Formation of barium hexaferrite

.k~ of Magneto's: and Magnetic Materials I 0 9 7 5 ) 144--152 ~ N k ~ o l l a . ' ~ c +Publishing Company

FORMATION OF BARIUM HEXAFEIEt~ITE A.M GADA1LA* and H.W. HENNICKE ~rxtab~ [fir Glas und Keramik, Tll Clausthal, West Germany ] h e re~ctis~ity of hematite produced by gyrolysis of iron salts at various temperatures was assessed by surface area . ~ent, dec;a'on microscopy, crystal size and by reacting it with BaCO 3 to form hexaferrite. The effect of heating ra~, soakSng temperature and time were studied. The mechanism of hexaferrite formation was investigated and the previously rep,.~tted unknown X-ray diffraction pattern was explained. It has been proven that the various compounds higher in llaO than the monoferrite were formed first, but the main intermediate phase was the moncferrite. Hexaferrite prep~:-ed at lower temperature for longer periods were found to be finer than that prepared at higher temperatures for shorter periods.

I. Introduction Although the technology of hexaferrite formation is now known, several points need further investigation to explain the conflicting data regarding the intermediate compounds formed during fabricatior~, specifications of raw materials and their effect on various production stages.

2. Previous work

, 1. TT~ereactivity of Fe203 Various conflicting data were reported on the reactivity of iron oxides. While Blum and Li [1 ] cotl:cluded that reactivity increases with partial size, Erzberger [2] reported the reverse. Jaworski et al. [3] compared two types of Fe203 prepared from sulphate and hydroxide and claimed to have similar physical properties. Using the thermobalance they found that sulphate ga~e oxide with higher reactivity and or, sintering, the hydroxide gave oxide producing higher dens~ies and ~-eater shrinkage. Recently however, Ratnam and h,gham [4] and Gal|a~her et al. [5] conclude~ that surface area and related parameters (tap density, average particle size..) d~es not represent the reactivity characteristics. Gallagher et al. [5] four:d that iron oxide prepared from nitrate at low temperatures has low reactivity, * A~oc~ate Prcfessor, Cairo University.

and at high tei,nperature large hard aggregates less suitable for sintering are produced. They found also that the oxide produced frora the sv',phate is obtained at relatively higher temperatures ~ I possesses a fresh more reactive surface.

2.2. The formation of barium hexaferi'ite Several research workers studied the ~ormation of hexaferrite from BaCO3 and Fe203 in oxygen, air and in vacuum using various techniques. They noticed the formation of an intermediate phase which was considered by Sadler et al. [6] as BaO and by various authors [7-10] as BaO" Fe203 . Accordingly they suggested that the monoferrite is formed first and at higher temperatures it reacts with the remaining Fe203 forming the hexaferrite. Suchet [ 11 ] showed also that the monoferrite, was formed between 700-900°C but below 800°C he showed the existence of a compound claimed to be BaO.2Fe203. Recently Bye and Howard [12] showed that apart from the monoferrite, some infrared absorption peaks appeared and could not be tel*fred to the known comFuultu~ n t l p t y m o a Cuniptex sequence u~ ~ evmtts. ..... At low temperatures, Haberey et at. [13] noticed the presence of an X-ray pattern other than that of the monoferrite and considered it to be of Fe3 04. They showed Fe304 and Fe203 to exist over a wide range of temperature (700-775°C) in air, and such a result is in contradiction to the phase rule. Bye and Howard [12] noticed also that the rate of disintegration of BaCO3 was far more rapid than the weight loss and . . . . .

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Fig. 2. Crystalline size of Fe203 produced from v~rious iron salts.

accordingly they suggested that the lbrmation of monoferrite involved a number of fast steps in which barium ferrites were formed but were not in sufficient quantity or crystallinity to appear using X-rays. Later they showed in polished sections [14] that 5BaOX Fe203, 3BaO. Fe203 , 2BaO. Fe203 and BaO. X Fe203 were present and that 3BaO.Fe203 was the msjor product.

for a constant time. Thermog~avimetric changes in these mixtures are shown in fig. 3. In all mixtures, a progressive weight loss occurred far below the dissociation temperature of pure BaCO3 , indicating that the iron oxide decreased its stability with various values depending on the reactivity of ferric oxide. With an active oxide the loss was rapid and a constant weight was reached at low temperature, if the specific area is the only factor determining the reactivity, one expects to find all the curves arranged with the highest value of surface area to the !eft followed by lower values irrespective to the origin of the iron oxide used. This order was not found and accordingly the curves were classified in (two) groups; (A) those showing reactivity in proportion to values of specific surface shown in fig. 1 and (B) those giving vflues unexpected from the point of view of surTace energy when superimposed on curves (A). It was found that while Fe203 produced from sulphate or oxalate has reactivity in proportion to specific surface, that produced from nitrate or hydroxide gave lower reactivity than expected. This effect was attributed to agglomeration which will be discussed later in view of the electron micrographs. Also the fact that the reactivity of oxide prepared from nitrate at 300°C was lower than 40O°C inspite of its very high surface area, will be discussed later. To compaxe particle size, particle size distribution, shape and tendency for agglomeration, some specimen were examined in an electron microscope. The micrographs obtained are shown in figs. 4 - 7 . It is evident that each particle of the salt produced at lc~wtemperatures agglomeration of small crystaUites wi'~hsmall pores between them and in some cases ~ts the cases of

3. Results and discussions

3.1. Reactivi~ of hematite lTroduced by pyrolysis o f iron salts Hydrated iron sulphate: nitrate, oxalate and hydroxide were fired at various temperatures above that required for producing alpha Fe203 [15]. The results obtained after measuring the specific surface area and X-ray diffraction broadening are shown in figs. 1 and 2 indicating that at low temperatures when iron oxide can be prepared from nitrate or oxalate, higher specific area was obtained but the rate of crystal growth was very high. (It should be noted th~Ltthe oxide produced from the oxalate gave crystal sizes higher tkan 1000 A and since these values are above the limit of accuracy of X-raye they were not shown in fig. 2.) In all cases the surface area drops at high temperatures due to increased particle size. To assess the reactivity, the oxide produced at various temperatures from different origins were mixed with BaCO 3 in the ratio to form BaO.6Fe203. For all mLxtures, mixing was carried out in an agate mortar

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e,'dde from oxalate at 400 and 700°C and from sulphate at 050°C, these pores are too small to admit N 2 m adsorption experiments, thus giving relatively lower sp,ecific surface areas. Nitrate and hydroxide gave ve~' fine c~slallites with high cohecive power, producing agglomerates with wide size distribution. A.: higher temperatures, gases evolving from sal~ha~.e c~,st,~Asshattered :hen; to fine particles with unifo~: size distribution. Such shattered particles will have sharp edges mad at the temperature used they will be rounded isy surface diffu,ion. On the other hand t1~e oxide produced from oxahte was retained in the erigin~ cE
fine Fe203, with very high surface area tend to floc and adhere to each other rather than to mix with BaCO 3, thus decreasing the number of contacts and the rate of reaction. During firing such agglomerates are expected to shrink giving loose contacts. Increasing the temperature of firing decreases the surface area to an optimum limit; thus decreasing the agglomeration tendency but giv':ng surface energy high enough for solid state reactions. It should be noted that for ultrafine grain ceramics, a wide range of particle size is not recommended since this will enhance exaggerated grain growth Accordingly iron oxide prepared from sulphate at moderate temperatures (650-850°C) is recommended.

3.2. Mechanism of hexaferrite formation After soaking BaCO3 and Fe203 (from various origins) at different temperatures, the main phases detected by X-rays were found to be monoferrite and extra lines decreasing in intensity with higher temperatures and disappearing above 1200°C. These lines could be considered 2BaO" 3Fe203 , the d..spacings of which were published recently [16]. Compounds higher in BaO than monoferrite were not detected. Using the internal standard method, the percentage of mono- and hexaferrite were calculated ;Jnd are shown

A.M. Gadadtla,H. W. Hen ~icke/Formation of barium hexalerrite

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Fig, 4. a-Fe203 produced from hydrated ferric nitrate at (a) 200°C (b) 500°C (c) 700°C (d) 1150°C. in fig. 8. It is evident that the formation of both depends on the raw materials and their ratio. High BaCO 3 content increased the rate of formation of monoferrite at low temperatures and of hexaferrite at high temperatures. These results disagree with

those published by Wullkopf [17], wb.o stated that the formation of monoferrite is not affected by the stoichiometric ratio of BaO/Fe203 and is 1.it,tie affected by grain size of raw materials. It should be noted that, the thermobalance gave

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Eig. 5. o Fe203 produced from hydrated iron sulphate at (a) 650°C for 1 hr (b) 850°C for 1 hr (c) 1150°C fear I hr. smc~oth ,:u:c:'es (fig. 3) with one peak oil DTG curves, ~,~ggesti,~g :h,~t on¢. reaction occurs: the formation of r~:'~;fcrrite. -Xfter a wei,~t k~ss correspoading ~to the

disappearence o f carbonate, the mixture should consist o f monoferrite and F e 2 0 x as well as small amounts o f hexaferfite, but the unexplained extra lines men-

A.M Gadalla, H. IV. tlennic,~e/Formation of barium hexalerrite

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Fig. 6. ot-]:c203 produced from hydrated iron oxide at (a) 400°C for 1 hr (b) 700°C for 1 hr (c) 1100°C for 1 hr.

tioned earlier were also present and were referred to 2BaO" 3Fe203. This confirms that the reaction is more complicated. To increase the sensitivity of the thermobalance, mixtures containing high BaCO 3 con-

tent (just to form tile monoferrite) were examined and the weight loss was not cor~tinuous as in fig. 3 (see fig. 9). DTG curves, which show the rate of weight change, gave three peaks separated by the tempera

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A.M. GadaUa, H. W. Hennicke/Forraation o f barium hexa[errite

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Fig. 11. Electron micrographsof hexaferrite prepared under various conditions. to its low crystalHnity. From 650-780°C BaO" F,,203 was the main product and at higher temperatures, the rate of formation of other intermediate compounds increased. 2BaO" Fe203 was detected above 780°C and 2BaO" 3Fe203 above 900°C. The latter phase disap-

peared above 1200°C due to its peritectic dissociation and above 1250°C the main phases were 2BaO'Fe203 . BaO-Fe203 and Fc203. In contrast to curves sho,,/n in fig. 3, constant weight corresponding to con lplele dissociation of c,a~-

152

A.M. C,aara!ia,H. W. Hennieke/Formation of barium hexaferrite

bo~r~te ~as not reached implying that the presence o f e x ~ F e 2 0 3 accelerated the reaction rates and lowerthe ~bo~v menlioned temperatures. It is also evident that ~ n oxide produced from the nitrate at 4 t ~ C .~: ~ ) r e reactive than that produced at 500"C f¢~nd earlier, F~g. 10 shows the effect o f the heating rate on w~qght ch~ng'-~ in mixtures forming the hexaferrite, indicating that higher reaction rates can be obtained with lk~ heating rates, This result does not agree with V~3akIer !7] who stated that heating and cooling rates o f 2 0 ~ ° ( ' ~ r did not affect the reactions and neglected this eff~:t in his studies.

rite at higher temperatures and produces fine crystals of hexaferrite in shorter time at lower temperatures. While 10wer preparation temperatures produced a fine graided pr0d0et suitable for isotropie magnets, higher temperatures produced a prodae~ suit~Ie for anisotropic or plastically bonded magnets.

Acknowledgement

The authors have pleasure in acknowledging the award of a fellowship to one of them (Dr. A.M. Gadalla) by the Alexander yon Humboldt Stiftung, which mad~ this investigation possible.

3.3. F.~tp~ affecting hexaferrite grain size "fn compare the hexaferrite prepared at various ~emperatures, it was found by X-rays thai the reaction v~s rt:ar~¢ complete at 1 IO0°C after 3 hrs, at 1200°C after 5 hn; and at 1300°C after 3 hrs. Electron micro~mapb,s for these three conditions are shown in fig. 11 (a), (b) and (e) respectively. Firing at low temperature for a tong time gave hexaferrite with very fine crystals and of nearly uniform size but strongly adhering to each other. Such material is expected to be reactive and wilt form small grains on sintering suitable for isotropic ma~.:~s. On the other hand, firing at high tem~ra~ures f~r shert periods gave well formed large ,:r.~~tals wh!ch can be easily broken by grinding into ~-~icles conststmg oi individual single crystals for anL~otropic or plastically bonded magnets. ~ s o fig. i 1 (d) shows that using BaO/Fe20 3 o f 1:5.6 gave az 1300°C hexaferrite with smaller grains after ordy one hour.

4. Conclusions

Reacti~Sty of iron oxide is not only a function o f cD,stat size and specific surface. The reactivity was a~e.~ed by the reaction with BaCO 3 in a thermobalance. Du,fing hexaferrite formation, various com~uwds higher in Bz,O thar~ the monoferrite are formed as well as 2BaO o3Fe20 3 which was present up to t 200°C. Fkgh BaO content increases the rate of formation of monoferrite at low temperatures and the hexafer-

References [1] S.L. Blum and P.C. Li, J. Amer. Ceram. Soc. 44 (1961) 611. [2] P. Erzberger, "Magnetic Ceramics", Proc. Brit. Ceram. Soe. 19 (1964). [3] J.M. Jaworski, G.A. lngham, W.S. Bowman and G.E. Alexander, J. Canad. Ceram. Soc. 38 (1969) 171. [4] D.V. Ratnam, G.A. lngham, N.F. Bright, R.H. Late and F. Rowland, J. Canad. Ceram. Soe. 36 (1967) 20. 15] P.K. Gallagher, D.W. Johnson, F. Schrey and D.J. Nitti, Amer. Ceram. Soc. Bull. 52 (11) (1973) 842. [6] A.G. Sadler, W.D. Wcstwood and D.C. Lewis, J. Canad. Ceram. Soc. 33 (19641 127. [7] G. Winkler, Reactivity of solids, Ed. G.M. Schwab (Elsevier Pub. Co., 1965). [8] J. Beretka and MJ. Ridge, J. Chem. Soc. (A) (1968) 2463. [9] H. S~.~'bleinand W. May, Ber. Deut. Ker am. Ges. 46 (1969) 126. [10] H. Sth~oleinand J. Willbrand, Reactivity of solids, Eds. J.S. Anderson, N.W. Roberts and F.S. Stone Chapman and Hall, London, 1972). [11] J. Suehet, Bull. Soc. France Ceram. 33 (1965) 33. [12] G.C. Bye and C.R. Howard, J. Appl. Chem. Biotechnol. 21 (1971) 319. [13] F. Haberey, M. V01icescuand A. Kocktel, Int. J. Magnetism. [In] C.M. Wilscm,G.C. Bye, C.R. Howard, J.H. Sharp, D.M. Tinsley and S.A. Wentworth-Rossi, Reactivity of solids, see ref. [5]. [15] A.M. Gadalla and H.W. Hennieke, Powder Met. Inter. 5 (1973) 196. [161 G. Sloccri, J. Ame:t.Ceram. Soe. 56 (1973) 489. [171 H. Wullkopf, Int..L Magnetism5 (1973) 147.