Common Inorganic Compounds

Why This Matters

Inorganic chemistry exams don't just ask you to identify compounds—they test whether you understand why compounds behave the way they do. When you see $\text{H}_2\text{SO}_4$ on a question, you're being tested on acid-base theory, oxidation states, and industrial applications all at once. The compounds in this guide represent the foundational categories you'll encounter repeatedly: acids and bases, ionic salts, molecular compounds, and oxidizing agents. Master these, and you'll have the building blocks for predicting reactivity, writing balanced equations, and explaining real-world chemistry.

Don't just memorize formulas and names—know what concept each compound illustrates. Can you explain why $\text{NaOH}$ is a strong base while $\text{NH}_3$ is weak? Why $\text{KMnO}_4$ oxidizes other species? These are the connections that separate students who struggle from those who ace the exam. Each compound below is grouped by its chemical behavior, so you're learning patterns, not just facts.

---

Strong Acids: Complete Dissociation and Industrial Workhorses

Strong acids dissociate completely in aqueous solution, releasing all their $\text{H}^+$ ions. This complete ionization is what makes them powerful proton donors and essential reagents in industrial processes.

Hydrochloric Acid ($\text{HCl}$)

• Monoprotic strong acid—dissociates completely into $\text{H}^+$ and $\text{Cl}^-$ ions in water
• Biological role in gastric juice where it maintains stomach pH around 1.5–3.5 for protein digestion
• Neutralization reactions with bases produce salt and water, a classic exam equation type

Sulfuric Acid ($\text{H}_2\text{SO}_4$)

• Diprotic acid and dehydrating agent—can donate two protons and removes water from compounds
• Industrial significance makes it the most-produced chemical globally, essential for fertilizers and batteries
• Highly corrosive due to both its acidity and its ability to generate heat when mixed with water

Nitric Acid ($\text{HNO}_3$)

• Strong acid and strong oxidizing agent—the nitrogen in $\text{HNO}_3$ has a +5 oxidation state, making it electron-hungry
• Industrial applications include fertilizer production (ammonium nitrate) and explosives manufacturing
• Passivation reactions with metals like aluminum form protective oxide layers, a testable concept

---

Strong Bases: Hydroxide Ion Donors

Strong bases dissociate completely to release $\text{OH}^-$ ions. Their reactivity stems from the hydroxide ion's ability to accept protons and attack electrophilic centers.

Sodium Hydroxide ($\text{NaOH}$)

• Complete dissociation in water releases $\text{Na}^+$ and $\text{OH}^-$ ions, making it a strong Arrhenius base
• Saponification reactions with fats produce soap, a classic application in organic-inorganic crossover questions
• Industrial scale uses include paper manufacturing, textile processing, and drain cleaners

---

Weak Bases and Amphoteric Compounds

Weak bases only partially accept protons in solution, establishing equilibrium. Understanding equilibrium position is key to predicting buffer behavior and pH calculations.

Ammonia ($\text{NH}_3$)

• Weak base behavior—accepts a proton to form $\text{NH}_4^+$, but the equilibrium lies far to the left
• Nitrogen cycle importance as a key intermediate in converting atmospheric $\text{N}_2$ to biologically usable forms
• Fertilizer production via the Haber process ($\text{N}_2 + 3\text{H}_2 \rightarrow 2\text{NH}_3$) is an industrial chemistry staple

Sodium Bicarbonate ($\text{NaHCO}_3$)

• Amphoteric character—can act as either acid or base depending on the reaction partner
• Decomposition reaction with acids produces $\text{CO}_2$ gas, the basis for baking leavening
• Buffer component in the bicarbonate buffer system, critical for blood pH regulation

---

Ionic Salts: Electrolytes and Biological Function

Ionic compounds dissolve in water to form electrolyte solutions. The ions released determine conductivity, osmotic pressure, and biological signaling.

Sodium Chloride ($\text{NaCl}$)

• 1:1 electrolyte—dissociates into $\text{Na}^+$ and $\text{Cl}^-$, creating a neutral solution (neither ion hydrolyzes significantly)
• Osmotic balance in cells depends on $\text{NaCl}$ concentration gradients across membranes
• Lattice energy of 787 kJ/mol explains its high melting point and solubility behavior

Potassium Chloride ($\text{KCl}$)

• Essential electrolyte—$\text{K}^+$ ions are critical for nerve impulse transmission and muscle contraction
• Fertilizer component provides potassium, one of the three primary plant macronutrients (N-P-K)
• Medical applications include IV fluids and treatment for hypokalemia

Calcium Carbonate ($\text{CaCO}_3$)

• Acid-base reaction with $\text{HCl}$ produces $\text{CO}_2$: $\text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2$
• Geological significance as the primary component of limestone, marble, and marine shells
• Thermal decomposition at high temperatures yields $\text{CaO}$ (quickite) and $\text{CO}_2$, a key industrial process

---

Molecular Compounds: Covalent Bonding and Unique Properties

These compounds feature covalent bonds and exhibit properties tied to molecular structure. Intermolecular forces—not ionic interactions—determine their physical behavior.

Water ($\text{H}_2\text{O}$)

• Hydrogen bonding network creates anomalously high boiling point, surface tension, and specific heat
• Universal solvent properties arise from its polarity and ability to stabilize ions through hydration
• Amphoteric behavior—can act as acid (donating $\text{H}^+$) or base (accepting $\text{H}^+$) depending on reaction partner

Carbon Dioxide ($\text{CO}_2$)

• Linear, nonpolar molecule—despite polar $\text{C}=\text{O}$ bonds, the symmetric geometry cancels dipole moments
• Acid anhydride behavior—dissolves in water to form carbonic acid: $\text{CO}_2 + \text{H}_2\text{O} \rightleftharpoons \text{H}_2\text{CO}_3$
• Greenhouse gas that absorbs infrared radiation due to asymmetric stretching vibrations

Hydrogen Peroxide ($\text{H}_2\text{O}_2$)

• Oxygen oxidation state of -1—intermediate between $\text{O}_2$ (0) and $\text{H}_2\text{O}$ (-2), making it both oxidizer and reducer
• Disproportionation reaction: $2\text{H}_2\text{O}_2 \rightarrow 2\text{H}_2\text{O} + \text{O}_2$, catalyzed by $\text{MnO}_2$ or catalase
• Bleaching mechanism involves oxidation of chromophores, destroying color-causing structures

---

Oxidizing Agents and Transition Metal Compounds

These compounds readily accept electrons from other species. High oxidation states on central atoms drive their oxidizing power.

Potassium Permanganate ($\text{KMnO}_4$)

• Mn in +7 oxidation state—the highest possible for manganese, making it a powerful electron acceptor
• Color change indicator—purple $\text{MnO}_4^-$ reduces to colorless $\text{Mn}^{2+}$ (acidic) or brown $\text{MnO}_2$ (neutral/basic)
• Titration standard for redox reactions, especially with oxalic acid and iron(II) compounds

Iron(III) Chloride ($\text{FeCl}_3$)

• Lewis acid behavior—the $\text{Fe}^{3+}$ ion accepts electron pairs, catalyzing Friedel-Crafts reactions
• Hydrolysis in water produces acidic solutions due to $\text{Fe}^{3+}$ coordinating water and releasing $\text{H}^+$
• Coagulant in water treatment—neutralizes colloidal particles, allowing them to aggregate and settle

Copper Sulfate ($\text{CuSO}_4$)

• Hydration behavior—anhydrous form is white; addition of water produces blue $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ (blue vitriol)
• Water detection test—color change from white to blue confirms presence of water in unknown samples
• Agricultural fungicide due to $\text{Cu}^{2+}$ toxicity to fungal cells

---

Quick Reference Table

---

Self-Check Questions

1. Which two compounds from this guide can act as both an acid and a base (amphoteric behavior), and what structural feature enables this?

2. Compare $\text{HCl}$ and $\text{HNO}_3$: both are strong acids, but only one is also a strong oxidizing agent. Which one, and why does its molecular structure enable oxidation?

3. If you needed to identify the presence of water in an unknown sample, which compound would you use and what observable change would confirm a positive result?

4. Rank $\text{NaOH}$, $\text{NH}_3$, and $\text{NaHCO}_3$ from strongest to weakest base. What determines the strength difference between them?

5. An FRQ asks you to explain why $\text{KMnO}_4$ is purple in solution but becomes colorless or brown after reacting with a reducing agent. What's happening to the manganese, and how does oxidation state relate to color?