The Soil Chemical Environment

1. The Soil Chemical Environment Reading: General background: Sparks,Chapter 1, pp. 1-28 Additional: Essington, Chapter 1 pp. 1-21.

2. Chemical Interactions in Soils

3. Description of complex interactions involves • Mineral Chemistry • Colloidal Chemistry • Physical Chemistry • Analytical Chemistry • Organic Chemistry • Biochemistry

4. Reactions occur at interfaces • Solid - liquid • Liquid - gas

5. System is open • fluxes of water and solutes • fluxes of gases • fluxes into biota and from decaying biota • energy fluxes

6. Equilibrium vs. Kinetics • Most soil chemistry is the study of reactions at equilibrium. • This is OK for prediction of fast reactions (e.g. ion exchange and many adsorption reactions), but soil is mostly a non-equilibrium system.

7. Kinetics • Very important for many reactions (e.g. precipitation and dissolution of most minerals). Hard to study in a mixed system like soil. • Kinetics can be used to describe the rates of all reactions but in many cases the rates are so slow that over the time of interest there is no reaction. Kinetics in mixed systems can be very complex.

8. Soil Chemistry and the Soil Solution • The soil solution is central to most soil chemistry • Soil chemistry is mostly about the interaction of ions and molecules in solution and their interaction with the gas phase and the solid phases. • Most concepts similar to those in aquatic chemistry courses.

9. Soil Solution Interaction with other components (Sparks Fig. 4.1)

10. Inorganic anions and cations in the soil solutions (Sparks Table 4.1)

11. Chemical systems in soils

12. Properties of the Elements • Web Elements has details on the properties and behavior of all elements • Web Elements http://www.webelements.com/

13. Chemical Elements in Soil

14. Abundance of the elements in soils • Major vs. minor and trace elements • In geochemistry and soil chemistry these the definitions are a bit fuzzy. • See also plant nutrition literature • Mostly “trace” and “minor” are used interchangeably • Trace or minor is sometimes defined as < 1% (10 g kg-1 ). Essington says trace is < 100 mg kg-1 • Some authors consider minor as including higher concentrations than trace

15. Major elements • Major: Al, Si, C, N, Fe, Ca, O, K, Ti (P in plants) • Remainder are minor (trace) elements see Table 2.4 in Sparks

17. Periodicity of the elements • Metals • Group 1 - alkali metals (+I), 1+ ions • Group 2 - alkaline earths (+II), 2+ ions • Transition elements • Non-metals • Group 6 chalcogens (- II) • Group 7 halogens (- I), 1- ions

18. Heavy metal?? • It is such a vague word, that is often used incorrectly. • Is boron a “heavy metal”?

19. Oxidation states and charge number • Oxidation state is indicated by Roman numerals. • E.g. Fe(III), Fe(II), Mn(IV), and Mn(II) • Charge number on an ion. • E.g. Ca2+ not Ca+2   • SO42- or with oxidation state S(VI)O42-

20. Acids and Bases See Sparks p. 65

21. ACIDS AND BASES • Arrhenius • Acid increases H+ concentration (activity) in solution. • First and least inclusive definition. • Base increases OH- activity. • Brönsted definition - typical definition used in environmental and soil chemistry • Acid is a proton donor. • Base is a proton acceptor • Includes Arrhenius acids and bases

22. Brönsted acids and Bases • Examples HCl + H2O = H3O+ + Cl- acid base conjugate acid conjugate base NH3 + H2O = NH4+ + OH- base acid CH3COOH + OH- = CH3COO- + H2O acid base

23. Lewis Acids and Bases • Lewis Acid – a broader definition. (Brönsted definition is subset of the Lewis definition) • Acid is an electron pair acceptor • Base is an electron pair donor. • Examples: • (see Brönsted acids and bases) e.g. H+ + H2O = H3O+

24. LeH+ is also a Lewis Acid • H+ is a proton --- can accept a pair of electrons from one of the two unshared pair of electrons in water to form a coordinate covalent bond. • water hydronium ion

25. Metal ion complexes • Lewis defines formation of complexes as acid base reactions. • e.g. Formation of Cu2+ amine complex ion. • Cu2+ + 2NH3 = Cu(NH3)22+ • Ammonia has one unshared pairs of electrons that can be donated to empty bonding orbitals in Cu2+.

26. Metal ion complexes • On the last slide NH3 is a ligand • The complex is also called and adduct (addition product). Not a term used very much in environmental chemistry.

27. Chelation • Chelation (multidentate complexes) • Fe(II) can accept 2 pairs of N electrons • Bidentate complex • This red complex is useful for detection of reduced Fe.

28. EDTA complexes

29. N atoms can donate pairs of electrons • Negatively charged O atoms can donate pairs of electrons. • Negative charges contribute to ionic bonding • Can form hexadentate complexes • e.g. Cu2+ and other first row transition metal cations. • Complexes with EDTA can be very strong

30. EDTA, is shown in 3-D at • EDTA • Fe DTPA, a similar complex is at: • Fe-EDTA

31. Systeme International (SI) Units

32. Based on mks system • Basic units; m, kg and s • Examples: • Concentration in a solid, mol kg-1 • Rate of reaction, mol L-1 s-1 • Writing units • mol L-1 s-1 notmol/L/ s

33. Essington Table 1.5

34. Some non SI (derived) units • We use many units that are not strictly part of the S.I. system • e.g. L for liter and ha. for hectare. • Land application - kg ha-1

35. Some equivalences • Mg m-3 = g cm-3 • mg kg-1 or mg L-1   =ppm • µg kg-1 or µg L-1   =ppb • c molc kg-1 same as meq/100g

36. Conversion of units • Example: Reporting 5 mg L-1 Ca extracted from a 10 g soil in 200 mL solution in units of cmolckg-1 = 0.5 mmolc kg-1

37. Essington Table 1.6

38. In class exercise • What is the common language unit for g T-1 (gram per ton)?

39. Summary • Soil solution is central to soil chemistry. • Soil chemistry concentrates on a small fraction of the periodic table. • Knowledge of periodicity is very useful. • We will mostly use the Brönsted definition of acid and bases. • Will use mostly SI units and a few derived from S.I