Mendeleev's Periodic Table

Mendeleev's Periodic Table

• Elements are classified by increasing atomic weights.
• Periodic law: Element properties are periodic functions of their atomic weights, properties repeat after regular intervals when elements are arranged by increasing atomic weights.
• Elements are arranged in horizontal rows (periods) and vertical columns (groups).
• Elements with similar properties are in the same group.
• Elements were arranged based on similarities in empirical formulae and compound properties.

Merits of Mendeleev’s Periodic Table

• The first comprehensive classification of elements.
• Corrected wrong atomic weights of some elements and placed them correctly.
• Vacant spaces were left for undiscovered elements, with predictions of some properties.
• Elements with similar properties were placed in the same group.

Drawbacks of Mendeleev’s Periodic Table

• Elements with dissimilar properties were in the same group.
• An exact position for hydrogen was not provided.
• An exact position for Lanthanoids and Actinoids, and for isotopes, was not provided.
• Did not strictly obey the increasing order of atomic weights.

Modern Periodic Law

• Physical and chemical properties of elements are periodic functions of their atomic numbers.
• Proposed by Moseley.

Blocks in the Modern Periodic Table

• Four blocks: s block, p block, d block, and f block.

s Block Elements

• Last electron enters the outermost s subshell
• Includes groups 1 and 2 elements.
• General outer electronic configuration is ns1 or ns2.
• Reactive metals with low ionization enthalpies.

p Block Elements

• Last electron enters the outermost p subshell.
• Includes groups 13 to 18 elements.
• Representative elements.
• General outer electronic configuration is ns2np1 to 6.

d Block Elements

• Last electron enters the penultimate d subshell.
• Includes groups 3 to 12 elements.
• Transition elements which show a transition from the most electropositive s block elements to the least electropositive p block elements.
• General outer electronic configuration is (n-1)d1 to 10ns0 to 2.

f Block Elements

• Last electron enters the anti-penultimate f subshell.
• Includes lanthanides of the 6th period and actinides of the 7th period.
• Inner transition elements.
• General outer electronic configuration is (n-2)f1 to 14(n-1)d0 to 1ns2.
• All metals.

Characteristics of Transition Elements

• All metals.
• Form colored compounds or ions.
• Show variable oxidation states and valencies.
• Show paramagnetism and catalytic properties.

Atomic Radius Trends

• Atomic size decreases from left to right in a period because electrons are added to the same shell, but effective nuclear charge increases.
• Atomic radius increases from top to bottom in a group because of increase in the number of shells and shielding effect.

Isoelectronic Species

• Atoms and ions with the same number of electrons, for example, O2-, F-, Ne, Na+, Mg2+ (all contain 10 electrons).
• Cations with greater positive charge have smaller radii.
• Anions with greater negative charge have larger radii.

Ionization Enthalpy

• Energy required to remove an electron from the outermost shell of an isolated gaseous atom in its ground state.
• Units are kJ/mol or J/mol.
• Affected by atomic size, nuclear charge, and shielding effect.

Ionization Enthalpy Trends

• Increases from left to right across a period due to decreasing atomic radius and increasing nuclear charge, alkali metals have the least and noble gases have the most.
• Decreases down a group due to increasing atomic radius and shielding effect, lithium has the least and francium has the most among alkali metals.

Ionization Enthalpy of Boron and Beryllium

• Boron's first ionization enthalpy is slightly less than beryllium's because of the completely filled orbitals in Be (1s22s2).

Ionization Enthalpy of Nitrogen and Oxygen

• Nitrogen's first ionization enthalpy is greater than oxygen's because N has a half-filled electronic configuration (1s22s22p3), which is more stable.

Electron Gain Enthalpy

• Heat change (enthalpy change) when an electron is added to the outermost shell of an isolated gaseous atom.
• Units are kJ/mol and it may be positive or negative.

Electron Gain Enthalpy Trends

• Becomes more negative from left to right across a period due to decreasing atomic radius and increasing nuclear charge.
• Becomes less negative down a group due to increasing atomic radius and shielding effect.

Electron Gain Enthalpy of Halogens

• The negative ∆egH decreases as follows: Cl> F > Br > I.

Electron Gain Enthalpy of Fluorine and Chlorine

• Chlorine has a higher negative electron gain enthalpy than fluorine.
• When an electron is added to F, it enters the smaller 2nd shell, causing more repulsion.
• For Cl, the incoming electron goes to the larger 3rd shell, so the electronic repulsion is low.

Electron Gain Enthalpy of Noble Gases

• Electron gain enthalpy values are zero due to their completely filled (stable) electronic configuration.