

Period 1 — 2 Elements
Period 2 — 8 Elements
Period 3 — 8 Elements

Hydrogen [H] and Helium [He].

The number of valence shell electrons in the elements at the end of Period 2 and Period 3 is 8.

• Ne = 2, 8
• Ar = 2, 8, 8

If an element has one electron in it's outermost energy level [shell] then it is likely to be metallic.

The similarities in the properties of a group of elements is because they have the same number of outer electrons.

The 18 vertical columns in the periodic table are called groups. Elements in the same group have the same number of valence electrons and similar chemical properties.

It will be non-metallic because being in Group 17 [or Group 7A], it has 7 valence electrons. Thus, its electronegativity is high and it has a tendency to accept electrons.

(a) The noble gas having duplet arrangement of electrons — Helium

(b) The noble gas having an electronic configuration 2, 8, 8 — Argon

(c) A metalloid in period 2 and in period 3 — Metalloid in period 2 is B (Boron) and in period 3 is Si (Silicon)

(d) The number of electron shells in elements of period 1, period 2 and period 3 — Period 1 — 1 shell, Period 2 — 2 shells, Period 3 — 3 shells

(e) The valency of elements in group 1 [IA] — +1

(f) The group whose elements have zero valency — Group 18 or 0 group

(g) An alkaline earth metal in period 3 — Magnesium [Mg]

(h) The non-metallic element present in period 3 other than sulphur and chlorine — Phosphorus [P]

(i) A non-metal in period 2 having electronic configuration 2, 6 — Oxygen [O]

(j) An electrovalent compound formed between an alkali metal and a halogen — Sodium chloride [NaCl]

(k) A covalent compound formed between an element in period 1 and a halogen — Hydrogen chloride [HCl]

(l) An alkali metal in period 3 which dissolves in water giving a strong alkali — Alkali metal is Na and it dissolves in water to give strong alkali - NaOH

(m) A metal in period 3 having valency 3 — Aluminium [Al]

(n) The bridge elements of period 3 of group 1 [IA], 2 [IIA] and 13 [IIIA] — Li & Mg, Be & Al, B & Si

(o) The periods which contain the inner transition elements — Period 6 and 7

(p) The formula of the hydroxide of the element having electronic configuration 2, 8, 2 —
Since valency of element X is 2+ and valency of OH is 1-

Formula of the compound:

$$\begin{aligned}\text{X}^{2+} \phantom{\nearrow} \text{OH}^{1-} \\[0.5em] \overset{\phantom{2}{2}}{\text{X}} \space {\searrow}\mathllap{\swarrow} \space \overset{1}{\text{OH}} \Rightarrow \underset{\phantom{1}{1}}{\text{X}} \space {\searrow}\mathllap{\swarrow} \space \underset{2}{\text{OH}} \\[0.5em]\end{aligned}$$

So, we get the formula as $\bold{X(OH)}\bold{_2}$

As the element with electronic configuration 2,8,2 is Mg, Hence the compound is Mg(OH)₂

(q) The valency of the element in period 3 having atomic number 17 — -1

(r) A non-metal in period 2 which is tetravalent — Carbon [C]

An element is the basic unit of matter. They are pure substances made up of one kind of atoms and cannot be broken down into simpler substances.

Elements need to be properly classified and the best classification is the one which puts together those elements which show maximum resemblance and separates the others.

The method of classification was discarded by early chemists since the elements behave both as metals and non-metals, show metallic and non-metallic nature and have variable valency.

Dobereiner arranged elements in increasing order of atomic weights - in groups of three called 'triads'.

Dobereiner's Law of Triads — Chemically analogous elements arranged in increasing order of their atomic weights formed well marked groups of three called 'triads' in which the atomic weight of the middle element was found to be generally the average [arithmetic mean] of the atomic weights of the other two elements.

Examples :

Atomic weight of sodium is the average (arithmetic mean) of Lithium and Potassium.

Dobereiner's method of classification of elements did not hold much weightage for future classification because —

1. It was a disordered collection of elements not holding true for all elements.
2. Properties of elements grouped into triads were found not to be similar.

Newland arranged elements in increasing order of atomic weights in series of eight.

Newland's Law of Octaves — Elements when arranged in increasing order of their atomic weights show resemblance in physical and chemical properties between the eighth element and the first element [similar to the eighth and first notes on a musical scale]

For example:

The eighth element from Lithium is Sodium and from Fluorine is Chlorine which show similarity in properties.

Elements when arranged in increasing order of their atomic weights show resemblance in physical and chemical properties between the eighth element and the first element similar to the eighth and first notes on a musical scale.

Mendeleeff arranged elements in increasing order of their atomic weights in the form of a table called 'periodic table'.
This arrangement had elements with similar properties appearing at regular intervals.
This phenomenon is called 'periodicity of properties' and the properties which repeat themselves at regular intervals are called 'periodic properties'.

Mendeleeff's Periodic Law — The physical and chemical properties of elements are periodic functions of their atomic weights.

Mendeleeff's arrangement had elements with similar properties appearing at regular intervals. This phenomenon is called 'periodicity of properties' and the properties which repeat themselves at regular intervals are called 'periodic properties'. Some examples of periodic properties are atomic radius, ionisation potential, electron affinity, etc.

The contributions of Mendeleeff towards the periodic table are:

1. Elements were arranged in increasing order of atomic weights in horizontal rows called 'periods' and vertical columns called 'groups'.
2. Elements which are similar with respect to their chemical properties are grouped together and have atomic weights of nearly the same value.
3. Elements in the same group had the same 'valency' and similar chemical properties.
4. Based on the periodicity of properties a number of gaps were left in the table for undiscovered elements i.e., elements now discovered e.g., Scandium, Gallium and Germanium originally called eka-boron, eka-aluminium and eka-silicon respectively.
5. The properties of undiscovered elements left in the vacant gaps were predicted.
6. Incorrect atomic weights of some of the arranged elements were corrected with the knowledge of the atomic weights of the adjacent elements.

Defects in Mendeleeff's Periodic Table :

1. Anomalous pairs of elements — Certain pairs of elements having higher atomic weights have been given positions before the elements having lower atomic weights.
   This defect disappears if elements were arranged according to their atomic numbers.
   e.g., Co [at. wt. 58.9, at. no. 27] was placed before Ni [at. wt. 58.6, at. no. 28].
2. Position of rare earths and actinides — Could be justified only if arranged according to their atomic numbers
3. Position of isotopes — Isotopes had to be placed in same position according to atomic numbers

Most of the defects in Mendeleeff's Periodic Table disappear if the basis of classification of elements is changed from atomic weights to atomic numbers.
Moseley arranged the elements in increasing order of atomic number in Seven horizontal rows called 'periods' and Eighteen vertical columns called 'groups'.
This forms the basis of the Modern Periodic Table.

In the Modern Periodic Table elements are arranged in increasing order of atomic number.

Modern Periodic Law states that the physical and chemical properties of elements are periodic function of their 'atomic numbers'.

Mendeleeff's Periodic Law states that the physical and chemical properties of elements are periodic function of their 'atomic weights'.

The basis of classification in Mendeleeff's Periodic Law was atomic weight of the element whereas in the Modern Periodic Law it is the atomic number of elements. This change in the Modern Periodic Law solves most of the defects in Mendeleeff's Periodic table.

Periods in the periodic table are the seven horizontal rows of elements arranged in increasing order of their atomic numbers.

Period number is the number assigned to each period i.e., 1, 2, 3, 4, 5, 6 or 7. It signifies the number of shells present in the atom of the element belonging to that period.

Period 1 is the shortest period. It has 2 elements [H and He]

Second period has 8 elements.

Third period also has 8 elements.

Period 1 — ¹H, ²He

Period 2 — ³Li, ⁴Be, ⁵B, ⁶C, ⁷N, ⁸O, ⁹F, ¹⁰Ne

Period 3 — ¹¹Na, ¹²Mg, ¹³Al, ¹⁴Si, ¹⁵P, ¹⁶S, ¹⁷Cl, ¹⁸Ar

The metallic, non-metallic and noble gas elements of periods 2 and 3 are listed below:

Period 2: Metallic — Li, Be; Metalloid — B; Non-metallic — C, N, O, F; Noble gas — Ne

Period 3: Metallic — Na, Mg, Al; Metalloid — Si; Non-metallic — P, S, Cl; Noble gas — Ar

Property trends of elements from left to right in a period

(a) Number of electron shells — remain the same.

(b) Valence electrons — increases by one.

(c) Transition from metallic to non-metallic character.

(d) Electronegativity increases.

Eighteen vertical columns in the periodic table are called groups.

Trends on moving down in a group of the periodic table :

• Number of shells — increases by one
• Metallic character increases
• Electronegativity decreases down a group
• Valence electrons — remain the same.
• Chemical properties — remain similar [or vary gradually]

(a) Alkali metals —

1. The elements of group 1 [IA] from ³Li [Lithium] to ⁸⁷Fr [Francium] are called alkali metals as they react with water forming strong alkalis.
2. They are light metals and are univalent with one valence electron in outer shell. This makes them strong reducing agents as they readily donate the one electron to achieve stable electronic configuration.
3. They are highly reactive electropositive metals, which are soft and can be cut with a knife.
4. They are good conductors of heat & electricity.
5. They form electrovalent compounds with non-metals. e.g. KCl, NaBr.

(b) Alkaline earth metals —

1. The elements of group 2 [IIA] from ⁴Be [Beryllium] to ⁸⁸Ra [Radium] are called Alkaline earth metals.
2. They are light metals and are divalent with two valence electrons in their outer shell. They are reducing agents as they can donate the two electrons to achieve stable electronic configuration.
3. They have properties similar to alkali metals but are slightly less reactive than alkali metals.
4. They are good conductors of heat & electricity.
5. They form electrovalent compounds with non-metals. e.g. CaCl₂, MgO.

(c) Halogens —

1. The elements of group 17 [VIIA] from ⁹F [Fluorine] to ⁵³I [Iodine] are called Halogens. They are called Halogens as Halogens means salt producer and these elements react with metals to form salts.
2. They are non-metals and are univalent with seven valence electrons in their outer shell. This makes them strong oxidising agents as they readily accept one electron to achieve stable electronic configuration.
3. They are highly reactive, electronegative non-metals.
4. They are bad conductors of heat & electricity.
5. They form covalent compounds with non-metals. e.g. HCl, PCl₃.

(d) Noble gases —

1. The elements of group 18 [0] from ²He [Helium] to ⁸⁶Rn [Radon] are called Noble gases.
2. They are non-metals and have eight valence electrons in outer shell [except Helium (He) which has only two]. They have stable electronic configuration hence their valency or combining capacity is zero.
3. Noble gases do not react with other elements & are therefore inert or unreactive. They are all monatomic gases.

(e) Transition and inner transition elements —

Transition elements —

1. The elements of group 3 to 12 [IB to VIIB & VIII] are called Transition elements.
2. They are all metals.
3. They lie at the centre of the periodic table between metals [highly electropositive] on the left and non-metals [highly electronegative] on the right.

Inner transition elements —

1. Lanthanide [Rare earths] and Actinides series are called inner transition elements.
2. They lie at the bottom of the periodic table in two horizontal rows [14 elements each].
3. Lanthanide series : from Cerium ⁵⁸Ce to Lutetium ⁷¹Lu : [Period - 6]
4. Actinide series : from Thorium ⁹⁰Th to Lawrencium ¹⁰³Lr : [Period - 7]
5. They are all metals and chemically very reactive [actinides are radioactive elements].

(f) Normal elements —

1. Elements of groups 1, 2, 13, 14, 15, 16, 17 are called normal elements.
2. Atoms of these elements have all inner shells complete except outermost shell.
3. They consist of metals, metalloids and non-metals.
4. The number of electrons in their outermost shell signifies the group to which they belong.

Li

Reason — The element having two electron shells with eight electrons in the outer valence shell is Neon and Neon is in Period 2. We need an element in the same period that is univalent & a strong reducing agent.
In Period 2, the elements are:

Li, Be, B, C, N, O, F, Ne

Among these, Lithium is univalent as it has valency 1 and a strong reducing agent because it is an electron donor.

Fluorine and Chlorine are strong oxidising agents as they are electron acceptors. Sodium, though a reducing agent and univalent but it is in Period 3.

I

Reason — Iodine is a solid halogen, univalent as it has one valency and acts as an oxidising agent as it is an electron acceptor.
Fluorine is not solid at room temperature; it is a gas. Silicon is not a halogen; it belongs to Group 14. Boron belongs to Group 13, not the halogens.

Both A and R are true but R is not the correct explanation of A.

Reason — A is true as noble gases already have a completely filled outermost shell (stable electronic configuration), so they neither gain, lose, nor share electrons easily.
R is true as noble gases exist as single atoms because they are chemically inert.

However, being monoatomic is a result of their stable electronic configuration, not the reason for their zero valency. The correct explanation for zero valency is their complete outermost shell, not their monoatomic nature.

Therefore, Both A and R are true and R is not the correct explanation of A is the correct option.

Chlorine is more electronegative than fluorine, but less electronegative than bromine.

Reason — Electronegativity order among halogens is:
F > Cl > Br

So chlorine is less electronegative than fluorine and more electronegative than bromine.
Fluorine is the most electronegative element.

(i)

(a) B is divalent.
(b) A is strong reducing agent.
(c) Y is unreactive.

(ii)

(a) X is a liquid.
(b) R is more electronegative.
(c) X is less non-metallic.

(iii) The numbers of electron shells are:

(a) Li — 2 (b) Mg — 3 (c) B — 4 (d) Ar — 3 (e) X — 4

(iv) The numbers of valence electrons are:

(a) Na — 1 (b) B — 2 (c) R — 7 (d) Ar — 8

(v) A is the stronger reducing agent.

(vi)

(a) Li, Na and Mg are normal elements.
(b) Li and Mg are bridge elements because they show diagonal relationship.
(c) Li and Na are alkali metals.

(vii)

(a) Valency of R = 1
(b) Valency of Y = 0

(viii)

(a) A is a metal.
(b) R and Y are non-metals.

(ix) C shows allotropy.

(x) F, R and X show similar chemical properties because they belong to the same group, Group 17, and have the same number of valence electrons, i.e. 7.

The metalloid with three valence electrons is boron, which belongs to Period 2. The non-metallic element in the same period which is diatomic and unreactive at ordinary temperatures is nitrogen, which exists as N₂.

Most electronegative element in group 17 is chlorine as electronegativity decreases down Group 17. Least non-metallic element is carbon as non-metallic character increases across a period from left to right.

C < O < F

1. Moseley
2. Lithium
3. Period 3
4. Fluorine
5. Group 2

1. Across a period from left to right in the Modern Periodic Table.
   No. of electron shells remains same; No. of valence electrons increases by one; Electronegativity increases; Character of elements changes from electropositive to electronegative.
2. Down a group in the Modern Periodic Table.
   No. of electron shells increases by one; No. of valence electrons remains same; Electronegativity decreases; Character of elements changes from electronegative to electropositive.

Mendeleeff's contributions to the periodic table laid the foundation for the Modern Periodic Table because he arranged elements in increasing order of their atomic weights in horizontal rows called 'periods' and vertical columns called 'groups'. His arrangement showed that:

1. properties of elements repeat at regular intervals.
2. elements with similar chemical properties are grouped together.
3. elements in the same group have the same valency and similar chemical properties.
4. gaps were left for undiscovered elements and their properties were predicted.
5. incorrect atomic weights of some elements were corrected using the atomic weights of adjacent elements.

The properties of elements are dependent on the electronic configuration of their atoms. As similar electronic configurations repeat after regular intervals of atomic numbers hence properties of elements are periodic functions of their atomic numbers and not atomic weights.
This was seen when Mendeleeff arranged the elements on the basis of their atomic weights. Certain defects were present in Mendeleeff's periodic table which got removed when elements were arranged as per their atomic numbers.

On moving across a period, the tendency to gain electron(s) increases due to an increase in nuclear pull and a decrease in the atomic size. Therefore, non-metallic character increases across a period and there is a transition from metallic to non-metallic elements.

Noble gases have their outermost or valence shell complete i.e., they have 2 electrons in outermost shell e.g., helium [electronic configuration 2] or they have 8 electrons in outermost shell e.g., neon [electronic configuration 2,8] or argon [electronic configuration 2,8,8 ]. So, they have stable electronic configuration. They do not gain, lose or share electrons and are unreactive or inert. Hence, they do not form compounds readily.

Elements of Group 1 [IA] are metals and react with water to form alkalis, hence are called alkali metals.

1. Group 2 [or Group II-A]
2. Group 17 [or Group VII-A]
3. Group 1 [or Group I-A]
4. Group 18 [or Group 0]
5. Group 17 [or Group VII-A]