Acids cause water hardness
lecture to the Liebig laboratory
The accompanying lecture on the Liebig Laboratory is intended to convey the necessary theoretical fundamentals that enable the participants of the course to carry out all laboratory work provided in the program safely and with specialist knowledge. The main focus is on imparting skills and abilities in the analysis of water samples (lime cycle), polyfunctional molecules such as amino acids (acid-base behavior and function as a chelating ligand) and redox processes (bleaching and disinfecting effects of strong oxidizing agents such as peroxides). Here, too, the imparting of important basic knowledge in chemical measurement analysis is a focus.
The accompanying lecture on the Liebig Laboratory will be given by PD Dr. Böttcher held from November 7th to 17th, 2011 (as shown in the timetable). The final exam for the Liebig Laboratory will take place on February 1, 2012 from 1 p.m. to 2.30 p.m. in the Liebig and Buchner lecture halls. For the writing of the exam it is necessary that you bring a (non-programmable) pocket calculator with you. All other materials will be made available to you.
We recommend the work as an accompanying textbook C. E. Mortimer, U. Müller: Chemistry. 9th edition, Thieme 2007, ISBN 978-3-13-484309-5. The "Mortimer" is also available electronically within the Munich university network: "E-Book"
The page numbers in this script refer to this edition.
The lime cycle
The learning objectives of this teaching unit include the following focal points: We discuss the circumstances that lead to so-called water hardness and learn how the main components of the hardness builders (calcium and magnesium ions) get into groundwater. Complexometric titration is dealt with as the standard method of classical water analysis. Furthermore, the term water hardness is defined, which is made up of the components "carbonate hardness" (CH, temporary hardness) and "non-carbonate hardness" (NCH, permanent hardness, including "sulfate hardness"). You will learn how the total hardness of industrial or drinking water can be determined. An important parameter is the total alkalinity of the water, which in principle can be determined by an acid-base titration.
Literature: Mortimer (2007) pp. 486-490, pp. 293-317, pp. 505-512.
General information on the lime cycle
The water hardness is mainly caused by the content of calcium and magnesium ions in water. Calcium and magnesium ions are released from rock materials in mountains (e.g. "dolomite", mixed compound of equal proportions of calcium and magnesium carbonate) by rainwater. This dissolution process is initially based on the actual solubility of these compounds in water (solubility product). The solubility is improved, however, because the rainwater has acidic components ("carbonic acid"), which are based on the one hand on the natural content of carbon dioxide in the air, but also on the basis of an increased degree of environmental pollution (including sulfur dioxide or similar acidic gases) . Complexometry is a suitable method for detecting metal ions in water.
Basics of complexometry
In this chapter you will get to know the basic principle of the method. It is a volumetric determination, whereby in principle the formation of very stable chelate complexes is used. The titrator here is a hexagonal chelating ligand: ethylenediaminetetraacetic acid. A great advantage of the method is that 1: 1 complexes are always formed, i.e. the metal ion to be determined and the ligand always react in a molar ratio of 1: 1. The main body of the ligand used is derived from ethylenediamine (1,2-diaminoethane) in that the four H atoms are each substituted by an acetic acid residue. In practice, "Titriplex III", the disodium salt of ethylenediaminetetraacetic acid, is often used as a standard solution. The equivalence point is indicated with so-called metal indicators. In principle, they represent organic dyes that are capable of forming a (stable) complex with the metal ion to be determined. During this complex formation, the indicator dye loses its own color. The ions of the metal to be determined, however, form much more stable complexes with edta, so at the equivalence point of the titration there is a displacement reaction of the indicator from the complex, so that the indicator loses its function as a ligand and is released into the solution in its original form. It takes on its own color again and there is a specific color change. Two important things are to be observed with complexometric titrations: since protons are released during complex formation from the titrator, the solution is in principle more "acidic". This can cause problems, as many complexes are only stable in certain pH value ranges. In this respect, it is often essential to buffer the solution. The metal indicators also require specific pH ranges in which the specific color change takes place. In practice, indicator buffer tablets are expediently used, i.e. these are the buffer components in combination with an indicator dye trituration. Nevertheless, you should not just "blindly" rely on the buffer tablet, but always check the pH value regularly during the titration!
Complexometry: different methods
Direct titration is often used, i.e. the solution of the metal ion is titrated directly with the Titriplex III solution against a metal indicator. The calcium determination (v 1.5) should be mentioned in this context. You will also get to know the method of back titration (e.g. determination of Al); it is based on the lower stability of the Mg or Zn complexes (with edta) compared to many other cations. The principle of back titration is as follows: a known volume of Titriplex III is added to the analysis sample in excess. The entire amount of the cation to be determined is bound. Then the rest of the unused standard solution is "back-titrated" with e.g. magnesium sulphate (or zinc sulphate) standard solution. The method is based on the different complex stability: lgβ(Al complex) = 16.70 vs. Ig3 (Mg complex) = 8.69. If this serious difference did not exist, a displacement reaction of the aluminum from its edta complex would take place.
Complexometry: very sensitive detection method!
In order to achieve accurate analysis results with complexometric titrations, it is imperative that you use extremely clean glass equipment. You should clean all glass appliances well with washing-up liquid and a brush and then rinse them several times with plenty of deionized (!) Water. Rinsing with tap water only would drag considerable amounts of foreign ions (Ca and Mg!) Into the vessels.
Complexometry: Principle of endpoint recognition
At the beginning of the titration, a suitable metal indicator dye is added (e.g. Eriochrome black T or Xylenol orange). This forms a (stable) complex with the cation to be determined, with the indicator's own color also changing. Finally, at the equivalence point, there is a displacement reaction of the indicator dye from the initially formed M-indicator complex, since the complex of the metal ion to be determined has a higher complex stability (competitive reaction around the central atom). When the indicator dye is released, its own color becomes visible again (equivalence point).
Determination of calcium
General: You work strictly according to the regulations of the internship script; check the pH value at the beginning and regularly during the titration! Work in very clean glass vessels! The starting point for calculating the calcium content in the sample is - as already noted - the formation of 1: 1 complexes, so the following applies for the calculation: 1 mL 0.1 m Titriplex III solution = 4.008 mg approx. For other metal ions (e.g. Fe) would apply analogously: 0.1 m Titriplex III solution = 5.585 mg Fe (or when using a 0.01 m standard solution corresponding to 0.5585 mg Fe, whereby the use of 0.01 m standard solutions is in principle more precise!).
Water analysis - water hardness
Where do the Ca and Mg ions in groundwater or service water come from? As mentioned above, for example, Ca and Mg ions (bound in the form of their carbonates) are released from the rock mineral dolomite by rainwater. On the one hand, the solubility of these salts in water is responsible for this: L (magnesium carbonate) = 2.6 × 10−5; L (calcium carbonate) = 4.7 x 10−9 minor. The solubility is improved by the "acid" rain (increased carbon dioxide or sulfur dioxide content in the air atmosphere), since the carbonates of all metal ions in the PSE generally dissolve in dilute mineral acids (including weak acids!). The content of calcium hydrogen carbonate determines the so-called "carbonate hardness" (CH, temporary hardness; is greatly reduced when heated, e.g. by the formation of scale, ie a considerable amount of calcium carbonate precipitates when the water is boiled. The content of other salts (except carbonates and hydrogen carbonates ) In water, such as sulfates or chlorides, there is no reaction when the water is heated (no precipitation of poorly soluble salts!) In this respect, this is the so-called part of the permanent (permanent) hardness, or we are talking about the proportion of non-carbonate hardness ( NCH). Completely desalinated water is obtained, for example, through ion exchange processes. These processes are only briefly referred to (see also in the old script for the lecture: Ion exchange processes). By boiling water, only partially desalinated water is obtained, if the precipitated calcium carbonate is removed by filtration. The total hardness of a water settles so as the sum of CH + NCH together. By definition, water hardness is given in so-called degrees of hardness and only relates to the content of calcium ions dissolved in the water. Please note, however, that some of this is also in the form of magnesium ions. A special complexometric separation (simultaneous determination of Ca in addition to Mg) allows the content of Ca or Mg to be precisely determined. However, you do not carry out this provision during the internship! The water hardness is given in ° dH ("degree of German hardness"): 1 ° dH = 10 mg CaO / L, for Ca means 1 ° dH = 7.15 mg Ca / L, for Mg means 1 ° dH = 4.34 mg Mg / L . For example, a mineral water (Gerolsteiner) contains an average of 348 mg Ca / L and 108 mg Mg / L. For both ions this would correspond to a total hardness of 73.55 ° dH. For experiment 1.6 (determining the total hardness of a given water), first prepare a 0.01 m Titriplex III solution (by diluting a commercially available standard solution, 0.1 m) and determine the factor of this standard solution. Exactly dried and precisely weighed calcium carbonate is used as the original titer; by dissolving it in hydrochloric acid, it provides you with a standard titer solution that is used to set the edta standard solution. With this standard solution you then determine the Ca content in accordance with the regulation for V 1.7 (the total contains Mg) and convert it to the corresponding degree of hardness.
Determination of the total alkalinity of drinking water
The total alkalinity of a water is understood to mean all components of a water that bind acid added above the transition point of methyl orange (pH 3.1–4.4) (acid-base titration with a 0.1 M hydrochloric acid standard solution). The determination of the alkalinity (corresponds to the acid consumption) is more exact than the specification of the so-called carbonate hardness.
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