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If a system at equilibrium is disturbed, then the system adjusts itself so as to minimise the disturbance. This is known as Le Chatelier s principle and is essential in determining how a change in concentration of a species, a change in total pressure of a gaseous reaction, or a change in temperature, will affect equilibrium.
Using the concentration levels of equation m as an example, an increase in concentration of either Fe3+ ions or SCN- ions will result in the equilibrium position moving to the right , using up the additional reactants to produce more FeSCN2+. A decrease in the concentration of FeSCN2+ will yield the same result. A decrease in the concentration of a reactant will result in the opposite, with the equilibrium position moving to the left , compensating for the lack of reactants. Again, an increase in FeSCN2+ will do the same.
When an ionic substance dissolves into water, it breaks up into ions which move independently throughout the solution. When a solution becomes saturated, ions still continue to break away from the crystals of solid and go into solution, but in addition an equal number of ion pairs from the solution precipitate out on to the solid.
This can be proven by adding radioactive solid to a saturated solution. Using lead nitrate as an example, when radioactive solid is added to the solution, initially there is no radioactivity within the solution. However, after some time, it can be observed that some radioactivity has transferred into the solution without changing the concentration of lead ions. Eventually the same fraction of radioactive ions will be present in the solution as in the solid, demonstrating the dynamic nature of the precipitation-dissolution equilibrium. There is a dynamic balance between dissolution and precipitation with both occurring at equal rates and thus inferring no overall change to the concentration of the solution.
Dissociation of chlorine solution into hypochlorite ion (and its subsequent hydrolysis into hypochlorous acid:
Pool chlorine is the chemical most often used to keep swimming pools free of micro-organisms. Although called chlorine , it actually takes the form of NaOCl if liquid or Ca(OCl)2 if solid. Both these compounds dissociate to form OCl- ions which then hydrolyse to create hypochlorous acid (HOCl) in an equilibrium reaction. HOCl and OCl- are referred to as free chlorine. Each species is able to kill micro-organisms by attacking cell walls and destroying internal cell structures through oxidation. Hypochlorous acid, however, is able to oxidise biological compounds in several seconds while a hypochlorite ion can take anywhere up to thirty minutes, largely due to its negative charge which retards cell entry.
HOCl and OCl- will eventually combine with ammonia-containing compounds to make chloramines such as NH2Cl and NHCl2 or be broken down by sunlight (see equation below), rendering them ineffective as a means of micro-organism elimination. Pools deficient in chlorine often experience high sunlight exposure or are frequently used, with humans introducing foreign residues and oils that contain ammonia.
Low chlorine levels can be compromised by adding NaOCl and keeping organic matter out of the pool. pH control is also necessary as if too basic, hypochlorous acid concentrations will drop. Table 5.3 pg 11. If chlorine levels have been left too low for too long and the pool begins to smell and irritate the skin and eyes, too much ammonia-based compound has formed and break point chlorination (BPC) will be required. Break point chlorination is the process of adding enough chlorine into the pool so as to facilitate the breakdown of chloramines into gas.
In order to achieve this, NaOCl must be added to the pool at a concentration level ten times the amount of the pool s current chloramine concentration (CC) in the formula; BPC=10CC. CC is calculated using a colorimetric DPD indicator test, similar to that of pH.
If chlorine levels are too high due to a surplus introduction of product, the concentration will eventually drop due to photolytic loss through sunlight however sodium thiosulfate or sodium sulphite can be added to facilitate a reduction.
The buffering system formed through the dissociation of sodium bicarbonate, and the three resultant hydrogen carbonate ion-based equilibria:
pH refers to the acidity or alkalinity of an aqueous solution and its management is imperative in the maintenance of pool water quality. If the water contained within a pool is too acidic, metal fittings and pump impellers corrode, plaster walls are etched, dissolved metals stain walls and eye irritation occurs. If the water contained within a pool is too basic; filters become overworked, water becomes cloudy, scale formation increases, pool walls are discoloured and chlorine activity is slowed allowing algae and micro-organisms to grow. Sweat, body oils, urine, faeces, vegetation and rain all affect the pH of a pool and as such its acidity and alkalinity are constantly changing. Fortunately, chemical systems can be imposed that are able to buffer pH change, and the most commonly used compound to create such a system is sodium bicarbonate.
When sodium bicarbonate dissociates into its ions a series of simultaneous equilibria are created, based on the hydrogen carbonate ion:
These three equilibria can be combined as demonstrated below:
Complying to Le Chatelier s Principle, if the concentration of hydrogen ions in solution increases (pH lowers) the reactions in equation ladida will move to the right, using up the excess H+ ions and raising the pH again. Similarly, if the concentration of H+ ions decreases, the reactions will move to the left, making more H+ ions and lowering the pH again.
Once all of the CO32- and HCO3- ions are used up, the pool cannot absorb any more hydrogen ions and the buffer slowly dissipates as gaseous CO2. In order to determine when the buffer needs to be replenished, colorimetric tests are conducted by adding indicator to pool water samples where the colour change corresponds to the amount of extra sodium bicarbonate that is required to be added to the pool water for its pH to return to the ideal range.
Equilibrium constants and pH:
For every equilibrium reaction there is a function of the concentrations of the species involved that has a constant value at equilibrium.
Equation x represents any equilibrium reaction where A and B are the concentration of reactants, C and D are the concentration of products, and a, b, c and d represents the number of each species involved. The equilibrium constant, K, of this reaction is represented by the expression;
The expression for acid dissociation constant is similar:
Using the equilibrium constants of the equilibria relevant to the carbonic acid-bicarbonate buffer system, the pH of the system can be determined. For calculation purposes, Equation 5.18 is ignored as it takes a relatively minor part in the buffering system.
These expressions can be rearranged to determine the equilibrium concentration of the hydrogen ion:
pH is equal to the negative log of hydrogen ion concentration therefore:
pH = -log[H+]
*where pK =
Equilibrium constant data tables determine that the value for pK in regards to Equation 5.16 and 5.17 is approximately 6.3 at 25 C therefore:
Dissolving and precipitation of calcium carbonate:
Concrete pools are often surfaced with a mixture of calcium carbonate and cement. Calcium carbonate is slightly soluble in water and its dissociation into ions can cause problems for pool maintenance. Le Chatelier s Principle dictates that as CaCO3 is a solid with constant concentration, it will have no effect on the equilibrium position and thus the amount of calcium carbonate that dissolves is dependent on the amounts of Ca2+ and/or CO32- already dissolved in the water. If only a small amount of Ca2+ and/or CO32- is dissolved, the reaction moves to the right causing the pool s lining to dissolve, while if too much Ca2+ and/or CO32- is present in the water, calcium carbonate will precipitate. This precipitation is known as scale and can be detrimental to the pool s piping systems.
Other equilibrium reactions affecting the precipitation of calcium carbonate include its reaction with hydrogen carbonate (formed as a result of dissolved carbon dioxide see Equation 5.28) and the reaction involving carbonate ions and hydrogen ions.
If hydrogen ion concentration were to increase in equation ii, thus lowering pH, the reaction would move to the right, using more carbonate ion and thus moving Equation z(caco3 to ca + co3) to the right dissolving more calcium carbonate. Equation i is exothermic therefore as temperature increases the reaction moves to the left to counteract the change resulting in calcium carbonate precipitation. It is a result of this that scale forms in hot water systems.
The concentration of Ca2+ ions can also be measured by forming complexes between the ions and a substance known as EDTA. In a basic solution of pH greater than 10, EDTA loses its H+ ions leaving the four oxygen ions with an unbonded pair of electrons able to form coordinate covalent bonds with Ca2+ ions. If EDTA is combined with a sample of pool water to complex the Ca2+ ions, the indicator Eriochrome Black T is used to determine when the ions have been fully complexed by changing the colour of the solution from red to blue. The number of moles of EDTA used is equivalent to the total number of moles of Ca2+ ions in the portion of the total amount of solution. The equilibrium constant for the Ca-EDTA complex is over 1010 with such a magnitude indicating that the reaction complexes all ions and is thus a reliable means of determining Ca2+ concentration.
The ideal range for Ca2+ ion concentration is between 200 and 400 parts per million (ppm) or milligrams per litre. If concentration is higher than this, the only way to reduce the concentration is by partially draining and refilling the pool water. If Ca2+ ion concentration is too low, calcium hypochlorite can be added which will also raise pH.
Solid ionic compounds in contact with saturated solutions of the same species share dynamic equilibrium. If these solids possess low solubility a quantitative relationship can be observed between the concentrations of the ions in the solution at equilibrium (see equations below).
This relationship between products is exactly the same as the expression for equilibrium constant discussed previously but with no concentration value applied to the solid. No value is applied in accordance with Le Chatelier s principle that increasing the amount of pure solid present in an equilibrium mixture does not alter the position of equilibrium as only the amount of solid is being changed, not the concentration. The equilibrium constant for reactions involving solids of low solubility is called the solubility product, Ksp, and is represented by the expression:
There is always a dynamic equilibrium between an ionic solid and its saturated solution however Equation abxy only applies to ionic solids of low solubilities (< 1mol/L).
In terms of pool maintenance, solubility product is specifically relevant to calcium carbonate precipitation. In pure water of pH 7, CaCO3 is poorly soluble however as discussed previously, the concentration of Ca2+ and CO32- ions affects the equilibrium position of the compounds dissociation reaction causing it to precipitate or dissolve.
= Ksp = 4.47x10-9
Knowledge of the pH of the pool can allow the concentration for both ion concentrations to be calculated through the expression:
Where Kh= 1.70x10-3 and represents the equilibrium constant for ,
Ka1= 2.5x10-4 and represents ,
Ka2= 5.61x10-11 and represents , and
kH= 29.76 mol/L and represents
In this expression, represents the partial pressure of CO2 and is equal to 3.5x10-4 atmospheres in ambient air while the H+ ion concentration is required to be calculated through the formula;
Once all these values have been substituted into the equation, the concentration of Ca2+ ions can be determined. The following example will calculate Ca2+ ion concentration for pool water at an optimum pH level of 7.4:
CO32- ion concentration can then also be calculated using the solubility product:
*TDS no solution other than pool draining. Conductivity.
Stabiliser takes years to go over 100 ppm apparently and is tied in with TDS, draining is only soln.