Reading Notes for Chapter 10

Chapter Summary:

Terminology Alert - Protons and Hydronium Ions:

Discussions of acids (and bases...) often use the term "proton" to refer to "H⁺". This can at times lead to confusion, because a "proton" is a subatomic particle and "H⁺" is an ion of hydrogen, but let's look closer. The most common isotope of hydrogen (atoms) has 1 proton in the nucleus and 1 electron outside the nucleus. When a hydrogen atom becomes a hydrogen cation, "H⁺", it loses 1 electron and we are left with... a proton! For this reason, we will often refer to the "H⁺" associated with acids (and bases...) as a "proton". We will use these terms ("proton" and "H⁺") interchangeably when talking about acids (and bases...). Application Note: You may have heard of some drugs that are "proton pump inhibitors" that can be used to treat excess stomach acid... a "proton pump" is the way your body provides H⁺ to your stomach!

A "proton" is very small, so although its charge is only +1, that charge is very dense because it is contained in a very small space. Because of this, it is very strongly attracted to negative charges, including the partial negative charge on the oxygen of a polar water molecule. In fact, if we have a "proton" floating around in water, it is very unlikely that it will just be wandering around all alone, it will always be associated with a water molecule. When "H⁺" sticks to a water molecule, we get "H₃O⁺", the "hydronium ion. When we are writing chemical equations, sometimes it's clearer (and easier) to write "H⁺(aq)", and other times it's clearer to write "H₃O⁺(aq)". Use whichever one more clearly communicates what you are trying to say.

Arrhenius, a functional/observational definition:

Identifying Acids & Bases:

Identifying acids and bases takes a little practice. Just because something has "H" doesn't mean it's an acid, and just because there's an "OH" doesn't mean it's a base. To some extent, "memorizing" a list like Table 10.4 in your textbook is a good start... read through that list a few dozen times and those names and formulas will start to stick. If you look closely at the acids and bases on that list, you'll probably start to see some trends, including:

• Metal hydroxides are usually bases. This often depends upon the solubility of the specific metal hydroxide... many hydroxides are insoluble so they may not appreciably change the hydroxide ion concentration when dissolved in water.
• Oxoanions with added "H" are usually acids. The names of these acids usually ends in "-ic acid" or "-ous acid". Nitric acid, carbonic acid, perchloric acid, sulfurous acid, phosphorous acid...
• Ammonia is an interesting compound... it might look like an acid (similar to H₂S(aq) or HBr(aq)), but it's actually a base.

Acid-Base Reactions - Neutralization:

When an acid reacts with a base, the resulting mixture is less acidic than the original acid and less basic than the original base. This type of reaction is called a neutralization.

When a neutralization occurs between an Arrhenius acid and an Arrhenius base, the reaction produces water and a salt.

For acids that have more than one proton (like phosphoric acid), the neutralization occurs 1 step at a time.

Bronsted-Lowry, a molecular definition:

Bronsted-Lowry definitions of acid and base recognize that the most important thing in an acid-base reaction is the movement of H⁺. A Bronsted-Lowry acid is a substance that supplies (loses) H⁺; a Bronsted-Lowry base is a substance that gains H+. Bronsted-Lowry acids are usually referred to as "H⁺-donors" or "proton-donors"; Bronsted-Lowry bases are usually referred to as "H⁺-acceptors" or "proton-acceptors".

Things that we called "acids" using the Arrhenius definition are still acids in the Bronsted-Lowry definition, and things we called Arrhenius "bases" are still Bronsted-Lowry bases. The two definitions are just looking at slightly different things when they assess whether something is an acid or a base.

Bronsted-Lowry actually helps us understand why ammonia is a base and H₂S is an acid. Energetically, it is much more favorable for NH₃ to accept a proton to become NH₄⁺(aq) than for it to lose a proton to become NH₂^-(aq); conversely, H₂S can lose a proton to become the (relatively) stable HS^-(aq), but it is much less likely to gain a proton because H₃S⁺(aq) is (relatively) unstable.

What's water?:

Consider the reactions just mentioned: ammonia or H₂S dissolving in water.

NH₃(aq) + H₂O(l)  -->  OH^-(aq) + NH₄⁺(aq)

H₂S(aq) + H₂O(l)  -->  H₃O⁺(aq) + HS^-(aq)

These reactions show ammonia gaining a proton (acting as a base) and H₂S losing a proton (acting as an acid), but what else is going on in these reactions? In the first reaction, water is losing a proton, so we would say that water is an acid. In the second reaction, water is gaining a proton, so we would say that water is a base. Um... so it's both??

In most situations, "acid" and "base" are relative descriptions of how a substance reacts during an acid-base reaction. If a substance reacts with something that is a stronger acid than itself, then it will act as the base in that reaction. Similarly, if a substance reacts with something that is a stronger base than itself, then it will act as the acid in that reaction.

Water is not a very strong acid or a very strong base, so its acid-base reactivity depends upon what it's reacting with. In a sample of pure water, water molecules act as both acid and base in a reaction called autoionization. In pure water, the concentration of H⁺(aq) {or H₃O⁺(aq) if you prefer} is equal to the concentration of OH^-(aq)

Acid and Base Strength:

Using the Bronsted-Lowry definitions, an acid is a proton-donor. The stronger the acid, the better it is at donating those protons. If an acid completely donates its protons, we call it a "strong" acid. Strong acids are completely ionized in solution.

Similarly for bases, a "strong" base is one that is completely ionized in solution. This is usually linked to solubility... a "strong" base is completely soluble.

At this point, it's probably easiest to just memorize the strong acids and bases because there aren't that many common ones. Table 10.2 in your textbook lists them. Read the list 43 times. Or make flash cards. Memorize.

For our purposes, any acid or base that is not listed as "strong" will be considered "weak". Weak acids (and bases...) do not completely dissociate. For example, if I have a solution of acetic acid (not on the list, so it's weak), that solution will have H₃O⁺(aq) and C₂H₃O₂^-(aq) ions floating around in solution, but there will also be a significant number of undissociated acetic acid molecules {HC₂H₃O₂(aq)} in the solution. These substances are all in equilibrium.

Equilibrium - We often write chemical reactions as if they only occur in one direction, but most chemical reactions can actually go both forward and backward. When the rate going forward is equal to the rate going backward, the observed concentrations do not change and the reaction is "at equilibrium". It is important to recognize that when a reaction is at equilibrium, the reaction has not stopped, the forward and reverse rates are just equal. Equilibrium is a dynamic process.

Measuring Acidity - pH:

REMINDER: When we are expressing the concentration of a solution in units of molarity, it is typical to simply enclose the solute in square brackets, "[ ]". Whenever you're talking about concentrations of substances in solution and see square brackets, it means that you're looking at a concentration in units of molarity. "[Fe⁺²]" means "the concentration of iron(III) ions in units of molarity".

If we think about acids and bases in a Bronsted-Lowry sense, it seems like the most important thing to watch is the protons. The concentration of protons (or "H⁺(aq)" or "H₃O⁺(aq)") tells us whether a solution is acidic, basic or neutral. Acidic solutions have [H₃O⁺] > 10^-7M. Basic solutions have [H₃O⁺] < 10^-7M. Neutral solutions have [H₃O⁺] = 10^-7M.

NOTE: Those concentrations are expressed as pretty exact numbers, but keep in mind that acid-neutral-base is a bit blurrier than that. If a solution has [H₃O⁺] = 1.5x10^-7M, it is still a "neutral" solution, even though the [H₃O⁺] is slightly greater than 10^-7M. In general, when we look at differences that we would consider significant, we're talking about differences in the power of 10. A solution with [H₃O⁺] = 10^-5M is definitely acidic. How do we deal with those powers of 10 using numbers that are a little easier to work with? Well...

Because [H₃O⁺] is usually a small number that is expressed in scientific notation, it can be a little cumbersome. To solve that, we use the "pH scale". This is just a mathematical change to make the [H₃O⁺] easier to express. Whenever we use "p" in this way, it means that we take the negative common logarithm of the quantity of interest. In this case, we're interested in the concentration of [H₃O⁺], or equivalently [H⁺], so we want to calculate p[H⁺]. To make the notation a little easier, we remove the square brackets and the "+" and are left with "pH".

A neutral solution has [H₃O⁺] = 10^-7M, so its pH = -log [H₃O⁺] = -log(10^-7) = 7.

For an acidic solution with [H₃O⁺] = 10^-5M, pH = -log [H₃O⁺] = -log(10^-5) = 5.

For a basic solution with [H₃O⁺] = 10^-10M, pH = -log [H₃O⁺] = -log(10^-10) = 10.

So solutions with pH<7 are acidic and solutions with pH>7 are basic.

CAUTION: pH tells you how acidic or basic a solution is. It does not tell you if an acid or base is "strong" or "weak". You can make a pH=6 solution using a strong acid and a pH=4 solution using a weak acid. The pH=4 solution is more acidic than the pH=6 solution, but the "strong" acid is still "strong" and the "weak" acid is still "weak". pH allows us to compare the acidity of a solution, it does not tell us how strong or weak an acid (or base) is. This can sometimes cause confusion because people (including me sometimes) will refer to a "strongly acidic solution" when talking about a solution with a (relatively) low pH, even if that solution is made with a weak acid!

An Exception... IF you prepare solutions of two acids that are exactly the same concentration, the stronger acid will have a lower pH. This is a very specific experiment, and it can be a useful one, but don't let this exception stick in your brain too hard. In 99% of situations, you cannot determine whether an acid is "strong" simply by looking at pH.

Buffers:

Buffers are everywhere, especially in biological systems. Why are they important? Let's look at an example...

Pure deionized water has [H₃O⁺] = 10^-7M, so its pH = -log [H₃O⁺] = -log(10^-7) = 7. Concentrated hydrochloric acid has a concentration of 12.0M. If 1.00mL of concentrated HCl(aq) is added to enough deionized water to make a liter of solution, the concentration of HCl(aq) in the resulting solution is:

C₁V₁=C₂V₂
(12.0M)(1.00mL) = C₂(1000.L)
C₂ = 0.012M

Since HCl(aq) is a strong acid, it completely dissociates in water so if [HCl] = 0.012M then [H₃O⁺] = 0.012M. This means the expected pH of the resulting solution would be

pH = -log [H₃O⁺] = -log(0.012) = 1.92 (!?!?!?!)

If pH were able to change that much with such small additions of acids or bases, it would be extremely hard for life to exist.

Buffers are mixtures of weak conjugate acid/conjugate base pairs that resist drastic changes in pH when (relatively) small amounts of acid or base are added.

Terminology Alert - A "conjugate acid/conjugate base pair" are two chemical species that differ by a single H⁺. For example, HCl and Cl^- are a conjugate acid/conjugate base pair; HCl is the conjugate acid of Cl^- and Cl^- is the conjugate base of HCl. Some other conjugate acid/conjugate base pairs: HC₂H₃O₂/C₂H₃O₂^-, H₂SO₄/HSO₄^-, HSO₄^-/SO₄^-2, H₃PO₄/H₂PO₄^-, H₂PO₄^-/HPO₄^-2, HPO₄^-2/PO₄^-3.

Buffers work because the added acid or base are "used up" when they react with the conjugate base or acid in the buffer. As long as the concentrations of the conjugate acid and conjugate base are similar (within about 10x of each other) and (relatively) high compared to the amount of added acid or base, the buffer will resist changes to pH. Remember, the conjugate acid of a buffer must be a weak acid.

Buffers are often referred to by their conjugate base component and a specific pH... "a pH 4.95 acetate buffer" is a buffer at pH 4.95 using a mixture of acetic acid and acetate ions. A "pH 6.65 carbonate buffer" is a buffer at pH 6.65 using a mixture of bicarbonate ions and carbonate ions.

Lewis Definitions - expanding the world of acids and bases:

There is one more definition of acids and bases that allows us to use the ideas of acid-base chemistry on a broader range of substances. Bronsted and Lowry defined acids and bases by following the proton (H⁺), but it can be argued that almost all chemical reactivity is a result of what the electrons are doing. G.N. Lewis favored looking at the electrons and chose to define acids and bases by describing what electrons are doing in acid-base reactions. In the Lewis Definitions, acids are electron pair acceptors and bases are electron pair donors. These definitions are complementary to and consistent with the Arrhenius and Bronsted-Lowry definitions, but the Lewis definitions allow us to explain even more things using acid-base chemistry as our framework. In fact, if you go all the way back to the top of this page, I mentioned that "AlCl₃ is also an acid even though there's no "H" in its chemical formula". Without any hydrogen in the chemical formula, it can be challenging to explain why AlCl₃ is an acid using the Bronsted-Lowry definitions, but the Lewis definitions give us a little easier picture - Al⁺³ is a small ion with a high positive charge, so we would expect it to attract electron pairs (from a Lewis base like water) very strongly.

Exploring Lewis acids and bases is a fascinating journey. I include the Lewis definitions here for completeness, but we will not be exploring Lewis acid-base chemistry in depth in this course.