Introduction to Periodic Table
The Symbol of Chemistry
One of the most important turning points in the history of science was the creation of the periodic table, which led to many important innovations. It is accurate to refer to the periodic table of elements as the “Symbol of Chemistry.” In the modern periodic table, 118 elements are arranged in tabular form based on their atomic number.
Historical Background
Early Classifications
Many elements, such as Gold, Silver, Iron, Phosphorus, Sulfur, Zinc, and Arsenic have been known since the pre-historic era. However, the first classification was made in the 18th century.
Antoine Lavoisier attempted to classify known elements as metals and nonmetals.
In 1829, Döbereiner grouped the elements into triads (a group of three) with similar properties, noting that the atomic weight of the middle element was roughly the average of the other two. Example: Lithium, Sodium, and Potassium.
In 1864, John Newlands observed periodicity in the 62 known elements arranged in increasing order of their atomic masses. He classified the elements into groups so that every eight elements resembled the first element in properties.
In the same year, Lothar Meyer developed his famous curves by plotting a graph between the atomic weight and atomic volumes of elements. These curves also showed periodicity.
In 1869, Russian chemist Dmitri Mendeleev arranged 63 elements by increasing atomic mass, aligning elements with similar properties into groups. He left gaps for undiscovered elements and predicted their properties accurately.
In 1913, Moseley determined the exact atomic numbers of known elements using X-ray emission, arranging the elements by atomic numbers instead of atomic masses. This led to the modern Periodic Law.
Modern Periodic Table – Main Features
Features
Presently, 118 elements are grouped in the table in ascending order of their respective atomic numbers.
There are seven horizontal rows called periods and eighteen vertical columns called groups.
Elements within the same group exhibit similar chemical properties because they have the same number of valence electrons. However, they show a gradual change in physical properties from top to bottom in a group.
Elements in a period show a gradual change in properties moving from left to right in periods. Atomic size gradually decreases as we move from left to right.
Metals, Non-Metals and Metalloids
Metals are elements which tend to lose electrons to form positive ions. Examples: Iron, copper, gold, silver, etc.
Non-metals are elements which tend to gain electrons to form negative ions. Examples: Chlorine, sulphur, phosphorus, etc.
Metalloids exhibit some properties of metals and some of non-metals. They separate the metals and nonmetals on a periodic table.
The stair-step line begins at boron (B) and extends down to polonium (Po) including Si, Ge, As, Sb and Te. Elements to the left are metals, to the right are nonmetals, and along the line are metalloids.
Blocks in Periodic Table
s-block elements: Groups 1 and 2, valence electrons in s-subshell.
p-block elements: Groups 13 to 18, valence electrons in p-subshell.
d-block elements: Transition metals.
f-block elements: Lanthanides and Actinides.
Families in Periodic Table
Alkali metals: Group 1 (Li, Na, K, Rb, Cs, Fr).
Alkaline earth metals: Group 2 (Be, Mg, Ca, Sr, Ba, Ra).
Transition elements: d-block and f-block elements.
Chalcogens: Group 16 (O, S, Se, Te, Po, Lv).
Halogens: Group 17 (F, Cl, Br, I, At, Ts).
Noble gases: Group 18 (He, Ne, Ar, Kr, Xe, Rn, Og).
Periodicity of Properties
Modern Periodic Law
The physical and chemical properties of elements are periodic functions of their atomic numbers.
Atomic Radius
Definition: Half of the distance between two identical atoms bonded together. Measured in picometers (pm) or Angstroms (Å).
Trend: Decreases across a period, increases down a group.
Ionic Radius
Definition: Distance from the nucleus of an ion to the outermost electron shell.
Cations are smaller than parent atoms; anions are larger.
Trend: Decreases across period, increases down group.
Ionization Energy
Definition: Energy needed to remove one electron from each atom in one mole of gaseous atoms.
Factors: Nuclear charge, size, electronic arrangement, shielding effect.
Trend: Increases across period, decreases down group.
Electron Affinity
Definition: Enthalpy change when 1 mole of electrons is added to 1 mole of gaseous atoms.
Factors: Size, nuclear charge, electronic configuration.
Trend: Becomes more negative across period, decreases down group.
Electronegativity
Definition: Measure of an atom’s attraction for bonding electrons.
Factors: Atomic size, effective nuclear charge.
Trend: Increases across period, decreases down group.
Metallic Character
Definition: Tendency to lose electrons.
Trend: Decreases across period, increases down group.
Reactions of Sodium and Magnesium
With Water
2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) (vigorous)
Mg(s) + 2H2O(l) → Mg(OH)2(aq) + H2(g) (slow)
Mg(s) + 2H2O(g) → MgO(s) + 2H2(g) (with steam)
With Oxygen
2Na(s) + O2(g) → 2Na2O2(s) (yellow flame)
2Mg(s) + O2(g) → 2MgO(s) (white flame)
With Chlorine
2Na(s) + Cl2(g) → 2NaCl(s)
Mg(s) + Cl2(g) → MgCl2(s)
Oxides and Chlorides of Period 3
Oxides: Ionic on left (Na2O), covalent on right (SO2).
Chlorides: Ionic on left (NaCl), covalent on right (PCl5).
Basic oxides: Na2O, MgO. Acidic oxides: SO2, P2O5. Amphoteric: Al2O3.
Neutral chlorides: NaCl, MgCl2. Acidic chlorides: AlCl3, SiCl4.
Oxidation Numbers
Na in Na2O: +1, Mg in MgO: +2, Al in Al2O3: +3, Si in SiO2: +4, P in P4O10: +5, S in SO3: +6.
In chlorides: Na in NaCl: +1, Mg in MgCl2: +2, Al in AlCl3: +3, Si in SiCl4: +4, P in PCl5: +5.