ICSE Chemistry- Grade 10- Chapter 1- “Periodic Properties and Variations of Properties – Physical & Chemical”

Chapter 1- “Periodic Properties and Variations of Properties – Physical & Chemical”

  

MENDELEEV’S PERIODIC LAW: “The physical and chemical properties of elements are periodic functions of their atomic weights.”

MODERN PERIODIC LAW: “The physical and chemical properties of elements are periodic functions of their atomic numbers.”

Cause of Periodicity: Similar electronic configuration or the same number of valence electrons.

LONG FORM OF PERIODIC TABLE

(i)       There are 18 VERTICAL COLUMNS (Groups) in the periodic table.

Ø  All the elements in the same group should have similar electronic configurations.

Ø  Elements in GROUP 1 and GROUP 2 placed to the extreme left of the periodic table are METALS.

Ø  Elements in GROUP 13 TO 17 placed to the right side of the periodic table are NON-METALS.

Ø  Elements in GROUP 18 are INERT GAS or NOBLE GAS elements.

Ø  Elements in GROUP 3 TO GROUP 12 are transition elements which generally exhibit intermediary properties.

(ii)     There are Seven Horizontal Rows (Periods) in the periodic table. 

Ø  Each period represents a Different Principal Energy Level (shell). In each element, valence electrons are placed in the principal energy level (valence orbit).

PERIODIC TRENDS IN PROPERTIES

1.         Atomic Radius:

Ø  Size INCREASES as you go down a GROUP as new shells are added.

Ø  Size DECREASES as you go across a PERIOD as the nuclear charge increases.

Isoelectronic Series: The size of the atom decreases as the number of protons increases.

Ø  Size of a cation is less than the parent Atom.

Ø  Size of the anion is more than the parent Atom.

2.         Metallic Character:

Ø  Metallic character DECREASES while non-metallic character INCREASES in a PERIOD.

Ø  Metallic character INCREASES while non-metallic character DECREASES in a GROUP.

3.         Ionization Energy Or Ionization Enthalpy Or Ionization Potential

“Ionization potential is defined as the minimum amount of energy that is required to remove the most a loosely bound electron from an isolated gaseous atom in its ground state so as to convert it into a gaseous cation.”

Ø  The ionization energy INCREASES IN A PERIOD as the atomic numbers increase.

Ø  The ionization energy DECREASES in a GROUP as the atomic size increases.

Order of successive Ionization Energy (IE):  IE₁ <IE₂ <IE₃

Factors influencing ionization energy

(i)       Atomic number                   

(ii)     Atomic size                          

(iii)    Electronic configuration

4.         Electron Gain Enthalpy Or Electron Affinity:

“Electron affinity an element may be defined as the energy released when a neutral isolated gaseous atom accepts an extra electron to form a gaseous anion, i.e. a negative ion.”

Ø  The electron gain enthalpy INCREASES on moving left to right in a period.

Ø  The electron gain enthalpy DECREASES on moving down the group.

5.         Electronegativity(EN):

Ø  INCREASES across the PERIOD from left to right.

Ø  DECREASES down the GROUP from top to bottom

GROUP 1: THE ALKALI METALS

·         Atomic Radii: Increases down the group from top to bottom.

·         Ionic Radii: Increases on moving down the group from top to bottom.

·         Soft metals:  Softness increases with increase in atomic number.

·         Melting and Boiling Points: Low melting and boiling points due to weak cohesive forces.

·         Flame Coloration: Reason of flame coloration is low ionization energy.

Element	Colour
Lithium (Li)	Crimson Red
Sodium (Na)	Golden Yellow
Potassium (K)	Pale Violet
Cesium (Cs)	Blue

·         The alkali metals are highly reactive elements: It increases from Li to Cs.

·         Hydrides. Alkali metals react with dry hydrogen to form ionic metallic hydrides.

LiH + H₂O            →      LiOH + H₂           

NaH + H₂O          →    NaOH + H₂

These hydrides are strong reducing agents and their reducing nature increases down the group.

·         Oxides. Alkali metals react with air and form BASIC OXIDES.                                   

 Basic Nature of Alkali Metal Hydroxides: Strong bases and highly soluble in water and stable to heat. Exception: Lithium hydroxide.

                                               

·         Reaction with water:  Alkali metals react with water and the other compounds containing acidic hydrogen atoms such as hydrogen halides (HX) liberate hydrogen gas.

2Na + 2H₂O → 2NaOH + H₂

2Na + 2HX → 2NaX + H₂      (where X = halogen)

·         Reaction with Halogens. Alkali metals react with halogens to form metal halides, which are ionic crystal solids having general formula M⁺X^-.

                                                               

·         Reaction with Non-Metals. Alkali metals, on heating, react with non-metals like sulfur and phosphorus to form sulfides and phosphides respectively.

2M + S     →    M₂S

3M + P     →     M₃P

GROUP 17 ELEMENTS – Halogen Family

·         Colour:   F₂ (Pale yellow), Cl₂ (greenish-yellow), Br₂ (dark red), I₂ (Violet)

·         Nature of Bonds: With metals - ionic halides. With non-metals - covalent compounds.

·         Ionisation Energy: Very high values of IE₁. These IE₁ values decrease in moving down the group.

·         Electronegativity: Most electronegative elements within their respective periods. Electronegativity decreases in moving down the group as given below:

F(4.0); Cl(3.0); Br(2.9); I(2.7)

·         Electron Affinity:  Highest electron affinities within their periods. Chlorine shows an exceptionally higher value than fluorine. The order of EA₁ is  Cl > F > Br > I

·         Hydrides:  

(i)       All halogens form hydrides with formula HX (HF, HCl, HBr, HI).

(ii)     Acidic nature of these hydrides is HI > HBr > HCl > HF

·         Oxidizing power: Strong oxidizing agents and the oxidizing power decrease down the group from fluorine to Iodine.

COMPARISON BETWEEN ALKALI METAL AND HALOGEN

S.No	Property	Alkali metal	Halogen
1	Name of element	Lithium(Li), Sodium(Na) Potassium(K), Rubidium(Rb) Cesium(Cs), Francium(Fr)	Fluorine(F), Chlorine(Cl) Bromine(Br), Iodine(I) Astatine(At)
2	Physical state	Solid	F,Cl–gas; Br-liquid; I-solid
3	Atomic number	Li – 3; Na – 11; K – 19; Rb -37; Cs – 55; Fr – 87	F- 9; Cl- 17; Br- 35; I -53; At -85
4	Electronic configuration	Li – 2,1 Na – 2,8,1 K – 2,8,8,1 Rb – 2,8,18,8,1	F- 2,7 Cl- 2,8,7 Br- 2,8,18,7 I -2,8,18,18,7
5	Valence electron	1	7
6	Ion formation	Forms cation (+ve ion)	Forms anion (-ve ion)
7	Electrovalence	+1	-1
8	Type of compounds	Electrovalent	Electrovalent and covalent
9	Atomic size	Highest in the period Increases from Li to Fr	Lowest in the period Increases from F to At
10	Thermal conductivity	Good conductor	Bad conductor
11	Electrical conductivity	Good conductor	Bad conductor
12	Ionization energy	Lowest in the period	Highest in the period
13	Electron affinity	Lowest in the period	Highest in the period
14	Electronegativity	Lowest in the period	Highest in the period
15	Metallic/non-metallic character	Metallic nature	Non – metallic nature
16	Nature of oxide	Basic	Acidic
17	Reducing agent	Strong reducing agent	Weak reducing agent

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