Weak Acid-Base Solutions
A solution of acetic acid has a pH of 2.54. What is the concentration of acetic acid in this solution?
0.462M
What is the Ka of the acid in problem 8? HA + H2O ↔ H3O+ + A-
Ka = (6.0x10-4M)( 6.0x10-4M)/ (0.20M) = 1.8x10-6
A 0.18M solution of the sodium salt of nicotinic acid (also known pharmaceutically as niacin) has a pH of 9.05. What is the value of ka for nicotinic acid? NaNic ↔ Na+ + Nic-
ka = kw/kb = 10-14/7.0x10-10 = 1.4x10-5
What percentage of C5H5NH+ reacts with water in a 0.10M solution of pyridinium chloride, C5H5NHCl? C5H5NH+ + H2O ↔ H3O+ + C5H5N
% ionization = (7.7x10-4M)/(0.10M) x100% = 0.77%
8. A 0.20 M solution of a weak acid, HA, has a pH of 3.22. What is the percentage ionization of the acid? HA + H2O ↔ H3O+ + A-
%ionization = ionized/not ionized x 100% = (6.0x10-4M)/(0.20M) = 0.30% ionization
How many grams of ammonium chloride would have to be dissolved in 500ml of 0.20M NH3 to prepare a solution buffered at pH 10.00? NH3 + H2O ↔ NH4+ + OH -
0.96g NH4Cl in 500ml of buffer solution
What mole ratio of NH4Cl to NH3 would buffer a solution at pH 9.25? NH3 + H2O ↔ NH4+ + OH -
1
The ka for HF is 6.8x10-4. What is the kb for F-? HF + H2O ↔ H3O+ + F- ka = 6.8x10-4
1.47x10-11
How many grams of sodium acetate, NaC2H3O2, would have to be added to 1.0L of 0.15M acetic acid (pKa 4.74) to make the solution a buffer for pH 5.00? HC2H3O2 + H2O ↔ H3O+ + C2H3O2-
22.3g in 1L soln
The decahydrate of sodium carbonate, Na2CO3∙ 10H2O, is known as washing soda. How many grams of Na2CO3∙ 10H2O should be dissolved in 1.00L of water at 25 C to have a solution with a pH of 11.62?
23.7g Na2CO3∙10H2O
For the titration of 25.00 ml of 0.1000 M acetic acid with 0.1000 M NaOH, calculate the pH: a. Before the addition of any NaOH solution, b. After 10.00 ml of the base has been added, c. After half of the HC2H3O2 has been neutralized, and d. At the equivalence point.
a. pH = -log(1.3x10-3) = 2.87 b. pH = - log(2.70x10-5M) = 4.57 c. pH = - log(1.80x10-5M) = 4.74 d. pH = 14.0 - 5.28 = 8.72
Phosphorous acid, H3PO3, is actually a diprotic acid for which ka1 = 1.0 X 10-2 and ka2 = 2.6 x 10-7. What are the values of [H+], [H2PO3-], and [HPO3-2] in a 1.0 M solution of H3PO3? What is the pH of the solution? H3PO3 + H2O ↔ H3O+ + H2PO3- ka1 = 1.0x10-2 H2PO3- + H2O ↔ H3O+ + HPO3-2 ka2 = 2.6x10-7
ka1= x = [H+] = [H2PO3- ] = 0.10M test => x = [H+] = [H2PO3- ]= 0.095M 0.095 = 0.10 within the accuracy of the measurements pH = -log(0.10) = 1.0 [H3PO3] = 1.0M - 0.10M = 0.90M ka2= [HPO3-2]= 2.6 x 10-7M
Tartaric acid, H2C4H4O6, is a diprotic acid for which Ka1 = 9.2 x 10-4 and Ka2 = 4.3 x 10-5. Calculate the molar concentrations of H+ and the two anions of tartaric acid in a solution that has a concentration of 0.10M. H2C4H4O6 + H2O ↔ H3O+ + HC4H4O6- ka1 = 9.2x10-4 HC4H4O6- + H2O ↔ H3O+ + C4H4O6-2 ka2 = 4.3x10-5
ka1= x = [H+] = [HC4H4O6-9.6x10-3M ka2= [H+] =6.42x10-4M ka2= x = [H+] = [HC4H4O6-]=3.8x10-5M
Meta-Periodic acid, HIO4, is an important oxidizing agent and a moderately strong acid. In a 0.10M solution [H+] = 3.8 x 10-2mole/L. Calculate the Ka and pKa for periodic acid HIO4 + H2O ↔ H3O+ + IO4-
ka=0.023 pka = -log ka = -log (0.023) = 1.63
Lactic acid, HC3H5O3, is responsible for the sour taste of sour milk. At 25 C. its Ka = 1.4 x 10-4. What is the Kb of its conjugate base, the lactate ion, C3H5O3-? HC3H5O3 + H2O ↔ H3O+ + C3H5O3-
kb = kw/ka = 10-14/1.4x10-4 = 7.1x10-11
Hydrogen peroxide, H2O2, is a weak acid with a Ka = 1.8 x 10-12. What is the value of Kb of its conjugate base? H2O2 + H2O ↔ H3O+ + HO2-
kb = kw/ka = 10-14/1.8x10-12 = 5.6x10-3
Ethylamine, CH3CH2NH2, has a strong, pungent odor similar to that of ammonia. Like ammonia, it is a Bronsted base. A 0.10M solution has pH of 11.86. Calculate the Kb and pKb for ethylamine. What is the percentage ionization of ethylamine in this solution? CH3CH2NH2 + H2O ↔ CH3CH2NH3+ + OH-
kb=5.6x10-4 pkb=3.25 % ionization= 7.2%
The compound para-aminobenzoic acid (PABA) is a powerful sun-screening agent whose salts were once used widely in sun-tanning and sun-screening lotions. The parent acid, which we may symbolize as H-Paba, is a weak acid with a pKa of 4.92 (at 25 C). What is the [H+] and pH of a 0.030M solution of this acid? H-Paba + H2O ↔ H3O + Paba-
pH = -log(0.010) = 2.0 % ionization = (0.010M)/(0.15M) x100% = 6.7%
What is the percentage ionization in a 0.15M solution of HF? What is the pH of the solution? HF + H2O ↔ H3O + F-
pH = -log(0.010) = 2.0 % ionization = (0.010M)/(0.15M) x100% = 6.7%
Calculate the concentrations of all the solute species involved in the equilibria in a 3.0M solution of H3PO4. Calculate the pH of the solution. H3PO4 + H2O ↔ H3O+ + H2PO4- ka1 = 7.1x10-3 H2PO4- + H2O ↔ H3O+ + HPO4-2 ka2 = 6.3x10-8 HPO4-2 + H2O ↔ H3O+ + PO4-3 ka3 = 4.5x10-13
pH = -log(0.146) = 0.84 ka1=0.146M ka2= 6.3x10-8M ka3= 1.9x10-19
What is the pH of a solution that contains 0.15M HC2H3O2 and 0.25M C2H3O2-? Use Ka = 1.8 x 10-5 for HC2H3O2. HC2H3O2 + H2O ↔ H3O+ + C2H3O2-
pH = -log(1.1x10-5M) = 4.97
Calculate the concentrations of all the solute species in a 0.15M solution of ascorbic acid (vitamin C). What is the pH of the solution? H2C6H6O6 + H2O ↔ H3O+ + HC6H6O6- ka1 = 6.7x10-5 HC6H6O6- + H2O ↔ H3O+ + C6H6O6-2 ka2 = 2.7x10-12
pH = -log(3.2x10-3) = 2.50 ka1=3.2x10-3M ka2=2.7x10-12M
Suppose 25.0ml of 0.10M HCl is added to a 250ml portion of a buffer composed of 0.25M NH3 and 0.20M NH4Cl. What is the pH of the buffer after the addition of the strong acid? We need to consider both the change in concentrations due to the addition of HCl and dilution. Dilution: C1V1 = C2V2
pH = 14 - log(2.0 x 10-5M) = 9.3
Calculate the pH of 0.20M NaCN. What is the concentration of HCN in the solution? CN- + H2O ↔ HCN + OH - (NaCN ↔ Na+ + CN- Since CN- is the conjugate base of a very weak acid, it is a relatively strong base. Na+, on the other hand, is the conjugate acid of an extremely strong base, NaOH, and is extremely weak. The CN- can be expected to react with water to control the pH in the salt solution.)
pH = 14.0 - 2.75 = 11.25 [HCN]=1.8x10-3M
Liquid, chlorine bleach is really nothing more than a solution of sodium hypochlorite, NaOCl, in water. Usually, the concentration is approximately 5% NaOCl by weight. Use this information to calculate the approximate pH of a bleach solution, assuming no other solutes are in the solution except NaOCl. (Assume the bleach has a density of 1.0 g/ml)
pH = 14.0 - 3.3 = 10.7
What is the pH of a 0.0050 M solution of sodium cyanide? NaCN ↔ Na+ + CN-
pH = 14.0 - 3.56 = 10.44
Calculate the pH of 0.12M Na2SO3. What are the concentrations of HSO3- and H2SO3 in the solution? This is the salt of a polyprotic acid: Na2SO3 ↔ 2Na+ + SO3-2 SO3-2 + H2O ↔ HSO3- + OH-
pH = 14.0 - 3.90 = 10.11 kb1=[HSO3- ]= 1.3x10-4M kb2=[H2SO3] = 8.3x10-13M
Sodium citrate, Na3C6H5O7, is used as an anticoagulant in the collection of blood. What is the pH of a 0.10M solution of this salt? Na3C6H5O7 ↔ 2Na+ + C6H5O7-3
pH = 14.0 - 4.30 = 9.70
A buffer is prepared containing 0.25M NH3 and 0.45M NH4+. Calculate the pH of the buffer using Kb for NH3. NH3 + H2O ↔ NH4+ + OH -
pH = 14.0 - 5.00 = 9.00
When 50 ml of 0.10M formic acid, HCHO2, is titrated with 0.10M sodium hydroxide, what is the pH at the equivalence point? (Be sure to take into account the change in volume during the titration.) Determining the pH value at the equivalence point is a salt problem. HCHO2 + NaOH ® NaCHO2 + HOH
pH = 14.0 - 5.78 = 8.22
What is the pH of a solution prepared by mixing 25 ml of 0.180M HC2H3O2 with 35.0 ml of 0.250M NaOH? HC2H3O2 + NaOH ® NaC2H3O2 + H2O
pH = 14.0 -1.15 = 12.85
What is the pH of a 0.50M solution of Na3PO4? In this solution. What are the concentrations of HPO4-, H2PO4-2, and H3PO4?
pH = 14.00 - 1.03 = 12.97 [HPO4-2] ≈ 0.094M [H2PO4-]=1.6x10-7M [H3PO4]=2.4x10-18M
Codeine, a cough suppressant extracted from crude opium, is a weak base with a pKb of 5.79. What will be the pH of a 0.020M solution of codeine? (Use Cod as a symbol for codeine). Cod- + H2O ↔ CodH + OH -
pH = 14.00 - 3.74 = 10.26
Calculate the pH of 0.15 M CH3NH3Cl. For methylamine, CH3NH2, Kb = 4.4 x 10-4. CH3NH3Cl ® CH3NH3+ + Cl- (Dissolved in water, this salt's cation, CH3NH3+, will react with water to change the pH of the salt solution.)
pH =-log (1.9x10-8) = 5.73
Many drugs that are natural Bronsted bases are put into aqueous solution as their much more soluble salts with strong acids. The powerful painkiller morphine, for example, is very slightly soluble in water, but morphine nitrate is quite soluble. We may represent morphine by the symbol Mor and its conjugate acid as H-Mor+, The pKb of morphine is 6.13. What is the calculated pH of a 0.20M solution of H-Mor+? H-MorNO3 ↔ H-Mor+ + NO3-
pH =-log (5.2x10-5) = 4.28
What are the concentrations of all the solute species in 0.150M lactic acid, HC3H5O2? What is the pH of the solution? This acid has Ka = 1.4 X 10-4. HC3H5O2 + H2O ↔ H3O+ + C3H5O2
x = [H+] = 4.6x10-3M pH = -log (4.6x10-3M) = 2.34
What is the pH of 0.15M hydrazoic acid, HN3? For HN3, Ka = 1.8 x 10-5. What percentage of the HN3 is ionized? N3H + H2O ↔ H3O+ + N3
x = [H+] = [N3-] = 1.6x10-3M pH = -log (1.6x10-3M) = 2.78 % ionization = 1.6x10-3M x 100% = 1.1% 0.15M
Suppose 25.00ml of 0.100M HCl is added to 125ml of an acetate buffer prepared by dissolving 0.100mole of acetic acid and 0.110mole of sodium acetate in 1L water. What are the initial and final pH values? What would be the pH if the same amount of HCl solution were added to 125ml of pure water?
△pH = 4.79 - 4.62 = 0.17 pH units [H+] = [HCl] = 0.017M => pH = -log(0.017) = 1.77
By how much will the pH change if 0.050mol of HCl is added to 1.00L of the buffer in problem 21? The acetic acid buffer solution contains 0.15M HC2H3O2 and 0.25M C2H3O2- and has a pH of 4.97; ka for acetic acid is 1.8 x 10-5 for HC2H3O2. HC2H3O2 + H2O ↔ H3O+ + C2H3O2-
△pH = 4.97 - 4.74 = 0.23 pH units
By how much will the pH change if 0.020moles of HCl is added to 1.00L of the buffer in problem 22? The buffer contained 0.25M NH3 and 0.45M NH4+ and had a pH of 9.00. NH3 + H2O ↔ NH4+ + OH -
△pH = 9.00 - 8.94 = 0.06 pH units