What is Inorganic Chemistry?
Inorganic chemistry is the branch of chemistry that deals with all the elements and their compounds, except for most carbon-based compounds. Only a few carbon compounds—like oxides, cyanides, and carbonates—are traditionally included in inorganic chemistry.
Simply put, it is the study of the elements in the periodic table and how they behave and combine with each other.
The Periodic Table
The periodic table is an organized arrangement of all known chemical elements in order of increasing atomic number. This arrangement is not random—it reflects the electron configuration of the atoms, which in turn determines their chemical properties.
Early Classification: Mendeleev and Lothar Meyer (1869)
In 1869, two scientists—Dmitri Mendeleev in Russia and Lothar Meyer in Germany—independently created the first comprehensive versions of the periodic table. They arranged the known elements according to their relative atomic masses (atomic weights).
They observed that when elements are placed in order of increasing atomic mass, similar properties reappear at regular intervals. This repeating pattern is called periodicity.
Mendeleev and Meyer placed elements with similar properties in vertical columns called groups, and horizontal rows were called periods. At the time, many elements (such as the noble gases, gallium, and germanium) had not yet been discovered, so there were gaps in their tables.
The Original Periodic Law
The early periodic law stated:
“The properties of elements are a periodic function of their relative atomic masses.”
Anomalies in the Early Table
However, this system had some problems. To keep chemically similar elements together, Mendeleev had to reverse the order of a few pairs based on their properties rather than strict atomic mass:
- Argon (39.9) and Potassium (39.1)
- Cobalt (58.94) and Nickel (58.69)
- Tellurium (127.6) and Iodine (126.9)
If he had followed atomic mass strictly, these elements would have been placed with very different elements (for example, potassium would have been separated from the other alkali metals).
These anomalies occurred because of isotopes. Elements with different isotopes can have average atomic masses that don’t perfectly match their chemical behavior. For instance, argon’s heavier isotopes are more abundant, raising its average mass above potassium’s.
Moseley’s Discovery and the Modern Periodic Law
In the early 20th century, Henry Moseley discovered that the atomic number (the number of protons in the nucleus) is the true basis for arranging elements. Atomic number determines the number of electrons and, therefore, the chemical properties of an element.
When elements are arranged by atomic number instead of atomic mass, all the previous anomalies disappear.
Modern Periodic Law:
“The properties of elements are a periodic function of their atomic numbers.”
This is the version we use today. Atomic number is the correct and reliable way to organize the periodic table.
Structure of the Modern Periodic Table
The periodic table is one of the most powerful tools in chemistry. It helps us understand trends in properties and predict how elements will behave.
In the table:
- Each box contains the element’s symbol, atomic number, and atomic mass.
- Additional information (electronegativity, melting/boiling points, oxidation states, etc.) is often included.
The modern long form of the periodic table (the one commonly used) has:
- Periods – horizontal rows
- Groups – vertical columns (elements in the same group have similar properties)
The table is divided into four main blocks (s, p, d, and f) based on the subshell in which the outermost or valence electrons are filling.
1. s-Block Elements
These elements fill their outermost electrons in the s-subshell. They are found in Group 1 (IA) and Group 2 (IIA). Examples: Lithium (Li), Sodium (Na), Potassium (K), Magnesium (Mg), Calcium (Ca). They are highly reactive metals (especially Group 1).
2. p-Block Elements
These elements fill their outermost electrons in the p-subshell. They range from Group 13 (IIIA) to Group 18 (VIIIA). Examples: Carbon (C), Nitrogen (N), Oxygen (O), Phosphorus (P), Sulfur (S), Chlorine (Cl), and the noble gases.
Main Group Elements: The s-block and p-block together are called the main group elements. They use only their outermost shell electrons for bonding. The noble gases (Group 18) have completely filled s and p subshells, making them very stable and chemically unreactive (though some can form compounds with fluorine and oxygen).
Note: Helium is an s-block element, even though it is placed at the far right with the noble gases.
3. d-Block Elements (Transition Metals)
These elements have electrons filling the d-subshell of the penultimate (second-to-last) shell. They are located in the middle of the periodic table. They can use both s and d electrons for bonding and often show multiple oxidation states and colorful compounds. Examples: Scandium (Sc), Titanium (Ti), Iron (Fe), Copper (Cu), Zinc (Zn), etc.
4. f-Block Elements (Inner Transition Metals)
These are the Lanthanides and Actinides. They are placed in two separate rows at the bottom of the table. They have electrons filling the f-subshell (and partially filled d-subshells). They are very similar to each other, making them difficult to separate. Examples: Lanthanum (La), Uranium (U), Thorium (Th), etc.
Periodic Trends and Properties of Elements
General Observations in the Periodic Table
- All elements in the same period have the same number of electron shells (energy levels), which may be partially or fully filled.
- The number of the main group (Group number) corresponds to the number of electrons in the outermost shell.
- The period number equals the principal quantum number (n) of the outermost shell and indicates the total number of occupied electron shells.
Periodic Trends in Physical Properties
Trends Down a Group (from top to bottom)
I. Atomic Radius (Atomic Size) Atomic radius is defined as half the distance between the nuclei of two identical atoms joined by a single covalent or metallic bond.
As you go down a group, atomic radius increases significantly. This happens because:
- A new electron shell is added.
- The outer electrons are farther from the nucleus.
- Inner electrons shield the outer electrons from the full attraction of the nucleus (shielding effect).
II. Ionic Radius Ionic radii follow a similar trend to atomic radii and increase down the group.
- Cations (positive ions) are smaller than their parent atoms because losing electrons increases the effective nuclear charge, pulling the remaining electrons closer. Example: Na (1.57 Å) → Na⁺ (0.97 Å)
- Anions (negative ions) are larger than their parent atoms because adding electrons increases electron-electron repulsion and reduces the effective nuclear charge. Example: Cl (0.99 Å) → Cl⁻ (1.81 Å)
III. Ionization Energy (I.E.) Ionization energy is the energy required to remove an electron from a gaseous atom or ion.
- First ionization energy: energy to remove the first electron.
- It decreases down a group.
Reason: Larger atomic size + greater shielding effect makes the outer electron easier to remove.
IV. Electron Affinity (E.A.) Electron affinity is the energy change when an electron is added to a gaseous atom.
- It is usually negative (energy is released) for non-metals.
- Halogens (Group VIIA) have the highest (most negative) electron affinities.
- There is a general decrease in electron affinity down the group (with some exceptions, e.g., Chlorine has higher E.A. than Fluorine due to less electron repulsion in the larger atom).
V. Electronegativity Electronegativity measures an atom’s ability to attract shared electrons in a covalent bond.
- It decreases down a group.
- The first element in each main group is the most electronegative.
Trends Across a Period (from left to right)
Atomic Radius Atomic radius decreases steadily across a period because:
- Nuclear charge increases (more protons).
- Electrons are added to the same shell, so shielding is minimal.
- Stronger attraction pulls electrons closer to the nucleus.
There is a sudden increase from halogens to noble gases because noble gases are measured differently (Van der Waals radius).
Ionization Energy Ionization energy generally increases across a period due to decreasing atomic size and increasing effective nuclear charge.
Exceptions (drops) occur at:
- Be → B and Mg → Al (removal of a higher-energy p-electron)
- N → O and P → S (extra stability of half-filled p-subshell in N and P)
Factors Affecting Ionization Energy
- Effective nuclear charge (higher → higher I.E.)
- Shielding/screening effect (higher → lower I.E.)
- Atomic radius (larger → lower I.E.)
- Stable full or half-filled sub-shells (require more energy to remove electrons)
Electron Affinity Generally increases (becomes more negative) across a period. Non-metals have higher electron affinities than metals.
Electronegativity Electronegativity increases across a period. Halogens are the most electronegative; alkali metals are the least.
Melting Point Melting points do not change uniformly. They depend on the type of bonding and structure:
- Increases sharply from Na to Mg (stronger metallic bonding).
- Silicon has a very high melting point due to its giant covalent structure.
- Phosphorus, Sulfur, and Chlorine have low melting points due to weak Van der Waals forces between molecules.
Periodic Trends in Chemical Properties (Period 3: Na to Ar)
As we move from left to right across Period 3, elements change from strongly metallic (Na, Mg) to weakly metallic (Al) to non-metallic (Si, P, S, Cl, Ar).
Hydrides
- NaH, MgH₂, AlH₃ → largely ionic or covalent, basic (form alkaline solutions with water).
- SiH₄, PH₃, H₂S, HCl → covalent, become increasingly acidic.
Chlorides
- NaCl, MgCl₂ → ionic.
- AlCl₃, SiCl₄, PCl₃/PCl₅, etc. → covalent.
- Ionic character decreases across the period.
- Hydrolysis increases across the period (AlCl₃ and later chlorides hydrolyze to give acidic solutions).
Hydroxides / Oxy-acids
- NaOH, Mg(OH)₂ → strongly basic.
- Al(OH)₃ → amphoteric.
- Si, P, S, Cl compounds → acidic (oxy-acids). Acidity increases with electronegativity.
Oxides
- Na₂O, MgO → basic (ionic).
- Al₂O₃ → amphoteric.
- SiO₂, P₄O₁₀, SO₂/SO₃, Cl₂O₇ → acidic. Solubility and basicity decrease across the period.
Diagonal Relationship
The first element of a group often shows similarities with the second element of the next group (diagonally placed) due to similar size, electronegativity, and polarizing power.
Common Pairs:
- Li and Mg
- Be and Al
- B and Si
Factors for Diagonal Similarity:
- Similar atomic and ionic sizes
- Similar electronegativity
- Similar polarizing power of ions
Anomalous Behaviour of the First Element in a Group
The first element in each group behaves differently from the rest because it has:
- Smallest size
- Highest ionization energy
- Highest electronegativity
- Highest electron affinity (in its group)
Anomalous Properties of Lithium (differs from other Alkali Metals)
- Forms covalent compounds (others are ionic).
- Reacts with nitrogen to form nitride.
- Reacts slowly with cold water.
- Forms hydrated salts.
- Forms monoxide when burnt in air (others form peroxides/superoxides).
- Its carbonate and hydroxide are less stable.
- Lithium salts can hydrolyze.
Anomalous Properties of Beryllium (differs from other Alkaline Earth Metals)
- Amphoteric oxides and hydroxides.
- Forms covalent compounds and fluoro-complexes.
- Does not react with water.
- Chloride dimerizes and dissolves in organic solvents.
- Carbide gives methane on hydrolysis (others give acetylene).
Anomalous Properties of Fluorine (differs from other Halogens)
- Always monovalent (no d-orbitals).
- Highest electronegativity → forces other elements to show maximum oxidation states.
- Strong hydrogen bonding in HF.
- Many solubility differences (e.g., fluorides often insoluble).
- Reacts with water to liberate oxygen.
- Behaves uniquely with alkalis and oxygen compounds.
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