Mendeleev's Periodic Table

• Elements cannot be broken down into simpler forms through chemical reactions.
• There are 118 known elements in the modern periodic table.

Discovery of Elements

• Initially, with few elements discovered, studying individual properties was feasible.
• As element discovery progressed, studying individual properties became nearly impossible.
• Scientists systematically classified elements in a tabular form, grouping similar elements.
• Grouping elements represents properties, leading to the periodic table.

Mendeleev's Periodic Table

• Dmitri Ivanovich Mendeleev, a Russian chemist, designed this table in 1869.
• Mendeleev's table was the first scientific and systematic classification of elements.
• Mendeleev is known as the "father of the periodic table."
• Mendeleev formulated the periodic law: "the physical and chemical properties of an element are the periodic function of their atomic mass/weight."
• Elements arranged by increasing atomic mass exhibit repeating properties at regular intervals.

Shortcomings of Mendeleev's Periodic Table

• Position of isotopes: Hydrogen has three isotopes (protium, deuterium, tritium) with different mass numbers, leading to an inconsistent position.
• Violation of periodic law: Certain elements were placed out of order of their atomic mass.
  • Argon (Ar) precedes potassium (K) despite having a greater atomic mass
  • Cobalt (Co) precedes nickel (Ni) despite having a smaller atomic mass.
• Separation of elements: Similar elements were separated (copper and mercury), while dissimilar elements were grouped together (copper, gold, silver with potassium & Cesium).
• Cause of periodicity: Mendeleev's table did not explain the underlying cause of periodicity.

Modern Periodic Table

• Henry Moseley, a British scientist, introduced this table in 1913 to address Mendeleev's shortcomings.
• Moseley formulated the modern periodic law: "the physical and chemical properties of an element are the periodic function of their atomic number."
• The modern periodic table classifies elements into groups, subgroups, periods, and blocks in a scientific manner.

Groups

• Groups are the vertical columns in the periodic table.
• Elements within the same group have similar electronic configurations and properties.
• The modern periodic table has 18 groups numbered I-VIII and zero,.
• Except for group eight and zero I-III group is divided into two sub-groups A and B.
• Group VIII consists of three vertical columns.
• The last group is called zero group and consists if inert gasses.
• According to the latest IUPAC recommendation in 1990, the elements are numbered from group 1-18.
• Groups 1 and 2 are on the extreme left and are known as alkali metals and alkaline earth metals, respectively.
• The right side of the periodic table consists of 6 groups (13-18).
• Groups 13, 14, 15, 16, 17, and 18 are known as Boron, Carbon, Nitrogen, Oxygen, Halogens and inert gases.
  • Nitrogen is known as Pnicogens.
  • Oxygen is known as Chalcogens.
• Groups 3-12 are called transition elements.

Periods

• Periods are the horizontal rows of the periodic table.
• There are 7 periods in the modern periodic table.
• Period 1 contains of two elements (H and He) and is called a very short period.
• Period 2 is short and consist of 8 elements, from Li to Ne,.
• Period 3 is short and consist of 8 elements, from Na to Ar,.
• Period 4 consists of 18 elements, from K to Kr, and is known as a long period.
• Period 5 consists of 18 elements, from Rb to Xe, and is known as a long period.
• Period 6 consists of 32 elements, from Cs to Rn, and is known as a very long period.
• Period 7 is incomplete, containing newly discovered elements (32 elements from Fr to Og).

Blocks

• Blocks are classified based on the atomic orbitals where the differentiating electron enters.

S-block Elements

• These elements have their differentiating electron entering the s-orbital of the outermost shell.
• The s-block consists of elements from Groups 1 and 2, known as alkali and alkaline earth metals, respectively.
• They are located on the extreme left side of the periodic table.
• Characteristics of s-block elements include:
  • General outermost electronic configuration of ns1-2 (ns1 for alkali metals, ns2 for alkaline earth metals)
  • They are reactive metals.
  • They are soft metals.
  • They have low melting and boiling points.
  • They form ionic compounds.
  • They are electropositive and form univalent (Na+, K+) and bivalent ions (Be++, Mg++).
  • The s-block elements impart characteristic colors to a flame.

P-block Elements

• The differentiating electron enters the p-orbital of their outermost shell.
• This block consists of 6 groups from Group number (13-18).
• Located on the extreme right side of the periodic table.
• General outermost electronic configuration is (ns²np¹⁻⁶).
• Consists of metals, non-metals, metalloids, heavy metals, and inert gasses.
• It mostly forms covalent compounds, however halogens form ionic compounds with alkali metals and alkaline earth metals.
• Mostly electronegative in nature and shows variable oxidation states.
• Examples of p-block elements are B, Al, C, Si, O, N, F, Cl, Ne, Si, Ge, Pb.

D-block Elements

• Differentiating electron enters the d-orbital of their outermost shell (penultimate shell).
• Consists of elements from Groups 3-12.
• It consist of metals.
• Examples of d-block elements are Fe, Co, Ni, Zn, Cu, Ag,.
• General outermost electronic configuration is (n-1)d¹⁻¹⁰ ns¹⁻².
• They are also called 'transition elements' because they represent a transition property between high-electropositive s-block and highly electronegative p-block elements.
• They have partially filled d-orbitals.
• They have variable oxidation states.
• They formed coloured compounds.
  • KMnO₄ (Potassium Permanganate) is purple.
  • K₂Cr₂O₇ is orange (Potassium dichromate).
• They have high melting and boiling points.
• They are successfully used as catalysts in various reactions.
  • Tron is used as catalyst in the manufracture of ammonia by Haber's process.
  • Divandium pentoxidel is used as a catalyst in the manufracture of sulfuric acid by contact process.
• They forms alloys.
  • Brass: (Cu+Zn)
  • Steel: (Fe+Cr+C).

F-block Elements

• Differentiating electrons enter the f-orbital of the anti-penultimate shell.
• There are 2 series of f-block elements, Lanthnides and actinides.
• Situated below the mainframe of the periodic table.
• f-block elements are also known as inner-transition elements.
• Lanthnides: The series of 14 elements occurring after lanthanum i.e. from ₅₈Ce to ₇₁Lu.
  • They are also called rare earth elements.
• Actinides: The series of 14 elements occurring after Actinium i.e. from ₉₀Th to ₁₀₃Lr
  • They are also known as radioactive elements.
• Their general outermost electronic configuration is (n-2) f¹⁻¹⁴ (n-1)d⁰⁻¹ ns².
• They have high melting and boiling points.
• They are all heavy metal.
• They form coloured salts.
• They have variable oxidation states.

Advantages of the Modern Periodic Table

• The long form is based on the modern periodic law, classifying elements by atomic number.
• Isotopes are justified by their atomic number.
• The table properly addresses Mendeleev's misfit points by choosing atomic number.
• The elements are easily studied via classification into four blocks.
• Elements within sub-groups are separated based on chemical properties.
• Metals and non-metals are separated.

Defects of Modern Periodic Table

• Inability to include Lanthanum (La) and Actinium (Ac) in its mainframe.
• Hydrogen's position remains controversial.
• Helium's position is not fully justified.

IUPAC Element Classification

• According to the latest IUPAC recommendation in 1990, the long form of the periodic table has been modified.
• The groups are numbered 1-18 from left to right across the periodic table.

Blocks and Corresponding Groups

• s-block: Groups 1 and 2
• p-block: Groups 13-18
• d-block: Groups 3-12
• f-block: No separate group

Periodic Trends and Periodicity

• Nuclear charge: The total positive charge on the nucleus.
  • Denoted by Zeff, it's always positive and depends on the number of proton in a nucleus.
• Effective nuclear charge: The actual charge experienced by the valence shell electron in an atom.
  • It is calculated by considering the effect inner.
• In multinuclear electron atoms are pulled by the nucleus and repelled by electron of inner shell, the outermost electron exprienced less attraction from the nucleus
• Shielding or screening effect: The decrease in nuclear attraction on the valence electron due to the presence of electron in the same shell or inner shell.
• Various properties of elements are atomic radii, ionic radii, ionization energy, electron affinity, electronegativity, and metallic character.
  • directly related to their electronic configuration; melting point, boiling point, and density are indirectly related
• Properties directly or indirectly related to electronic configuration are called atomic properties whose variation along a group and period is periodic

Atomic Radii (AR)

• Defined as the distance between the center of the nucleus and the outermost shell of electrons.
• Atoms are tightly held to the nucleus, therefore the size decreases the ionic radii.

Variation of Atomic Radii in the Periodic Table

• In a group, moving from top to bottom, the number of shells increases causing an increase in the atom size.
• In a period, moving from left to right, nuclear charge increases, and electrons are more strongly attracted due to which the size of atom decreases.

Ionic Radii (IR)

• Defined as the effective distance from the center of the nucleus of an ion to the point up to which it effects its electron cloud
• Size of cation is always smaller than respective parent atom whereas the size of an anion is larger than the respective parent atom

Isoelectronic ions

• Isoelectronic ions or species have the same number of electrons and decreases with increase in nuclear charge.

Variation of Atomic/Ionic Radius

• In a group, ionic radius increases because of increase in number of shells by moving from top to bottom,.
• In a period, nuclear charge increases as electron are more tightly held as you go from left to right.

Ionization Energy/Ionization Potential

• The minimum energy needed to remove the most loosely bound electron (valence electron) from an isolated gaseous atom in its ground state.
• Expressed in electron volts (eV) or kJ/mol.

Successive Ionization Energy

• The energy required to remove subsequent electrons.

Factors affecting ionization energy (I.E)

• As atomic size increases, ionization energy decreases because larger the distance of outer electron from the nucleus is.
• ionization energy increase with an increase in nuclear charge because there is more nuclear attraction.
• The element configration influences ionization energy because the more stable the electronic.

Variation of ionization in a periodic table

• the ionization group energy gradually goes on decreasing as you move from top to bottom
• On moving from top to bottom in a group because of the following reason
  • There is increased atomic size and increasing the number of shell in the electron.
  • Increase in nuclear charge which causes the increasing sheilding effect on the outermost electron.

Variation in a Period

• In general ionization energy increases with increase in atomic number
• combined effect of increase nuclear charge is equal decrease in atomic size.

Electro Affinity (E A

• Electro affinity is is the minimum amount of energy released where an electron is added to an isolated gaseous item
• EA is measured in term of eV per atom or KJ/mol
• In EA the value of EA is always negative

Factors Affecting EA

• Atomic size: As the size of atom increases the distance between the nucleus and the incoming electron increases.
• Nuclear charge: Greater the Nuclear charge, greater will be the electron affinity.
• Electronic Configuration: Electrons having stable E.C like half-filled and completely filled orbital have E A

Variation of E.A in a Periodic Table

• In a group: On moving from top to bottom in a group, the atomic size and nuclear charge increases,.

Intresting Trends of E.A

• Halogens has the highest E.A- The eC of the is to required the halogens in stable tendency

Electronegativity

• Is the tendency of an atom in a molecule to attract the shared-pair of e-toward itself.

Factor Affecting Electronegativity

• Atomic size: the tendency of atom to effect the pairing of the shell increases with decreasing size
• No. of inner shell: Atom with larger number of inner shells has lower value of electronegativit

Variation of Electronegativity

• Electronegativit: an element decreases downt eh group this is because of increasing number of sheels consequently the shell
• In variation in periods, it increases from left to right.

Metallic Character

• Is can be defined as the tendency of an element ti loose electrons ti form positive charged Cation.
• Fore.g - block element have high metallic character.

Variation of Metallic Character

• Increases tendary of electron from top to bottom