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 connections separate students who struggle from those who ace the exam. Each compound below is grouped by chemical behavior, so you're learning patterns, not just facts.

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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 that dissociates completely into $\text{H}^+$ and $\text{Cl}^-$ 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: $\text{HCl} + \text{NaOH} \rightarrow \text{NaCl} + \text{H}_2\text{O}$

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

• Diprotic acid and dehydrating agent that can donate two protons and removes water from compounds
• The most-produced industrial chemical globally, essential for fertilizer manufacturing (phosphoric acid production) and lead-acid batteries
• Highly corrosive due to both its strong acidity and the large exothermic heat of mixing with water (always add acid to water, not the reverse)

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

• Strong acid and strong oxidizing agent because nitrogen sits at its +5 oxidation state, making it electron-hungry
• Industrial applications include fertilizer production (ammonium nitrate) and explosives manufacturing
• Passivation of metals like aluminum and chromium forms a thin protective oxide layer that resists further corrosion

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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 $\text{OH}^-$, making it a strong Arrhenius base
• Saponification reactions with fats produce soap, a classic application in organic-inorganic crossover questions
• Industrial uses include paper manufacturing (Kraft process), textile processing, and drain cleaners

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Weak Bases and Amphoteric Compounds

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

Ammonia ($\text{NH}_3$)

• Weak base behavior: accepts a proton from water to form $\text{NH}_4^+$ and $\text{OH}^-$, 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
• Haber-Bosch process: $\text{N}_2 + 3\text{H}_2 \rightleftharpoons 2\text{NH}_3$ (note the equilibrium arrow; this reaction is reversible and requires high temperature, high pressure, and an iron catalyst)

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

• Amphoteric character: the $\text{HCO}_3^-$ ion can act as either an acid (donating its remaining proton) or a base (accepting a proton), depending on the reaction partner
• Reaction with acids produces $\text{CO}_2$ gas: $\text{NaHCO}_3 + \text{HCl} \rightarrow \text{NaCl} + \text{H}_2\text{O} + \text{CO}_2$
• Buffer component in the bicarbonate buffer system, critical for maintaining blood pH near 7.4

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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 that 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 (801 °C) and the energy required to dissolve it

Potassium Chloride ($\text{KCl}$)

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

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

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

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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 an anomalously high boiling point (100 °C), high surface tension, and high specific heat capacity
• Universal solvent properties arise from its polarity and ability to stabilize ions through hydration shells
• Amphoteric behavior: can act as an acid (donating $\text{H}^+$) or a base (accepting $\text{H}^+$) depending on the reaction partner, with $K_w = 1.0 \times 10^{-14}$ at 25 °C

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

• Linear, nonpolar molecule: despite having polar $\text{C}=\text{O}$ bonds, the symmetric geometry cancels the dipole moments, giving a net dipole of zero
• 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 its IR-active asymmetric stretching and bending vibrational modes

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

• Oxygen in the -1 oxidation state: intermediate between $\text{O}_2$ (0) and $\text{H}_2\text{O}$ (-2), making it capable of acting as both an oxidizing agent and a reducing agent
• 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 the enzyme catalase
• Bleaching mechanism involves oxidation of chromophores, destroying the conjugated structures responsible for color

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Oxidizing Agents and Transition Metal Compounds

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

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

• Mn in the +7 oxidation state: the highest possible for manganese, making it a powerful electron acceptor
• Color change as a redox indicator: purple $\text{MnO}_4^-$ reduces to colorless $\text{Mn}^{2+}$ in acidic solution or brown $\text{MnO}_2$ in neutral/basic solution
• Titration standard for redox reactions, especially with oxalic acid ($\text{C}_2\text{O}_4^{2-}$) and iron(II) compounds. The self-indicating endpoint (purple to colorless) means no separate indicator is needed.

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

• Lewis acid behavior: the electron-deficient $\text{Fe}^{3+}$ ion accepts electron pairs from donors, making it useful as a catalyst in Friedel-Crafts reactions
• Hydrolysis in water produces acidic solutions because $\text{Fe}^{3+}$ coordinates water molecules and polarizes them, releasing $\text{H}^+$: $[\text{Fe}(\text{H}_2\text{O})_6]^{3+} \rightleftharpoons [\text{Fe}(\text{OH})(\text{H}_2\text{O})_5]^{2+} + \text{H}^+$
• Coagulant in water treatment: neutralizes the charge on colloidal particles, allowing them to aggregate and settle out

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

• Hydration behavior: anhydrous $\text{CuSO}_4$ is white; adding water produces blue $\text{CuSO}_4 \cdot 5\text{H}_2\text{O}$ (copper(II) sulfate pentahydrate, historically called blue vitriol)
• Water detection test: the color change from white to blue confirms the presence of water in unknown samples
• Agricultural fungicide (Bordeaux mixture) due to $\text{Cu}^{2+}$ toxicity to fungal cells

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Quick Reference Table

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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. Both $\text{HCl}$ and $\text{HNO}_3$ are strong acids, but only one is also a strong oxidizing agent. Which one, and what about its structure enables oxidation?

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

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

5. Why is $\text{KMnO}_4$ purple in solution but becomes colorless or brown after reacting with a reducing agent? Connect the color change to what's happening to the manganese oxidation state and its $d$-electron configuration.