When added to soil, urea fertilizer granules will typically dissolve in the soil water after a short time (minutes to hours), depending on the soil water content and temperature. Even applications at the soil surface will dissolve, typically within a few days, from water near the soil surface and from dews. The dissolved urea then reacts (catalyzed by soil urease enzyme) to form NH4+ and HCO3−, as described by eqn [4]. Concurrently, bicarbonate produced in eqn [4] forms carbon dioxide in eqn [5]. Below pH 8.2, eqn [5] will continue to consume all of the bicarbonate, provided that CO2 diffuses freely from the soil, which would be the case for surface applications of urea. In the case that eqns [4] and [5] go to completion, two H+ are consumed and two NH4+ ions released for each urea hydrolyzed (or one H+ consumed for each NH4+ released).
The ratio of one H+ consumed for each NH4+ released results in an effective cation exchange site being formed and occupied by an NH4+ ion, as shown in eqn [6]. The H+ on the right side of eqn [6] is consumed in eqns [4] or [5]. The term ‘H+ − Soil’ in eqn [6] refers to the pH-dependent charge in the soil, which can be measured as titratable acidity.
Some of the soil cations are not adsorbed on the cation exchange sites. These cations (typically Ca2+, Mg2+, K+, NH4+, and Na+) exist in the soil solution and exchange freely with the same cation species on the exchange sites. Since the equilibria described by eqn [1] only apply to the ions in the soil solution, the cation exchange equilibria, shown in eqn [7] for the case of Ca2+, affects the amount of ammoniacal N in the soil solution, and therefore the proportion that will exist in solution as NH3.
The consumption of H+ causes soil pH near the dissolved urea to rise. The amount of pH rise depends on the soil's H+ buffering capacity, which will be discussed below. If the soil pH rises above 7, a significant amount of NH3 can form, which depends primarily on the soil pH, temperature, and the concentration of NH4+ in the soil solution, as described by the equilibrium in eqns [7] and [1]. Based on equilibrium in eqn [1], any movement of NH3 away from the reaction site will release H+, thereby lowering pH.
Eqns [1], [4], [5], [6], and [7] are the predominant ones below a soil pH of 8.2. The process of nitrification, in which NH4+ is oxidized to NO3−, will add two H+ to the soil for each NH4+ oxidized. This process will not be discussed any further here. Its effect on NH3 volatilization is not very pronounced in the first week following application, since there is often a lag time of a few days in the buildup of a nitrifier population following N application.
Above pH 8.2, eqns [8] and [9] become important, and CO2 formation from reaction [5] stops, as shown graphically in Figure 3. The predominant species is HCO3− (99.2% HCO3 and 0.8% CO32−) around pH 8.2. Any HCO3− from eqn [4] remains in the soil solution or is changed to CO32− as pH is raised further, as shown by eqn [8]. If there are sufficient Ca2+ ions in the soil solution to exceed the solubility product of CaCO3, then solid CaCO3 is formed in the soil, as shown in eqn [9].
Figure 3. The distribution of inorganic C species in a dilute solution as affected by pH. (Reproduced with permission from Koelliker JK and Kissel DE (1988) Chemical equlibria affecting ammonia volatilization. In: Bock BR and Kissel DE (eds) Ammonia Volatilization from Urea Fertilizers, pp. 37–52. Bulletin Y-206. National Fertilizer Development Center. Muscle Shoals, AL: Tennessee Valley Authority.)
The formation of solid CaCO3 removes CO32− from the soil solution, causing a new equilibrium to be established in eqn [8], releasing H+ that slows the rise in pH. This effect is shown in Figure 4. In this study, soils were titrated with either NH4OH or urea + urease enzyme to raise soil pH. Up to pH 8.2, soil pH was raised identically by NH4OH and urea that had completely hydrolyzed. Above pH 8.2, urea was less effective than NH4OH in raising soil pH. The lower effectiveness of urea was due in part to the urea-C not forming CO2 but remaining as HCO3− and CO32− in the soil solution or by precipitating as CaCO3.
When added to soil, urea fertilizer granules will typically dissolve in the soil water after a short time (minutes to hours), depending on the soil water content and temperature. Even applications at the soil surface will dissolve, typically within a few days, from water near the soil surface and from dews. The dissolved urea then reacts (catalyzed by soil urease enzyme) to form NH4+ and HCO3−, as described by eqn [4]. Concurrently, bicarbonate produced in eqn [4] forms carbon dioxide in eqn [5]. Below pH 8.2, eqn [5] will continue to consume all of the bicarbonate, provided that CO2 diffuses freely from the soil, which would be the case for surface applications of urea. In the case that eqns [4] and [5] go to completion, two H+ are consumed and two NH4+ ions released for each urea hydrolyzed (or one H+ consumed for each NH4+ released).
The ratio of one H+ consumed for each NH4+ released results in an effective cation exchange site being formed and occupied by an NH4+ ion, as shown in eqn [6]. The H+ on the right side of eqn [6] is consumed in eqns [4] or [5]. The term ‘H+ − Soil’ in eqn [6] refers to the pH-dependent charge in the soil, which can be measured as titratable acidity.
Some of the soil cations are not adsorbed on the cation exchange sites. These cations (typically Ca2+, Mg2+, K+, NH4+, and Na+) exist in the soil solution and exchange freely with the same cation species on the exchange sites. Since the equilibria described by eqn [1] only apply to the ions in the soil solution, the cation exchange equilibria, shown in eqn [7] for the case of Ca2+, affects the amount of ammoniacal N in the soil solution, and therefore the proportion that will exist in solution as NH3.
The consumption of H+ causes soil pH near the dissolved urea to rise. The amount of pH rise depends on the soil's H+ buffering capacity, which will be discussed below. If the soil pH rises above 7, a significant amount of NH3 can form, which depends primarily on the soil pH, temperature, and the concentration of NH4+ in the soil solution, as described by the equilibrium in eqns [7] and [1]. Based on equilibrium in eqn [1], any movement of NH3 away from the reaction site will release H+, thereby lowering pH.
Eqns [1], [4], [5], [6], and [7] are the predominant ones below a soil pH of 8.2. The process of nitrification, in which NH4+ is oxidized to NO3−, will add two H+ to the soil for each NH4+ oxidized. This process will not be discussed any further here. Its effect on NH3 volatilization is not very pronounced in the first week following application, since there is often a lag time of a few days in the buildup of a nitrifier population following N application.
Above pH 8.2, eqns [8] and [9] become important, and CO2 formation from reaction [5] stops, as shown graphically in Figure 3. The predominant species is HCO3− (99.2% HCO3 and 0.8% CO32−) around pH 8.2. Any HCO3− from eqn [4] remains in the soil solution or is changed to CO32− as pH is raised further, as shown by eqn [8]. If there are sufficient Ca2+ ions in the soil solution to exceed the solubility product of CaCO3, then solid CaCO3 is formed in the soil, as shown in eqn [9].
Figure 3. The distribution of inorganic C species in a dilute solution as affected by pH. (Reproduced with permission from Koelliker JK and Kissel DE (1988) Chemical equlibria affecting ammonia volatilization. In: Bock BR and Kissel DE (eds) Ammonia Volatilization from Urea Fertilizers, pp. 37–52. Bulletin Y-206. National Fertilizer Development Center. Muscle Shoals, AL: Tennessee Valley Authority.)
The formation of solid CaCO3 removes CO32− from the soil solution, causing a new equilibrium to be established in eqn [8], releasing H+ that slows the rise in pH. This effect is shown in Figure 4. In this study, soils were titrated with either NH4OH or urea + urease enzyme to raise soil pH. Up to pH 8.2, soil pH was raised identically by NH4OH and urea that had completely hydrolyzed. Above pH 8.2, urea was less effective than NH4OH in raising soil pH. The lower effectiveness of urea was due in part to the urea-C not forming CO2 but remaining as HCO3− and CO32− in the soil solution or by precipitating as CaCO3.
Figure 4. The change in soil pH of Kahola silt loam as affected by the addition of NH4OH (squares) or urea (circles) allowed to hydrolyze completely. (Reproduced with permission from Kissel DE, Cabrera ML, and Ferguson RB (1988) Reactions of ammonia and urea hydrolysis reaction products with soil. Soil Science Society of America Journal 52: 1793–1796.)
Figure 4. The change in soil pH of Kahola silt loam as affected by the addition of NH4OH (squares) or urea (circles) allowed to hydrolyze completely. (Reproduced with permission from Kissel DE, Cabrera ML, and Ferguson RB (1988) Reactions of ammonia and urea hydrolysis reaction products with soil. Soil Science Society of America Journal 52: 1793–1796.)
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