A Short Revision in Inorganic Chemistry for Undergraduates

 

Among the undergraduate Bachelor's degree courses I studied was chemistry, zoology, physiology, mathematics and medicine 

I still have some of these textbooks on: 

1. inorganic chemistry. 

2. organic chemistry. 

3. physical chemistry. 

4. zoology. 

No less than than 40 books on medicine 

I have been reading some on chemistry over the last two weeks for revision as an academic exercise to jolt back my memory. 

I just like to share with younger students today what we learnt in chemistry way back over 60  years ago in this revision journey. Come along with me in this short revision journey. 

Let me structure this revision in the following way:

1. Inorganic Chemistry – A review of periodic trends, chemical bonding, coordination compounds, acids and bases, redox reactions, and main group & transition metal chemistry.

2. Organic Chemistry – A refresher on fundamental concepts like hybridization, isomerism, reaction mechanisms (substitution, addition, elimination, rearrangement), functional groups, and important reactions.

3. Physical Chemistry – Revisiting thermodynamics, kinetics, quantum chemistry, spectroscopy, electrochemistry, and surface chemistry.

We shall discuss botany and zoology as separate subjects  with a review on zoology with animal classification, physiology, evolutionary biology, and selected topics in genetics, anatomy, and ecological adaptations.

Let me start with Inorganic Chemistry and move sequentially.

 

Inorganic Chemistry: Revision Part 1

Let’s begin with Periodic Trends—a fundamental concept in inorganic chemistry that helps in understanding element properties.

1. The Periodic Table & Periodic Trends

The periodic table is arranged in increasing atomic number, with elements grouped by similar chemical properties. The key periodic trends include:

• Atomic Radius: Decreases across a period (left to right) due to increased nuclear charge but increases down a group due to additional electron shells.
• Ionization Energy: Increases across a period as electrons are held more tightly and decreases down a group as electrons are farther from the nucleus.
• Electron Affinity: Becomes more negative across a period as atoms gain electrons more easily; however, noble gases have near-zero electron affinity.
• Electronegativity: Increases across a period and decreases down a group; fluorine (F) has the highest electronegativity.

2. Chemical Bonding

• Ionic Bonding: Transfer of electrons (e.g., NaCl)

• Covalent Bonding: Sharing of electrons (e.g., H₂, CH₄)

• Metallic Bonding: Delocalized electrons in a lattice of positive ions (e.g., Cu, Fe)

• Coordinate Covalent Bonding: A single atom donates both electrons in a bond (e.g., NH₃ → BF₃ complex)

3. Acid-Base Theories

• Arrhenius Theory: Acids produce H⁺ in water, bases produce OH⁻.

• Bronsted-Lowry Theory: Acids donate protons, bases accept protons.

• Lewis Theory: Acids accept electron pairs, bases donate electron pairs.

Let us now continue with Inorganic Chemistry, covering Coordination Chemistry and Redox Reactions, both of which are crucial topics in this field.

4. Coordination Chemistry

Coordination compounds consist of a central metal ion surrounded by ligands (molecules or ions that donate electron pairs). These compounds are fundamental in bioinorganic chemistry, catalysis, and industrial applications.

Key Terms in Coordination Chemistry are: 

• Coordination Number: The number of ligand donor atoms attached to the central metal (e.g., in [Cu(NH₃)₄]²⁺, Cu²⁺ has a coordination number of 4).

• Ligands: Can be classified based on the number of donor atoms:

1. Monodentate: Bind through a single donor atom (e.g., NH₃, Cl⁻, H₂O).
2. Bidentate: Bind through two donor atoms (e.g., ethylenediamine, C₂O₄²⁻).
3. Polydentate (Chelating Ligands): Bind through multiple donor atoms (e.g., EDTA⁴⁻).   

Common Geometries of Coordination Complexes

• Linear (CN = 2) → Example: [Ag(NH₃)₂]⁺
• Tetrahedral (CN = 4) → Example: [ZnCl₄]²⁻
• Square Planar (CN = 4) → Example: [Pt(NH₃)₂Cl₂] (cisplatin)
• Octahedral (CN = 6) → Example: [Fe(CN)₆]³⁻

Crystal Field Theory (CFT)

This theory explains the color and magnetic properties of coordination compounds by describing how d-orbitals split in an electric field created by ligands.

• Octahedral Splitting: d-orbitals split into two sets:

  1. Higher energy: ${\mathrm{<br>}}_{}$ ($d_{x^2-y^2},d_{z^2}$)

  2. Lower energy: $t_{2g}$ (${}_{}$dxy,dxz,dyzd_{xy}, d_{xz}, d_{yz})

• Tetrahedral Splitting: The energy levels are reversed; $e_g$ orbitals are lower than $t_{2g}$
  

• High-spin vs. Low-spin Complexes:

1. Weak field ligands (e.g., Cl⁻, H₂O) → high-spin (more unpaired electrons, paramagnetic).
2. Strong field ligands (e.g., CN⁻, CO) → low-spin (fewer unpaired electrons, often diamagnetic).  

This explains why [Fe(H₂O)₆]³⁺ is pale violet (high-spin) while [Fe(CN)₆]³⁻ is yellow (low-spin) due to different energy gaps in d-orbital splitting.

Applications of Coordination Chemistry

• Biological systems: Hemoglobin (Fe), Vitamin B₁₂ (Co), and chlorophyll (Mg).
• Medicinal chemistry: Cisplatin for cancer therapy.
• Catalysis: Zeigler-Natta catalysts in polymer production.

5. Redox Reactions & Electrochemistry

Redox (reduction-oxidation) reactions involve the transfer of electrons.

Oxidation & Reduction

• Oxidation: Loss of electrons (increase in oxidation number).
• Reduction: Gain of electrons (decrease in oxidation number).

• Redox Example:

$$Zn+C{u}^{2+}\to Z{n}^{2+}+Cu\mathrm{<br>}\mathrm{<br>}\mathrm{<br>}$$

Zn is oxidized (loses electrons, reducing agent).
Cu²⁺ is reduced (gains electrons, oxidizing agent).

Balancing Redox Equations (Ion-Electron Method)

1. Split into half-reactions (oxidation & reduction).
2. Balance atoms (except H & O).
3. Balance O with H₂O, then balance H with H⁺.
4. Balance charge with electrons (e⁻).
5. Equalize electron transfer and combine half-reactions.

Example: Balancing the acidic reaction

$$MnO_4^- + Fe^{2+} → Mn^{2+} + Fe^{3+}$$

Half-reactions:

$$MnO_4^- + 8H^+ + 5e^- → Mn^{2+} + 4H_2O$$
$$Fe^{2+} → Fe^{3+} + e^-$$

Multiply Fe half-reaction by 5 and add them.

Final balanced equation:

$$MnO_4^- + 5Fe^{2+} + 8H^+ → Mn^{2+} + 5Fe^{3+} + 4H_2O$$

Electrochemical Cells

Redox reactions drive electrochemical cells, categorized into:

1. Galvanic (Voltaic) Cells – Spontaneous reaction producing electrical energy (e.g., Daniell cell).
2. Electrolytic Cells – Non-spontaneous reaction driven by an external voltage (e.g., electrolysis of water).

Standard Electrode Potential ($E^\circ$)

• Measures a species' tendency to gain/lose electrons.

• Standard Hydrogen Electrode (SHE) is set at 0.00V.

• A positive $E^\circ$ means a species is a strong oxidizing agent (e.g., $E^\circ_{Cu^{2+}/Cu} = +0.34V$).

• A negative ${\mathrm{E0 means\; a\; species\; is\; a\; strong reducing\; agent<br>}}^{}$$E^\circ_{Zn^{2+}/Zn} = -0.76V$), e.g.

Nernst Equation (for non-standard conditions):

$$E = E^\circ - \frac{0.0591}{n} \log Q$$

where $Q$ is the reaction quotient, and $n$ is the number of electrons transferred.

Next Steps in Inorganic Chemistry

Now that we have covered Periodic Trends, Bonding, Coordination Chemistry, and Redox Reactions, the next topics could be:

1. Main Group & Transition Metal Chemistry (properties & trends of Groups 1-18, d-block elements).

2. Industrial & Environmental Applications (e.g., Haber process, metallurgy).

Inorganic chemistry overlaps with physical chemistry in areas like electrochemistry, bonding theories, thermodynamics, and reaction kinetics. As we continue, I will highlight these connections so you can see the interplay between these disciplines.

Let’s move forward with Main Group & Transition Metal Chemistry, followed by Metallurgy and Industrial Applications.

6. Main Group & Transition Metal Chemistry

The periodic table is divided into main group elements (s- and p-block) and transition metals (d- and f-block). Their chemistry is crucial in industry, medicine, and materials science.

6.1 Main Group Elements (s- and p-block)

Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr)

• Highly reactive metals, react vigorously with water:

  $$2Na+2H_{2}O\to 2NaOH+H_2$$
  

• Form strong bases (e.g., NaOH, KOH).

• Exist as ionic compounds (e.g., NaCl, K₂CO₃).

• Low ionization energy → easily lose electrons to form +1 oxidation state.

Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

• Less reactive than alkali metals, but still react with water (except Be).

• Important in biological systems (e.g., Ca²⁺ in bones, Mg²⁺ in chlorophyll).

• Form oxides and hydroxides:

  $$CaO+H_{2}O\to Ca(OH{)}_{\mathrm{2<span\; style="font-weight:\; bold;"></span>}}$$

• Common compounds: Mg(OH)₂ (milk of magnesia), CaCO₃ (limestone), BaSO₄ (used in X-rays).

Group 13: Boron Group (B, Al, Ga, In, Tl)

• Boron (B) is a metalloid, others are metals.

• Aluminum is amphoteric:

  $$Al(OH{)}_3+NaOH\to Na[Al(OH{)}_4](\text{acts as an acid})$$

$$Al(OH{)}_3+HCl\to AlC{l}_3+H_{2}O(\text{acts as a base)}$$

Group 14: Carbon Group (C, Si, Ge, Sn, Pb)

• Carbon: Basis of organic chemistry; forms strong covalent bonds.

• Silicon (Si): Found in sand (SiO₂), semiconductors.

• Tin & Lead: Show +2 and +4 oxidation states due to the inert pair effect.

Group 15: Nitrogen Group (N, P, As, Sb, Bi)

• Nitrogen (N₂) is an inert gas due to its strong triple bond.

• Phosphorus (P) forms allotropes:

  • White P (reactive, stored in water).

  • Red P (stable, used in matches).

• Ammonia (NH₃) is a key base:

  $$N{H}_3+H^{+}\to N{H}_4^{}$$

Group 16: Oxygen Group (O, S, Se, Te, Po)

• Oxygen (O₂): Strong oxidizing agent.

• Sulfur (S): Forms oxides like SO₂ (acid rain precursor).

• Sulfuric Acid (H₂SO₄): Industrially significant; acts as an oxidizer and dehydrating agent.

Group 17: Halogens (F, Cl, Br, I, At)

• Highly reactive nonmetals, exist as diatomic molecules (F₂, Cl₂).

• Fluorine (F₂): Most electronegative element, strong oxidizer.

• Chlorine (Cl₂): Used in water purification.

• Form halide salts:

  $$Na+C{\mathrm{l }}_{}\to NaCl$$
  

Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

• Chemically inert, but Xe can form compounds (e.g., XeF₄).

• Argon (Ar) used in welding to provide an inert atmosphere.

6.2 Transition Metals (d-block)

Transition metals show variable oxidation states, complex formation, and catalytic properties.

General Properties

• Partially filled d-orbitals lead to colorful compounds.

• Multiple oxidation states:

  • Fe → Fe²⁺, Fe³⁺

  • Cu → Cu⁺, Cu²⁺

• Good catalysts: Fe in Haber process, Pt in catalytic converters.

Notable Transition Metals

• Iron (Fe): Found in hemoglobin, steel.

• Copper (Cu): Electrical wiring, corrosion-resistant.

• Silver (Ag): Photography, antimicrobial properties.

• Platinum (Pt): Catalysts, jewelry.

• Chromium (Cr): Stainless steel, pigments

Aluminum oxide (Al2O3) and aluminium hydroxide (Al (OH) are amphoteric salts, meaning they can react with both acids and bases. 

Here's how: Aluminum Oxide (Al₂O₃):

• Reaction with acids (acting as a base):

  • Al₂O₃(s) + 6HCl(aq) → 2AlCl₃(aq) + 3H₂O(l) 
  • Al₂O₃(s) + 3H₂SO₄(aq) → Al₂(SO₄)₃(aq) + 3H₂O(l) 
• Reaction with bases (acting as an acid):

  • Al₂O₃(s) + 2NaOH(aq) + 3H₂O(l) → 2NaAl(OH)₄ 
  • Al₂O₃(s) + 2NaOH(aq) → 2NaAlO₂(aq) + H₂O(l) 

Aluminum Hydroxide (Al(OH)₃):

• Reaction with acids (acting as a base):
  • Al(OH)₃(s) + 3H⁺(aq) → Al³⁺(aq) + 3H₂O(l)
• Reaction with bases (acting as an acid):
  • Al(OH)₃(s) + OH⁻(aq) → [Al(OH)₄]⁻(aq) 

7. Metallurgy & Industrial Applications

Metallurgy is the process of extracting metals from ores and refining them.

Metallurgical Processes

1. Ore Concentration (removing impurities).

2. Froth flotation for sulfide ores.

3. Magnetic separation for iron ores.

4. Extraction (reducing metal ions to metals).

   1. Reduction of oxides:

      $$F{e}_2{O}_3+3CO\to 2Fe+3C{O}_2$$

   2. Electrolysis for reactive metals (Al, Na, Mg).

5. Refining (purifying the metal).

• Electrolytic refining for Cu, Ag.

• Zone refining for semiconductors (Si, Ge).

Connections to Physical Chemistry

Many inorganic concepts involve physical chemistry principles:

• Crystal Field Theory (CFT) (Quantum Mechanics).

• Electrochemistry (Redox & Nernst Equation).

• Metallurgy (Thermodynamics & Gibbs Free Energy).

• Transition Metal Spectroscopy (Absorption of light in d-orbitals).

This sums up an undergraduate  course in inorganic chemistry in a nutshell