Cellular Ceramics / p2

2.9 Cellular Concrete 205

Figure 9 a) TEM image showing crystals of tobermorite. b) Enlargement of overlapping crystals of tobermorite.

AAC formulations have bulk compositions that are significantly more silica rich than their low-temperature counterparts. Typical formulations based on quartz flour and fly ash are given in Tab. 1. Quartz flour AAC slurries normally contain a greater proportion of lime relative to Portland cement vis vis an equivalent AAC made from Class F fly ash, which usually contains a greater proportion of cement and less lime.

Table 1 Formulation of a quartz AAC and a fly ash AAC (in wt %)

In both cases, the starting materials contain enough lime and Portland cement that the slurries will hydrate at room temperature and develop green strength, but at these temperatures the additives (quartz sand flour or Class F fly ash) will mostly act as inert fillers. However at elevated steam temperatures of 160–180 C, these additives will combine with the lime-rich C–S–H phase in the green body to form tobermorite (Ca5Si6O17·10 H2O). Densities of autoclave-cured masonry are in the

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200–1000 kg m–3 range which is a direct consequence of varying the air content of the mixture [11].

Autoclave curing at 180–190 C and saturated steam pressure causes the C–S–H formed at ambient temperatures in the green block to undergo varying degrees of crystallization. Sato and Grutzeck [19] proposed that crystallization occurs most easily from a lime-rich C–S–H precursor that forms during the initial setting of the cellular concrete. They suggest that poorly crystallized lime-rich C–S–H contains shorter silicate units [(Si2O7)6–] that are more easily rearranged during higher-tem- perature curing than a comparable C–S–H already containing considerably longer silicate chains [(Si3O9)n6–]. Very fine silica such as a colloidal silica is so reactive (pozzolanic) that it causes tobermorite-like C–S–H with a Ca/Si ratio closer to unity to form in the green cake. This tends to inhibit further crystallization at autoclave temperatures. Conversely, the industry standard, a fine-grained silica sand (silica flour) is relatively unreactive. The C–S–H that forms tends to form around the cement grains and coat the quartz grains. As a consequence the C–S–H has a Ca/Si ratio near 1.7. Once this is autoclaved the C–S–H and remaining anhydrous ingredients combine to form a very crystalline tobermorite, similar to that in Figures 6–9. Clearly, small amounts of impurities as well as changes in silica source and granularity tend to alter the shape and size of the tobermorite crystals. The effect of such changes on transport properties is unknown.

The calcium aluminate hydrates that form and coexist with C–S–H are crystalline at all temperatures: Ca2Al2O5·8 H2O which in shorthand becomes C2AH8, Ca4Al6O7·13 H2O (C4AH13), and hydrogarnet Ca3Al2O6·6 H2O (C3AH6) and are thus much easier to study than C–S–H by X-ray diffraction and electron microscopy. The final phase that forms in a hydrated cement is portlandite Ca(OH)2 (CH). Portlandite forms as a byproduct of the decomposition and formation of C–S–H from tricalcium silicate Ca3SiO5 (C3S) and dicalcium silicate Ca2SiO4 (C2S). For more information on the nature of these and other intermediate phases found in cementitious materials the reader is referred to Taylor [1] and Lea [2].

2.9.5

Portland Cement

Because Portland cement and its hydration products are central to the production of cellular concrete, it is necessary to spend some time discussing its history, chemistry and properties. However, this is not an easy matter, because the chemistry of Portland cement hydration is daunting. Even after 100 years or so of commercial use in the USA there are still questions as to what actually happens in terms of dissolution, nucleation, and growth when Portland cement is hydrated [1, 2].

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2.9.5.1

History

Portland cement was formulated, evaluated, and first commercialized in the UK in the mid 1800s. It was found that a mixture of certain argillaceous limestone could be fired at 1400–1500 C to form a mixture of calcium compounds (phases) that would react with water, and given time would harden to form a rock-hard product much like the famed Portland stone quarried in the southern parts of England. Its name was meant to imply that Portland cement would be as good as or perhaps even better than Portland stone. As it turned out, it was as good as it was supposed to be and became immediately popular. Engineers began to use it to repair landmark structures and the rest is history. A revolutionary new material was born. The technology was brought to the USA in the early 1900 by entrepreneurs. It was referred to as a synthetic cement in contrast to what was then available in the USA, that is Rosendale and other natural cements made from argillaceous limestone deposits in Rosendale, NY and Louisville, KY. Once it was realized that Portland cement could be made from limestone and kaolinite clay, cement plants sprouted up all across the world [1,2].

2.9.5.2

Fabrication of Portland Cement

Like many commercially available cements the needed raw materials are often fired in a furnace of some type to dehydrate, vitrify, or otherwise make a substance that contains a large amount of latent free energy. In other words, once treated it will react with a solution of some sort and form hydrated solids that have the unique ability to intergrow and produce monolithic solids with useful physical and mechanical properties. Portland cement is currently made in rotary kilns equipped with a host of preheaters, coolers, recyclers, grinders, and filters needed to make them energy efficient and environmentally compliant. The kilns are fired at one end and fed with raw materials at the other. The limestone and clay move through the kiln and gradually become dewatered, decarbonated, sintered, and then react at 1450–1500 C in the presence of a small amount of a liquid phase [2]. The now-agglomerated reaction product (5–20 mm) is called clinker. The clinker falls through a grate where it is cooled by incoming air. The clinker is then ground in ball mills containing steel balls and 3–5 wt % gypsum added to control premature setting. Grain size is such (ca. 20–50 mm) that surface areas typically are in the 2000–3000 cm2 g–1 range. The Portland cement is bagged as it is or blended with a variety of active and, more recently, inert fillers. The active additives are known as pozzolanic materials. They are typically “glassy” in nature and are reactive with the lime in the pore solution that forms during hydration. Typically Class F fly ash and granulated blast furnace slag are added to the ground clinker. Natural pozzolans such as diamateous earth, Trass from Germany, pozzolana from Italy, and manufactured products such as metakaolinite and condensed silica fume (an industrial byproduct essentially equivalent to silica fume produced by various manufacturers such as Cabot Corporation in the USA and Degussa in Europe) can also be added [1,2]. Recently, there has been a

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move afoot to base standards governing the manufacture of Portland cement on performance rather than on chemistry alone. For example, clinker can be ground extremely fine, and this makes it more reactive in the first few days and weeks. Standards require a minimum compressive strength at 28 d. These cements react to such an extent that they are often stronger than they need to be. Therefore, the manufacturer can add 5–10 wt % limestone as inert filler. Unfortunately for the material scientist these “alterations” go unnoticed because the manufacturer is not required to label such ingredients on their bags.

Producing Portland cement clinker is a true “ceramic” process. In fact Portland cement clinker is the largest volume ceramic produced in the world. In the simplest case ground Portland cement clinker contains four major phases. Listed in order of decreasing abundance they are: tricalcium silicate (Ca3SiO5, C3S, ca. 60 wt %), dicalcium silicate (Ca2SiO4, C2S, ca. 20 wt %), tricalcium aluminate (Ca3Al2O6, C3A, ca. 10 wt %), and tetracalcium aluminoferrite (Ca4Al2Fe2O10, C4AF, 10 wt %). Although C3A is the most reactive phase (and the reason gypsum is added to the clinker during grinding; see Eq. 3 below) it plays a rather minor role in the overall property development of Portland cement based materials. By far the most important and consequently the most studied of the four phases is C3S. When C3S is mixed with water it undergoes dissolution, the solution reaches saturation/supersaturation, and nucleation begins, followed by rapid growth of highly insoluble hydrates.

2.9.5.3

Hydration

Once Portland cement is mixed with water the resulting viscoelastic mixture undergoes hydration and will harden with time. The anhydrous grains undergo surface dissolution, the solution that forms becomes supersaturated with respect to calcium ions, which in turn causes hydrated phases to nucleate and grow. The calcium silicate hydrates that form are collectively known as calcium silicate hydrate (C–S–H). Once the paste hardens it is called hydrated cement paste. The strength of the hardened cement paste will continue to increase until all of the anhydrous starting materials are consumed. This process can take as long as a year or considerably longer if particles are large and water content is small. Portland cement, like many cements can be mixed and then placed, since the reaction is not immediate (induction period) as in the case of mixing two homogeneous solutions such as Na2SO4 and BaCl2 to instantaneously precipitate BaSO4. Epoxies and other cements exhibit this same characteristic. In the case of Portland cement, the first C–S–H that precipitates from solution does so within minutes of mixing. Although it does not have the necessary morphology or layer structure needed to join the grains together [20], it does coat the grains and cause the reaction to become diffusion-controlled; it is suggested that this change is the cause of the induction period. The induction period of cement slurry is normally 3–5 h. The end of the induction period is signaled by a vigorous increase in the rate of heat evolution of the sample. The reaction that occurs is highly exothermic and has a very distinct footprint on the output of a conduction calorimeter. Figure 10 shows an example of a conduction calorimetric output for hydration of C3S with water at 21 C for 24 h [21].

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Fig. 10 Heat flux of hydrating tricalcium silicate as a function of time, determined by using an isothermal conduction calorimeter. Stage I = surface wetting, State II = induction period, State III = acceleratory period, State IV = deaccelatory period, State V = normal diffusion-controlled reaction. After Kondo and Daimon (published by Young [21]).

Portland cement continues to hydrate for days, months, and years. Water is important to the setting and curing process, and more harm than good comes from allowing a concrete to dry out. When Portland cement is hydrated the reaction is relatively slow and the development of the hydrated phases required to convert a slurry into a solid is much more gradual than the hydration of Plaster of Paris (CaSO4·0.5 H2O), for example, which goes to completion in approximately 15 min. Portland cement will harden in approximately 8 h. The first 3–4 h are considered the induction period. After entering what is called the acceleratory period, it normally takes 4-5 h for cement paste to harden to the touch. During the acceleratory period reaction rates accelerate as a second C–S–H having a layer structure begins to form. Compared to the initial sorosilicate-like C–S–H that forms initially, the surface area of the second hydrate is large, often greater than 100 m2/g [17]. In addition, the C–S–H has a layer structure which contains Ca2+ ions, protons, and water molecules. Its morphology is often described as fibrillar to foil-like [16]. The developing hydrates intergrow and interlock, and when connectivities are great enough the hydrating grains become “glued” together, and thus cause the cement paste to harden. As larger quantities of C–S–H are formed water pockets become isolated from one another and the rate of reaction becomes diffusion-controlled. As the degree of joining increases so does the rigidity of the paste.

Hydration reactions for the four clinker phases in Portland cement are given below [5]:

In this case the excess water will come off during drying because the gypsum crystals are much larger and do not have interlayers that are capable of containing chemically bound interstitial water. The only water molecules present are those within the structure itself, hydrogen-bonded to the Ca2+ ions in the gypsum lattice. Even the largest grains have been converted to gypsum and just a few remnant hemihydrate cores normally remain [22].

The currently accepted model for the structure of C–S–H was proposed by Taylor [1]. It proposes that the C–S–H that forms has a rudimentary layer structure, similar to that found in a phyllosilicate (i.e., clay), but without the hexagonal silicate structure of a phyllosilicate (i.e., C–S–H does not have Q3 connectivity, as determined by 29Si MAS NMR). Rather it is proposed that the C–S–H consists of a central layer of octahedrally coordinated Ca2+ ions (CaO22– sheet) upon whose surface are bonded varying amounts of dimeric and longer polymeric silicate chains (dreierketten). Two such layers stacked next to each other would define an interlayer. The dreierketten have a repeat distance of three. The central silicate tetrahedron in a chain of three is the so-called bridging tetrahedron. It has two unpaired oxygen atoms that extend into the interlayer space. It is in this layer and in the vicinity of the oxygen ions that Ca2+ ions and protons can substitute for one another, leading to what can be considered to be a solid-solution series. The interlayer also contains varying amounts of molecular water. The Ca/Si ratio of the C–S–H will vary in response to the composition of the coexisting solution. Taylor proposed that the initial silicate units that attach themselves to the CaO2 layer are predominantly dimeric (Si2O76–), but with time the silicate dimers tend to polymerize to form increasingly longer chains. Grutzeck has published a model for the reaction of C3S and water which describes the earliest reactions in more detail [20]. It purports that there are two major C–S–H phases that form during the first 8–10 h of curing at ambient temperature. The first phase that forms has a sorosilicate-like structure with a Ca/Si ratio close to two that

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closely resembles a calcium-deficient jaffeite (tricalcium silicate hydrate Ca3Si2O7·3 Ca(OH)2 with loss of some Ca(OH)2) similar in structure to a dimeric jennite C–S–H described by Richardson [16]. It is proposed that the sorosilicate-like C–S–H develops as a result of an equilibration process that occurs during the induction period, and that the initial hydrate formation during the first few minutes after mixing is a highly localized event that occurs in the vicinity of the anhydrous grains as a consequence of rapid dissolution and buildup of critically supersaturated concentration gradients around each grain. The C–S–H that precipitates is a highly disorganized, thermodynamically unstable monomeric hydrate that coats the grains and causes the reaction to become diffusion-controlled (i.e., causes the induction period). The hydrate now controls the solubility of the C3S. It dissolves incongruently and thus begins to reorganize itself to form more stable dimers surrounded by double chains of octahedrally coordinated Ca atoms. The jaffeite structure contains molecular Ca(OH)2, which is mobile and could conceivably be replaced by water molecules. It is proposed that the C–S–H will continue to exist until the coexisting solution phase becomes deficient in Ca ions, causing it to decompose. It forms via a phase transformation of the sorosilicate as a consequence of radical changes in coexisting pore solution chemistry as a consequence of Ca(OH)2 precipitation. It is proposed that the jaffeite-like structure converts to a dimeric jennite-like structure, which has a layer structure that is similar in principle to that exhibited by a Ca-defi- cient tobermorite. Additional lime is lost by sorosilicate forming the jennite structure. With time jennite is joined by tobermorite-like C–S–H with gradually lengthening dreierkette as the system drives towards equilibrium. Although Grutzeck’s model is empirical [20], it can be used to visualize a mechanism responsible for much of the kinetic data reported in the literature.

The detailed chemistry of Portland cement is far too complex and as such well beyond the scope of this chapter, but at the same time the general reaction is easy enough to visualize. The matrix phase in a cellular concrete forms as a result of hydrate formation around the entrained gas bubbles. The matrix is not glassy nor is it theoretically dense as in traditional ceramics. It is porous because the developing foils and fibrils that form on the surfaces of the anhydrous grains intergrow and interlock, and with time the hydrates replace both the anhydrous phases and the water that held the anhydrous starting materials in suspension during mixing and placing [20]. For further information on the chemistry of cement, see Refs. [1, 2, 5].

2.9.6

Properties of Calcium Silicate Hydrate in Cellular Concretes

Unlike high-fired ceramics, cellular concretes have an extremely large degree of latitude in terms of the ingredients that can be used to make them. They can be produced from a variety of materials with a variety of foaming agents. Cellular concretes can be manufactured from Portland cement, but they can also be manufactured from Portland cement mixed with waste materials such as Class F fly ash and ground blast-furnace slag. These additives are waste products that often have to be

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land-filled. Because of their pozzolanic character they actually aid the long-term performance of the cellular concrete. However, one should be aware of possible deleterious chemical interactions that might occur before proceeding with full-scale use of a given cellular concrete. One should at least evaluate the proposed combination experimentally. An example of a worst case scenario follows. A contractor was using a cellular concrete fill material as a base material for an asphalt-covered parking lot in Texas. The fill specifications included Portland cement and Class F fly ash. The contractor noted that a nearby coal burning power plant was literally giving its fly ash away so he used it in formulating his mixture. The mixture was allowed to cure for a day or so and then was coated with asphalt. Everything went smoothly.

However, the contractor neither understood nor realized that the fly ash was the product of flue gas desulfurization, and in addition to fly ash it also contained a great deal of gypsum and hannebachite. After sitting in the Texas sun for a month or so the sulfate and sulfide reacted with the calcium aluminate and the calcium silicate hydrates in the fill material and additional water adsorbed from the ground to form the expansive minerals known as ettringite and/or thaumasite. The asphalt covering heaved and buckled and had to be replaced.

2.9.6.1

Cast-in-Place or Precast Cellular Concrete

The strength of a low-temperature cellular concrete will decrease as the amount of porosity (void space occupied by air bubbles and evaporable water) increases, as proposed in 1972 by Hoff [4], who also provided an equation for calculating strength as a function of theoretical porosity of preformed foamed cellular concretes made with Type III and Class G oil-well cements. More recent data taken from a Cellular Concrete LLC brochure on cellular concrete [10] was used to produce the graphs in Figs. 11 and 12. The trends as a function of mixture density are similar to those published by Hoff. Curve-fitting programs were used to fit the data points to a curve. The equations are given, as is the goodness of fit.

The trend of the strength versus dry density curve of the cured cellular concrete tends to follow the same trend as that of AAC (see Fig. 13). However, the fact that the strength curve tracks the lower limit of the AAC curve suggests that a properly formulated low-temperature cellular concrete is not as strong as an a similarly formulated autoclaved equivalent. Hoff suggests that 280 kg m–3 may be the practical lower limit for such products cured at ambient temperature. Lower densities lead to such low-strength products that they tend to break when handled. Thermal conductivity increases with increasing density (Fig. 12), as does the permeability to and ability to absorb water. Water adsorption is on the order of 14–20 wt % after long-term immersion. The layer of matrix between each bubble is quite impermeable because it consists of a relatively dense C–S–H microstructure. Matrix thickness between walls in the cellular concretes pictured in Fig. 1 are on the order of 0.1 mm. Adding coarse-grained quartz or limestone sand to the mixture increases its density without effecting its resistance to water-penetration and freeze/thaw resistance. Adding a similar amount of Class F fly ash will compromise the isolated cellular structure