Re: Ph of water-heads -- lots of water tested

Hi,
well I put on a link to the website where I ordered mine for $79, and that was pretty much the cheapest price I could find. I just sent you an email back, I hope you get it soon. Buffered solutions contain a weak acid and it's conjugate base. This resists changes in pH when strong acids or bases are added to the solution. It's better stated by Benjamin Abelow, M.D. in his book "Understanding Acid-Base" (which I'm thinking of getting, but have not read in total yet) where he says:
"A buffer is a chemical substance that helps stabilized the [H+] (and hence the pH) of a solution. The word buffer comes from an archaic usage of the word buff, which means “to limit the shock of” – in this case, the shock of a pH change. A buffered solution is a solution that has a buffer dissolved in it. If you add strong acid to a buffered solution, pH falls much less than it would if no buffer were present; if you add strong base to the solution, pH rises much less. All body fluids are buffered solutions. The term buffering refers to the process by which buffers act to minimize pH changes."

The remainder of the explaination I got from a section of the book that was available on line, and is why I think I want to get it. It's pretty long. At the bottom of this post you should find the link to the pH100. Ok, here is the rest of Abelow's introduction to buffers:

"In the first two of these chapters, we discuss the mechanism by which buffering takes place. In the next two chapters, we describe the particular chemical substances that act as buffers in the body. In the final chapter of the section, we examine how the different buffers work together to stabilize the pH of the extracellular fluid.

How Buffers work: Part 1

Understanding how buffers work is one of the more complex topics in acid-base. As a result, many students will find this chapter and the next quite challenging. Take your time. You need not memorize anything, but it is important to follow the progression of ideas closely. The effort you devote now will allow you to master all of acid-base, including areas of great clinical importance.

Introduction:

A buffer consists of two parts and, for this reason, is often referred to as a buffer pair. Even if we don’t say “pair,” the word is implied – just as it is implied when we speak of (a pair of) pants. What are these two parts? In human body fluids, the important buffers all consist of (1) a weak acid and (2) the conjugate base of that acid. As described earlier (p. 9), a weak acid is one that has only a modest tendency to dissociate.

For the rest of this chapter, we will work with a hypothetical weak acid, symbolized “HA.” It will be our demonstration model. HA consists of a proton, H+, and a negatively charged structure, designated “A-“ (for anion), which remains after the proton dissociates. HA is the weak acid and A- is its conjugate base. Thus, buffered solutions contain both dissolved HA and A-. Buffered solutions are generally prepared by mixing both HA and a salt of A- (e.g., NaA, which dissociates into Na+ and A-) in water.

How do buffers work? The brief answer is as follows: When [H+] starts to rise (say, due to the addition of a strong acid, like HCl), some of the dissolved A- combines with free protons via the reaction A- + H+ -> HA. This reaction takes some (but not all) of the added free protons out of solution, thereby minimizing the rise in [H+]. Conversely, when [H+] starts to fall (say, due to the addition of a strong base, like NaOH), some of the dissolved HA dissociates and liberates, or “releases,” free protons via the reaction HA -> A- + H+. This reaction replaces some (but not all) of the removed protons, thereby minimizing the fall in [H+]. Both A- and HA are needed because each plays a specific role. A- removes protons when [H+] rises and HA donates protons when [H+] falls. Together, A- and HA minimize changes in [H+] in either direction. We will now explore these processes in greater detail, beginning with underlying concepts and then applying these concepts to the buffering situation.

Underlying concepts:

When a weak acid (HA) is first dissolved in water, it begins to dissociate into free protons (H+) and anions (A-), As the concentrations of H+ and A- rise, the rate of randome collisions between these two ions increases. These collisions allow H+ and A- to interact chemically, re-forming molecules of undissociated HA. Soon an equilibrium is reached, which can be represented like this:

HA <--> H+ + A-

Or as two separate but simultaneously occurring reactions:

HA -> H+ + A- (called the “dissociation reaction”)
And
HA <- H+ + A- (called the “association reaction”)

At equilibrium, the rates of these two reactions are equal. Each molecule of HA that dissociates is immediately replaced by the combination (“association”) of an H+ and A-. And each H+ and A- that associate into HA are immediately replaced by the dissociation of an HA. Thus, [HA], [A-], and [H+] all remain constant. Chemical equilibria are sometimes called dynamic equilibria to emphasize that fact that, even though concentrations do not change, both the association and dissociation reactions never stop.

The rates of the association and dissociation reaction are described by the law of mass action. This law states that rate is proportional to the product of the concentrations of the reactants. For the association reaction (H+ + A- -> HA), the rate is proportional to [H+] x [A-]. For example, if [H+] doubles but [A-] stays the same, the reaction rate doubles. If both [H+] and [A-] double, the reaction rate quadruples. If both [H+] and [A-] are halved, the reaction rate is quartered (because ½ x ½ = ¼). This “multiplicative” relationship holds because the association of two ions requires that the ions collide with each other, and the probability of a random collision between two dissolved ions is proportional to the product of their concentrations.

To better grasp this concept, consider a box containing red and yellow marbles. If the box is shaken, the frequency of collisions between red and yellow marbles will be proportional to the product of their numbers. For example, if the box contains ten red marbles and one yellow marble, and we add another yellow marble, the frequency of red-yellow collisions will double. The same basic idea holds for collisions between H+ and A- in solution.

What about the dissociation reaction (HA -> H+ + A-)? Here the reaction rate is proportional simply to [HA]. Thus, if [HA] doubles, the reaction rate doubles. If [HA] is halved, the reaction rate is halved. This proportionality is also explained by the law of mass action, although the reason is less obvious. According to Bronsted-Lowry theory, the release of a proton from HA involves its transfer to water (HA + H2O -> H3O+ + A-). Because this transfer requires a collision between HA and H2O, the rate of the dissociation reaction is proportional to [HA] x [H2O]. (When we speak of concentration, we are usually referring to solutes, but it is perfectly correct to talk about the concentration of solvents, like water, as well.) However, the concentration of water is so high, about 55.5 moles per liter, that it remains essentially constant even if small amounts of water are consumed or produced during a chemical reaction (e.g., the percentage change in [H2O] is minute even if, say, 2 millimoles of water are produced). Since [H2O] does not change, we can simplify and say that the rate of dissociation reaction is proportional to [HA].

We can summarize the rules governing raction rates like this:
Dissociation reaction rate proportional to [HA] ->
HA <--> H+ + A-
<- Association reaction rate proportional to [H+] x [A-]

Going further. Actually, the reaction rate of ions is slightly lower than would be expected based on concentrations alone. Why? Neighboring ions in solution exert electrical attractions and repulsions on eachother; this slightly restricts their movements, reducing the rate of collisions. Thus, the more precise term “activity” (which takes into account both concentration and the reduction in ionic mobility) is sometimes used to designate ionic reactivity. For example, it is said that the rate of association reaction is proportional to (Activity of H+) x (Activity of A-). However, in dilute solutions such as body fluids, the average distance between ions is large. This keeps inter-ionic forces to a minimum, so activity is very close to estimates of reactivity based solely on concentration. It is therefore common to speak of reaction rates as being a function of concentrations, not activities.

Buffering in action:

Let’s now use the above ideas to explain how buffers stabilize pH. Consider a buffered solution (i.e., one containing HA and A-) to which we add a strong acid, like hydrochloric acid, HCl. The Hcl immediately and completely dissociates (HCl -> H+ + Cl-). The rise in [H+] increases the product of [H+] x [A-], so the rate of the association reaction increases. H+ and A- are converted into HA. Thus, some of the free protons that were released from the HCl are removed. However, as the buffering reaction proceeds, [H+] and [A-] fall, and [HA] rises. As a result, the rate of the association reaction decreases while the rate of the dissociation reaction increases. When these rates become equal, a new equilibrium is reached and no more H+ is removed. Buffering stops.

At this new equilibrium, [H+] is higher than it was before the strong acid was added, but it is much lower than it would have been if no buffer were present. Thus, bufferes are said to resist or mitigate changes in pH but – this is important – not to prevent them entirely.

ph100 $79 link