MEV 013: Unit 02 - Environmental Chemistry-II

 UNIT 2: ENVIRONMENTAL CHEMISTRY – II

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2.0 Introduction

Environmental chemistry extends into the behavior of acids, bases, redox reactions, and equilibria in natural systems such as soil, water, and air. These chemical principles govern crucial phenomena including nutrient cycling, pollutant transformation, acid rain formation, and water buffering capacity. This unit explores the fundamentals of acid–base chemistry, ionic equilibrium, and redox processes with specific relevance to environmental systems.

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2.1 Objectives

After studying this unit, you will be able to:

• Understand the nature of acid–base reactions in environmental systems.
• Calculate and interpret pH, pOH, and ionic product of water.
• Explain hydrolysis and buffer actions in natural waters.
• Apply the common ion effect in chemical equilibria.
• Understand the environmental implications of oxidation and reduction processes.

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2.2 Acid–Base Reactions

Arrhenius Concept:

• Acids produce H⁺ ions in aqueous solutions.
• Bases produce OH⁻ ions.

Bronsted–Lowry Concept:

• Acids donate protons (H⁺); bases accept protons.

Lewis Concept:

• Acids accept electron pairs; bases donate electron pairs.

Environmental Relevance:

• Soil acidity, lake acidification, and acid rain.
• Buffer systems maintain the pH of biological and aquatic systems.

Example:
Carbonic acid (H₂CO₃) in water maintains oceanic pH:
H2CO3⇌H++HCO3−\text{H}_2\text{CO}_3 ⇌ \text{H}^+ + \text{HCO}_3^-H2​CO3​⇌H++HCO3−​

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2.3 Ionic Product of Water (Kw)

Water undergoes self-ionization:

H2O (l)⇌H+(aq)+OH−(aq)\text{H}_2\text{O (l)} \rightleftharpoons \text{H}^+ (aq) + \text{OH}^- (aq)H2​O (l)⇌H+(aq)+OH−(aq)

The ionic product at 25°C is:

Kw=[H+][OH−]=1.0×10−14K_w = [\text{H}^+][\text{OH}^-] = 1.0 \times 10^{-14}Kw​=[H+][OH−]=1.0×10−14

This equilibrium is temperature dependent and forms the foundation for pH calculations.

Implications in Environment:

• Water bodies’ ability to maintain neutrality.
• pH changes due to acid deposition or pollution affect biodiversity.

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2.4 pH and pOH

pH: Negative logarithm of hydrogen ion concentration

pH=−log⁡[H+]\text{pH} = -\log[\text{H}^+]pH=−log[H+]

pOH: Negative logarithm of hydroxide ion concentration

pOH=−log⁡[OH−]\text{pOH} = -\log[\text{OH}^-]pOH=−log[OH−] pH+pOH=14\text{pH} + \text{pOH} = 14pH+pOH=14

Environmental Relevance:

• pH < 5.5 in rainfall indicates acid rain.
• Aquatic organisms are sensitive to pH changes.

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2.5 Hydrolysis

Hydrolysis refers to the reaction of ions with water, especially salts.

Types:

• Cationic Hydrolysis: e.g., Al³⁺ + H₂O → Al(OH)²⁺ + H⁺
• Anionic Hydrolysis: e.g., CO₃²⁻ + H₂O → HCO₃⁻ + OH⁻
• Amphoteric Hydrolysis: Both ion types contribute.

Environmental Role:

• Controls pH of natural waters.
• Important in soil chemistry, especially for nutrient availability.

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2.6 Buffer Solutions

Buffer solutions resist changes in pH when small amounts of acid or base are added.

Types:

• Acidic Buffer: Weak acid + salt (e.g., CH₃COOH + CH₃COONa)
• Basic Buffer: Weak base + salt (e.g., NH₄OH + NH₄Cl)

Henderson-Hasselbalch Equation:

pH=pKa+log⁡([Salt][Acid])\text{pH} = pK_a + \log \left( \frac{[\text{Salt}]}{[\text{Acid}]} \right)pH=pKa​+log([Acid][Salt]​)

Environmental Importance:

• Carbonate buffer system in oceans.
• Buffering in wastewater treatment and blood pH.

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2.7 Common Ion Effect

The common ion effect is the suppression of ionization of a weak acid or base when a strong electrolyte containing a common ion is added.

Example:
Adding NaCl to HCl suppresses ionization of HCl.

Environmental Relevance:

• Influences solubility of minerals in groundwater.
• Important in remediation and precipitation reactions.

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2.8 Oxidation and Reduction

Oxidation: Loss of electrons or gain of oxygen.

Reduction: Gain of electrons or loss of oxygen.

Redox Reaction Example:

Fe2++MnO4−→Fe3++Mn2+\text{Fe}^{2+} + \text{MnO}_4^- → \text{Fe}^{3+} + \text{Mn}^{2+}Fe2++MnO4−​→Fe3++Mn2+

Oxidizing Agents: Accept electrons (e.g., O₂, Cl₂, KMnO₄)

Reducing Agents: Donate electrons (e.g., Fe²⁺, H₂)

Environmental Applications:

• Redox potential (Eh) governs chemical behavior of elements like Fe, Mn, As in groundwater.
• Redox reactions control degradation of pollutants in water and soils.
• Oxygen demand in water bodies determines their health and self-purification capacity.

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2.9 Let Us Sum Up

This unit focused on the chemical behavior of acids, bases, and redox species in environmental systems. Key principles such as the ionic product of water, pH, hydrolysis, buffer systems, and the common ion effect were explored with relevance to real-world applications like water treatment, soil chemistry, and atmospheric interactions. Redox reactions, which govern transformation and mobility of pollutants, were also discussed.

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2.10 Glossary

·         pH-Measure of hydrogen ion concentration

·         Hydrolysis-Chemical reaction involving water

·         Buffer Solution-Maintains stable pH upon acid/base addition

·         Common Ion Effect-Suppression of ionization due to shared ion

·         Redox Reaction-Combined process of oxidation and reduction

·         Oxidizing Agent-Species that gains electrons

·         Reducing Agent-Species that loses electrons

·         Ionic Product (Kw)-Product of H⁺ and OH⁻ ion concentrations in water

·         Bronsted-Lowry Acid-Proton donor

·         Lewis Base-Electron pair donor