SOIL FERTILITY DECOMPOSITION Nutrient dynamics and energy flow

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SOIL FERTILITY DECOMPOSITION Nutrient dynamics and energy flow Powered By Docstoc
					BIO/EVS 481--TERRESTRIAL ECOLOGY FALL 2004

SOIL FERTILITY & DECOMPOSITION Nutrient dynamics and energy flow are a major focus of ecosystem ecology. Understanding processes related to soils are fundamental to the management of ecological resources and is perhaps the most important research area in ecosystem ecology. Soils control what types of plants grow in an area, they affect rates of carbon storage in ecosystems, and they determine the fertility of a site. Soil processes are also receiving much attention in light of global change. With atmospheric levels of CO2 rising rapidly as a result of industrialization, there is a tremendous amount of interest in determining how this added carbon will affect our climate and ecosystems. In fact, the study of carbon accumulation in ecosystems is one of the most actively researched and funded areas of ecological research today. It has direct relevance on the functioning of ecosystems and how they may respond to anthropogenic changes in the Earth's carbon cycle. Today in lab, we will be investigating three factors including soil nitrogen availability, cation exchange capacity, and decomposition rates. Each of these measures is critical for a variety of ecosystem functions and have greater implications to ecosystem health and productivity. For our study, we will be revisiting Hitchcock Nature Area and again comparing the forest and prairie on opposing hill aspects. I. Litter Decomposition

Decomposition is a key ecosystem process linking the vegetation (or animal) and soil nutrient pools. Therefore, it is essential in the recycling of plant essential nutrients. Decay has also received much attention for its role in global climate change. Soil organic matter represents a large but relatively available C pool globally. Although it is assumed and largely accepted that plants and oceans will be able to consume some of the anthropogenically produced atmospheric carbon dioxide, the role of soils remains known to a lesser degree. In northern latitudes, boreal forest and tundra have sequestered large amounts of C in the soil because decomposition rates are very slow; however, global warming is projected to be greatest near the poles and may enhance decomposition. If this scenario plays out, increased C released from decomposition may counter any increased C assimilated in NPP. Our decomposition samples are those that you constructed earlier in the semester using tree litter from Hitchcock and placed in fiberglass window screen litterbags. Nine litterbags were placed in each of 2 plots in the forest and 2 plots in the prairie, very near the transects that you used earlier in the semester. Litterbags were placed at Hitchcock September 15 and collected December 8. When the litterbags were placed in the field, several bags were immediately returned to the laboratory and oven dried to be used as a correction factor for any mass lost due to transportation or handling and oven versus air drying (called “travelers”). While the use of tree litter in both forest and prairie sites is imperfect as ideally we would use tree litter in the forest and grass litter in the prairie, it does represent a uniform litter type allowing as to more directly gauge the role of environment in decomposition rates between the two sites. Additionally, the litter used was not fresh litter but likely last year’s senesced leaves already partially decomposed. Nor was the timeframe for this experiment ideal. If time was not a constraint, we would have used fresh litter, allowed the experiment to continue for at least 2 years, and periodically sampled bags to determine percent mass remaining at multiple time intervals. Regardless, this experiment still provides an adequate means of comparing the effects of environment on litter decomposition. Procedure: The litterbags have already been returned to the lab, the contents emptied into paper bags and placed in the oven overnight. You must find your litter samples. Remove non-leaf material: place sample in 1mm sieve and shake. If obvious plant material falls through the sieve, return it to sample, or if obvious non-plant material remains in sample, remove it. Weigh the sample and record on class data sheet.

BIO/EVS 481--TERRESTRIAL ECOLOGY FALL 2004

After all of the samples have been compiled for the class, we will complete the calculations necessary to determine decay rates (k) for the two sites. First, because we are interested in decay rates expressed on an annual basis, we must convert the time period from number of days in the field (88 days) to number of years in the field (88/365=0.241 yrs). Second, we must determine the correction factor from the “travelers” which is then multiplied by each samples final dry mass. Third, average initial and final masses (independently) by site. Fourth, plot the natural log of litter mass against # years in the field- this plot will only consist of 2 points for our experiment, the initial mass (at time 0) and the final mass (at time 0.241 years). Fifth, add the trendline (linear regression curve) and determine the slope which will equal k, the decay rate, or the proportion of the litter that decomposes over the period of a year. The higher decay rate of the two sites should correspond with the site that has the lower average percent mass remaining. II. Cation Exchange Capacity Cation exchange capacity (CEC) is defined as the capacity of a soil to hold exchangeable cations on negatively charged sites at the surface of soil minerals and soil organic matter. In addition to affecting soil mineral weathering, leaching, and buffering soil pH, this measure can provide an index to the fertility of soils as many soil cations are essential plant nutrients. Given the fact that the soils at our two sites at Hitchcock are very similar in texture and because CEC is determined largely by two factors, soil clay and soil organic matter, our determination of CEC should reflect the difference in the amounts of soil organic matter. As we discussed in lecture, many soil colloids have a net negative charge. As a result, cations are attracted to and in a sense bound to soil colloids. In fact, the ability of a soil to bind cations is proportional to the number of negative charges per unit weight of soil. There are a number of ways to measure CEC. Different labs tend to use modifications of the same basic technique. Most are based on the principle of cation displacement. This refers to a process in which the cations occupying the exchange complex are displaced by 'flooding' the soil with excess cations from reagents such as barium chloride (Ba2+ is the displacing cation). CEC is expressed in units that may appear rather unusual, unless you have previous experience in soil science. In the older literature, the units used were meq/100 g. The newer standard units are centimoles of positive charge per kilogram of dry soil, written as cmoles+ kg-1. These units are meant to indicate how many cations can be held per unit weight of soil. You will determine the CEC of your sample by using a cation displacement type procedure. This involves displacing the cations already on the complex with cations of your choice. You do this by mixing the soil with a solution rich in cations. In our lab we will do this by mixing with a barium chloride solution. By 'flooding' the soil with barium you replace whatever cations were already on the soil with Ba2+ ions. The ions that you displace end up in the soil solution. By centrifuging the sample, you separate the soil solution from the soil. So when you pour off the supernatant you pour off the displaced cations. Therefore, in the first part of the procedure you displace the cations on the soil complex with Ba2+ cations. And you have now saturated the cation exchange complex on the soil with Ba2+ ions. If you then had a way to measure how many Ba2+ ions were on the soil exchange complex you could determine the CEC. How can we do this? The first step in this process is to wash the soil with a neutral compound like deionized water. Because it is electrically neutral, it will not displace any of the barium ions on the exchange complex, but it will wash away any Ba2+ ions that may still be present in the small amount of soil solution left after centrifuging. In the final part of the procedure, you want to wash the Ba2+ ions off the soil and in effect

BIO/EVS 481--TERRESTRIAL ECOLOGY FALL 2004

'count them' or quantify them to develop an estimate of CEC. To do this you simply flood the soil with another cation. In this lab we will use Mg2+. Now Mg2+ occupies the entire cation exchange complex and the barium ends up in the supernatant. This is why you save your supernatant in the final steps of the procedure. This liquid contains the Ba2+ ions that you will measure via titration. Procedure: For each of 4 samples: • Add 2.0 g of sieved soil to a centrifuge tube, cap and weigh (w1) • Add 40 mL of buffered barium chloride solution, cap tube and shake vigorously for 5 minutes • Centrifuge the mixture and discard the liquid • Add 40 mL distilled water, cap tube, and shake thoroughly, centrifuge, discard liquid • Re-weigh capped tube (w2) • Pipette 20 mL of magnesium sulfate into tube, cap tube and shake vigorously for 5 minutes • Centrifuge and transfer liquid (Keep the solution!) to a stoppered flask • Transfer 5 m: of the final extract into a 125 mL flask, add 10 mL of 2 M NH4OH, and titrate with standard EDTA solution using two drops of the indicator solution. The end-point is shown by a change in color from red to blue. Record this end-point (A1 mL). • Titrate 5 mL of the original magnesium sulfate solution under similar conditions, recording the end-point (B mL) • Calculations: 1. The titration end-point for the soil extract (A1) must be corrected for the volume of liquid retained by the centrifuged soil: Corrected end-point (A2): = A1 × (100 + w2 - w1)/100 2. CEC = 4 (B - A2) (units are meq/100 g soil) III. Nitrogen extraction Recall that terrestrial ecosystems are largely limited by the availability of nitrogen and/or phosphorus. For reasons discussed in class, we can speculate that plant production at Hitchcock (and especially the prairie) is limited by N and not P. Nitrogen primarily exists as two inorganic, plant-available forms in the soil: nitrate (NO3-) and ammonium (NH4+). Of the two forms, ammonium is typically in higher concentrations in the soil solution because it is produced first by microbes (it is a reactant in the nitrification process) and only certain nitrifying bacteria are able to produce nitrate. For that reason and the fact that extracting and analyzing ammonium in soil is a simpler process, we will be determining the concentration of ammonium in the two sites. This will provide us a snapshot of soil N availability for which to compare the two sites. Our timing is less than ideal as the plants and soil organisms are currently largely dormant. This measure can be improved upon by measuring the rate of inorganic N production (mineralization). This is done in the field by installing a capped tube in the soil which prohibits plant uptake of soil N but permitting soil organisms to continue to be active and consume soil organic matter and therefore mineralizing N (other methods are available but this is the most widely used). Because the soil microbes in the tube are also using some of that mineralized N for their own uptake and assimilation, the difference in soil inorganic N when the tube was installed and at some time later (perhaps a month) when the soil from the tube is removed, equals net N mineralization which is a good index of plant available N. Many techniques used to determine nutrient analyses in solution involve colorimetric or spectrophotometric analyses. The basis for such methods is the development of a particular color (wavelength) in the solution. Color development is based on the nutrient concentration and the appropriate reagents added to the solution. The intensity of the color in the samples (the amount of color development) is proportional to the nutrient concentration and can be compared to a standard curve in which a known range of concentrations is prepared for spectrophotometric analysis. The basis for spectrophotometric analysis of samples is Beer’s Law which states that log (Io/I)=abc where:

BIO/EVS 481--TERRESTRIAL ECOLOGY FALL 2004

Io = intensity of light being received by the solution I = intensity of light after passing through samples Io/I = abosorbance (A) a = absorbance coefficient (or extinction coefficient) b = pathlength (= diameter of cuvette) c = concentration of the sample Therefore, the absorbance is proportional to the concentration of the nutrient being analyzed. The basic set up of a spectrophotometer consists of a broad spectrum light, which after passing through a monochrometer (to achieve the appropriate wavelength for the particular color or nutrient) is passed through the sample. The intensity of the light is then measured by a photodetector and the absorbance then calculated by Beer’s Law. The standard curve that is developed from known concentration yields an equation which can be used to determine concentrations of unknown samples from their absorbance (A) values. For example, if the absorbance is graphed as a function of the concentration in the standard curve, the regression equation would be in the form A = bc +a (where b=slope & a=y intercept). This equation is then transformed to calculate the concentrations of your soil solutions (c): c = (A-a)/b. There are a few limits to this approach. First, the chemical composition of the samples must be at equilibrium. Also, sample concentrations must be within the range of concentrations of the standard curve, otherwise dilutions will have to be done (either by increasing pathlength or by simple dilution with the appropriate matrix prior to the addition of reagents). Lastly, the wavelength at which the light intensity is to be measured is best at the wavelength which yields peak absorbance. For ammonium analysis, we will use the indophenol blue method which measures both NH3 and NH4+ by converting NH4+ to NH3 under basis conditions (pH=11) (NH3 + H2O ↔ NH4+ + OH-). This reaction enables the determination of NH3 concentrations as follows: phenol + NH3 (+ hypochlorite + sodium nitroprusside) indophenol blue. The blue color (indophenol blue) is produced by the reaction of NH3 and phenate in the presence of a strong oxidizing agent (hypochlorite). The sodium nitroprusside is a catalyst used to quicken the reaction. Although analytical equipment is available (unfortunately not to us) to perform these reactions automatically, we will complete them manually and analyze absorbance on a Spec 20. Before we perform the above chemistry, we must first extract the NH4+ from the soil. To do this, we add a KCl solution to the soil. The K+ ions will displace NH4+ ions in the soil leaving the NH4+in soil solution. Procedure: • Sieve enough soil for 4 samples (no more than 2 cups). • Weight (record exact mass) approximately 12 g soil into each of 4 125 mL flasks • Add 50 ml 2 M KCl to each flask, cover with parafilm, and shake vigorously for 30 minutes (keep track of this time). • Filter the soil-KCl suspension through a Whatman 42 filter into another 125 mL flask. • Pipette a 3 mL aliquot of the filtrate in a 25 mL volumetric flask. • Add 1 mL EDTA reagent and mix. Let stand for 1 minute. • Add 2 mL phenol-nitroprusside reagent followed by 4 mL of the buffered hypochlorite reagent and immediately dilute the flask to volume with distilled water and mix well. Place the water in the water bath (40˚C) and allow it to remain for 30 minutes. • Remove the flask and allow to cool (~ 10 minutes) and determine the absorbance of the colored complex at a wavelength of 636 nm. • Determine the NH4+-N concentration of the sample by reference to a calibration (standard) curve plotted from results obtained with a 25 mL standard sample containing 0, 2, 4, 6, 8, 10, and 12 µg of NH4+-N. To prepare this curve, add 3 mL of 2 M KCl to a series of 25 mL volumetric flasks. Then add 0, 1, 2, 3, 4, 5, 6 mL of the 2 mg NH4+-N L-1 solution to the series of flasks and measure the intensity of blue color developed with these standards by the procedure described for analysis of the extract.

BIO/EVS 481--TERRESTRIAL ECOLOGY FALL 2004

CEC Materials: Centrifuge and tubes (24) Flasks (24 125mL) Pippettes (50 mL) Burettes (4 10 mL) with stands & clamps Graduated cylinder (100 mL) Reagents: Distilled water Triethanolamine solution: Dilute 90 mL of triethanolamine to 1 L and adjust to pH 8.1 with 2 M HCl (about 150 mL of acid required). Dilute to 2 liters and cover. 1 M Barium chloride solution (BaCl2·H2O): 244 g/L (1 liter needed) Buffered barium chloride solution: mix equal volumes (1 L of each) of barium chloride and triethanolamine solutions. (2 L needed) 0.05 N Magnesium sulfate solution (MgSO4·7H2O): 6.2 g/L (1 liter needed) 0.01 M EDTA solution: 3.723 g/L (1 liter needed) Calmagite indicator: 0.1 g reagent/100 mL water (100 mL needed) 2 M NH4OH (70.1 g/L) (1 liter needed) Ammonium extraction and analysis Materials: Spec 20 & cuvettes (40) Flasks (24 125mL) Water bath (40˚C) Volumetric flasks (24 25mL) Pippette (10 mL) Graduated cylinder (50 mL) Plastic weigh boats and spoons Parafilm Whatman 42 filter paper and funnels Reagents: 2 M KCl: 150 g KCl/L (1 liter needed) Stock ammonium standard solution: dissolve 0.3821 g of dry ammonium chloride (NH4Cl) in a 1 L volumetric flask, dilute to volume with deionized water. Store in fridge. Immediately before use, dilute 4 mL of solution to 200 Ml Phenol-nitroprusside reagent: dissolve 7 g of phenol and 34 mg of sodium nitroprusside (disodium penacyanonitrosylferrate, Na2Fe(CN)5NO·2H2O) in 80 mL of deionized water, dilute to 100 mL. Mix well, store in dark colored bottle in fridge. Buffered hypochlorite reagent: dissolve 1.480 g sodium hydroxide (NaOH) in 70 mL of deionized water, add 4.98 g sodium monohydrogen phosphate (Na2HPO4) and 20 mL of sodium hypochlorite (NaOCl) solution (~5% NaOCl). pH should be between 11.4-12.2, add NaOH if need to raise pH. Dilute final volume to 100 mL. EDTA reagent: dissolve 6 g pf ethylenediaminetetraacetic acid disodium salt (EDTA disodium) in 80 mL deionized water, adjust to pH 7, mix well, dilute final volume to 100 mL.