Article pubs.acs.org/jced
Measurement and Correlation of the Solubility of Enrofloxacin in Different Solvents from (303.15 to 321.05) K Meng-jiao Liu, Hua-lin Fu,* Dai-ping Yin, Yan-li Zhang, Chao-cheng Lu, Hang Cao, and Jian-yu Zhou Department of Pharmacy, College of Veterinary Medicine, Sichuan Agricultural University, Yaan, Sichuan 625014, People’s Republic of China ABSTRACT: The solubility of enrofloxacin in ethanol, water, 2-propanol, acetonitrile, and acetic ether was measured by a static analytical method from (303.15 to 321.05) K. The concentrations of enrofloxacin in saturated solvent were analyzed by UV spectrometry. The experimental solubility data for pure solvents decreases in the order acetonitrile > acetic ether >2-propanol > ethanol > water. The experimental solubility data was correlated by the modified Apelblat equation, the ideal solution equation, and the Clarke−Glew equation. The results show that the mole fraction solubility of enrofloxacin in pure solvents increased with increasing temperature; the Clarke−Glew two parameters equation (Clarke−Glew 2 equation) can provide a more accurate mathematical model than the modified Apelblat equation, the ideal solution equation, and the Clarke−Glew three parameters equation (Clarke−Glew 3 equation). In addition, the discussion of Gibbs energy, enthalpy, and entropy for enrofloxacin in different solvents indicated that the dissolution process of enrofloxacin was endothermic and entropy-driven.
1. INTRODUCTION Enrofloxacin (C19H22FN3O3, CAS 93106-60-6, Figure 1), the first fluoroquinolone approved for veterinary use, has a
For this reason, the solubility of enrofloxacin in various solvents is needed. Crystallization process is a multifarious mass and energy transfer process between different phase, its main driving force comes from the crystalline properties of multiphase systems in nonequilibrium thermodynamics. Thus, the thermodynamics plays a crucial role for the development of the crystallization process. In the present work, a static method was employed to measure the solubility of enrofloxacin in several different solvents at the temperature ranging from (303.15 to 321.05) K. The experimental solubility data was correlated by the modified Apelblat equation, the ideal solution equation and the Clarke− Glew equation. The thermodynamic parameters of the dissolution process for this system, including the Gibbs energy, enthalpy, and entropy obtained by van’t Hoff analysis, were discussed as well.
Figure 1. Structural formula of enrofloxacin.
chemical name 1-cyclopropyl-7-(4-ethyl-1-piperazinyl)-6-fluoro-1,4-dihydro-4-oxo-3-quinolinecarboxylic-acid. Plenty of clinical studies have been conducted about enrofloxacin. Its good potency against bacteria1−4 and pharmacokinetic properties5,6 suggest that it could be an excellent antimicrobial agent for the treatment of bacterial infections in animals. Drug crystallization is a decisive factor to determine the quality and yield for a formulation product. The crystallization process requires the solubility data to determine high yield and low energy, which is essential for pharmacy, fine chemical industry, and some solid products. Solid−liquid equilibrium data in a crystalline system are not only selected as the basis of the solvent system during crystallization purification and crystallization methods but also the key factor of maximum productivity and theoretical yield for a crystallization process. © 2014 American Chemical Society
2. EXPERIMENTAL SECTION 2.1. Materials. The materials table is depicted in Table 1. 2.2. Apparatus and Procedure. The equilibrium solubility was measured similar to our previous work.7,8 The reliability test of this prodedure, which have been validated in our previous work, is detailed in the literature.8 An excess of enrofloxacin was taken into a 150 mL erlenmeyer flask, which was placed in a thermostatic shaker (HZS-H water bath oscillator, China). The temperature was corrected by a mercury Received: March 6, 2014 Accepted: May 5, 2014 Published: May 9, 2014 2070
dx.doi.org/10.1021/je5002158 | J. Chem. Eng. Data 2014, 59, 2070−2074
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Table 1. Characteristics of Chemicals Used in This Study chemical name ethanol acetonitrile enrofloxacin 2-propanol aceticether water a
the mass fraction purity
source Chengdu Kelong Chemical poReagent Co., Ltd. Tianjin Kermel Chemical poReagent Co., Ltd. Zhejiang Guobang Pharmaceutical Co., Ltd. Tianjin Shentai Chemical poReagent Co., Ltd. Tianjin Kermel Chemical poReagent Co., Ltd. double-distilled water prepared by the laboratory
x1p =
analysis method
m1/M1 + (m − m1)/M p
(1)
where m represents the mass of 1 mL mixture, m1 is the mass of solute in 1 mL mixture, and xp1 is the mole fraction solubility of the solute. M1 and Mp represent the molar mass of solute, pure solvent, respectively. The relative deviation (RD) between the calculated and the experimental values are obtained from eq 2.
≥99.7% ≥99.9% ≥99%
m1/M1
UVa
≥99.7% ≥99.5%
RD =
x1p − x1cal x1p
(2)
where xcal denotes the back-calculated solubility data by 1 different equations. The root-mean-square deviation (RMSD) defined in eq 3 was used to evaluate the agreement between the experimental data and the model predictions.
Ultraviolet spectrophotometry.
thermometer (uncertainty of ± 0.1 K). Until solid−liquid equilibrium was reached, the samples were filtered through a microporous membrane (0.22 μm) before analysis. A 5 mL mixture was removed by a precise preheated pipet (uncertainty of 0.05 mL) to weigh the mass. A 2 mL sample transferred to volumetric flask to a suitable dilution with 70% ethanol aqueous solution. According to the standard curve (Figure 2), the
⎡ N (x p − x cal)2 ⎤1/2 1 ⎥ RMSD = ⎢∑ 1 ⎢⎣ i = 1 N − 1 ⎥⎦
(3)
where N is the number of solubility data. To evaluate the accuracy and predictability of the four models, the mean value of RMSDs was denoted as overall RMSD (ORMSD) and is given by eq 4. 1 N−1
ORMSD =
N
∑ RMSD (4)
1
where N is the number of RMSD.
3. RESULTS AND DISCUSSION 3.1. Correlation Models of Enrofloxacin. The modified Apelblat equation which could be used as a classical correlation model was described below. B ln x1p = A + + C ln(T /K) (5) T /K
Figure 2. Standard curve for the dependence of enrofloxacin concentration on UV absorbance.
where xp1 is the mole fraction solubility of the solute in different solvents and T is the experimental temperature in Kelvin. A, B, and C are the parameters of the model. The temperature dependence of the solubility of enrofloxacin in the selected solvents can be correlated by the following ideal solution equation:9
content of enrofloxacin was measured by UV spectrophotometer detector (WFZ UV-2000, Unico, China) at 271 nm wavelength. All of the experiments were repeated three times and the mean values were used to calculate the mole fraction of the solubility. For pure solvents, the mole fraction solubility xp1 of the solute was obtained by eq 1.
ln x1p = A +
B T /K
(6)
xp1
where is the mole fraction solubility of enrofloxacin and T is the corresponding temperature in Kelvin. A and B are the parameters of the equation.
Table 2. Parameters of the Clarke−Glew 3 equation for enrofloxacin in the selected solvents from (303.15 to 321.05) K.a parameters
a
pure solvent
A0 ± SE
A1 ± SE
A2 ± SE
RMSD
ORMSD
ethanol water acetic ether acetonitrile 2-propanol
304.65 ± 50.80 28.38 ± 77.10 163.43 ± 50.78 −267.96 ± 50.78 415.29 ± 79.95
−297.19 ± 50.84 −16.68 ± 77.17 −156.47 ± 58.48 275.20 ± 50.83 −408.01 ± 80.02
−300.84 ± 48.58 −37.68 ± 73.73 −161.35 ± 55.88 245.90 ± 48.56 −407.49 ± 76.46
1.5540 2.3425 1.7807 1.5589 2.4325
1.9337
Standard uncertainties u are u(T) = 0.05 K, u(A0, A1, A2, SE) = 0.01, u(RMSD) = 0.0001, u(ORMSD) = 0.0001. 2071
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containing active groups such as >N−, −COOH, −F, enrofloxacin perhaps involve various forces such as electrostatic force, van der Waals bond, hydrophobic interaction, hydrogen bond, and stereoscopic effect in the dissolving process.11 Acetonitrile is weakly alkaline, but enrofloxacin molecule contained acidic groups −COOH. So the two molecules could produce a salt linkage between the molecules; the carbon oxygen double bond of acetic ether could form hydrogen bond with enrofloxacin; in the remaining three solvents, van der Waals force is generated, which increased with increasing molecular weight of the solvents. In fact, the essence of the modified Apelblat equation and the Clarke−Glew 3 equation, the ideal solution equation and the Clarke−Glew 2 equation are exactly alike, so resulting in the same results. By a series calculation, it can be seen that most of the RMSD from eqs 5 to 8 do not exceed 6%, and the average of RMSD (ORMSD) are 1.9337, 1.9337, 4.3543, and 4.3543, respectively (In our paper, we only list the results of eqs 7 and 8 in Tables 2 and 3.). Obviously, the ORMSD from eqs 5 and 7 are lower than eqs 6 and 8, but nonlinear regression of eqs 5 to 7 indicate that each of the fit parameters have a very large uncertainty, in other words, the system of eqs 5 to 7 are over parametrized, and the result of eq 8 is quite perfect. Otherwise, the ORMSD of Clarke−Glew 2 is less than 5, which is in the range of allowable error. Take all these into consideration, we can draw that the Clarke−Glew 2 equation can provide a more accurate mathematical model than the modified Apelblat equation, the ideal solution equation and Clarke−Glew 3 equation. 3.3. Thermodynamic Properties of Solutions. In our work, the temperature was in a narrow range (less than 20 K), the heat capacity change, which could be considered as zero. The solution system could regard as an ideal solution. The relationship between solubility and Gibbs free energy is described as follows.
The Clarke−Glew three parameters equation (Clarke−Glew three parameters equation)10 which was linked with the relationship between temperature and solubility in different solvents was shown as follows: −ln x1p = A 0 + A1(Tref /T ) + A 2 ln(T /Tref )
(7)
xp1
where is the mole fraction solubility of the solute in different solvents, T and Tref are the experimental temperature and reference temperature of 25 °C in Kelvin, respectively. A0, A1, A2 are the parameters (uncertainty of ± 0.01) and listed in Table 2, together with its standard error (SE, uncertainty of ± 0.01), RMSD and ORMSD (uncertainty of ± 0.0001). Within a temperature range of less than 20 K the change of heat capacity can probably be neglected. For the change in heat capacity is supposed zero, only parameters A0 and A1 are thermodynamically meaningful. The temperature dependence of the solubility of enrofloxacin in the selected solvents could be calculated as the following Clarke−Glew two parameters equation (Clarke−Glew 2 equation): −ln x1p = A 0 + A1(Tref /T )
(8)
xp1
where is the mole fraction solubility of enrofloxacin, T is the corresponding temperature in Kelvin, Tref is the selected temperature of 298.15 K, and A0 and A1 are the parameters (uncertainty of ±0.01) of the equation. The parameters A0 and A1 are listed in Table 3, together with its standard error (SE, uncertainty of ± 0.01), RMSD and ORMSD (uncertainty of ± 0.0001). Table 3. Parameters of the Clarke−Glew 2 Equation for Enrofloxacin in the Selected Solvents from (303.15 to 321.05) Ka parameters pure solvent
A0 ± SE
ethanol water acetic ether acetonitrile 2-propanol
−9.93 ± 1.17 −11.02 ± 0.57 −5.28 ± 0.73 −10.83 ± 0.98 −10.80 ± 1.62
A1 ± SE
RMSD
ORMSD
± ± ± ± ±
5.0806 2.4160 3.1654 4.2087 6.9010
4.3543
17.65 22.75 12.38 17.85 18.44
1.23 0.59 0.76 1.03 1.69
ΔGs /RT = −ln aenrfloxacin = −ln x1p − ln f1p
(9)
where aenrfloxacin is the activity of enrofloxacin in saturated aqueous solutions and fp1 is the activity coefficient. For ideal solution, fp1 = 1. And we could use the following ideal solution equation to calculate the thermodymic parameters. According to van’t Hoff analysis,12,13 the apparent enthalpy change of solution could be related to the temperature and the mole fraction solubility as the following equation:
a
Standard uncertainties u are u(T) = 0.05 K, u(A0, A1, SE) = 0.01, u(RMSD) = 0.0001, u(ORMSD) = 0.0001.
3.2. Experimental and Ideal Solubility. The solubilities of enrofloxacin in different solvents are presented in Table 4 in addition to the ideal solubilities in mole fraction of the solute (xcal 1 ) calculated from eqs 5 to 8. And the solubility versus temperature was graphically plotted in Figure 3. From Table 4 and Figure 3, the conclusion can be draw that the solubility value of enrofloxacin in different solvents increase with increasing temperature, which denotes the processes of enrofloxacin dissolving in the selected solvents are endothermic. However, the extent of solubility increase in pure solvents is different. It can be observed from Figure 3 that the solubility follows the order: acetonitrile > acetic ether > 2-propanol > ethanol > water. Compared to the polarity of the solvents, water > acetonitrile > ethanol > 2-propanol > acetic ether, we can determine that the polarity of the solvents is not just the crucial factor of the solubility of enrofloxacin in the selected solvents. This phenomenon may be resulted from the solvent molecular structures and space conformation between solute and solvent. Owing to the solute enrofloxacin molecules
⎛ ∂ln x1p ⎞ ΔHs = −R ⎜ ⎟ ⎝ ∂(1/T ) ⎠P
(10)
xp1
where is the mole fraction solubility, R represents the universal gas constant (8.314 J·k−1·mol−1) and T is the experimental temperature in Kelvin. The ΔHs can be obtained from the slope of the solubility curve in a so-called van’t Hoff plot where ln xp1 is plotted versus 1/T. The values of ΔHs would be valid for the reference temperature (Tref = 298.15 K) for the selected solvents. Equation 10 can also be written as eq 11. ⎛ ⎞ ∂ln x1p ΔHs = −R ⎜ ⎟ ⎝ ∂(1/T − 1/Tref ) ⎠P
(11)
The ΔHs (uncertainty of ± 0.01) can be obtained from the slope of the solubility curve where ln xp1 is plotted versus 1/T − 2072
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Table 4. Experimental Mole Fraction Solubility of Enrofloxacin (xp1) in Pure Solvents at Standard Atmospheric Pressure (p = 0.1 a MPa) in the Temperature Ranging from (303.15 to 321.05) K, and Calculated Solubility (xcal 1 ) Obtained from eqs 5 to 8. T/K
103 xp1b
303.15 306.35 309.15 312.15 314.95 318.55 321.05
0.6314 0.7161 0.7883 0.8974 1.1016 1.4018 1.6393
303.15 306.35 309.15 312.15 314.95 318.55 321.05
1.0530 1.1623 1.2316 1.3641 1.6076 1.8556 2.0384
303.15 306.35 309.15 312.15 314.95 318.55 321.05
1.1344 1.4124 1.7669 2.0786 2.4140 2.7013 3.0573
c 103 xcal.1 1
d 103 xcal.2 1
T/K
103 xp1b
0.6339 0.7047 0.7926 0.9215 1.0846 1.3784 1.6599
0.5905 0.7079 0.8271 0.9741 1.1316 1.3668 1.5545
303.15 306.35 309.15 312.15 314.95 318.55 321.05
0.0115 0.0153 0.0179 0.0217 0.0272 0.0335 0.0418
1.0519 1.1472 1.2536 1.3963 1.5620 1.8326 2.0680
1.0126 1.1500 1.2826 1.4385 1.598 1.8243 1.9965
303.15 306.35 309.15 312.15 314.95 318.55 321.05
0.7321 0.7706 0.869 1.0032 1.2455 1.6509 1.9005
1.1279 1.4412 1.7401 2.0751 2.3892 2.7742 3.0155
1.1953 1.4359 1.6806 1.9831 2.3077 2.7934 3.1816
Ethanol
c 103 xcal.1 1
d 103 xcal.2 1
0.0118 0.0147 0.0179 0.0221 0.0268 0.0345 0.0410
0.0117 0.0147 0.0180 0.0223 0.0270 0.0344 0.0406
0.7194 0.7849 0.8771 1.0225 1.2171 1.5884 1.9631
0.6534 0.7897 0.9291 1.1023 1.2892 1.5704 1.7963
Water
Acetic ether
2-Propanol
Acetonitrile
a
Standard uncertainties u are u(T) = 0.05 K, ur(x1) = 5% bxp1 denotes the experimental solubility data. cxcal.1 denotes back-calculated solubility data 1 denotes back-calculated solubility data by variant 2 of the by variant 1 of the modified Apelblat equation and the Clarke−Glew 3 equation. dxcal.2 1 ideal solution equation and the Clarke−Glew 2 equation.
Figure 4. Mole fraction solubility (ln xp1) of enrofloxacin versus temperature (1/T − 1/Tref) in different solvents: yellow square, water; blue circle, ethanol; red triangle, 2-propanol; ×, acetic ether; black diamond, acetonitrile.
Figure 3. Mole fraction solubility (xp1) of enrofloxacin at different temperatures in pure solvents: yellow square, water; blue circle, ethanol; red triangle, 2-propanol; ×, acetic ether; black diamond, acetonitrile.
ΔSs =
1/Tref. Tref is the reference temperature of 298.15 K. Figure 4 shows the linear ln xp1 − (1/T − 1/Tref) curves of enrofloxacin in different solvents. The Gibbs energy change (ΔGs, uncertainty of ±0.01) for the dissolution process can be calculated as eq 13 according to Krug et al.14 ΔGs = −RTref × intercept = ΔHs − Tref ΔSs
ΔHs − ΔGs Tref
(13)
Table 5 shows the standard thermodynamic parameters for the dissolution process in the selected solvents. From Table 5, it can be draw that the standard molar enthalpy and entropy for all cases is positive, which is indicated that the dissolution process is endothermic and entropy-driven in the experimental temperature range. The high values of the enthalpy suggest that more energy is needed for overcoming the cohesive force between the solute and the solvent, which also indicating a strong temperature dependence of solubility.15 Otherwise, Gibbs energy could determine the degree of difficulty in the
(12)
In which, the intercept can be obtained by plotting ln xp1 versus 1/T − 1/Tref in Figure 4. Thus, the standard molar entropy of solution (ΔSs, uncertainty of ± 0.01) can be calculated as eq 12 from eq 13. 2073
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Table 5. Dissolution Enthalpy and Entropy of Enrofloxacin in the Selected Solvents from (303.15 to 321.05) Ka
a
pure solvent
ΔHs/KJ·mol−1
ΔGs/KJ·mol−1
ΔSs/J·K−1·mol−1
% ζTS
% ζH
ethanol water acetic ether acetonitrile 2-propanol
43.76 56.39 30.69 44.26 45.71
19.15 29.09 17.60 17.41 18.93
82.53 91.59 43.90 90.04 89.82
35.99 32.62 29.90 37.76 36.94
64.01 67.38 70.10 62.24 63.06
Standard uncertainties u are u(T) = 0.05K, u(ΔHs) = 0.01, u(ΔGs) = 0.01, u(ΔSs) = 0.01, u(%ζTS) = 0.01, u(%ζH) = 0.01. (4) Watts, J. L.; Salmon, S. A.; Sanchez, M. S.; Yancey, R. J. In vitro activity of premafloxacin, a new extended-spectrum fluoroquinolone, against pathogens of veterinary importance. Antimicrob. Agents Ch. 1997, 41 (5), 1190−2. (5) Kaartinen, L.; Salonen, M.; Alli, L.; Pyorala, S. Pharmacokinetics of enrofloxacin after single intravenous, intramuscular and subcutaneous injections in lactating cows. J. Vet. Pharmacol. Ther. 1995, 18 (5), 357−62. (6) Mengozzi, G.; Intorre, L.; Bertini, S.; Soldani, G. Pharmacokinetics of enrofloxacin and its metabolite ciprofloxacin after intravenous and intramuscular administrations in sheep. Am. J. Vet. Res. 1996, 57 (7), 1040−3. (7) Zhang, C.; Zhao, F.; Wang, Y. Thermodynamics of the solubility of ciprofloxacin in methanol, ethanol, 1-propanol, acetone, and chloroform from 293.15 to 333.15 K. J. Mol. Liq. 2010, 156 (2−3), 191−193. (8) Zhou, J.; Fu, H.; Cao, H.; Lu, C.; Jin, C.; Zhou, T.; Liu, M.; Zhang, Y. Measurement and correlation of the solubility of florfenicol in binary 1, 2-propanediol+ water mixtures from 293.15 to 316.25 K. Fluid Phase Equilib. 2013, 360, 118−123. (9) Ren, Y.; Duan, X. Equilibrium Solubilities of 2-Cyanoguanidine in Water+ (Ethane-1, 2-diol or N, N-Dimethylformamide) Mixtures at Different Temperatures: An Experimental and Correlational Study. J. Chem. Eng. Data 2013, 58 (11), 3282−3288. (10) Lorimer, J. W.; Cohen-Adad, R. Thermodynamics of Solubility; In Experimental Determination of Solubilities; Hefter, G. T., Tomkins, R. P. T., Eds.; John Wiley & Sons: Chichester, U.K., 2003. (11) Zhang, C.; Zhao, F.; Wang, Y. Thermodynamics of the solubility of sulfamethazine in methanol, ethanol, 1-propanol, acetone, and chloroform from 293.15 to 333.15 K. J. Mol. Liq. 2011, 159 (2), 170− 172. (12) Chaires, J. B. Possible origin of differences between van’t Hoff and calorimetric enthalpy estimates. Biophys. Chem. 1997, 64 (1−3), 15−23. (13) Atkins, P.; De Paula, J. Physical Chemistry, 8th ed.; W.H. Freeman and Company: New York, 2006. (14) Krug, R. R.; Hunter, W. G.; Grieger, R. A. Enthalpy-entropy compensation. 1. Some fundamental statistical problems associated with the analysis of van’t Hoff and Arrhenius data. J. Phys. Chem. 1976, 80 (21), 2335−2341. (15) Schröder, B.; Santos, L. M. N. B.; Marrucho, I. M.; Coutinho, J. A. P. Prediction of aqueous solubilities of solid carboxylic acids with COSMO-RS. Fluid Phase Equilib. 2010, 289 (2), 140−147. (16) Perlovich, G. L.; Kurkov, S. V.; Bauer-Brandl, A. Thermodynamics of solutions: II. Flurbiprofen and diflunisal as models for studying solvation of drug substances. Eur. J. Pharm. Sci. 2003, 19 (5), 423−432. (17) Delgado, D. R.; Holguín, A. R.; Almanza, O. A.; Martínez, F.; Marcus, Y. Solubility and preferential solvation of meloxicam in ethanol + water mixtures. Fluid Phase Equilib. 2011, 305 (1), 88−95.
dissolving process. The greater the Gibbs energy is, the more difficult it is to dissolve. In the selected solvents, the Gibbs energy of ethanol, water, acetic ether, acetonitrile and 2propanol are 19.15, 29.09, 17.60, 17.41, and 18.93, respectively. This is coincident with the analysis of their molecular structures and space conformations. For comparing the relative contribution to the standard Gibbs energy by enthalpy (%ζH, uncertainty of ± 0.01) and entropy (%ζTS, uncertainty of ± 0.01) in the solution process, eq 14 and eq 15 were employed by the literature.16,17 %ζH = 100
|ΔHs| |ΔHs| + |Tref ΔSs|
%ζTS = 100
|Tref ΔSs| |ΔHs| + |Tref ΔSs|
(14)
(15)
The calculated values of %ζH and %ζTS are listed in Table 5, which indicated that the values of enthalpy were the main contribution to the Gibbs energy of the enrofloxacin dissolution process, because the values of %ζH are ≥ 62.24%.
4. CONCLUSION The solubilities of enrofloxacin in different solvents were measured in the temperature range from (303.15 to 321.05) K by a static analytical method. From the experimental solubility data, we could make a conclusion that for pure solvents the mole fraction solubility decreases in the order acetonitrile > acetic ether > 2-propanol > ethanol > water. The Clarke−Glew 2 equation can show better agreement with the experimental data than the modified Apelblat equation, the ideal solution equation and the Clarke−Glew 3 equation. The experimental solubility can be used as fundamental data in the purify process of enrofloxacin in industry, and the Clarke−Glew 2 model could be applied for related process design. In addition, the discussion of Gibbs energy, enthalpy, and entropy for enrofloxacin in the selected solvents indicated that the dissolution process was endothermic and entropy-driven.
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AUTHOR INFORMATION
Corresponding Author
*Tel: +86-0835-2885614. E-mail:
[email protected]. Notes
The authors declare no competing financial interest.
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REFERENCES
(1) Brown, S. A. Fluoroquinolones in animal health. J. Vet. Pharmacol. Ther. 1996, 19 (1), 1−14. (2) Neu, H. C. Quinolone antimicrobial agents. Annu. Rev. Med. 1992, 43, 465−86. (3) Hannan, P. C.; Windsor, G. D.; de Jong, A.; Schmeer, N.; Stegemann, M. Comparative susceptibilities of various animalpathogenic mycoplasmas to fluoroquinolones. Antimicrob. Agents Ch. 1997, 41 (9), 2037−40. 2074
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