Unit 2: Fundamentals of environmental chemistry
Atomic structure and electronic configuration
All matter is made of atoms, which consist of:
- Protons (p⁺): Positively charged, in the nucleus.
- Neutrons (n⁰): No charge, in the nucleus.
- Electrons (e⁻): Negatively charged, orbit the nucleus in specific energy levels or orbitals.
Electronic configuration is the arrangement of electrons in these orbitals (s, p, d, f). The electrons in the outermost shell are called valence electrons, and they determine the chemical behavior of an atom.
Example: Oxygen (O) has 8 electrons. Its configuration is 1s²2s²2p⁴. It has 6 valence electrons (in the 2s and 2p orbitals). To become stable (like Neon, 1s²2s²2p⁶), it needs to gain 2 electrons, making it highly reactive.
Periodic properties of elements
These are trends in the periodic table that predict how elements will behave.
- Ionization Potential (or Ionization Energy): The energy required to remove one electron from an atom.
Trend: Increases across a period (left to right) and decreases down a group.
- Electron Affinity: The energy change that occurs when an electron is added to a neutral atom.
Trend: Generally increases (becomes more negative/favorable) across a period.
- Electronegativity: The ability of an atom to attract shared electrons in a chemical bond.
Trend: Increases across a period and decreases down a group. Fluorine (F) is the most electronegative element, followed by Oxygen (O) and Nitrogen (N).
Environmental Relevance: Electronegativity is crucial. The large difference in electronegativity between Oxygen and Hydrogen is what makes the H-O bonds in water (H₂O) highly polar. This polarity is responsible for almost all of water's unique life-supporting properties.
Types of chemical bonds
- Ionic Bonds: Complete transfer of one or more valence electrons from one atom (a metal) to another (a non-metal). This creates charged ions that are held together by electrostatic attraction.
Example: Na + Cl → Na⁺Cl⁻ (Table Salt). In water, these easily dissociate back into ions (Na⁺ and Cl⁻).
- Covalent Bonds: Sharing of one or more pairs of valence electrons between atoms (usually non-metals).
Example: O₂ (non-polar covalent, equal sharing), H₂O (polar covalent, unequal sharing due to Oxygen's high electronegativity). Most organic pollutants are held together by covalent bonds.
- Coordinate Bonds (or Dative Bonds): A type of covalent bond where both shared electrons come from the same atom (the "donor").
Example: The bond between a metal ion (like Fe²⁺) and a ligand (like H₂O) in an aquatic solution.
- Hydrogen Bonds: A weak electrostatic attraction between a hydrogen atom (which is covalently bonded to a highly electronegative atom like N, O, or F) and a nearby electronegative atom.
Example: The bonds *between* water molecules. This "stickiness" gives water its high boiling point, high heat capacity, and surface tension.
Mole concept, molarity and normality
Mole Concept
A mole (mol) is a specific quantity, just like a "dozen" means 12.
1 mole = Avogadro's Number (6.022 x 10²³) of particles (atoms, molecules, ions).
The Molar Mass (g/mol) of a substance is the mass of one mole of that substance. It's numerically equal to its atomic/molecular weight.
Molarity (M)
This is the most common unit of concentration in chemistry.
Formula: Molarity (M) = Moles of solute / Liters of solution
Example: A 1 M solution of NaOH contains 1 mole of NaOH (40.0 g) dissolved in 1 Liter of water.
Normality (N)
This unit of concentration is based on the "reactive capacity" of a substance.
Formula: Normality (N) = Gram equivalent weight / Liters of solution
The equivalent weight is the molar mass divided by the number of reacting units (e.g., H⁺ for acids, OH⁻ for bases, or e⁻ for redox).
- For HCl (1 H⁺), 1 M = 1 N.
- For H₂SO₄ (2 H⁺), 1 M = 2 N.
- For NaOH (1 OH⁻), 1 M = 1 N.
In environmental reports, concentrations are often given in mg/L (milligrams per liter). For water, 1 mg/L is equal to 1 ppm (part per million). You must be able to convert mg/L to Molarity.
Conversion: Molarity (mol/L) = [Concentration (mg/L)] / [Molar Mass (g/mol)] / 1000
Acids, bases and salts, solubility products
- Acids: Substances that release H⁺ ions in solution (e.g., H₂SO₄, HNO₃, H₂CO₃). They have a pH < 7.
- Bases: Substances that release OH⁻ ions in solution or accept H⁺ ions (e.g., NaOH, NH₃). They have a pH > 7.
- Salts: Ionic compounds formed from the reaction of an acid and a base (e.g., NaCl, CaCO₃).
Solubility Product (Ksp)
This describes the equilibrium of a "sparingly soluble" salt dissolving in water.
Example: Lead(II) chloride (PbCl₂), a toxic solid.
Reaction: PbCl₂(s) ⇌ Pb²⁺(aq) + 2Cl⁻(aq)
Solubility Product Expression: Ksp = [Pb²⁺][Cl⁻]²
Environmental Application: Ksp tells us the
maximum amount of a substance that can be dissolved.
- If the product of ion concentrations [Pb²⁺][Cl⁻]² is less than Ksp, more solid can dissolve.
- If the product is greater than Ksp, the solid will precipitate out of the solution.
This controls the concentration of toxic metals in water and the formation of minerals like limestone (CaCO₃).
Solutes and solvents
- Solvent: The substance that does the dissolving (present in the largest amount). In the environment, the "universal solvent" is water.
- Solute: The substance that gets dissolved (e.g., salt, sugar, O₂, CO₂, pollutants).
"Like dissolves like" is the general rule:
- Polar solvents (like water) dissolve polar solutes (like salts, acids, sugars).
- Non-polar solvents (like oil or fat) dissolve non-polar solutes (like DDT, PCBs).
Application (Biomagnification): Non-polar pollutants like DDT are not water-soluble, so they don't get flushed out of organisms. Instead, they dissolve in and accumulate in body fat. This is why they build up in food chains.
Redox reactions
Redox stands for Reduction-Oxidation. These reactions involve the transfer of electrons and are fundamental to all life processes and environmental transformations.
- Oxidation: Loss of electrons (LEO - Loss of Electrons is Oxidation). An atom's oxidation state increases.
- Reduction: Gain of electrons (GER - Gain of Electrons is Reduction). An atom's oxidation state decreases.
Mnemonic: OIL RIG (Oxidation Is Loss, Reduction Is Gain).
Environmental Examples:
- Photosynthesis: Carbon in CO₂ is reduced (gains e⁻) to form sugar (C₆H₁₂O₆).
- Respiration: Sugar is oxidized (loses e⁻) back to CO₂.
- Decomposition (Aerobic): Organic matter is oxidized by O₂.
- Decomposition (Anaerobic): In the absence of O₂, other substances are reduced (e.g., NO₃⁻ → N₂, SO₄²⁻ → H₂S). This is why swamps (reducing environments) smell like rotten eggs (H₂S).
Concepts of pH and pE
These two "master variables" control most of environmental chemistry, especially in water.
pH (Acidity)
A measure of the hydrogen ion (H⁺) concentration. It's a logarithmic scale (a change of 1 pH unit is a 10x change in acidity).
Formula: pH = -log[H⁺]
- pH < 7: Acidic (high [H⁺])
- pH = 7: Neutral
- pH > 7: Basic or Alkaline (low [H⁺])
Environmental Importance: Most aquatic life can only survive in a narrow pH range (e.g., 6.5-8.5). Acid rain (pH < 5.6) can devastate lakes. pH also controls the solubility and toxicity of heavy metals.
pE (Redox Potential)
A measure of the electron activity in a system. It's analogous to pH, but for electrons instead of protons.
Formula: pE = -log[e⁻]
- High pE: Oxidizing environment. Electrons are scarce. (e.g., a fast-flowing, well-aerated river). O₂, NO₃⁻, and Fe³⁺ are common.
- Low pE: Reducing environment. Electrons are abundant. (e.g., a waterlogged swamp bottom). H₂S, CH₄, and Fe²⁺ are common.
pH and pE together determine the fate of pollutants. A chemical's form (species) changes depending on the pH and pE of the environment.
Example (Arsenic):
- In oxidizing (high pE) surface water, Arsenic exists as Arsenate, As(V), which is less toxic and sticks to soil particles.
- In reducing (low pE) groundwater, it becomes Arsenite, As(III), which is highly toxic and mobile, leading to well water contamination.