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    Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling
    Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling
    Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling
    Ebook83 pages46 minutes

    Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling

    By R. H. Hopkins

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    About this ebook

    This text contains a detailed guide to the biochemical aspects of brewing beer, including a wealth of detailed information on subjects such as mashing, boiling, and cooling. The perfect book for serious brewers with an interest in the scientific side of the process, this antique text is timeless the value of its information and constitutes a great addition to collections of brewing literature. The sections of this book include: The Mashing Liquor, The Mashing Process, Influence of Temperature on the Quantity and Quality of Extract, Influence of pH on the Quantity and Quality of Extract, The Various Mashing Systems, Boiling with Hops, The Bitter Acids and Resins of Hops, The Precipitation of Proteins during Wort Boiling, Influence of Various Conditions on Wort Boiling, et cetera. We are proud to republish this antiquarian text now complete with a new introduction on brewing beer.
    LanguageEnglish
    PublisherRead Books Ltd.
    Release dateJul 7, 2014
    ISBN9781473394216
    Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling

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      Biochemistry Applied to the Brewing Processes - Mashing, Boiling, Cooling - R. H. Hopkins

      THE BREWING PROCESSES, MASHING, BOILING, COOLING

      The Mashing Liquor

      It is a well-known fact that the composition of the liquor has a pronounced influence on the character of the wort and the beer, and also that this influence is associated with the hardness of the water. The latter influences the pH of the mash, with the result that the enzymic processes are either delayed or accelerated, whilst the salt ions themselves may and almost certainly do exert minor independent influences on the extraction of matters from the malt and on the enzyme actions during mashing. Further, protein coagulation during boiling is influenced by pH and probably also, but in less degree, by the salt ions independently.

      The hardness of water consists of (1) temporary hardness caused by the calcium and magnesium bicarbonates, and so called because it may be eliminated on aeration or boiling (in fact, some magnesium and a trace of calcium carbonate are not eliminated on boiling), and (2) permanent hardness, caused by all other salts of the same metals.

      (1) Temporary hardness.—The bicarbonates are dissociated as follows:

      Ca(HCO3)2 → Ca++ + 2HCO3–

      The bicarbonate ion will, as explained in section 5, react with water, forming undissociated carbonic acid:

      HCO3– + H2O → H2CO3 + OH–

      Therefore, in spite of the presence of free carbonic acid, the water will have an alkaline reaction, the pH of such natural waters being 7·2–7·5.

      This value of pH will be found alike in very hard and rather soft waters, the mixture of bicarbonate ions and carbon dioxide having a pronounced buffer action. Only if one of these substances is completely eliminated may considerable change in the pH occur, and this is exactly the reason for the characteristic effect of a hard brewing liquor on the enzymic processes.

      From the titration curve of a moderately hard water as it appears on Fig. 1 (page 22) it will be seen that a change of pH from that of the water to the optimal pH of a mash, which will be about 5·3, will require the addition of more than 5 c.c. N/10 (strong) acid per 100 c.c. The weak acids of the malt are quite incapable of effecting this change, giving as they do a pH of 5·6–5·8 with distilled water (with hard water a pH amounting to 6·3 may be attained). These differences, which may seem small, are nevertheless of great significance, which will be dealt with in detail in sections 39 and 45.

      (2) Permanent hardness.—The permanent hardness, mainly caused by gypsum, which in certain types of water may be present in large amounts, influences the reaction of the mash in an indirect way, the Ca ions reacting with the secondary phosphates of the mash. The inorganic phosphates present in malt or liberated during mashing, if regarded as potassium salts, will occur as KH2PO4 and K2HPO4, since the pH of the mash is about 5·5. The phosphate ions present will include H2PO4–, HPO4– –, and PO4– – –, although the last named ion will be present in only relatively minute proportions, and the equilibrium

      will be set up, the proportions depending on pH. When calcium salts are heated with inorganic phosphates at pH 5·0–6·5, the insoluble phosphate precipitated is Ca3(PO4)2. This involves the removal of PO4––– ions from the system, which is followed by their replacement according to the laws of chemical equilibria and mass action, i.e. the above reactions proceed from left to right. This involves the splitting off of hydrogen ions as indicated in the reactions, and a fall in pH is observed. The whole reaction may conveniently be expressed,

      3CaSO4 + 4K2HPO4 → Ca3(PO4)2 + 2KH2PO4 + 3K2SO4

      All calcium salts will function in this way, but magnesium salts are less effective owing to the appreciable solubility of the magnesium phosphate formed. The effect of permanent hardness is seen to be the reverse of that of temporary hardness.

      For the brewing of light lager beers of the Pilsener type a soft water is necessary, and if the natural water at disposal is hard it must be softened artificially. This is done in practice either by boiling or by aeration (seldom), when the following reaction takes place:

      or by addition of lime water:

      To accelerate the formation of a crystalline precipitate of CaCO3 the treated water passes a so-called catalysator, which is a container with coarsely powdered marble to act as inoculation crystals.

      If the last reaction is written as an ionic reaction it leaves but, as mentioned in the beginning of the section, a considerable quantity of HCO3– ions may be present in equilibrium with free carbonic acid at as high a pH as 7·5 and even more. To remove the HCO3– ions effectively the concentration of OH– ions must therefore be rather high, corresponding to a pH of about 9. Thus it will be seen that water softened in this way is much more alkaline than the original hard water. The fact that such alkaline water nevertheless gives a more acid mash is easily understood when the low buffer content of the softened water is taken into consideration, even a small quantity of weak acid is capable of changing the pH from 9 to about 6.

      In general, the value of pH of a water, by itself, is of little significance, the amount and nature of the dissolved salts are important.

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