PERIODIC TABLE

SCHEME OF WORK

Week 1: Periodic Table

Week 2:Periodic Table

Week 3: Chemical Reactions

Week 4: Chemical Reactions

Week 5: Mass and Volume Relationships in Reactions

Week 6-7: MID TERM

Week 8: Acid Based Reactions

Week 9: Acid Based Reactions

Week 10: Water

Week 11: Water

Week 12: Revision

WEEK 1: 

TOPIC: PERIODIC TABLE

• Periodic Law
• Blocks of Elements
• Families of Elements
• Properties
• Diagonal Relationships
• Ionization Energy and electron affinity changes
  
  
   

Introduction

The periodic table is the table which shows the arrangement of elements in the order of increasing atomic number.

Historical Development of the Periodic Table
In Chemistry, classification of elements began with Lavoiser in 1787, whose classification was essentially into metals and non-metals. Other scientists like Doberainer, Newland and Lother Meyer also attempted to classify elements based on their properties. In 1869, a Russian scientist Dmitri Mendeleev prepared the first periodic table where he arranged elements based on the atomic masses of the elements.

Mendeleev development of the periodic table was based on the periodic Law. His periodic law stated that the properties of the element varied periodically with their relative atomic masses. He arranged the elements in increasing order of relative atomic mass. With this, elements having similar properties kept recurring at regular intervals or periods forming families of related elements.

With discovery of the electronic structure of the electronic structure of teh atoms, it became clear that elements do vary regularly not with their relative atomic mass, but with their atomic number. Hence, in the modern periodic table, classification is based on atomic number, rather than relative atomic mass. Therefore the modern periodic law states that ‘the properties of elements are periodic function of their atomic number’.

FEATURES OF THE PERIODIC TABLE

The modern periodic table is divided into eight vertical columns known as GROUPS and seven horizontal rows known as PERIODS.

 

GROUPS:

The vertical columns of elements or groups are numbered from I to VIII (or 0). Elements in the same group have the same number of electrons in the valence shell. Hydrogen can be placed in group I or VIII because it can donate its one electron like the group I elements or accept like group VII elements, but for convenience and because of its simple valence electron, it is placed in group 1. In group VIII, which is also group 0, Helium (He) has two electrons while the other elements have eight valence electrons. Besides the eight groups, there are also the transition groups of elements. These lay between group II and group III in the periodic table.

PERIODS:

The horizontal rows of elements or periods are numbered 1 to 7. Elements in the same period have the same number of electron shells. Among the elements in period six and seven are the elements of Lanthanides and Actinides series known as inner transition metals.

BLOCKS OF ELEMENTS

The modern periodic table is divided into four main blocks based on the type of atomic orbital being filled with electrons. These blocks are:

s-block, p-block, d-block, and f-block.
This classification is important for understanding the electronic configuration and chemical behavior of elements.

• S-BLOCK ELEMENTS

The s-block elements are those in which the last electron enters the s-orbital of the outermost energy level. They are located on the left side of the periodic table and include Group 1 (Alkali metals) and Group 2 (Alkaline earth metals), as well as hydrogen and helium. The general electronic configuration of s-block elements is ns¹ for alkali metals and ns² for alkaline earth metals, where n represents the principal quantum number.

Alkali metals, such as lithium, sodium, and potassium, are soft, silvery-white metals that are highly reactive. Their reactivity increases down the group because the atomic size increases, making it easier to lose the outer electron. They react vigorously with water to form alkaline hydroxides and hydrogen gas. Alkaline earth metals, such as magnesium and calcium, are harder and less reactive than alkali metals. They also form basic oxides and hydroxides, but react less violently with water.
 

Properties of s-block Elements

1. Highly Electropositive: They easily lose electrons to form positive ions (cations).
2. Strong Reducing Agents: They readily donate electrons in chemical reactions.
3. Low Ionization Energy: Ionization energy decreases down the group, making them more reactive.
4. Metallic Character: Metallic nature increases down the group.
5. Flame Colors: They impart characteristic colors to flames: lithium (red), sodium (yellow), potassium (lilac), calcium (brick red), and barium (green).
6. Formation of Basic Oxides: Their oxides and hydroxides are strongly basic in nature.
   
    

Important compounds of s-block elements include NaOH (used in soap making), CaCO₃ (used in cement and chalk), and NaCl (common salt). These elements are widely used in industries: sodium in street lights, calcium in building materials, and magnesium in alloys and fireworks.

• P-BLOCK ELEMENTS

The p-block elements are those in which the last electron enters the p-orbital of the outermost energy level. They are found on the right side of the periodic table, occupying Groups 13 to 18. These elements include a wide variety of elements such as metals, non-metals, and metalloids. Their general electronic configuration is ns² np¹–⁶, where n is the principal quantum number.

This block includes important families such as the boron family (Group 13), carbon family (Group 14), nitrogen family (Group 15), oxygen family (Group 16), halogens (Group 17), and noble gases (Group 18). Examples include boron, carbon, nitrogen, oxygen, fluorine, and neon.
 

Properties of p-block Elements

1. Diverse Nature: They include metals (e.g., Al), non-metals (e.g., O, N), and metalloids (e.g., Si).
2. Variable Oxidation States: Most elements show different oxidation states due to the availability of both s and p electrons for bonding.
3. Electronegativity Trend: Electronegativity increases across the period and decreases down the group.
4. Covalent Bond Formation: Many p-block elements form covalent compounds because of high electronegativity.
5. Acidic Oxides: Non-metals in this block form acidic oxides (e.g., SO₂, CO₂), while metals form amphoteric or basic oxides.
6. Presence of Allotropes: Elements like carbon (diamond, graphite) and phosphorus (white, red) show allotropy.

Important compounds of p-block elements include CO₂ (carbon dioxide), NH₃ (ammonia), HCl (hydrogen chloride), and SO₂ (sulfur dioxide). Noble gases in Group 18 are inert and used in lighting and welding.

• D-BLOCK ELEMENTS (TRANSITION METALS)
  The d-block elements are those in which the last electron enters the d-orbital of the penultimate shell. They are found in the center of the periodic table and occupy Groups 3 to 12. These elements are also called transition elements because they form a transition between the s-block and p-block elements. Their general electronic configuration is (n−1)d1–10ns1–2.
  Examples include iron (Fe), copper (Cu), zinc (Zn), chromium (Cr), and manganese (Mn). Transition metals are important industrial metals because of their strength and ability to form alloys.
   

Properties of d-block Elements

1. Metallic Nature: All d-block elements are metals.
2. Variable Oxidation States: They show multiple oxidation states due to the involvement of (n-1)d and ns electrons in bonding.
3. Colored Compounds: Many of their compounds are colored because of d-d electronic transitions (e.g., CuSO₄ is blue).
4. Catalytic Properties: Transition metals and their compounds often act as catalysts (e.g., Fe in Haber process).
5. Formation of Complex Ions: They easily form complex compounds with ligands (e.g., [Fe(CN)₆]³⁻).
6. Good Conductors: They conduct heat and electricity well due to the presence of free electrons.
   Examples of uses include iron in steel production, copper in electrical wiring, and platinum as a catalyst.

• F-BLOCK ELEMENTS

The f-block elements are those in which the last electron enters the f-orbital of the atom. They are placed separately at the bottom of the periodic table and are also called inner transition elements. This block consists of two series: the lanthanides (atomic numbers 57–71) and the actinides (atomic numbers 89–103).

Examples of lanthanides include cerium (Ce), neodymium (Nd), and europium (Eu), while examples of actinides include thorium (Th), uranium (U), and plutonium (Pu). Most actinides are radioactive and some are used as nuclear fuels.

Properties of f-block Elements

1. Metallic Nature: All f-block elements are metals.
2. Variable Oxidation States: They show different oxidation states because of similar energies of 4f/5f, 5d, and 6d orbitals.
3. Radioactivity: Most actinides are radioactive (e.g., uranium and thorium).
4. Formation of Colored Compounds: Many compounds are colored due to f-f transitions.
5. High Atomic and Ionic Sizes: They show lanthanide contraction (gradual decrease in size across lanthanides).
6. Complex Formation: They form complex ions due to small size and high charge.

Uses include uranium and plutonium as nuclear fuels, thorium in gas mantles, and lanthanides in strong magnets and electronics.

FAMILIES OF ELEMENTS

Elements in the same group maybe said to belong to a family since they show similar properties because their atoms have the same numer of valence electrons. At the same time, certain properties of the element in the same group show a gradual change with an increase in atomic number. Such gradual change of property within a group is known as GROUP TREND

• GROUP 1:
  Group 1 elements are univalent elements and are known as alkali metals because they form strong alkaline solutions (bases) when they react with water. These elements are found in Group 1 of the Periodic Table and include: Lithium (Li), Sodium (Na), Potassium (K), Rubidium (Rb), Cesium (Cs), and Francium (Fr). All Group 1 elements have one electron in their outermost shell. 
  Alkali metals are very reactive and so do not occur free in nature. They are found in combined form as salts, eg. NaCl (table salt), KNO3  (potassium nitrate)
  
  General Characteristics of Group 1 Elements

These are features that describe alkali metals in general:

1. They are soft metals and can be cut with a knife.
2. They have a shiny metallic luster when freshly cut.
3. They are good conductors of heat and electricity.
4. They have low density (Li, Na, and K float on water).
5. They are very reactive and do not occur free in nature.
6. They are good reducing agents since they can easily donate one electron to form cations.
   
   Physical Properties of Group 1 Elements

These are measurable features of the elements:

1. Atomic size: Increases down the group due to addition of shells.
2. Density: Generally increases down the group, though Li, Na, and K are less dense than water.
3. Melting and boiling points: Decrease down the group.
4. Electrical conductivity: All are good conductors.
5. Softness: Increases down the group (Li is hardest, Cs is softest).
    

Chemical Properties of Group 1 Elements

These describe how alkali metals react with other substances:

1. Reaction with Oxygen: Forms oxides, peroxides, and superoxides:

4Li+O2⟶2Li2O

2K+O2⟶K2O2

1. Reaction with Water: They react vigorously with cold water to liberate hydrogen gas and form alkali, hence are known as ALKALI METALS.
                                                           2Na+2H2O⟶2NaOH+H2↑
2. Reaction with Halogens: Forms ionic halides like NaCl, KCl.
3. Reaction with Acids: Violent reaction, liberating hydrogen.
4. Reactivity trend: Increases down the group (Li < Na < K < Rb < Cs).
5. The oxides of group 1 elements dissolve in water to give very strong alkalis. Example:
   K2O(s) + H2O(l)         ⟶          2KOH(aq)
   NOTE: Reactivity and Atomic radius increases down the group while melting point decreases down the group. Density also increases down the group but with some irregularities.
   
    

• GROUP 2: 
  
  Group 2 elements are divalent elements and are called alkaline earth metals because they form alkaline hydroxides and occur naturally in the earth’s crust. These elements belong to Group 2 of the periodic table and include: Beryllium (Be), Magnesium (Mg), Calcium (Ca), Strontium (Sr), Barium (Ba), and Radium (Ra). All Group 2 elements have two elements in their outermost shell. 
  
  General Characteristics of Group 2 Elements

These are general features of alkaline earth metals:

1. They are harder and denser than alkali metals.
2. They have a shiny metallic luster when freshly cut.
3. They are good conductors of heat and electricity.
4. They have higher melting and boiling points than Group 1 metals.
5. They are less reactive than alkali metals, but reactivity increases down the group.

Physical Properties of Group 2 Elements

These are measurable features of the elements:

1. Atomic size: Increases down the group (Be < Mg < Ca < Sr < Ba < Ra).
2. Density: Increases down the group, though not regularly.
3. Melting and boiling points: Generally decrease down the group.
4. Electrical conductivity: All are good conductors.
5. Hardness: Harder than alkali metals, but softness increases down the group.

Chemical Properties of Group 2 Elements

These describe their chemical behavior:

1. Reaction with Oxygen: Forms oxides:
   2Ca+O2⟶2CaO
2. Reaction with Water:
   • Be and Mg react slowly with water (Mg reacts with steam).
   • Ca, Sr, Ba react readily with cold water:
     Ca+2H2O⟶Ca(OH)2+H2↑
3. Reaction with Acids: Reacts with dilute acids to liberate hydrogen gas.
4. Reducing Property: They ionize by donating their two valence electrons, hence, they are good reducing agents.
5. Formation of Hydroxides: Forms strong bases that are less soluble than alkali hydroxides.
6. Reactivity trend: Increases down the group (Be < Mg < Ca < Sr < Ba).

Uses of Group 2 Elements

• Magnesium (Mg): Used in fireworks, flares, and making light alloys.
• Calcium (Ca): Used in cement and as a reducing agent.
• Barium (Ba): Used in X-ray contrast solutions.
• Radium (Ra): Formerly used in luminous paints (now mostly discontinued).
  
   
• GROUP 3:
  The Group 3 elements of the Periodic Table are Boron (B), Aluminium (Al), Gallium (Ga), Indium (In), and Thallium (Tl). They are trivalent elements. These elements belong to Group 13 in the modern periodic table. They have three electrons in their outermost shell (ns²np¹), which determines their chemical behavior. The first element, boron, is a metalloid, while the others are metals. This group shows a mixture of properties because of the gradual transition from non-metal to metal as we go down the group.
  
  Characteristics of Group 3 Elements
• They all have three valence electrons.
• Boron is a non-conductor, while others are good conductors of heat and electricity.
• Metallic character increases down the group, making boron the least metallic and thallium the most.
• Their atomic and ionic sizes increase down the group, while ionization energy decreases.
  
  Properties of Group 3 Elements
• Physical Properties:
  
   
  1. They are good conductors of heat and electricity (except boron, which is a poor conductor).
  2. Melting points decrease down the group due to weakening metallic bonds.
  3. Boron is hard and brittle, while aluminium is light and malleable.
     
      
• Chemical Properties:
  1. Oxidation States: Common oxidation state is +3; Tl also shows +1 due to the inert pair effect.
  2. Electropositivity: Increases down the group as metallic character increases.
  3. Reducing Power: Increases down the group because elements lose electrons more easily.
  4. Nature of Oxides: Boron oxide (B₂O₃) is acidic, aluminium oxide (Al₂O₃) is amphoteric, and others are more basic as we go down the group.
  5. Formation of Complex Compounds: Aluminium forms many complexes (e.g., [AlF₆]³⁻).
  6. Reactivity: Aluminium reacts with acids and alkalis due to its amphoteric nature, while boron does not react easily because it is non-metallic.
     
     Uses of Group 3 Elements
• Boron: Used in borosilicate glass and as a neutron absorber in nuclear reactors.
• Aluminium: Used for making utensils, foils, and aircraft parts.
• Gallium: Used in LEDs and solar cells.
• Indium: Used in touchscreens and semiconductors.
• Thallium: Used in optical glasses and electronic devices.
  
   
• GROUP 4: 
  Group 4 elements belong to Group 14 of the Modern Periodic Table and include Carbon (C), Silicon (Si), Germanium (Ge), Tin (Sn), and Lead (Pb). These elements have four valence electrons in the outer shell, with a general electronic configuration of ns²np². Carbon and silicon are non-metals/metalloids, germanium is a metalloid, while tin and lead are metals. As we move down the group, metallic character increases while non-metallic character decreases.

Characteristics of Group 4 Elements

• All have valency of 4, but they can exhibit +2 and +4 oxidation states.
• Carbon and silicon are covalent, while tin and lead show metallic properties.
• Electronegativity decreases down the group.
• The ability to form catenation (bonding with itself) is strongest in carbon and decreases down the group.
  
  Properties of Group 4 Elements
  Physical Properties
• Carbon exists in various forms (diamond, graphite, fullerene) and is a good covalent network solid.
• Silicon is a hard, brittle solid and a semiconductor.
• Tin and lead are soft metals with good conductivity.
• Melting points are high for carbon and silicon but decrease as we move to metals.
  
   

Chemical Properties

• Oxidation States: They show +4 and +2 oxidation states. +4 is common in lighter elements (C, Si), while +2 becomes stable in heavier elements (Pb) due to the inert pair effect.
• Nature of Oxides: Carbon dioxide is acidic; silicon dioxide is acidic; tin and lead oxides are amphoteric.
• Catenation: Strong in carbon, weak in lead.
• Reducing/Oxidizing Nature: +4 compounds are oxidizing for heavier elements; +2 compounds are reducing.
• Reactivity: Carbon reacts with oxygen to form CO and CO₂, silicon reacts slowly, while tin and lead react with acids and alkalis.
• Reaction with Water: Carbon does not react with water in any form, but silicon and tin react with steam at red heat to form +4 state oxides and hydrogen.
   

Uses of Group 4 Elements

• Carbon: Used in fuels, steel production, graphite for electrodes, and diamonds in cutting tools.
• Silicon: Used in glass, ceramics, and as a semiconductor in electronics.
• Germanium: Used in semiconductors and fiber optics.
• Tin: Used in coating food cans and in alloys like bronze.
• Lead: Used in batteries and radiation shielding (though toxic).

• GROUP 5:
  Group 5 elements belong to Group 15 of the Modern Periodic Table and include Nitrogen (N), Phosphorus (P), Arsenic (As), Antimony (Sb), and Bismuth (Bi). These elements have five valence electrons in their outer shell with the general electronic configuration ns²np³. The group shows a gradual change from non-metallic to metallic character as we move down the group: nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids, while bismuth is a metal.
  
  Characteristics of Group 5 Elements
• All elements in the group have valency of 3 or 5 due to the presence of five outer electrons.
• They can form both covalent and ionic bonds.
• Nitrogen and phosphorus exist as discrete molecules (N₂, P₄), while others form giant structures.
• The group shows increasing metallic character down the group.
  
  
  
  
  Properties of Group 5 Elements

Physical Properties

• Nitrogen is a colorless, odorless gas, while phosphorus is a solid that exists as white, red, and black allotropes.
• Arsenic and antimony are brittle solids, and bismuth is a soft metal.
• Melting and boiling points increase down the group, except nitrogen which is gaseous at room temperature.
  
   

Chemical Properties

• Oxidation States: Common oxidation states are –3, +3, and +5. Stability of +5 decreases while +3 increases down the group due to the inert pair effect.
• Nature of Oxides:
  • Nitrogen oxides are acidic.
  • Phosphorus oxides are acidic.
  • Arsenic and antimony oxides are amphoteric.
  • Bismuth oxide is basic.
• Reactivity: Nitrogen is relatively inert due to its triple bond but forms compounds like ammonia (NH₃). Phosphorus is very reactive, especially white phosphorus, which catches fire in air.
• Reducing/Oxidizing Behavior: Nitrogen compounds can be strong oxidizing agents since they are electron acceptors, while heavier elements act as reducing agents.
  
  Uses of Group 5 Elements
• Nitrogen: Used in making ammonia for fertilizers and in liquid form as a coolant.
• Phosphorus: Used in matches, fertilizers, and detergents.
• Arsenic: Used in pesticides and semiconductors.
• Antimony: Used in flame-proofing materials and alloys.
• Bismuth: Used in medicines and cosmetics.

• GROUP 6:
  Group 6 elements belong to Group 16 of the Modern Periodic Table and include Oxygen (O), Sulphur (S), Selenium (Se), Tellurium (Te), and Polonium (Po). These elements are called Chalcogens because they are often found in ores (from the Greek word chalcos, meaning ore). They have six valence electrons in their outermost shell with a general electronic configuration of ns²np⁴. The group shows a transition from non-metallic to metallic nature as we move down the group: oxygen and sulphur are non-metals, selenium and tellurium are metalloids, and polonium is a metal.
  
  Characteristics of Group 6 Elements
• All elements have valency of 2, 4, or 6, depending on the compound.
• Oxygen is a colourless gas, while sulphur is a yellow solid.
• Non-metallic character decreases down the group, while metallic character increases.
• They have a strong tendency to form covalent compounds.
  
  Properties of Group 6 Elements

 Physical Properties

• Oxygen exists as a diatomic gas (O₂), while sulphur exists as a solid in various allotropes (e.g., rhombic and monoclinic sulphur).
• Melting and boiling points increase down the group.
• They exist as solids at room temperature except for oxygen.
• Density and atomic size also increase as we move down the group.
  
   

        Chemical Properties

• Oxidation States: Common oxidation states are –2, +4, and +6. The negative oxidation state (–2) is more stable in oxygen, while the positive oxidation states (+4, +6) are more common in heavier elements like sulphur, selenium, and tellurium.
• Nature of Oxides:
  • Oxides of oxygen (O₂) are neutral.
  • Oxides of sulphur and selenium are acidic.
  • Tellurium oxides are amphoteric.
  • Polonium oxides are basic.
• Reactivity: Oxygen is highly reactive, supporting combustion. Sulphur reacts with metals and non-metals, forming compounds like SO₂ and SO₃.
• Electronegativity: Oxygen is the most electronegative after fluorine, and electronegativity decreases down the group.
• Reaction with Hydrogen: Oxygen and Sulphur combine directly with hydrogen to yield water and hydrogen sulphide respectively.
  
  Uses of Group 6 Elements
• Oxygen: Used for respiration, welding, and in hospitals for patients.
• Sulphur: Used in making sulphuric acid (H₂SO₄), matches, and vulcanization of rubber.
• Selenium: Used in photocells and glass manufacturing.
• Tellurium: Used in alloys and semiconductors.
• Polonium: Used as a heat source in space satellites (radioactive).
  
  
   
• GROUP 7:
  Group 7 elements belong to Group 17 of the Modern Periodic Table and include Fluorine (F), Chlorine (Cl), Bromine (Br), Iodine (I), and Astatine (At). These elements are called Halogens (meaning “salt formers”) because they readily form salts with metals. They have seven valence electrons in their outermost shell with a general electronic configuration of ns²np⁵. All halogens are highly reactive non-metals, and reactivity decreases down the group.
  
  Characteristics of Group 7 Elements
• Each halogen needs one electron to complete an octet, making them strong oxidizing agents. They ionize to form univalent anions.
• They exist as diatomic molecules (F₂, Cl₂, Br₂, I₂).
• Physical states vary: Fluorine and chlorine are gases, bromine is a liquid, and iodine and astatine are solids.
• Color deepens down the group: fluorine is pale yellow, chlorine is greenish-yellow, bromine is reddish-brown, iodine is violet.
  
   

Properties of Group 7 Elements
Physical Properties

• Low melting and boiling points compared to metals, but these increase down the group.
• Poor conductors of heat and electricity.
• They have a strong, irritating smell.
  
   

      Chemical Properties

• Oxidation State: Commonly –1 in compounds, but positive states like +1, +3, +5, +7 occur in oxoacids and interhalogen compounds.
• Reactivity: Fluorine is the most reactive halogen; reactivity decreases from F → Cl → Br → I → At.
• Nature of Oxides: Mostly acidic.
• Displacement Reaction: A more reactive halogen displaces a less reactive one from its compound (e.g., Cl₂ can displace Br₂ and I₂ from their salts).
• Halogenation: They react with metals to form halides (e.g., NaCl, KBr).
• Hydrogen Compounds: All form hydrogen halides (HF, HCl, HBr, HI), which dissolve in water to give acids.
  
  Uses of Group 7 Elements
• Fluorine: Used in toothpaste (fluoride) and in Teflon production (polytetrafluoroethylene - PTFE).
• Chlorine: Used in disinfectants, bleaching powder, and water treatment.
• Bromine: Used in dyes and flame retardants.
• Iodine: Used as an antiseptic and in iodized salt.
• Astatine: Rare and radioactive, used in medical research.
  
  
  
   
• GROUP 8:
  Group 8 elements belong to Group 18 of the Modern Periodic Table and are known as Noble Gases or Inert Gases. They include Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), and Radon (Rn). These gases are called noble because they are generally unreactive under normal conditions due to having a complete outer electron shell. Their general electronic configuration is ns²np⁶ (except helium, which is 1s²).
  
  Characteristics of Noble Gases
• They exist as monoatomic gases at room temperature.
• They have very low boiling and melting points because they are held together by weak Van der Waals forces.
• They are colorless, odorless, and tasteless.
• They are the least reactive elements in the periodic table.
  
  Properties of Noble Gases

Physical Properties

• Gases at room temperature.
• Extremely low density (Helium is lighter than air).
• Poor conductors of heat and electricity.
   

   Chemical Properties

• They are generally chemically inert because of a stable electronic configuration (octet).
• Helium and neon do not form compounds.
• Xenon and krypton can form compounds with highly electronegative elements like fluorine and oxygen under special conditions (e.g., XeF₂, XeO₃).
  
  Uses of Noble Gases
• Helium: Used in balloons, airships, and as a coolant in superconductors.
• Neon: Used in neon advertising signs.
• Argon: Used in electric bulbs to prevent oxidation of filaments.
• Krypton: Used in special light bulbs and lasers.
• Xenon: Used in camera flash lamps and anesthesia.
• Radon: Used in cancer treatment (radioactive).
  
  
  PERIODIC TABLE SUMMARY (GROUP 1 - 8)
  
• TRANSITION ELEMENTS
  Transition elements are the elements found in Groups 3 to 12 of the Modern periodic table. They are also called transition metals because they form a transition between the highly reactive s-block metals and the p-block elements. Their general electronic configuration is:    (n−1)d1−10ns1−2 
  
  Characteristics of Transition Elements

1. Variable oxidation states – They can lose different numbers of electrons (e.g., Fe²⁺ & Fe³⁺).
2. Formation of colored compounds – Due to d-d electron transitions.
3. Good conductors of heat and electricity – Like copper and silver.
4. High melting and boiling points – Strong metallic bonding.
5. Form complex ions – e.g., [Fe(CN)6]3.
6. Exhibit catalytic properties – e.g., Fe in Haber process, V₂O₅ in Contact process.
7. Magnetic properties – Many are paramagnetic due to unpaired d-electrons.
   
   Examples of Transition Elements

• First transition series: Sc, Ti, V, Cr, Mn, Fe, Co, Ni, Cu, Zn.
• Second transition series: Y, Zr, Nb, Mo, Tc, Ru, Rh, Pd, Ag, Cd.
• Third transition series: Hf, Ta, W, Re, Os, Ir, Pt, Au, Hg.
  
  Uses of Transition Elements
• Iron (Fe): Steel production.
• Copper (Cu): Electrical wiring.
• Chromium (Cr): Stainless steel, electroplating.
• Manganese (Mn): Steel hardening.
• Nickel (Ni): Alloy production.
• Platinum (Pt): Catalysts, jewelry.
  
  
   
• Lanthanides (Lanthanons)
  
  Lanthanides are a group of 15 chemical elements with atomic numbers 57 to 71, starting from Lanthanum (La) to Lutetium (Lu). They belong to the f-block of the periodic table and are also called rare earth elements and they resemble each other greatly. Their general electronic configuration is:      [Xe]4f1−145d0−16s2
  
  Characteristics of Lanthanides

1. Silvery-white metals – Soft and can be cut with a knife.
2. High reactivity – Especially when finely divided.
3. Lanthanide contraction – Gradual decrease in atomic and ionic sizes across the series due to poor shielding of 4f electrons.
4. Variable oxidation states – Mostly +3, but some show +2 and +4.
5. Strong reducing agents – Due to low standard electrode potentials.
6. Formation of colored ions and complexes – Due to f–f electronic transitions.
7. High melting and boiling points – Metallic bonding is strong.
   
   Uses of Lanthanides

• Neodymium (Nd): Used in making powerful magnets.
• Cerium (Ce): Used in catalytic converters and polishing powders.
• Lanthanum (La): Used in optical lenses and camera glass.
• Europium (Eu): Used in fluorescent lamps and TV screens.
• General: Used in nuclear reactors and high-strength alloys.
  
  
   
• Actinides

Actinides are a series of 15 elements with atomic numbers 89 to 103. They are found in the seventh period starting from Actinium (Ac) to Lawrencium (Lr). The famous Uranium is in this group. They belong to the f-block of the periodic table, specifically the 5f series, and most are radioactive. Their general electronic configuration is:   [Xe]4f145d106s26p65f1−146d0−17s2

The elements with atomic numbers from 93 to 103 are known as artificial elements. This is because they do not occur naturally but were formed during nuclear reactions.

Characteristics of Actinides

1. Radioactive nature – All actinides are radioactive, and some are highly unstable.
2. Metallic character – They are soft metals with silvery appearance.
3. Variable oxidation states – Common states are +3, +4, +5, and +6.
4. Formation of colored ions – Due to f–f electronic transitions.
5. Strong tendency to form complexes – Especially with oxygen and halogens.
6. Most are synthetic – Except Actinium, Thorium, Protactinium, and Uranium, which occur naturally.
7. High radioactivity and toxicity – Require careful handling.

Uses of Actinides

• Uranium (U): Used as a fuel in nuclear reactors and in atomic bombs.
• Thorium (Th): Used in nuclear reactors and in strong alloys.
• Plutonium (Pu): Used in nuclear weapons and reactors.
• Americium (Am): Used in smoke detectors.
• General: Used in nuclear energy production and research.