Chapter 17

The effectiveness of a buffer depends on two factors:

(1) the relative amounts of buffer acid and base, and (2) the absolute concentrations of buffer acid and base.

buffer

- resists changes in pH by neutralizing added acid or added base. -Contains either significant amounts of a weak acid and its conjugate base. - significant amounts of a weak base and its conjugate acid. ***Blood has a mixture of H2CO3 and HCO3−.

Selective Precipitation

A solution containing several different cations can often be separated by addition of a reagent that will form an insoluble salt with one of the ions but not the others (selective precipitation). A successful reagent can precipitate with more than one of the cations, as long as their Ksp values are significantly different.

The Effect of Common Ion on Solubility

Addition of a soluble salt that contains one of the ions of the "insoluble" salt decreases the solubility of the "insoluble" salt.

Qualitative Analysis

An analytical scheme that utilizes selective precipitation to identify the ions present in a solution is called qualitative analysis. Wet chemistry A sample containing several ions is subjected to the addition of several precipitating agents. Addition of each reagent causes one of the ions present to precipitate out.

How Acid Buffers Work: Addition of Base HA(aq) + H2O(l) A−(aq) + H3O+(aq)

Buffers work by applying Le Châtelier's principle to weak acid equilibrium. Buffer solutions contain significant amounts of the weak acid molecules, HA. These molecules react with added base to neutralize it. HA(aq) + OH−(aq) → A−(aq) + H2O(l) You can also think of the H3O+ combining with the OH− to make H2O; the H3O+ is then replaced by the shifting equilibrium.

Calculating pH Changes in a Buffer Solution

Calculating the new pH after adding acid or base requires breaking the problem into two parts: 1. A stoichiometry calculation for the reaction of the added chemical with one of the ingredients of the buffer to reduce its initial concentration and increase the concentration of the other Added acid reacts with the A− to make more HA. Added base reacts with the HA to make more A−. 2. An equilibrium calculation of [H3O+] using the new initial values of [HA] and [A−]

The Effect of pH on Solubility

For insoluble ionic hydroxides, the higher the pH, the lower the solubility of the ionic hydroxide. And the lower the pH, the higher the solubility Higher pH = increased [OH−] Mg(OH)2(s) Mg2+(aq) + 2 OH−(aq) For insoluble ionic compounds that contain anions of weak acids, the lower the pH, the higher the solubility. CaCO3(s) Ca2+(aq) + CO32−(aq) H3O+(aq) + CO32− (aq) HCO3− (aq) + H2O(l)

Group 1

Group 1 cations are Ag+, Pb2+, and Hg22+. All of these cations form compounds with Cl− that are insoluble in water. As long as the concentration is large enough Precipitated by the addition of HCl

Group 2

Group 2 cations are Cu2+, Bi3+, Cd2+, Pb2+, Hg2+, As3+, Sb3+, and Sn4+. All these cations form compounds with HS− and S2− that are insoluble in water at low pH. Precipitated by the addition of H2S in HCl

Group 3

Group 3 cations are Fe2+, Co2+, Zn2+, Mn2+, and Ni2+ precipitated as sulfides, as well as Cr3+, Fe3+, and Al3+ precipitated as hydroxides. All of these cations form compounds with S2− that are insoluble in water at high pH. Precipitated by the addition of H2S in NaOH.

Group 4

Group 4 cations are Mg2+, Ca2+, and Ba2+. These cations form compounds with PO43− that are insoluble in water at high pH. Precipitated by the addition of (NH4)2HPO4

Group 5

Group 5 cations are Na+, K+, and NH4+. All of these cations form compounds that are soluble in water; they do not precipitate. They are identified by the color of their flame.

Common Ion Effect

In a weak acid solution, adding a salt containing the anion NaA, the conjugate base of the acid (known as the common ion), shifts the position of equilibrium to the left. This lowers the H3O+ ion concentration and causes the pH to be higher than the pH of the acid solution.

Making an Acidic Buffer Solution

It must contain significant amounts of both a weak acid and its conjugate base. If a strong base is added, it is neutralized by the weak acid (HC2H3O2) in the buffer. NaOH(aq) + HC2H3O2(aq) H2O(l) + NaC2H3O2(aq) If the amount of NaOH added is less than the amount of acetic acid present, the pH change is small. If a strong acid is added, it is neutralized by the conjugate base (NaC2H3O2) in the buffer. HCl(aq) + NaC2H3O2 HC2H3O2(aq) + NaCl(aq) If the amount of HCl is less than the amount of NaC2H3O2 present, the pH change is small.

Precipitation

Precipitation will occur when the concentrations of the ions exceed the solubility of the ionic compound. If we compare the reaction quotient, Q, for the current solution concentrations to the value of Ksp, we can determine if precipitation will occur. Q = Ksp, the solution is saturated, no precipitation. Q < Ksp, the solution is unsaturated, no precipitation. Q > Ksp, the solution would be above saturation, the salt above saturation will precipitate. Some solutions with Q > Ksp will not precipitate unless disturbed; these are called supersaturated solutions.

How Acid Buffers Work: Addition of Acid HA(aq) + H2O(l) A−(aq) + H3O+(aq)

The buffer solution also contains significant amounts of the conjugate base anion, A−. These ions combine with added acid to make more HA. H+(aq) + A−(aq) → HA(aq) After the equilibrium shifts, the concentration of H3O+ is kept constant.

If Ka1 >> Ka2, there will be two equivalence points in the titration.

The closer the Ka values are to each other, the less distinguishable the equivalence points are. Titration of 25.0 mL of 0.100 M H2SO3 with 0.100 M NaOH

solubility product, Ksp

The equilibrium constant for the dissociation of a solid salt into its aqueous ions And its solubility product constant is Ksp = [Ag+][Cl−].

Formation Constant

The reaction between an ion and ligands to form a complex ion is called a complex ion formation reaction. Ag+(aq) + 2 NH3(aq) Ag(NH3)2+(aq) The equilibrium constant for the formation reaction is called the formation constant, Kf.

Complex Ion Formation

Transition metals tend to be good Lewis acids. They often bond to one or more H2O molecules to form a hydrated ion. H2O is the Lewis base, donating electron pairs to form coordinate covalent bonds. Ag+(aq) + 2 H2O(l) Ag(H2O)2+(aq) Ions that form by combining a cation with several anions or neutral molecules are called complex ions. For example, Ag(H2O)2+ The ions or molecules that act as Lewis bases are called ligands. For example, H2O

buffer capacity

is the amount of acid or base a buffer can neutralize. **the amount of acid or base that can be added to a buffer without causing a large change in pH. The buffer capacity increases with increasing absolute concentration of the buffer components. As the [base]:[acid] ratio approaches 1, the ability of the buffer to neutralize both added acid and base improves. Buffers that need to work mainly with added acid generally have [base] > [acid]. Buffers that need to work mainly with added base generally have [acid] > [base].

molar solubility

is the number of moles of solute that will dissolve in a liter of solution. The molarity of the dissolved solute in a saturated solution Molar solubility is related to Ksp. But you cannot always compare solubilities of compounds by comparing their Ksp values. To compare Ksp values, the compounds must have the same dissociation stoichiometry.

buffer range

is the pH range in which the buffer can be effective.

Solubility

the amount of solute that will dissolve in a given amount of solution at a particular temperature.