Acid-Base Chemistry

An Introduction • For thousands of years people have known that

vinegar, lemon juice and many other foods taste sour. However, it was not until a few hundred years ago that it was discovered why these things taste sour - because they are all acids. The term acid, in fact, comes from the Latin term acere, which means "sour". While there are many slightly different definitions of acids and bases, in this lesson we will introduce the fundamentals of acid/base chemistry.

• In the seventeenth century, the English writer and amateur chemist Robert Boyle first labeled substances as either acids or bases (he called bases alkalies) according to the following characteristics:

• Acids taste sour, are corrosive to metals, change litmus (a dye extracted from lichens) red, and become less acidic when mixed with bases.

• Bases feel slippery, change litmus blue, and become less basic when mixed with acids.

• While Boyle and others tried to explain why acids and bases behave the way they do, the first reasonable definition of acids and bases would not be proposed until 200 years later.

• In the late 1800s, the Swedish scientist Svante Arrhenius proposed that water can dissolve many compounds by separating them into their individual ions. Arrhenius suggested that acids are compounds that contain hydrogen and can dissolve in water to release hydrogen ions into solution. For example, hydrochloric acid (HCl) dissolves in water as follows:

• HCl H+(aq) + Cl-(aq)H2O

• Arrhenius defined bases as substances that dissolve in water to release hydroxide ions (OH-) into solution. For example, a typical base according to the Arrhenius definition is sodium hydroxide (NaOH):

• NaOH Na+(aq) + OH-(aq)• The Arrhenius definition of acids and bases explains a

number of things. Arrhenius's theory explains why all acids have similar properties to each other (and, conversely, why all bases are similar): because all acids release H+ into solution (and all bases release OH-). The Arrhenius definition also explains Boyle's observation that acids and bases counteract each other. This idea, that a base can make an acid weaker, and vice versa, is called neutralization.

H2O

Neutralization: As you can see from the equations, acids release H+ into solution and bases release OH-. If we were to mix an acid and base together, the H+ ion would combine with the OH- ion to make the molecule H2O, or plain water:

H+ (aq) + OH-(aq) H2O

• The neutralization reaction of an acid with a base will always produce water and a salt, as shown below:

 

 

Acid Base Water Salt

HCl + NaOH H2O + NaCl

HBr + KOH H2O + KBr

Definitions of acids and bases• Arrhenius

acid: generates [H+] in solutionbase: generates [OH-] in solutionnormal Arrhenius equation: acid + base <---> salt + waterexample: HCl + NaOH <---> NaCl + H2O

• Bronsted-Lowery:acid: anything that donates a [H+] (proton donor)base: anything that accepts a [H+] (proton acceptor)normal Bronsted-Lowery equation: acid + base <---> acid + baseexample: HNO2 + H2O <---> NO2- + H3O+Each acid has a conjugate base and each base has a conjugate acid. These conjugate pairs only differ by a proton. In this example: HNO2 is the acid, H2O is the base, NO2- is the conj. base, and H3O+ is the conj. acid.

• Lewis:acid: accepts an electron pairbase: donates an electron pairThe advantage of this theory is that many more reactions can be considered acid-base reactions because they do not have to occur in solution.

Salts

A salt is formed when an acid and a base are mixed and the acid releases H+ ions while the base releases OH- ions. This process is called hydrolysis. The pH of the salt depends on the strengths of the original acids and bases:

This is because the conjugate base of a strong acid is very weak and cannot undergo hydrolysis. Similarly, the conjugate acid of a strong base is very weak and likewise does not undergo hydrolysis.

Acid Base Salt pH

strong strong pH = 7

weak strong pH > 7

strong weak pH < 7

weak weak depends on which is stronger

Water• We typically talk about acid-base reactions in aqueous-

phase environments -- that is, in the presence of water. The most fundamental acid-base reaction is the dissociation of water:

• In this reaction, water breaks apart to form a hydrogen ion (H+) and a hydroxyl ion (OH-). In pure water, we can define a special equilibrium constant (Kw) as follows:

• Where Kw is the equilibrium constant for water (unitless)

[H+] is the molar concentration of hydrogen[OH- is the molar concentration of hydroxide

• An equilibrium constant less than one (1) suggests that the reaction prefers to stay on the side of the reactants -- in this case, water likes to stay as water. Because water hardly ionizes, it is a very poor conductor of electricity.

pHUnder the Brønsted-Lowry definition, both acids and bases are related to the

concentration of hydrogen ions present.  Acids increase the concentration of hydrogen ions, while bases decrease the concentration of hydrogen ions (by accepting them).  The acidity or basicity of something therefore can be measured by its hydrogen ion concentration.

In 1909, the Danish biochemist Sören Sörensen invented the pH scale for measuring acidity.  The pH scale is described by the formula:

pH = -log [H+]

.  When measuring pH, [H+] is in units of moles of H+ per liter of solution.

For example, a solution with [H+] = 1 x 10-7 moles/liter has a pH equal to 7 (a simpler way to think about pH is that it equals the exponent on the H+ concentration, ignoring the minus sign). The pH scale ranges from 0 to 14. Substances with a pH between 0 and less than 7 are acids (pH and [H+] are inversely related - lower pH means higher [H+] and a stronger acid). Substances with a pH greater than 7 and up to 14 are bases (the higher the pH, the stronger the base). Right in the middle, at pH = 7, are

neutral substances, for example, pure water.

• pOH gives us another way to measure the acidity of a solution. It is just the opposite of pH. A high pOH means the solution is acidic while a low pOH means the solution is basic.

pOH = -log[OH-]

pH + pOH = 14.00

Strong Acids: These acids completely ionize in solution so they are always represented in chemical equations in their ionized form. There are only seven (7) strong acids:

HCl, HBr, HI, H2SO4, HNO3, HClO3, HClO4

To calculate a pH value, it is easiest to follow the standard "Start, Change, Equilibrium“

Example Problem: Determine the pH of a 0.25 M solution of HBr.

Weak Acids: These are the most common type of acids. They follow the equation:

HA(aq) <---> H+(aq) + A-(aq)

The equilibrium constant for the dissociation of an acid is known as Ka. The larger the value of Ka, the stronger the acid.

Example Problem: Determine the pH of .30 M acetic acid (HC2H3O2) with the Ka of 1.8x10-5.

Strong Bases: Like strong acids, these bases completely ionize in solution and are always represented in their ionized form in chemical equations. There are only seven (7) strong bases:

LiOH, NaOH, KOH, RbOH, Ca(OH)2, Sr(OH)2, Ba(OH)2

• Example Problem: Determine the pH of a 0.010 M solution of Ba(OH)2.

Weak Bases: They follow the equation:

Weak Base + H2O <---> conjugate acid + OH-

example: NH3 + H2O <---> NH4+ + OH+

Kb is the base-dissociation constant:

Ka x Kb = Kw = 1.00x10-14

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. It is at this point where the pH = pKa of the weak acid.

Henderson-Hasselbalch Equation 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.

Where:pH is the log of the molar concentration of the hydrogenpKa is the equilibrium dissociation constant for an acid[base] is the molar concentration of a basic solution [acid] is the molar concentration of an acidic solution

This equation is used frequently when trying to find the pH of buffer solutions.

• Buffers: • A buffer is a compound that limits the

change in hydrogen ion concentration (and so pH) when hydrogen ions are added or removed from the solution. It may be useful to think of the buffer as being like a sponge. When hydrogen ions are in excess, the sponge mops up the extra ions. When in short supply the sponge can be squeezed out to release more hydrogen ions!

• All buffers are in fact weak acids or bases. Figure 3 shows how as hydrogen ions are added to a buffer solution they combine with A- (the conjugate base) and the reaction is pushed to the left. This creates more HA whilst removing the excess H+ from the solution. Similarly, as hydrogen ions are removed from solution by addition of a strong base the reaction moves to the right restoring the hydrogen ion concentration and reducing the quantity of HA.

• The effects of buffers can also be illustrated graphically. If a strong acid is added slowly to a buffer solution and the hydrogen ion concentration [H+] is measured then a plot similar to the one in figure 4 will be generated. Notice that during the highlighted portion of the curve a large volume of acid is added with little change in [H+] or pH.

• As we shall see later buffers are crucial in maintaining hydrogen ions within a narrow range concentrations in the body.

NORMAL pH • There is a normal pH value in each body

compartment (i.e. extracellular fluid, plasma, intracellular fluid etc). Intracellular pH is difficult to measure and may vary in different types of cells and in different parts of cells.

• pH of the plasma (i.e. pH of the plasma of whole blood = conventional "blood" pH) is controlled at 7.4 (7.35 - 7.45).

• There are three mechanisms which diminish pH changes in body fluid: buffers; respiratory; renal.

THE BUFFER SYSTEMS OF THE BODY • (a) Proteins are the most important buffers in the body.

They are mainly intracellular and include haemoglobin. The plasma proteins are buffers but the absolute amount is small compared to intracellular protein. Protein molecules possess basic and acidic groups which act as H+ acceptors or donors respectively if H+ is added or removed.

• (b) Phosphate buffer (H2PO4- : HPO4

2-) is mainly intracellular. The pK of this sytem is 6.8 so that it is moderately efficient at physiological pH's. The concentration of phosphate is low in the extracellular fluid but the phosphate buffer system is an important urinary buffer.

• (c) H2CO2 : HCO3- is not an important true buffer system

because normal blood pH (7.4) is so far from its pK (6.1). H2CO3 and HCO3

- are involved in pH control but they are not acting as a buffer system

Examination of "Buffering" Properties of HCO3-:H2CO2 System

Most texts state that the HCO3- : H2CO2 system is an efficient

physiological buffer because the components of the pair are controlled separately. As it is not a chemical buffer of any reasonable efficiency at the blood pH use of the term "buffer" in respect to HCO3

- : H2CO2 action introduces considerable confusion. This is illustrated in the following example.

Plasma has a [HCO3-] of approximately 24meq/l and [H2CO2] of

1.2meq/l, hence: