Carbon dioxide equilibrium with water

Buffers are extremely important in biological systems. The pH of arterial blood is about 7.4. The pH of the blood in your veins is just slightly less. If the pH of hlood drops to 7.0, or rises above 7.5, life-threatening problems develop. To maintain its pH within a narrow range, blood contains a number of buffer systems. The most important buffer system in the blood depends on an equilibrium between hydrogen carbonate ions and carbonate ions. Dissolved carbon dioxide reacts with water to form hydrogen carbonate ions. 

The CO2 in the blood also participates in the carbonic acid-bicarbonate buffer equilibrium. Carbon dioxide reacts with water in the blood to form carbonic acid ... 

One very important buffer solution is human blood An equilibrium between carbonic acid (H2CO3) and its conjugate base bicarbonate (HCOsi helps blood to maintain a relatively constant pH of around 7.4. The carbonic acid buffer system is created by carbon dioxide (CO2) dissolved in blood carbon dioxide reacts with water (H2O) to form carbonic acid. Since the amount of carbon dioxide in the blood depends on the rate at which you breathe, your blood pH is influenced by your breathing rate. Your body can... 

Cement mortar will be attacked by waters that have an excess of free carbon dioxide compared with that of waters that are in a lime-carbonic acid equilibrium. There is a two-step mechanism with a carbonization process according to... 

Water often is a reagent in an aqueous equilibrium. For example, when carbon dioxide dissolves in water, it reacts with a water molecule to form carbonic acid ... 

The melting point of carbon dioxide increases with increasing pressure, since the solid-liquid equilibrium line on its phase diagram slopes up and to the right. If the pressure on a sample of liquid carbon dioxide is increased at constant temperature, causing the molecules to get closer together, the liquid will solidify. This indicates that solid carbon dioxide has a higher density than the liquid phase. This is true for most substances. The notable exception is water. 

Once cooled, the effluent from the process separates into a liquid-water phase and a gaseous phase, the latter containing mainly carbon dioxide along with oxygen which was in excess of the stoichiometric requirements, and nitrogen. The separation is carried out in multiple stages in order to minimize erosion of valves, as well as to maximise the separation due to phase-equilibrium constraints. 

Carbonic acid (H2CO3) is formed reversibly whenever carbon dioxide dissolves in water. All carbonated beverages contain carbonic acid in equilibrium with C02 and water. 

Carbonate iron sediments are formed when iron is precipitated in the presence of dissolved carbonic acid or as a result of interaction of the primary sediments with organic matter in the course of diagenesis. The determination of the stability of primary iron carbonates with respect to iron oxides and hydroxides was made on the basis of using the functional dependence of and on pH, calculated by the method of Garrels and Christ (1968) for the system carbonate-water. According to this dependence, in conditions of FeCOj in equilibrium with water the value of partial pressure of carbon dioxide decreases in proportion to increasing pH. 

Carbon dioxide dissolved in water leads to the formation of carbonic acid and a consequent increase in H. Ponnamperuma (1967) has calculated that water at 25°C, in equilibrium with the normal concentration of CO2 in the earths s atmosphere (0.03% by volume), will attain a pH of 5.63. The weathering action of this weak acid over geologic time is well known to geologists (Krauskopf, 1967). Ponnamperuma s calculations also indicate that increased atmospheric CO2 concentrations will result in further decreases in pH, down to pH 3.97 with one amosphere of CO2. Respiratory CO2 concentrations in soil atmospheres can be 10 to 100 times greater than the normal 0.03% in the earth s atmosphere (Stotsky, 1972). Thus, pH values considerably lower than 5.63 can be achieved through respiration. Similarly, respiratory activity in shallow waters and tidal flats, especially at night when photosynthetic CO2 assimilation is halted, can cause a marked decrease in pH (Oppenheimer and Master, 1965). 

Some of the limits mentioned above can be relaxed by carefully using various SOLMINEQ.88 options. For example, if the formation temperature is not known, the calculated value from one of the chemical geothermometers can be used (21). If an element has not been analysed for, an approximate value can be obtained by assuming equilibrium with the appropriate formation mineral using the mineral dissolution and precipitation option. If carbon dioxide was known to have been lost between the time of sampling and analysis but the pH was measured in the lab and at the time of sampling, the amount of carbon dioxide in the water at the time of sampling can be recalculated. Other techniques and options in SOLMINEQ.88 can be used, but each assumption of this type should reduce confidence in the results. 

When carbon dioxide dissolves in water, it undergoes a multistep equilibrium process, with A overaii = 4.5X10 , which is simplified to the following ... 

Carbonated water. CO2+ 2H2O H30 + HCO3. The carbon dioxide in the atmosphere is sufficient to bring water that is otherwise pure to a pH of 5.6. In surface soils, respiration creates a higher CO2 level, that in equilibrium with water creates a pH < 5.0. 

Carbonic acid (H2CO3) is formed when carbon dioxide dissolves in water. Although carbonic acid itself is always in equilibrium with carbon dioxide and water, it has several important stable derivatives. Carbonate esters are diesters of carbonic acid, with two alkoxy groups replacing the hydroxyl groups of carbonic acid. Ureas are diamides of carbonic acid, with two nitrogen atoms bonded to the carbonyl group. The unsubstituted urea, simply called urem, is the waste product excreted by mammals from the metabolism of excess protein. Carbamate esters (urethanes) are the stable esters of the unstable carbamic acid, the monoamide of carbonic acid. 

Carbon dioxide CO2 gas is a natural component of natural waters. It is also generated by biological oxidation of organic matter, especially in polluted water. Carbon dioxide is very soluble in water above pH 5 and forms bicarbonate and carbonate buffers by exposure to atmosphere. In the case of surfece waters it can enter water from the atmosphere when its concentration in water is less than that in equilibrium with CO2 in the atmosphere in accordance with Henry s law. Free CO2 is dissolved carbon dioxide gas in water. 

Determine the Henry constant for carbon dioxide (1) in water (2) with the help of the following phase equilibrium data at 50 "C ... 

The revision reduces the dew-point temperature predicted by the Parrish et al. conelation. As an example, the equilibrium water dew point of a methane gas containing 10 mole percent carbon dioxide contacted with 99.9 wt% TEG at 100°F and 1,500 psia is estimated to be -33°F based on the Parrish correlation, Figure 11-15. The revised dew point, based on equation 11-1 and Figure 11-16 is about 8°F, which is 15 degrees lower than the original prediction. If the gas consists primarily of methane, a smaller cotrection is indicated (about six degrees for these conditions). 

The most important property of the dissolved solids in fresh waters is whether or not they are such as to lead to the deposition of a protective film on the steel that will impede rusting. This is determined mainly by the amount of carbon dioxide dissolved in the water, so that the equilibrium between calcium carbonate, calcium bicarbonate and carbon dioxide, which has been studied by Tillmans and Heublein and others, is of fundamental significance. Since hard waters are more likely to deposit a protective calcareous scale than soft waters, they tend as a class to be less aggressive than these indeed, soft waters can often be rendered less corrosive by the simple expedient of treating them with lime (Section 2.3). 

The effect of pH on the corrosion of zinc has already been mentioned (p. 4.170). In the range of pH values from 5 -5 to 12, zinc is quite stable, and since most natural waters come within this range little difficulty is encountered in respect of pH. The pH does, however, affect the scale-forming properties of hard water (see Section 2.3 for a discussion of the Langelier index). If the pH is below the value at which the water is in equilibrium with calcium carbonate, the calcium carbonate will tend to dissolve rather than form a scale. The same effect is produced in the presence of considerable amounts of carbon dioxide, which also favours the dissolution of calcium carbonate. In addition, it is important to note that small amounts of metallic impurities (particularly copper) in the water can cause quite severe corrosion, and as little as 0-05 p.p.m. of copper in a domestic water system can be a source of considerable trouble with galvanised tanks and pipes. 

Table 21.22 Saturated solubilities of atmospheric gases in sea-water at various temperatures Concentrations of oxygen, nitrogen and carbon dioxide in equilibrium with 1 atm (lOI 325 N m ) of designated gas...

A carbonated beverage is made by saturating water with carbon dioxide at 0°C and a pressure of 3.0 atm. The bottle is then opened at room temperature (25°C), and comes to equilibrium with air in the room containing... 

It is necessary to draw attention to the variable pH of water which may be encountered in quantitative analysis. Water in equilibrium with the normal atmosphere which contains 0.03 per cent by volume of carbon dioxide has a pH of about 5.7 very carefully prepared conductivity water has a pH close to 7 water saturated with carbon dioxide under a pressure of one atmosphere has a pH of about 3.7 at 25 °C. The analyst may therefore be dealing, according to the conditions that prevail in the laboratory, with water having a pH between the two extremes pH 3.7 and pH 7. Hence for indicators which show their alkaline colours at pH values above 4.5, the effect of carbon dioxide introduced during a titration, either from the atmosphere or from the titrating solutions, must be seriously considered. This subject is discussed again later (Section 10.12). 

Before dealing with these, it is necessary to refer briefly to the stability of thiosulphate solutions. Solutions prepared with conductivity (equilibrium) water are perfectly stable. However, ordinary distilled water usually contains an excess of carbon dioxide this may cause a slow decomposition to take place with the formation of sulphur ... 

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