Clay Minerals (1984) 19, 865-875
MAGNESIA
FROM
A. S. B H A T T I ,
SEAWATER:
D. D O L L I M O R E
A REVIEW
ANt) A. D Y E R *
University of Toledo, Department of Chemistry, Toledo, Ohio 43606, USA, and *University of Salford,
Department of Chemistry & Applied Chemistry, Salford M5 4 WT, UK
(Received 7 November 1983; revised 18 May 1984)
ABSTRACT: The process developed in the UK to produce magnesium hydroxide from
seawater is described, together with the heat treatment that the hydroxide receivesto produce the
active oxide. Some of the characteristics required of the dolomite used in the process are also
discussed. Impuritiesintroduced by the seawater are noted and the means by which they can be
reduced explained.
Magnesium oxide has long been an important industrial material. The major tonnage outlet
for magnesia has been the refractories industry in which it is employed as 'dead-burned'
grain magnesite for the construction and maintenance of the hearths or bottoms of
open-hearth steel furnaces. Magnesia also is used in the form of calcined refractory
magnesite brick and as a chemical addition in the manufacture of chrome, olivine and other
refractories.
The industrial importance of magnesia is due to its chemical behaviour and this is largely
determined by the position of magnesium in the periodic table. Magnesium oxide is a weak
base, particularly at high temperatures, and therefore magnesium oxide is barely attacked
by alkalis but readily by acids. All strong mineral acids dissolve or corrode magnesium
oxides, even at room temperatures. Only phosphoric, boric and, to a certain degree, silicic
acids have little effect on magnesium oxides, because salts with these acids are insoluble in
water and stable substances. For the same reason, hydrofluoric acid is not particularly
aggressive towards magnesium oxide, and the insoluble magnesium fluoride causes a
certain degree of protection of the magnesium oxide surface against further attack.
Commercial magnesium oxides are generally used in the form of fine grains and powders
(an exception is hot-pressed MgO) and, as such, are susceptible to attack by atmospheric
moisture and carbon dioxide, forming hydrates and carbonates respectively. Larger
crystals of calcined oxides, especially after hot-pressing, are resistant for all practical
purposes to water vapour and to carbon dioxide.
Magnesium oxide can only be melted at 2800~ in a fully oxidizing atmosphere, and
there is very little volatilization below 3200~ A slightly reducing atmosphere suffices at
these temperatures, however, to reduce it to the metal which is then oxidized by air in the
upper part of the furnace to produce thick clouds of magnesium oxide, the density of which
increases from 3.0 to 3.58 g/cm 3 for the sintered product.
DIFFERENT
TYPES
OF MAGNESIAS
In general, only two classifications of calcined magnesia are in common useage:
'caustic/calcined' magnesia, calcined at low temperatures (350-450~
and 'dead-burned'
9 1984 The Mineralogical Society
866
A. S. Bhatti et al.
magnesia, calcined at much higher temperatures (>1300~
Temperature, rate of
calcination and impurities materially influence the properties of magnesia obtained by
heating the carbonate, hydroxide and other compounds. The quantitative effects of these
factors on ultimate crystal size and reaction rates have not always been completely
evaluated in commercial products.
Magnesium oxide does not occur naturally to any great extent. It is made by calcining
magnesium salts and there are two main types available. These are:
1. calcined natural magnesite (MgCO3) or brucite (Mg(OH)2);
2. calcined precipitated magnesium hydroxide, utilizing seawater or natural brines as the
main source of Mg 2+ ions.
The first type has industrial applications based on its relatively non-reactive properties,
e.g. as a filler in rubber or polyester resins. The second type is widely used as a filler where
more reactivity is required; it is this type that will be discussed here.
Fig. 1 shows the calcination of magnesium hydroxide and the type of relationship
between surface area and magnesium oxide content of the products. By controlling the
calcination conditions, a porous magnesium oxide with large surface area can be formed.
The surface area affects its rate of reaction with acids. Active magnesium oxides will react
rapidly with water in the air to form the hydroxide and with carbon dioxide to form the
carbonate.
Magnesium is present as a soluble salt in seawater which contains about 1.3 • 10-3
kg/litre Mg 2+ ions, associated with chloride and sulphate ions (Hall & Spencer, 1973). This
is equivalent to 2 • 10 -3 kg/litre magnesium oxide and, in principle, can be easily extracted
by the addition of a suitable alkali. In the early 1930s small-scale processes designed to
extract chemical magnesia from concentrated brine and seawater were investigated. Earlier
attempts to operate a small plant had not been successful as the gelatinous magnesium
hydroxide precipitates proved too slow in settling and filtering. In 1938 the first large-scale
commercial plant for producing refractory magnesia was built at Hartlepool, by the
Mg(OH) z
70
h~, ~' Mc:JO~(OH)2 ~y + yH20
810
9'0
% MgO
FIG. 1. Schematic representation of variation in surface area vs. M g O content of heated
Mg(OH)2.
Magnesia from seawater
867
Steetley Company, using a then unique process of extracting magnesia using seawater and
dolomite. Because of their consistent composition and quality, magnesias from seawater
often are preferred to natural magnesias.
The trace element content of the product is controlled by careful raw material selection
and, in some cases, certain trace elements can be added to the magnesium hydroxide
precipitates, before calcination, to improve the sintering characteristics of the magnesias.
By controlling the preparation and precipitation conditions, the amounts of chloride,
sulphate and other ions can be controlled. These, in turn, affect the sintering characteristics
of the magnesia (Green, 1969; Guilliat & Brett, 1971; Banerjee, 1969; Riddell, 1969;
Laminy, 1969)--e.g. a very low content of C1 ions would indicate that the particular
magnesia will calcine more quickly than one with a higher concentration.
In a review of the production and properties of seawater magnesia, Hall & Spencer
(1973) outlined the physical and chemical characteristics required of magnesia powders for
use in oxygen steel-making vessels. These include:
(i) ahigh density (3.35-3-40 • 103 kg/m3);
(ii) a low boron content (B203 < 0.05%);
(iii) The correct CaO/SiO 2 molar ratio to give the refractory phase dicalcium silicate;
(iv) a CaO content less than the solubility limit in MgO.
Magnesia, like other solids, especially oxides, has physical properties which are
determined by its previous history and heat treatment. The physical properties of such
solids, as distinct from their chemical properties, may be affected profoundly by the
pretreatment that they have received. The term loosely applied in this connection is
'activity'. Although the chemical constitution remains the same, the more 'active' a solid is,
the more rapidly it takes part in chemical reaction. An early detailed investigation of the
factors governing the activity of solids was made by Huttig (1941, 1942), with particular
reference to the changes produced on heating.
The preparation of active magnesia
The usual way of preparing active solids is to utilize thermal decomposition or
calcination (Sing, 1949; Gregg & Sing, 1951); the best known active solids prepared by
this method are lime or plaster of paris. Consider a chemical reaction of the type
Solid A --, Solid B + Gas C
in which, providing the correct conditions of temperature and time of heating are chosen,
the solid B is obtained in a very active form. The essential requirement for solid B to be
more 'active' than A is that A and B should have different structures, i.e. the above
equation represents a true chemical decomposition. The mechanism of activation has been
discussed by Gregg & Sing (1951), Hill (1951) and Sing (1949): the decomposition of solid
A first gives rise to product B, which is in the form of a pseudo-structure of A in which the
metal and oxygen ions are still occupying much the same positions as they held in the
original structure. RecrystaUization of the pseudo-structure produces the stable solid B.
Care must be taken when preparing active solids that the conditions are not too
vigorous, for then an inactive solid is obtained due to sintering; the active sample may be
produced temporarily, but the stringent environment can cause the small crystallites
to adhere with a subsequent reduction in internal area. An investigation by Huttig
(1942) into the sintering of metals and metal oxides demonstrated the importance of
868
A. S. Bhatti et al.
calcination temperature in decreasing the activity of the solid, but in this investigation the
solids were always morphologically stable over the temperature range used. Huttig's metal
oxides were preheated to a high temperature and then finely ground; his metals were also
in a finely-divided form with an active solid prepared by a decomposition of the type shown
in the above equation, In equilibrium, however, the sintering process is altered by the
occurrence of a polymorphic transition, since according to Hedvall (1937) such a
transformation should give rise to additional activation.
Sing (1949) and Gregg & Sing (1951) studied aluminium hydroxide and hydrated ferric
oxide, two systems showing different activation-temperature curves. With ferric oxide
activation was greatest with the unheated sample, and subsequent heating caused a steady,
unbroken fall in activity; when aluminium hydroxide was examined, unheated samples
were not very active and alumina (with a surface area of ~300 mZ/g) was produced only
when the sample was heated to 300-500~ Magnesium hydroxide can readily be obtained
as a stable, crystalline solid, and on heating it loses water to give magnesium oxide; the
only solid components in this system are magnesium hydroxide (brucite) and magnesium
oxide which does not undergo any subsequent changes up to its melting point (2800~
Some workers have studied the changes that occur on heating magnesium hydroxide
while others have examined the closely allied system of magnesium carbonate. Thus
Budnikov (1930) used the rate of dissolution in water as a criterion for the activity of
magnesium carbonate at various temperatures, and discovered that the most 'active'
sample was that obtained at 700~
Lamb & West (1946) studied the changes in
adsorption of nitrous oxide on samples prepared by removing different amounts of water
from magnesium hydroxide. However, this work did not restrict the time and temperature
o f heating within definite limits and, in addition, did not extend outside the range
300-500~
the rate of sintering cannot therefore be estimated from their data. Again,
Fricke & Luke (1935) and Fricke (1936) showed from the heats of solution in acids and
other measurements that the total energy content of magnesium hydroxide and magnesium
oxide varied considerably with the method and temperature of preparation; a variation of
11.3 k J/mole was noted between the energy contents of the most active and least active
sample. Similar studies by later workers (Ginque, 1949; Torgoson & Schamer, 1948;
Taylor & Wells, 1938) confirmed this effect. Taylor & Wells (1938) noted no change in
crystal structure, so they attributed the variations in heat content mainly to differences in
surface properties.
Adsorption on solid magnesium oxide has been used as a basis for measurements of its
surface area. Adsorption of dyes and other substances from the liquid phase is fraught with
difficulties due to possible preferential adsorption of the solvent, as emphasized by McBain
& Dunn (1948). Zettlemayer & Walker (1947) compared the adsorption of iodine from
carbon tetrachloride solution with the results obtained by nitrogen adsorption on a number
of active magnesium oxide samples. Agreement between the two methods was not good,
iodine adsorption in all cases being lower. The active magnesium samples used by these
workers were obtained from a large-scale plant and the experiments they performed were
aimed at characterizing the particular samples rather than finding the exact changes in
activity with varying treatments. Some calcination studies were carried out on magnesium
oxide by Eubank (1951) who reached similar conclusions to those of other investigators,
i.e. the change from the brucite to the periclase structure occurs rapidly above 350~ and
that magnesium oxide only exists in one of these forms. However, he did notice a slight
distortion of the unit cell of the periclase structure at lower temperatures.
Magnesia from seawater
869
More detailed research (Glasson, 1963) has shown that the unit-cell dimension of the
periclase varies from 4.24 A to a limiting value of 4.21/~ as the calcination temperature is
increased from 400 to 1000~
the final density of the periclase being 3.58. The
contraction of the unit cell is accompanied by expulsion of traces of hydroxyl water.
Magnesium oxide from basic magnesium carbonate gave similar results. Changes in
surface area during calcination of magnesium hydroxide and magnesium carbonate
(magnesite) have been followed by a gravimetric nitrogen gas sorption technique (Glasson,
1956). Magnesium oxide of maximum surface area is obtained at about 80-90% hydroxide
decomposition and 90-100% carbonate decomposition. Similar studies on Steetley
dolomite (Glasson, 1964b) gave analogous results. In the dolimes formed, the average
crystallite sizes (from X-ray line-broadening) of the magnesium oxide were correspondingly much smaller than those of the calcium oxide. Steetly and Coxhoe dolomites gave
products of correspondingly similar activities at lower calcination temperatures, but above
700~ the Coxhoe dolimes were less active, as impurities such as ferric oxide promoted
sintering. Further studies have been made on the hydration of calcium and magnesium
oxides (Glasson, 1958, 1960, 1961, 1963)and dolomitic lime (Glasson, 1964a), and their
reactions with seawater have been studied by surface area, optical- and electronmicrographic (Glasson et al., 1968) and radio-isotopic (Glasson et al., 1968) techniques.
Nearly every commercial process that has been developed for the production of
magnesium metal or magnesia requires either a high-calcium or dolomitic lime as the basic
raw material (Boynton, 1966). Thus the Dow Seawater Process uses hydrated time,
MgC12 (from seawater) + Ca(OH) 2 --, Mg(OH) 2 + CaC12,
followed by either
(A) Mg(OH)2 + 2HC1 --, MgC12 + 2H20 and MgC12 + electrolysis -, Mg + C12
or
(B) Mg(OH) z + heat --, MgO + H20 (Seawater Chemical Magnesian Process)
The Dow Natural Brine Process uses dolomitic lime,
CaC12, MgC12 + CaO.MgO + 2CO z --, 2CaCO 3 + 2MgC12,
followed by electrolysis of MgC12. The Steetley Process also uses dolomitic lime and is
described now in detail.
PROCESS
FOR
PRODUCING
MAGNESIA
FROM
SEAWATER
The chemistry of the process is relatively simple. First, dolomite is decomposed in a rotary
kiln at 1350-1400~ The calcination reaction is
CaMg(CO3) 2 ~
dolomite
MgO. CaO + CO2
dolime
The dolime is then hydrated
MgO. CaO
+ H20 ~
MgO. Ca(OH) 2
870
A. S. Bhatti et al.
and finally reacted with the seawater to precipitate magnesium hydroxide
MgO. 2Ca(OH) 2 +
MgSO 4
CaSO 4
MgC12
--, M g O . 2 M g ( O H ) 2 + CaC1 z
magnesium salts
spent
seawater
seawater
Some of the MgO from the dolomite will hydrate during the process, but most of it comes
through unchanged, because when the dolomite is fired to 1350-1400 ~ C the MgO forms a
deadburnt layer around itself and becomes almost inert.
The production process
The calcined dolomite is slaked to a fine dry powder in pay hydrators. The hydrated
product is made into a slurry to facilitate handling by mixing with seawater and is classified
to remove impurities, particularly fuel ash and any non-calcined or over-calcined,
unreactive materials. The seawater is first pretreated to remove calcium bicarbonate-achieved by adding a small proportion of sulphuric acid to change the pH to 3.4-4.0 to the
centre of the tank by rotating r a k e s - - a n d then pumped out as a thick slurry. This slurry is
washed with fresh water to reduce the sodium chloride content.
The main problem with this process has always been that the precipitates settle badly
and are very difficult to filter. It has been found that magnesium hydroxide derived from
Fe2+ ~ 13%
Fe2~ < 0 1%
CaCOa.MgCO3
CaCO3. MgCO~
l oo~
t
CaO.MgO
CaO
,l HzO
I H2o
Ca(OH)2. MgO
Ca(OH)2
I H20 + Mg2+
I MgZ+
2Mg(OH)2 + Ca2+
Mg(OH)2 + Caz{
Rotary vacuum filters
Paste of 50% Mg(OH)z + 50% H20
Multiple hearth furnace
MgO
FIG. 2. Two methods of producing magnesiumoxides.
Magnesia from seawater
871
the dolomite acts as nuclei for precipitated magnesium hydroxide and improves the settling
rate. The settled slurry still contains 0.1 to 0.12 kg/litre of solids and, even after
steam-heating to 90 ~ C before filtering, only produces a filter cake of 30-35% solid content
(Hedvall, 1937).
A process of seeding has been developed (Budnikov, 1930) where up to 85% of the
product is fed back to the reaction tanks. These recycled magnesium hydroxide particles
act as nuclei for the precipitation of further magnesium hydroxide and the resulting
precipitate settles down to an increased solids concentration of 0.3-0.4 kg/litre; this can be
filtered more rapidly, and to a higher solids concentration (50%), without heating.
In 1955, organic flocculants became available and they are effective at very low
concentrations (1-10 p.p.m.), thus making their use economically viable. Those based on
polyacrylamide are particularly effective in enhancing the settling rate of magnesium
hydroxide precipitates but do not result in a filter cake with a lower free water content.
The slurry is filtered on rotary vacuum disc filters, and additions of lime, silica, and iron
are made in order to control the ratios of the impurities in the magnesia. The filter cake,
containing 50% solids, is then either fed directly into rotary kilns or calcined in a
multi-hearth Herreshoff furnace and formed into pellets by passing through pelletizing
rolls. The rotary kilns 'dead-burn' the product at temperatures in excess of 1600~ to
produce refractory grade magnesia.
Magnesias prepared using dolomite contain ~ 1.3% Fe203; magnesias containing
<0.1% F%O 3 are prepared using calcium carbonate. The two methods are shown
schematically in Fig. 2.
The Herreshofffurnaee
The paste from the rotary vacuum filters is fed to the top of the furnace and it slowly
travels along the hearths until lightly calcined magnesia is discharged from the bottom. To
produce magnesium oxide of high purity, high density and large grain size, a carefully
controlled temperature profile is required.
The upper hearths effect the first stage, i.e. removal of free water, the middle hearths
remove combined moisture and volatiles, and the lower hearths condition (calcine and
sinter) the oxide. The temperature profile ranges from 400~ at the top to 900~ at the
bottom of the furnace. The exhaust gases carry with them some of the oxide (20%) as dust
which is drawn from the furnace and cleaned in a bag filter unit which strips out the oxide
and allows the clean gases to pass to atmosphere, the collected dust being returned to the
middle section of the furnace.
Impurities from dolomite
Careful control of raw dolomite quality must be maintained because the impurities in the
calcined dolomite are precipitated in the magnesium hydroxide. In order to achieve this
control, it is sometimes necessary to blend rock from different areas of the quarry, and the
crushed dolomite is washed to remove clays, and hence reduce potential AlzO 3 and SiO 2
contamination. Hydrocycloning removes coarse grit, which has a high silica content (Hall
& Spencer, 1973).
Dolime produced initially from dolomite must be of good reactivity, as any unreacted
dolime will appear as a lime contamination in the product; the calcination conditions of the
872
A. S. Bhatti et al.
dolomite must therefore be such that all carbon dioxide is removed, while avoiding
overburning which would result in poor slakeability. Gilpin & Heasman (1952)listed some
of the compounds formed during calcination of the dolomite, with their stabilities in
seawater.
Many of the impurities are difficult to remove because they are very finely distributed,
and inevitably the dolomite introduces SiO 2, Fe203, A120 3 and CaO impurities, with
perhaps lower levels of others such as MnO and B203. Various methods of chemically
purifying the magnesia are being investigated and these could result in extremely pure
grades (Gilpin & Spencer, 1972).
Impurities from seawater
Lime. Seawater contains calcium bicarbonate which reacts with calcium hydroxide to
precipitate calcium carbonate.
Ca(HCO3) 2 + Ca(OH) 2 --, 2CaCO 3 + 2H20
This, being insoluble, appears as a lime impurity in the product at a concentration of about
3% (Gilpin & Heasman, 1963). Earlier methods of minimizing this impurity have been
abandoned in favour of the newer acid pretreatment of seawater whereby sulphuric acid is
injected to lower the pH of the seawater from about 8.5 to 3.5-4.0. The sulphuric acid
reacts with the calcium bicarbonate to form calcium sulphate and carbon dioxide. The
carbon dioxide is removed in a wooden desorption tower. Here, the water is broken into
small droplets and the carbon dioxide escapes; with this method the lime content of the
product can be reduced by 0 . 4 - 0 . 5 % (Gilpin & Heasman, 1963).
Silica. Most of the silica impurity comes from sand suspended in seawater. Although
seawater inlets are sited in sheltered (Hall & Spencer, 1973) areas, some very fine sand is
inevitably drawn into the system and in heavy storms silica contamination rises. A way of
reducing the silica contamination would be to absorb the silica onto a precipitate of either
magnesium hydroxide or calcium carbonate in the hydro-treaters prior to the main
reaction.
Boron. Boron is present in small amounts in seawater as boric acid (Svendup &
Johnson, 1942). The precipitates of magnesium hydroxides have a high capacity for
absorbing boron and the final concentration in the oxide can be as high as 0.4% (seawater
contains 15 p.p.m, expressed as B203). This boron contamination can be reduced by using
an excess of lime in the precipitation but this also increases the lime contamination;
however, it has been used without increasing the lime contents beyond acceptable limits.
Alternative methods have been studied, and one commonly used is to remove the boron
at the dead-burning stage by volatilization. The rate can be accelerated by the addition of
certain alkali metal salts, e.g. potassium hydroxide. The use of such methods may reduce
the boron content of seawater magnesias to 0.03% (Gilpin & Spencer, 1972).
Calcium sulphate. As the precipitation of magnesium hydroxide proceeds, the
concentration of Ca 2+ increases until after ~75% of the magnesium hydroxide has been
removed; the theoretical solubility limit of calcium sulphate is then exceeded and hence the
possibility arises of gross calcium sulphate contamination. Fortunately, calcium sulphate
forms stable super-saturated solutions and this prevents the calcium sulphate contamination problems.
Magnesia from seawater
873
Calcination--effect of time
The effect of the time of heating on magnesium hydroxide can be demonstrated by
quoting some of our own results. Adsorption-desorption measurements of nitrogen at 77.4
K were carried on the parent as well as on the thermal decomposition products of the
hydroxide, which were formed by heating at 350~ for 1, 2, 3, 4, 5, 6 and 7 h in both wet
and dry nitrogen. The BET surface areas of the products are shown as a function of the
calcination time in Fig. 3. The most striking result is that obtained by heating in dry
nitrogen for 6 h; the area increases about 13-fold over that of the parent material. Heating
at 350~ for longer than 6 h causes a gradual decrease in surface area, obviously due to
sintering forces (Laminy, 1969). Calcination of the magnesium hydroxide is also affected
by the presence of a constant water vapour pressure. Fig. 3 shows that the specific surface
areas of the various products obtained by decomposition in the presence of water vapour
are appreciably lower than the corresponding areas obtained in dry nitrogen. However, the
general trend of variation of area with respect to the calcination time is approximately the
same in dry and wet nitrogen. In both cases, the area increases with time of calcination to
reach a maximum. For decomposition in the presence of water vapour, the maximum is
400
9 dry nitrogen
0 wet nitrogen
300
to
200
~u
03
1 O0
I
1
l
2
I
3
I
4
I
5
I
6
I
7
C a l c i n a t i o n t i m e (h)
F]G. 3, The development of surface area in dry and 'wet' nitrogen at 350~ The 'wet' nitrogen
was obtained by passing it through a series of Dreschel bottles filled with water and glass rods.
A. S. Bhatti et al.
874
shifted to a longer calcination time. The efficiency of removal of the water vapour formed
on decomposition of Mg(OH) 2 affects the energy of activation which therefore varies
appreciably from sample to sample of Mg(OH) 2 with different surface areas and porosity.
For the decomposition of Mg(OH)2 in wet nitrogen the water vapour concentration
gradients along the capillary passages in the product will be smaller than in dry nitrogen,
giving lower rates of nucleation of the newly formed magnesium oxide crystallites.
Differences of similar type are found for magnesium and calcium hydroxides decomposed
in air with vacuum conditions (Glasson, 1956, 1963).
Calcination--effect of temperature
Our own results can be used also to show the effect of temperatures on the activity of
magnesium oxide produced. Results of experiments on the adsorption of nitrogen at 77.4
K on samples of magnesium hydroxide calcined in air for 5 h at various temperatures are
shown in Fig. 4.
The appearance of a maximum in the specific area-temperature of dehydration curves
has been frequently found in the decomposition of many crystalline hydroxides used as
parent materials in the preparation of active solids. This maximum generally is attributed to
the existence of more than one process during thermal treatment, namely, decomposition,
recrystallization and sintering. It is known that when dehydration occurs at lower
temperatures the former two processes predominate; whereas above these temperatures
sintering becomes predominant. At temperatures near that producing the maximum surface
area, the processes of activation and sintering overlap (McBain & Dunn, 1948).
300
200
%
1 O0
I
200
J
I
400
I
600
Calcination t e m p e r a t u r e (~
F1G. 4. The effect of calcination on the adsorption properties of magnesium oxide.
Magnesia from seawater
875
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