# Ka and ph relationship alkalinity

### Difference Between pKa and pH | Definition, Values, Relationship

Learn what pH, pKa, Ka, pKb, and Kb mean for acids and bases, plus A low pH value indicates acidity, a pH=7 is neutral, and a high pH value indicates alkalinity . The pH pKa and pKb are related by the simple relation. 1 is highest acidic and 14 is highest alkaline. 7 is neutral, pH is measured by specific sensors/electrodes and in room temperature along with proper calibrations. In simple terms, pH is the concentration of acid protons [H+]. Bicarbonate in particular, is the strongest buffer (largest Ka value) and the effect of other buffers .

This is when you add a salt to a weak acid or base which contains one of the ions present in the acid or base. To be able to use the same process to solve for pH when this occurs, all you need to change are your "start" numbers.

Add the molarity of the ion which comes from the salt and then solve the Ka or Kb equation as you did earlier. Find the pH of a solution formed by dissolving 0. H3O2 in a total volume of 1.

## Difference Between pKa and pH

Acid-Base Titrations An acid-base titration is when you add a base to an acid until the equivalence point is reached which is where the moles of acid equals the moles of base. For the titration of a strong base and a strong acid, this equivalence point is reached when the pH of the solution is seven 7 as seen on the following titration curve: For the titration of a strong base with a weak acid, the equivalence point is reached when the pH is greater than seven 7.

The half equivalence point is when half of the total amount of base needed to neutralize the acid has been added. In an acid-base titration, the base will react with the weak acid and form a solution that contains the weak acid and its conjugate base until the acid is completely gone.

To solve these types of problems, we will use the weak acid's Ka value and the molarities in a similar way as we have before. Before demonstrating this way, let us first examine a short cut, called the Henderson-Hasselbalch Equation. This can only be used when you have some acid and some conjugate base in your solution.

If you only have acid, then you must do a pure Ka problem and if you only have base like when the titration is complete then you must do a Kb problem. Determine the pH of the solution. This equation is used frequently when trying to find the pH of buffer solutions.

### Calculating a Ka Value from a Known pH - Chemistry LibreTexts

A buffer solution is one which resists changes in pH upon the addition of small amounts of an acid or a base. The buffer capacity is the amount of acid or base the buffer can neutralize before the pH begins to change to an appropriate degree. This depends on the amount of acid or base in the buffer. Why exactly is this true? Here's a simple non-quantitative way to think about it: Carbon dioxide has a specific solubility in water as carbonic acid H2CO3.

Regardless of pH, at a given salt concentration this carbonic acid concentration is always the same. Further, at any given pH there is an exact mathematical relationship between H2CO3 and both bicarbonate and carbonate. These relationships are, in fact, those used to derive equation 2 the first term in equation 2 is bicarbonate, and the second term is carbonate.

For example, at a pH of about 9. At higher pH this multiplier rises, and there is consequently more bicarbonate and carbonate present. More bicarbonate and carbonate results in higher alkalinity, as is shown in Figure 1. A second interesting feature of Figure 1 is that at any pH, the alkalinity of seawater is much higher than that in fresh water.

The reason in simple terms is that the multipliers described above are larger in salt water. The more quantitative reason is that the dissociation constants are higher in salt water.

## Calculating a Ka Value from a Known pH

Higher dissociation constants force a higher concentration of bicarbonate and carbonate to be present for a given concentration of carbonic acid. Hence, they result in a higher alkalinity.

A third feature of this relationship involves the pH of seawater as the ambient CO2 level rises. If CO2 is allowed to double Figure 2the pH drops by 0. Consequently, in the future, the pH of seawater may actually drop into the upper 7's from the 8. The theoretical relationship between carbonate alkalinity and pH for seawater in equilibrium for preindustrial air green; ppm carbon dioxidecurrent air blue; ppm carbon dioxide and possible future air red; ppm carbon dioxide using equations 2 and 3.

An important point to keep in mind is that the relationship will be altered slightly if the tank is not in equilibrium with the air. Specifically, reef tanks are often not in equilibrium with the air, making the internal pCO2 for the tank something different than the surrounding air. For example, tanks using limewater can have a pH value of 8.

Looking at Figure 2, this puts them off of the theoretical relationship for seawater in ambient air. The fundamental explanation is that the tank is deficient in CO2. In effect, the tank has an internal pCO2 that is more like that for the preindustrial air with ppm CO2 Figure 2.

In this case, driving more CO2 from "normal air" into the water would lower the pH to about 8. Again, that set of values falls off of the theoretical curve shown in Figure 2. In this case, the tank has an artificially high internal pCO2 of more than twice "normal air". Driving more CO2 from the tank into "normal air" would raise the pH to about 8.

A third way that reef tanks can present unusual combinations of pH and alkalinity is if the tank is in an environment where the ambient CO2 is far from normal. Rarely would such a situation involve reduced CO2, but homes and businesses are frequently elevated with respect to CO2. Such levels as those represented by the ppm line in Figure 2 are frequently encountered by aquarists, especially those living in newer, "tighter" homes and some have proven this fact to themselves with carbon dioxide detectors.

Aquarists that experience chronic low pH despite adequate alkalinity and aeration may do so because their homes have such elevated levels of carbon dioxide. Many of these aquarists have found that the pH of their tanks rises substantially by simply leaving a window near the tank open to permit better exchange with exterior, "normal" air. Finally, pCO2 fluctuates within a reef tank every day because of the activities of the organisms present.

### Testing Alkalinity In Boiler Water

Some are producing CO2 as a waste product of metabolism, including all organisms in the dark. Those that photosynthesize consume CO2 during the day.

As a consequence, the pCO2 rises during the night and declines during the day. This change in pCO2 is largely responsible for the pH fluctuation over the course of a day. For all of these reasons, a tank may move between the red and green lines of Figure 2 or further in extreme cases without the alkalinity changing at all.

Typical diurnal pH fluctuations in a reef tank and in some natural lagoons, for that matter are about 0. For tanks with a larger fluctuation than about 0. This minimization is best accomplished by maximizing the gas exchange between the tank and "normal" air through better circulation, better aeration through devices such as skimmers, having part of the tank system, such as a refugium, on a reverse photocycle so some organisms are always photosynthesizing, or by more rapidly exchanging the room air with exterior air.

One can also impact the diurnal pH fluctuation by adding high pH additives like limewater or other high pH alkalinity additives during the nightly pH minimum, and by adding low pH additives like sodium bicarbonate during the daily pH maximum. The magnitude of the alkalinity itself, of course, can influence pH stability, and that is the focus of the next section.

What is "Buffering" Buffer and buffering are terms that are thrown around indiscriminately in the world of reefkeeping, and the actual meaning of these terms is often lost.

Many aquarists refer to any alkalinity supplement as a buffer, but this isn't the case. For example, neither sodium bicarbonate nor sodium carbonate, taken alone, is a true buffer.

A buffer is something that helps minimize pH changes in the presence of added acid or base. No buffer can completely stop the pH from changing when acid or base is added. The change in pH, however, is made smaller when an appropriate buffer is used. A buffer is almost always comprised of two different chemical entities.

Find the Ka of an acid (Given pH) (0.1 M Hypochlorous acid) EXAMPLE

Bicarbonate and carbonate together, for example, form a buffer in the pH range from about 8 to 11 in seawater, though the buffering is best between about 8.

Here's what is happening on a chemical level. When a base such as OH- is added to the system in an effort to raise pHsome of the bicarbonate is converted to carbonate. This process effectively "uses up" some of the OH- that was added, and the pH does not rise as much as it would without the "buffer".

At about pH 8. At lower pH, there is less CO, and at pH 8. Consequently, seawater is not especially well buffered against substantial pH drops when the pH is already less than 8.

It is, however, well buffered against substantial pH rises.