Sanitary Chemistry Lecture Notes

Chemistry Building Blocks

• Acid solutions are used to lower pH levels.

• Chemistry's foundation lies in balancing charges (positive and negative particles).

• Examples of balanced charges influencing properties: temperature, melting point.

Atomic Structure

• Nucleus: Contains protons and neutrons.

• Proton: Positively charged particle with a charge of +1.602 \times 10^{-19} C.

• Neutron: No charge (neutral).

• Electron: Negatively charged particle with a charge of -1.602 \times 10^{-19} C; lightest of the three.

• A compound must obey the law of balanced charges to be a true compound.

• White gold is an example of a mixture, not a compound, as its constituents aren't chemically bonded.

• 1 \frac{g}{mol} = 1 \text{ amu} = 1 \text{ Dalton}

• Helium, a noble gas, wasn't evident in early periodic tables.

• The modern periodic table organizes elements by atomic number, based on Mendeleev's original concept.

Periodic Table Groups

• Main-group elements: Found in the eight A groups (1A-8A).

• Transition elements: Located in the ten B groups.

• Inner transition elements: Two series (lanthanides and actinides) placed between Groups 3B and 4B.

• Elements within a group share similar chemical properties; elements within a period have differing properties.

• Group 1A: Alkali metals (excluding hydrogen, which is a non-metal).

• Group 2A: Alkaline Earth Metals.

• Group 3A: Boron Family.

• Group 4A: Carbon Family.

• Group 5A: Nitrogen Family.

• Group 6A: Oxygen Family (Chalcogens).

• Group 7A: Halogens.

• Group 0 (8A): Noble gases.

Element Distinctions

• Metals: Shiny solids at room temperature that conduct heat and electricity well; malleable and ductile.

• Nonmetals: Generally gases or dull, brittle solids at room temperature; poor conductors of heat and electricity.

• Metalloids: Semi-metals.

Periodic Trends

• Atomic Size: Increases with the number of electron shells.

  • Atomic size increases significantly down a group.

  • Atomic size generally decreases across a period (increasing proton number pulls electrons closer).

• Ionization Energy: Energy needed to remove an electron from an atom.

  • Ionization energy decreases down a group.

  • Ionization energy generally increases across a period.

• Electron Affinity: Energy released when an electron is added to an atom.

  • Electron affinity decreases down a group for non-metals (like halogens).

  • Electron affinity increases across a period.

• Electronegativity: An atom's tendency to attract electrons.

  • Electronegativity decreases down a group.

  • Electronegativity increases across a period. (Pauling Scale)

• Each period adds an energy level.

Molecular Perspective

• Molecule: A structure of two or more chemically bound atoms acting as a unit.

Calculations and Examples

• Number of electrons calculations:

  • Example 1: 1(K) + 3(G) + 2 = 26

  • Example 2: 1(S) + 3(a) + 1 = 24

  • Example 3: 1(A) + 4(G) + 2 = 32

  • Example 4: 1(A) + 2(G) + 2 = 18

• The positive charge of Ca^{+2} is reduced to minus.

Atomic Mass and Molecular Weight

• Atomic masses of elements:

  • C = 12.01

  • H = 1

  • N = 14.01

  • O = 16

• TNT (C7H5N3O6) Molecular Weight Calculation:

  • MW_{TNT} = 7(12.01) + 5(1) + 3(14.01) + 6(16) = 227.07 \frac{g}{mole}

• Moles Calculation Example:

  • Moles = \frac{650g}{227.07 \frac{g}{mole}} = 2.864 \text{ moles}

• Number of Molecules Calculation:

  • 2.86 \text{ moles} \times 6.022 \times 10^{23} \frac{\text{molecules}}{\text{mole}} = 1.725 \times 10^{24} \text{ molecules}

Chemical Formulas

• Sodium hypochlorite: NaOCl ; oxidation state -2

• Slaked lime: Ca(OH)_2 ; oxidation state +2

• Quicklime: CaO ; oxidation state +2

• Ammonium Sulfate: (NH4)2SO4 ; ions NH4^+ and SO_4^{2-}

• Barium Phosphate: Ba3(PO4)_2

• Stannic oxide: SnO_2; tin ion Sn^{4+}

Calculations

• Percentage Calculations:

  • 24.96

  • 68.65

  • 56.2

  • 20.03

  • 33.3

  • 19

• Ferrous and Ferric compounds

  • 27.9/18.6

  • 86.2

• Atomic Mass Example: Calculate Molecular Weight of Fe2(SO4)_3

  • MW{Fe2(SO4)3} = 2(55.8) + 3(32.1) + (3)(4)(16) = 399.9 \frac{g}{mole}

• Another Formula

  • Fe(SO_4) formula = C

• Equivalent Weight Calculation Example:

  • Equivalent = MW/valence

  • Normality = # equivalent = 30

• Normality Calculation

  • NA VA = NB VB

Equations and Concepts

• If equivalence is not given

• Reducing agent

• Oxidizing agent

• Mole calculations:

  • Moles = \frac{mass}{MW}

  • Example: Moles = \frac{35}{MW} = 0.53 \text{ moles}

• Normality Equation for H2SO4:

  • N = \frac{\text{# of } H2SO4 \text{ solute}}{V_{sol}}

  • Normality = 0.816 N = 0.13 Molar

• Mole Calculation Example

  • Moles = \frac{mass}{MW} = 0.125 \text{ moles}

  • MW = 23 + 16 + 1 = 40 \frac{g}{mole}

Chemical Reactions

• Stoichiometry: Reactants and products have numerical relationships.

• Acid-Base Reactions (Neutralization):

  • HCl + H2O = H3O^+ + Cl^-

  • Acids release hydrogen ions and anions when added to water.

• Gas Producing Reactions: Reactions proceed to completion as gas escapes.

• Ammonia Stripping Example: Nitrogen removal from wastewater by raising pH above 10.8, converting NH4^+ to NH3 gas.

Physical States of Matter - Gases

• Gases vs. Liquids and Solids:

  1. Gas volume is highly pressure-sensitive.

  2. Gas volume is highly temperature-sensitive.

  3. Gases have low viscosity.

  4. Gases have low densities under normal conditions.

  5. Gases are miscible.

Water Impurities - Turbidity

• Turbidity definition: Cloudiness due to suspended particles that interfere with light passage.

• Common Impurities:

  • Suspended Impurities:

    • Colloids (negatively charged particles)

    • Suspended inorganic matter

    • Suspended organic matter

  • Living matter

Turbidity Measurement Methods

• Jackson Candle Turbidimeter: Measures depth at which a candle is no longer visible through the water column.

• Turbidimeter (Nephelometer): Measures scattered light at a 90° angle.

• Secchi Disk: Measures maximum depth at which a black and white disk is visible.

• Turbidity Tube (Transparency Tube): Combines Jackson candle and Secchi disk principles.

Gas Laws

• Boyle's Law: Volume varies inversely with pressure at constant temperature (isobaric condition).

• Charles's Law: Volume varies directly with absolute temperature at constant pressure.

• Generalized Gas Law: Combination of Boyle's and Charles's laws.

• Dalton's Law of Partial Pressure: Partial pressure is proportional to gas amount in a mixture.

• Raoult's Law: Describes vapor pressure of ideal solutions based on volatility and mole fraction.

  • Ideal Solution: Properties are a molar average of component properties.

• Henry's Law: Gas weight dissolved in liquid is proportional to the gas's pressure above the liquid.

  • Applications: Aeration for gas removal, oxygenation of sewage, industrial wastewater gas removal.

• Graham's Law: Gas diffusion rates are inversely proportional to the square root of their density.

• Gay-Lussac's Law of Combining Volumes: Gas volumes in reactions are related by whole numbers.

Turbidity Environmental Significance

• Aesthetics: Turbidity suggests contamination.

• Filterability: Turbidity reduces filter lifespan and increases purification costs.

• Disinfection: Particles shield pathogens from disinfectants.

Historical and Current Turbidity Standards

• Original Standard: JTU (Jackson Turbidity Unit), measured using Jackson tube turbidimeter.

  • 1 \frac{mg}{L} SiO_2 = 1 \text{ unit of turbidity}

• Current Standard: Uses formazin polymers, measuring light scattered at right angles.

  • NTU - Nephelometric Turbidity Unit

• Common Standard: Secchi Disk use (black and white disk)

Turbidity Levels

• Noticeable to average consumer: 5 JTU

• Clear lake: 25 JTU

• Muddy water: Over 100 JTU

pH Measurement

• pH is a measure of hydrogen ion concentration.

• pH Indicators: Change color in specific pH ranges. Examples:

  • Thymol Blue: 1.2-2.8 and 8.0-9.2

  • Methyl Orange: 3.0-4.4

  • Bromphenol Blue: 3.0-4.6

  • Congo Red: 3.0-5.0

  • Bromcresol Green: 3.8-5.4

  • Methyl Red: 4.4-6.2

  • Bromthymol Blue: 6.0-7.6

  • Phenol Red: 6.8-8.4

  • Phenolphthalein: 8.2-10.0

  • Thymolphthalein: 9.3-10.5

• pH meters use glass electrodes to measure hydrogen-ion activity.

pH Scale

• 0 to 14 scale: 7.0 is neutral, below 7 is acidic, above 7 is basic.

• Examples:

  • Battery acid: ~1

  • Human blood: 7.4

  • Drain cleaner: 12.8

• Acids taste sour and react with bases; bases taste bitter, feel slippery, and react with acids.

Application of Turbidity Data

• Water Supply: Determine need for coagulation/filtration; monitor groundwater quality.

• Domestic/Industrial Water Treatment: Assess suspended solids removal; adjust chemical dosages.

pH Definition and History

• pH = -log[H+], where [H+] is the hydrogen ion concentration.

• Søren Sørenson (1868-1939) introduced pH around 1909.

Litmus Paper

• Oldest pH indicator: Dye mix from lichens (e.g., Roccella tinctoria).

• Blue litmus turns red in acid, red litmus turns blue in base, neutral is purple.

Acidity and Aquatic Life

• Rapid pH changes stress fish; changes >0.3 units/day harmful.

• Examples of water samples with pH levels: Trout, Bass, Perch, Frogs etc.

Acidity Measurement and Treatment

• Treatment: Remove corrosive CO2 via aeration or neutralization with lime/NaOH.

• CO2 removal is key in water softening.

• Neutralize mineral acidity in industrial wastes before discharge.

Alkalinity

• Alkalinity: Capacity to neutralize acids.

• Sources: Bicarbonates (from CO2 reacting with soil), salts of weak acids (borates, silicates, phosphates), humic acids, ammonia/hydroxides, and salts like acetic acid or H_2S in polluted water/Algae.

• Components: Hydroxide, carbonate, and bicarbonate.

Alkalinity and Hardness Calculations

• Given T=80mg/L and A=80mg/L

  • Carbonated Hardness A = 80mg/L

  • Non-carbonated Hardness = T-carbonate = 112-80 = 32mg/L

• If T=80mg/L and A=112mg/L:

  • Carbonate = T = 80mg/L

  • Non-carbonate = T - carbonate = 80 - 80 = 0

• Calculations with P and T

  • If: P = 10mg/L, T = 52mg/L

  • Bicarbonate = T - 2P = 52 - 2(10) = 32mg/L

  • Carbonate = 2P = 20mg/L

  • Hydroxide = 0

• Comparing the results the value should be less than 26Mg/L.

Hardness as CaCO_3

• Equation:
  \text{Hardness as } CaCO3 = 127 \frac{mg}{L} \text{ as } CaCO3

• Calculations:

  • Ca hardness = 65.9mg

  • Mg hardness = 37.16

  • Total hardness:
    \text{Total hardness} = \text{CA hardness} + \text{Mg hardness}
    \text{Total hardness} = 107.085 mg

• From toothpaste to simpler hydrocarbons.

Organic Compounds and Functional Groups

• Many organic compounds contain heteroatoms (atoms other than C or H) like N, O, S, P, and halogens.

• Functional Group: Specific arrangement of atoms (often with C=C or C-heteroatom bonds) that reacts characteristically.

Hydrocarbons

• Simplest organic compounds: Contain only H and C.

• Chains - aliphatic compounds. Rings - aromatic compounds.

Classification of Hydrocarbons

• Alkanes, Alkenes, Alkynes.

Alkanes

• Contain only single bonds.

• General formula: CnH{2n+2}

• Saturated hydrocarbons (max # of bonds).

• Aliphatic hydrocarbons from petroleum (methane, ethane, etc.).

• Paraffins: Longer aliphatic waxes.

• Methane (CH4) is the simplest; explosive when mixed with air.

Alkenes

• Have at least one C=C bond.

• Also called ethylene series.

• General formula: CnH{2n}

• Unsaturated hydrocarbons.

• Olefins – multiple carbon-carbon bonds.

• Formed as byproducts of petroleum breakdown (e.g., ethylene, propylene).

• Hydrogenation: Addition of hydrogen gas under controlled conditions.

Alkynes

• Have at least one C \equiv C.

• General formula: CnH{2n-2}

• Named similarly to alkenes with -yne suffix.

• More reactive than alkanes.

Aromatic Hydrocarbons (Benzene Series)

• Contain one or more benzene rings + alternating single and double bonds.

• Benzene (C6H6) is the parent compound.

• More rings = less soluble and more persistent.

Alcohols

• Hydrocarbons with at least one hydroxyl (-OH) group.

• Example: Methyl alcohol (methanol) - CH_3OH

Straight Chain Alkanes

• Methane: CH_4

• Ethane: C2H6

• Propane: C3H8

• Butane: C4H{10}

• Pentane: C5H{12}

• Hexane: C6H{14}

• Heptane: C7H{16}

• Octane: C8H{18}

• Nonane: C9H{20}

• Decane: C{10}H{22}

Halogenated Hydrocarbons

• Contain at least one halogen atom (e.g., chloromethane).

Carboxylic Acids

• Highest oxidation state for organic radical.

• Further oxidation forms CO_2 and water.

Functional Groups

• Formic Acid: Formula: HCOOH, Proper name ending oic acid.

• Acetic Acid: Formula: CH_3COOH, Example in IUPAC name:Ethanoic acid

• Propionic Acid: Formula: CH3CH2COOH, Example in IUPAC name: Propanoic acid

• Butyric Acid: Formula: CH3CH2CH_2COOH, Example in IUPAC name: Butanoic acid

Other Organic Compounds

• Glycols: Two H atoms substituted by hydroxyl groups (antifreeze, anesthetics).

• Ethers, ketones, aldehydes, esters, cyanides.

• Carbonyl Group: C=O grouping.

• Aldehyde: Contains carbonyl group with C bonded to H.