The identification of substances and ions present in a aqueous solution sample.
Solutions that contain a weak acid or a weak base and one of its salts. A means of stabilizing the pH of aqueous solutions such as blood plasma, seawater, etc. (i.e. buffers).
A solution in which the pH resists changes when small amounts of strong acids or bases are added.
12.1 Buffer Action Summary
A buffer is a mixture of a weak conjugate acid-base pair that stabilizes the pH of a solution by providing both a source and a sink for protons.
A buffer made of a weak acid and its salt, pH>7.
A buffer made of a weak base and its salt, pH<7.
Calculating the pH of a buffer solution (acid example)
Identify the weak acid and its conjugate base. Write the proton transfer equilibrium between them and rearrange the expression for Ka to give [H₃O⁺]. Calculate the pH.
(Use tabled values and approximate conjugate acid/base concentrations by their initial values).
Calculating the pH change of a buffered solution (acid example)
Calculate the new concentration of acid (initial concentration minus how much reacts [based on rxn stoichiometry]). Do the same for the conjugate base. Calculate the pH from such.
Buffers are often made with equal amounts of conjugate acids/bases. This means that, since the acids/bases are so weak, and dissociate so little, in calculating Ka, the concentrations can be set equal to their initial concentrations so that they cancel out.
tl;dr when [HA]ini=[A⁻]ini, pH=pKa
pH = pKa + log [base]ini/[acid]ini
12.2 Designing a Buffer Summary
The pH of a buffer solution is close to the pKa of the weak acid component when the acid and base have similar concentrations.
The maximum amount of acid or base that can be added before the buffer loses its ability to resist large changes in pH.
Buffer capacity relies on...
A more concentrated buffer is more resistant to pH changes. The relative concentrations of weak acid/base also matter; when the weak base is at least 10% of the weak acid, the buffer better resists the addition of acid, and when the weak acid is at least 10% of the weak base, the buffer better resists the addition of base.
When the acid is 10× as abundant as the base ([acid] = 10[base], the pH is given by
pH = pKa + log [base]/10[base] = pKa + log 1/10 = pKa - 1
When the base is 10× more abundant than the acid, then pH = pKa + log 10[acid]/[acid] = pKa + 1
Buffers act most effectively when the pH is within a range of ±1 unit of pKa.
12.3 Buffer Capacity Summary
The capacity of a buffer is determined by its concentration and pH. A more concentrated buffer can react with more added acid or base than can a less concentrated one. A buffer solution is generall most effective when the pH is in the range pKa ± 1.
The unknown sample in a titration.
The solution of known concentration in a titration.
When the amount of OH⁻ or H₃O⁺ added as titrant is equal to the amount of H₃O⁺ or OH⁻ initially present in the analyte.
A plot of the pH of the analyte solution against the volume of titrant added during a titration.
Physiological Buffers (Blood)
Biological systems are highly reliant on pH buffers, ex. in the blood. Blood uses HCO₃⁻/H₂CO₃ in a ration of ~20:1, with most of the carbonic acid in the form of CO₂.
Alkalosis: when the pH of blood rises above normal levels.
Acidosis: when the pH of blood falls below normal levels.
Reliant on the ratio of HCO₃⁻ to H₂CO₃ present.
How blood pH is maintained
H₂CO₃: exhalation as CO₂
HCO₃⁻: excretion in urine
12.4 Strong Acid-Strong Base Summary
in the titration of a strong acid with a strong base or vice versa, the pH changes slowly initially, changes rapidly through pH=7 at the stoichiometric point, and then changes slowly again.
At the halfway point of an acid-base titration, since pH = 7...
[HA] = [A⁻] and thus pH = pKa
12.5 Strong Acid-Weak Base and Weak Acid-Strong Base Titrations Summary
Halfway to the stoichiometric point, the pH is equal to the pKa of the acid. The pH is greater than 7 at the stoichiometric point of the titration of a weak acid and strong base. The pH is less than 7 at the stoichiometric point of the titration of a weak base and strong acid.
A device that uses a special electrode to measure H₃O⁺ concentration.
A water-soluble organic dye with color that depends to pH.
How indicators work
An indicator changes color with pH because it is a weak acid that has one color in its acidic form and another color in its conjugate base form. Because it is a weak acid, it takes place in a proton transfer equilibrium and has its own Ka/pKa.
End Point of an Indicator
When the concentrations of its acid and base forms are equal, i.e. [acid]=[base]; the color change occurs when pH=pKin (in=indicator).
12.6 Acid-Base Indicators Summary
Acid-base indicators are weak acids that change color close to pH = pKin; an indicator should be chosen so that its end point is close to the stoichiometric point of the titration.
Acids with more than one donatable H⁺.
12.7 Stoichiometry of Polyprotic Acid Titrations Summary
The titration of a polyprotic acid has a stoichiometric point corresponding to the removal of each acidic hydrogen atom. The pH of a solution of a polyprotic acid undergoing a titration is estimated by considering the primary species in solution and the proton transfer equilibrium that determines the pH.