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pH = -log10[H+] Equation 1
Where [H+] is the concentration of H+ ions in moles per liter. Because H+ ions associate with water molecules to form hydronium (HO+) ions, pH also is often expressed in terms of the concentration of hydronium ions. If acid is added to water, however, an excess of HO+ ions is formed H+ (acid) + HO (water) yields HO+ (hydronium ions). When the concentration of HO+ exceeds the concentration of OH-, the solution becomes acidic. In an acidic solution, the concentration of hydronium (HO+) ions can range from 1 to 1 x 10-6 moles/liter, depending on the strength and amount of the acid. Therefore, acid solutions have a pH ranging from 6 (weak acid) to 0 (strong acid). Inversely, when the concentration of OH- exceeds the concentration of HO+, the solution becomes basic. In a basic solution, the concentration of hydroxyl (OH-) ions can range from 1 to 1 x 10-6 moles/liter. Therefore, basic solutions can have a pH ranging from 8 (weak base) to 14 (strong base).
Pure water ionizes slightly into hydronium and hydroxide ions. It has been found by experiment that one liter of pure water contains only one ten millionth of a mole of hydronium ions and one ten millionth of a mole of hydroxide ions. A substance which contains more hydronium ions than hydroxide ions is acidic; a substance which contains more hydroxide ions than hydronium ions is basic; and a substance such as pure water which contains an equal number of hydronium ions and hydroxide ions is neutral. The dissociation constant, K, changes with temperature, and this must be taken into account in interpreting data involving H+ and OH- ions.
For example, many water treatment operations are carried out at high temperatures, and samples from the system are usually cooled prior to analysis. The H+ and OH- concentrations measured on the cooled sample, even though different from those in the hot system, are usually used for control purposes.
The pH of a solution can be measured by titration, which consists of the neutralization of the acid (or base) by a measured quantity of base (or acid) of known concentration, in the presence of an indicator (the color depends on the pH). The pH of a solution can also be determined directly by measuring the electric potential arising at special electrodes immersed in the solution.
In chemistry, a buffer solution is a substance that inhibits the solutions change in pH, or hydrogen-ion concentration. Such substances consist either of the pairing of a weak acid and a related salt of the acid, or of the similar pairing of a weak base and a salt of the base. Fluids in living organisms are strongly buffered, and sea water and soil substances are other examples found in nature. Buffers are used in chemistry and serve as references in the measurement of pH.
An indicator is a natural or synthetic substance that changes color in response to the nature of its chemical environment. Indicators are used to provide information about the degree of acidity of a substance pH or the state of some chemical reaction within a solution being tested or analyzed. One of the oldest indicators is litmus, a vegetable dye that turns red in acid solutions and blue in basic ones. Other indicators include alizarin, methyl red, bromocresol and phenolphthalein, each one being useful for a particular range of acidity or a certain type of chemical reaction.
Alkalinity is the capacity of water to accept protons. It is the sum of all the bases present. This is call total alkalinity.
Total alkalinity is easily measured from an alkalinity titration from the known amount of acid added to the solution. Total alkalinity is a conserved quantity. It is not affect by pressure or temperature changes. Although the individual species in the total alkalinity may be affected by temperature and pressure changes, the total alkalinity is not. In most natural waters, the concentration of borate, silic acid, bisulfide, organic acids, and the hydrogen and hydroxyl ions are small compared to the bicarbonate and carbonate ions. In these situations the total alkalinity is approximately equal to the carbonate alkalinity. This facilitate find some of the terms in the equation by equilibrium expressions as follow.
If the alkalinity is low, it indicates that even a small amount of acid can cause a large change in our pH. Consider the pond owner whose pH was 8.0. He was told that 7.0 were better so he puts in chemicals to lower it. The next day, it is back to 8.0 so he adds more chemicals. The following day it tests at 7.5. He feels good because it is finally starting to come down and dumps in some more stuff. All of a sudden he finds that the pH is 5.0. His bio-converter bacteria were destroyed and his fish are dying of ammonia poisoning compounded with pH shock. Each treatment kept reducing the alkalinity until it was so low that the final addition caused a major pH transition.
Alkalinity is related to the amount of dissolved Calcium, Magnesium, and other compounds in the water and as such, alkalinity tends to be higher in harder water. Lime leaching out of concrete ponds is a primary source of alkalinity but it is also slowly increased by evaporation which concentrates the source compounds. Alkalinity is naturally decreased over time through bacterial action which produces acidic compounds that combine with and reduce the alkalinity components.
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