Chapter 15 • Acids/Bases with common ion – This is when you have a mixture like HF and NaF. – The result is that you have both the acid and it’s conjugate base so you have to keep track of both initial concentrations. – When doing a calculation make sure you include both initial concentrations
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Chapter 15 Acids/Bases with common ion –This is when you have a mixture like HF and NaF. –The result is that you have both the acid and it’s conjugate.
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Chapter 15
• Acids/Bases with common ion– This is when you have a mixture like HF and
NaF. – The result is that you have both the acid and
it’s conjugate base so you have to keep track of both initial concentrations.
– When doing a calculation make sure you include both initial concentrations
• Calculate the pH for a aqueous solution containing 1.0 M HF and 1.0 M NaF Ka = 7.2 x 10-4
– HF F- + H+
– Chem. [init.] [equil.]– HF 1.0 1.0 – x = 1.0– F- 1.0 1.0 + x = 1.0– H+ 0 + x
– 7.2 x 10-4 = (1.0)x / 1.0 x = 7.2 x 10-4
– pH = - log (7.2 x 10-4) = 3.14
• Buffered Solutions– Resists a change in pH by reacting with any
H+ or OH- added to it.– Done by having a weak acid and its conjugate
base or a weak base and its conjugate acid in the same solution (common ion)
– Ex. HNO2 / NaNO2 (NO2-)– Adding acid reacts with the NO2- to make
HNO2, adding base reacts with the HNO2 to make NO2-
• Henderson – Hasselbalch equation– From the definition of Ka we can create this
equation by taking the - log and simplifying
– pH = pKa + log ([A-] / [HA]) A- is the base HA is the acid
– This equation is an easy way to calculate the pH if you know the equilibrium concentrations of the acid and its conjugate base
• Buffering capacity
• Calculate the pH for a solution of 0.75M lactic acid (HC3H5O3) Ka = 1.4 x 10-4 and 0.25M sodium lactate (conj. Base).– Since the Ka and the conc. are not very close
we can assume very little acid changes to base or base changes to acid
– So the equilibrium conc. are equal to the initial conc. and we can use the Hend-Hassel eq.
• Weak Acid / Strong Base overview– Acid only Ka problem– Before eq. pt. Ka problem– ½ way pt. pH = pKa
– Eq. pt. Kb problem– Past eq. pt. strong base only
• Calculating Ka
– We can use a titration set of data and the pH to determine an unknown Ka value.
– Just determine the final concentrations of the weak acid and conjugate base then insert into the equation
– pH = pKa + log [A-]/[HA]
• Weak Base/Strong Acid– Before any acid Kb problem– Before eq. pt. Kb problem– ½ way pt. pOH = pKb– At eq. pt. Ka problem– Past eq. pt. focus on strong acid conc.
• Acid/Base Indicators– Weak acids/bases that change color at
specific pH points
– Usually use the ratio [I-]/[HI] = 1/10 to determine when an indicator changes color
– Phenolphthalein is most common indicator and is clear in acid solution and turns pink in base solution
• Solubility Equilibria– The equilibrium for a partially soluble ionic
substance in water– Solubility is the amount of solid that
dissociates– PbCl2 (s) Pb2+ + 2Cl-
• Ksp– Equilibrium constant for solubility– Ksp = [Pb2+][Cl-]2
• Calculate the Ksp for Bi2S3 that has a solubility of 1.0 x 10-15 M.– Bi2S3 (S) 2Bi3+ + 3S2-
– Ksp = [Bi3+]2[S2-]3
– 1.0 x 10-15 M Bi2S3 = 2.0 x 10-15 M Bi3+
– 1.0 x 10-15 M Bi2S3 = 3.0 x 10-15 M S2-
– Ksp = (2.0 x 10-15 )2(3.0 x 10-15 )3
– Ksp = 1.1 x 10-73
• Calculate the solubility for Cu(IO3)2 with a Ksp = 1.4 x 10-7
– Cu(IO3)2 Cu2+ + 2IO3-
– Ksp = [Cu2+][IO3-]2
– 1.4 x 10-7 = (x)(2x)2 = 4x3
– x = 3.3 x 10-3 M
• Common Ion Effect– Presence of one of the ions that are formed
by the ionic substance
– AgCl with Cl- (from NaCl) this would impact the solubility of the AgCl since some Cl- is already present
• Calculate the solubility of CaF2 in a 0.025 M NaF solution if the Ksp = 4.0 x 10-11