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Fe (II)'nin oksidasyonuna organik asitlerin etkisi

Başlık çevirisi mevcut değil.

  1. Tez No: 55933
  2. Yazar: GÜLSÜN ÖZGÜL
  3. Danışmanlar: DOÇ.DR. CUMALİ KINACI
  4. Tez Türü: Yüksek Lisans
  5. Konular: Çevre Mühendisliği, Environmental Engineering
  6. Anahtar Kelimeler: Belirtilmemiş.
  7. Yıl: 1996
  8. Dil: Türkçe
  9. Üniversite: İstanbul Teknik Üniversitesi
  10. Enstitü: Fen Bilimleri Enstitüsü
  11. Ana Bilim Dalı: Belirtilmemiş.
  12. Bilim Dalı: Belirtilmemiş.
  13. Sayfa Sayısı: 67

Özet

ÖZPT Yeraltı ve yüzeysel sularda yüksek konsantrasyonlarda demir bulunması, evsel ve endüstriyel amaçlı kullanımları olumsuz yönde etkilediğinden bu tür sulardan demirin giderilmesi gerekmektedir. Bu amaçla bu çalışmada, Fe(II)'nin oksijenle oksidasyonunda reaksiyon ürünü olan Fe(III)'ün katalitik etkisi ve bu katalitik etkinin kinetiği incelenmiştir. Tannik asit, asetik asit, glikozun Fe(II)'nin oksidasyon hızına ve ardından Fe(III) konsantrasyonlarının oksidasyon hızına katalitik etkileri araştırılmıştır. Birinci bölümde, yapılan çalışmanın amaç ve kapsamı açıklanmış ve önemi vurgulanmıştır. İkinci bölümde, demir giderilmesinin önemi, sulu ortamlarda demirin kimyası detaylı bir şekilde verilmiştir. Üçüncü bölümde, havalanma ile Fe(II) oksidasyonunun kinetiği ve Fe(III)'ün katalitik etkisi incelenmiştir. Bu konuda mevcut bilgi düzeyi bilimsel çerçevede sunulmuştur. Dördüncü bölümde doğal sularda bulunan organik ve inorganik maddelerin Fe(II)'nin oksidasyon kinetiğine olan etkileri ayrıntılı bir. şekilde anlatılmıştır. Beşinci bölümde, deneysel çalışmada kullanılan yöntem ve düzeneklerle ilgili ayrıntılı bilgi verilmiştir. Fe(II) konsantrasyonu 25 mg/1 alınarak ve her aşamada Fe(III) konsantrasyonu artırılarak Fe(II)'nm oksijen ile oksidasyonunun kinetiği incelenmiştir. Fe(III) konsantrasyonunun 0-1500 mg/1 arasındaki konsantrasyon değerlerinde Fe(III)'ün katalitik etkisi kontrollü olarak doldur-boşalt tipi 2 litrelik reaktörlerde incelenmiştir. Bu deneylerde, oksijenin kısmi basıncı, pH ve sıcaklık sabit tutularak ve sadece Fe(III) konsantrasyonları değiştirilerek deneyler tekrar edilmiştir. İkinci safhada, tannik asit, asetik asit, glikoz ilave edilerek Fe(II)'nin oksijen ile oksidasyonunun kinetiği araştırılmıştır. Ortamda Fe(III) konsantrasyonu 0-1500 mg/1 arasındaki değerlerde tutularak, Fe(in)'ün katalitik etkisi kontrollü olarak doldur- boşalt tipi 2 litrelik reaktörlerde incelenmiştir. Altıncı bölümde, deney sonuçlan değerlendirilmiştir. pH= 6.7, sıcaklığın 25°C, alkalinitenin 2.1 0"2 eq/l ve yüksek Fe(D) başlangıç konsantrasyonu değerlerinde (25 mg/1) otokatalitik etki nedeniyle oksidasyon hızının önemli ölçüde arttığı tesbit edilmiştir. Tannik asidin oksidayonu yavaşlattığı, asetik asit ve glikozun ise oksidasyon hızım etkilemedikleri gözlemlenmiştir.

Özet (Çeviri)

SUMMARY THE EFFECT OF ORGANIC ACIDS ON FERROUS OXIDATION In oxygen-free aquatic environments, such as groundwaters and hypolimnetic waters of eutrophic lakes, iron exists predominantly in the ferrous state, Fe(II). For those groundwaters used for domestic and industrial purposes, removal of iron is desirable because it can form rust (iron oxide) deposits, causing staining of plumbing fixtures, laundered goods, and manufactured products, as well as imparting a metallic taste to the water. Conventional water treatment for the removal of iron consists of aeration of the raw water, providing for the oxidation of ferrous iron Fe+2+l/402 + lT = Fe+3+ 1/2H20 (1) The resultant ferric iron hydrolyzes to form the highly insoluble ferric hydroxide Fe+3 +3H20 = Fe(OH)3(S) + 3H+ (2) which is subsequently removed by sedimentation and filtration. The net reaction is Fe+2 + l/402 + 5/2 H20 = Fe(OH)3(S) + 2ff (3) In natural surface waters, iron serves as a nutrient for algae, higher plants, and other forms of aquatic life. During the spring and fall over turns in eutrophic lakes, the anoxic hypolimnetic water containing ferrous iron is exposed to oxygen and Fe(II) is normally oxidized to Fe(III), as in Equation 3. The availability of iron as a nutrient is dependent upon the rate of Fe(II) oxidation and the stability of the resultant Fe(m); the concentration of total soluble ferric iron in equilibrium with ferric hydroxide at pH 8 is approximately 0.2 ug/1. [The solubility product of ferric hydroxide, Kso, is approximatly 10“36.1] -d[Fe(II)]/ dt = k[Fe(H)] (p02) [Off ] (4) where k =8.0 1013 mole”2 atm“1 min”1 at 20.5° C. (At 25° C, fc= 1.36 1014 mole“2 atm”1 min“1.). The oxidation of ferrous iron proceeds relatively rapidly at neutral pH values; the half- time of the reaction is 4 minutes at pH 7 and a partial pressure of oxygen of 0.21 atm. Thus, under oxygen -rich conditions, Fe(II) would be rapidly oxidized, making iron removal a relatively simple process and hmiting iron availability as a nutrient for aquatic growth in natural surface waters. Occurrence of iron levels in water supplies exceeding the 0.3 mg/1 limit is common. The oxidation of the ferrous iron is affected by several factors such as Fe(II) and oxygen concentration, pH, temperature, organic matter and other ions present in the solution. Ferrous iron oxidation is also affected and is accelerated in the presence of the ferric hydroxide. The recent studies have demonstrated the catalytic effect of the XIferric hydroxide but the effect becomes noticeable at Fe(III) concentrations exceeding 5-10 mg/1 (Sankaya,1980; Tamura et.al.,1976; Robinson et.al.,1981). It has been reported that the oxidation rate increases linearly with Fe(III) concentrations up to 100 mg/1 (Tamura et.al.,1976). In the presence of Fe(III), the oxidation of ferrous iron by the atmospheric oxygen occurs by two simultaneous processes. One of these is the homogeneous reaction taking place in solution, and the other one is the heterogeneous reaction occurring on the surface of the ferric hydroxide precipitates. Thus: Fe(II) + 02 - ^-> Fe(ffl) + O^ (homogeneous) (5) Fe(II) + 02 - ^-> Fe(HI) + 0~2 (heterogeneous) (6) The reaction rate under the constant pH and O2 concentration is given as (Sankaya,1980; Tamura et.al.,1976; Sung etal., 1980): -d[Fe(H)]/dt = (k + k'[Fe(HI)]) [Fe(fl)] (7) in which the first term indicates the homogeneous and the second one heterogeneous reaction rates. The explicit form of the constants are(Tamura et.al.,1976; Sung etal, 1980): k = k0[OH-]2[O2] (8) k' = ks,0[O2]K/[H+] (9) k' = ksK/[H+] (10) ks=ks>0[O2] (11) where ko and K,a are the real rate constants, k, is the surface rate constant and K is the equilibrium constant for the adsorption of Fe(ü) on Fe(III) hydroxide. The numerical values of the constants are (Tamura et.al.,1976); 14 -3 -1 ko = 2.3x10 M s k =73 M-1 s”1 S,0 K=10^-6Mmg“1 One of the aims of this study İs to determine the catalytic effect of Fe(III) on the ferrous iron oxidation by atmospheric oxygen at Fe(lTI) concentrations beyond 100 mg/1 (Tamura et.al.,1976). In the experimental study in which a large amount of Fe(lII) hydroxide was added at the beginning of the oxygenation of a small amount of Fe(II) the concentration of XllFe(lII) hydroxide is almost constant throughout the experiment. So equation (7) can be written as -d [Fe (II) ] /dt = kc [Fe(II)] (12) where k^ is a constant for a constant pH and O2 concentration. However, certain natural organic compounds, usually referred to as humic material, exert a stabilizing effect on iron in aquatic systems, i.e., High concentrations of iron are often associated with organic matter of natural origin in many natural aquatic systems. Previous investigators (Theis and Singer, 1973; Komolrit, 1962: Ghosh et al. 1967) have indicated that the oxidation of ferrous iron is severely retarded in many natural waters which contain humic substances, including groundwaters, surface waters, and municipal wastewaters. It has been demonstrated (Theis and Singer ) that many model organic compounds which possess the structural features of humic substances are capable of significantly altering the rate of oxidation of ferrous iron from that reported by Stumm and Lee (1961) for simple aqueous media. Gallic acid, tannic acid and pyrogallol completely inhibited Fe(II) oxidation! Glutamic and tartaric acids behaved analogously to glutamine, while vanillic acid, vanillin, phenol, resorcinol, syrincig acid, and histidine had no effect on the oxidation rate. Morgan and Stumm (1964) suggested that the inhibition of Fe(II) oxidation by organic species involved the cyclic oxidation of Fe(II) followed by the reduction of Fe(HI) by the organic substances. Their scheme is given in equations la through lc. Fe(n) + 02 -> Fe(HI) la Fe(III) + Organic ->. Fe(II) + Oxidized Organic lb Fe(II) + 02 -> Fe(IH) lc This report includes six chapters as following: In the first chapter, the importance and the general objectives of the study are defined in detail. In the second chapter, the chemistry of aqueous iron and character of iron precipitates are given. In the third chapter, the kinetics of ferrous iron oxidation and its oxygenation products in aqueous systems are explained. In the fourth chapter, the inhibition of Fe(II) oxidation by organic species involved the cyclic oxidation of Fe(II) followed by the reduction of Fe(III) by the organic substances are explained. In the fifth chapter, the experimental techniques for this subject are explained. In the sixth chapter, the results of the experiment are discussed. XlllThe study was carried out in two stages: In the first stage, the kinetics of ferrous iron oxidation by atmospheric oxygen has been studied in batch systems of 2 1 volume. Fe(III) concentrations were varied within the range of 0-1500 mg/1, keeping partial pressure, pH, and temperature constant. A modified 1.10 phenanthroline method which enables the analysis of ferrous iron in the presence of the high concentrations of Fe(IH) has been applied in the study. To find out the effect of tannic acid, acetic acid and glicose on the ferrous iron oxidation these substances were added in batch systems of 21 volume. All of the other conditions were kept same as the first stage. The experimental results obtained high (25 mg/1) Fe(II) concentrations are given. No Fe(III) was initially added in these experiments. The linear relationship between the Fe(II) concentration. and time on a semi logarithmic plot shows a first order kinetics. This is in agreement with the results given in the literature. Reproducibility of the results is good in the homogeneous systems. When initial Fe(II) concentration was increased keeping all other experimental conditions the same, the reaction is accelerated due to the Fe(III) formed as a result of oxidation. Average homogeneous reaction rate constant, k for initial Fe(II) concentration of 25 mg/1 is found as 0.10907mm”1. The results of the experiments with initial FeQT) concentration of 25mg/l and with varying initial Fe(III) concentrations are given. Catalytic rate constant, k^ is calculated from the slopes of the lines on semi logarithmic plots of Fe(II) versus time. Time needed for the completion of the reaction is about 46 minutes when no Fe(III) is present in the reactor. The reaction time is reduced to about 26 minutes when 200 mg/1 of Fe(III) is added initially into the reactor. It has been found that k^t increases linearly with increasing Fe(lTI) concentration up to about 500 mg/1 and the rate of the increase decreases beyond this value, lc^ reaches its maximum value at an Fe(m) concentration of about 500 mg/1. This means that there is no additional catalytic effect of Fe(III) on the ferrous iron oxidation at Fe(IÜ) concentrations beyond 500 mg/1. The kinetics of the catalytic reaction can be represented by linear expression in the concentration range of Fe(III) 0-75 mg/1 and the second order polynomial in the range 0-1500mg/l. Applying the curve fitting techniques to the data, the following equations are obtained between the k^t and Fe(III) concentration. In the linear range, k^t is given as. ^ = k + k [Fe(III)]0 (13) kcat= 0.00 12[Fe(in)]+0. 1076 [Fe(n)] = [Fe(II)]o e*0-T121T*01076^ XIVwhere [Fe(III)]0 indicates the initial Fe(m) concentration. Oxygenation rate in terms of kcat and Fe(II) is. d\Fe(II)] r, dt -MFg(//)l (14> Assuming Fe(III) is constant, integration of equation 14 yields [Fe(ir)] = [Fe(II)]oe-k-' (15) in which [Fe(II)]0 is the initial Fe(II) concentration. Substituting the value of k^ obtained from equation 13 into equation 15 yields: In the second order polynomial range, kcat is found as: [Fe(II)] = [Fe(II)]o e50006^^01687*1 where [Fe(III)] and [Fe(II)]o indicate the concentration of Fe(III) and initial concentration of Fe(IT), respectively. The results of the second stage are tannic acid completely inhibited the oxygenation of ferrous iron over the 105 minute observation period and acetic acid, glicose had no observable effect on the oxidation reaction. XV

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