1. WATER CHEMISTRY Chapter 1 Basics of Aquatic Chemistry

2. Chemical equilibrium Size and shape of water molecules; space between molecules Conc’n of water in environmental solutions Polarity Hydrogen bonds Dissolution Dissociation Salts, ions Hydrophobic/hydrophilic Conc’n of solutes in environmental solutions Units for solute conc’ns – mol/L, ppm, ppb, equiv/L pX Ionic strength, conductivity, TDS, hardness Chemical (re)activity Introductory Terms and Concepts

3. Course Overview Focus on equilibrium composition of aqueous solutions: what is in a given system; what we have to do to alter it in a desired way – Equilibrium: Applies to specific reactions (not individual species); Condition of no net production/destruction of chemicals via that reaction – Composition: Emphasis is on speciation, i.e., the distribution of the components among different chemical forms H 2 S vs SO 4 2  NH 4 + vs. NH 3 Cr 3+ vs. CrO 4 

4. Example: What are the molar concentrations of (a) H 2 O in pure water, and (b) Ca 2+ in a solution containing 20 mg/L Ca 2+ ?

5. Water: Molecular Structure and Key Properties Molecular structure and size

6. Water: Molecular Structure and Key Properties Dipolar, facilitating interactions with other dipoles or charged molecules (ions)

7. Water: Molecular Structure and Key Properties Structure facilitates formation of hydrogen bonds, leading to unusually high density and cohesiveness (high boiling point, surface tension) Up to four H bonds can form; three typically exist at normal temperatures

8. Characterizing Water/Solute Interactions Dissolution: Surrounding of individual molecules of solute by water molecules – H bonds between H 2 O molecules and solute-solute bonds between dissolving molecules must break, both of which oppose dissolution – Formation of bonds between H 2 O and solute favors dissolution – Balance between net changes in bonding energy and randomizing effect of molecular kinetic energy controls dissolution; statistical process, so some dissolution always occurs; designation as “hydrophobic” or “hydrophilic” is always relative to some implicit baseline

9. Characterizing Water/Solute Interactions Some compounds dissociate (i.e., split apart) when they dissolve – Compounds that dissociate extensively and release ions in the process are called salts – Positive ions are cations; negative ions are anions; all ions are hydrophilic

10. Expressing Solute Concentrations Dimensions and Units – Mass/volume (e.g., mg/L, mg/m 3, mol/L [molar, M]) – Mass/mass (e.g., mg/kg, mass fraction, ppm m, ppb m, mol/kg [molal, m]) – Mol/mol (mole fraction) – Volume/volume (e.g., volume fraction, ppm V ); used primarily for gases – Note: subscripts on ppm, ppb, etc., often implicit – mass for liquids and solids, volume for gases; ppm in water is ~mg/L The “p-convention”: pX ≡  log 10 (X)

11. Expressing Solute Concentrations Composite or Surrogate Concentrations – Used to characterize mixtures of species that behave similarly or when the individual species are unknown – Examples Total or Dissolved Organic Carbon (TOC or DOC), Alkalinity (ALK) Total Organic Halogen (TOX) Hardness (Hd 2+ ) Free Available Chlorine (FAC) Ammonia-N (NH 3 -N) – Note several ways to represent concentrations in terms of elements: DIC, NH 3 -N, AsO 4 as As, TOTCr

12. Example (HW 1-3) The concentration of silver (Ag) in seawater is 50 parts per trillion, and the volume of seawater in the world is 1370×10 6 km 3.  sw =1.025 kg/L, AW Ag =107.9. a)How much Ag is in the oceans of the world? b)What volume of seawater would have to be processed to recover 1 kg of Ag if the process were 100% efficient? c)What is p[Ag] in the ocean?

13. Example (HW 1-16a) A solution contains 30 mg/L NH 4 + and 5 mg/L NO 3 . a)Express these concentrations as mg/L NH 4 -N and NO 3 -N, respectively. Which species makes a larger contribution to the total N concentration in the water?

14. Concentrations Expressed As an “Equivalent” Amount of a Different Species – Usage – Basis of the equivalency generally inferred from context Examples – Charge – Acquisition/release of H + in acid/base reactions – Acquisition/release of e  in oxidation reactions

15. Ionic Strength I or  : A composite, calculated (not measurable) property of the solution (not of individual species) – Always calculated using c values in mol/L, but commonly reported without dimensions – Includes contributions from all ions, but not neutral species – Indicator of the ability of the solution to neutralize the electrical field surrounding an ion (by attraction of like-charged and repulsion of oppositely-charged ions) – High ionic strength reduces an ion’s “sphere of influence” and shields it from interactions with the rest of solution

16. Composite Parameters for Overall Solute Concentrations Conductivity, Specific Conductance  microSiemens,  S) – In situ measurable indicator of overall ionic content of a solution – Includes contributions from all ions, weighted by their charge and mobility Total Dissolved Solids (TDS; mg/L) – Non-specific indicator of overall mineral content – Measurable offline – Includes contributions from all solutes that do not easily volatilize (evaporate)