Acid Base Solubility Equilibria Lesson 17 by Owen Borville 10.31.2025
A system at equilibrium will shift in response to being stressed. The addition of a reactant or product can be an applied stress.
After the initial concentration in M is given for the reactant, the change in x for reactants and products can be calculated using an equilibrium table. When there is an addition in the products, equilibrium is directed toward the reactant.
AB (aq) ⇌ A(aq) + B(aq)
Initial concentration (M)
Change in concentration (x)
Mr=concentration of reactant
Mp=concentration of product
Equilibrium concentration of reactant AB = Mr-x
Equilibrium concentration of product A = x
Equilibrium concentration of product B = Mp + x
Therefore, these equilibrium concentrations are substituted into the equilibrium expression to give:
[(x)(Mp + x)]/[(Mr - x)]
If the x is expected to be very small, it can be considered negligible and the expression becomes:
[(x)(Mp)]/[(Mr)]
A buffer is a solution that contains a weak acid and its conjugate base (or a weak base and its conjugate acid).
Acid (aq)(reacts with added base) ⇌ H+(aq) + conjugate base (aq)(reacts with added acid)
Buffer solutions resist changes in pH. To calculate the pH of a buffer knowing the initial concentration of reactant acid and conjugate base, and knowing the equilibrium constant,
Ka = (x)(Mp + x)/(Mr - x)
When additional acid is added to the same buffer, the added acid reacts with the conjugate base.
so Ka = (x)(Mp + x)/(Mr - x)
If ionization is suppressed by the presence of a common ion, x can be neglected. The pH of a buffer solution can often be calculated by using a derived expression, the Henderson-Hasselbalch equation: pH = pKa + log[conjugate base]/[weak acid]
pKa is the Ka of an acid (strength)
Preparing a buffer solution with a specific pH: Buffers must have conjugate acid and base concentrations within a factor of 10.
10>=[conjugate base]/[weak acid]>=0.1
The pH of a buffer cannot be more than one pH unit different than the pKa of the weak acid it contains.
To make a buffer with a specific pH: (1) Pick a weak acid whose pKa is close to the desired pH. (2) Substitute the pH and pKa into the equation to obtain the necessary [conjugate base]/[weak acid] ratio.
pH = pKa + log[conjugate base]/[weak acid]
Strong Acid-Strong Base Titrations: The reaction between a strong acid and strong base: Net ionic equation: OH-(aq) + H+(aq) => H2O(l)
A strong acid-strong base titration involves neutralizing a strong acid with a strong base (or vice-versa) to determine an unknown concentration. The reaction forms a neutral salt and water, and the equivalence point occurs at a pH of 7. A pH indicator, such as phenolphthalein, is used to signal the endpoint, where the indicator changes color.
Weak Acid-Strong Base Titrations: A weak acid and strong base titration is a process to determine the concentration of a weak acid by reacting it with a strong base. At the equivalence point, the moles of acid and base are equal, and the resulting solution is basic, or pH greater than 7 because the weak acid has been converted into its conjugate base, which hydrolyzes water to produce hydroxide ions. Common indicators for this type of titration are phenolphthalein, which changes from colorless to pink in the basic pH range of the equivalence point.
Strong Acid-Weak Base Titrations: involves adding a strong acid to a weak base to determine the base's concentration. The pH starts high and decreases as the acid is added, with an acidic equivalence point less than pH 7 because the reaction produces a conjugate acid that creates a weakly acidic solution. A distinct titration curve is formed, including a buffer region where the pH changes slowly until the equivalence point is reached.
Acid-Base Indicators can be used to visualize the equivalence point in a titration: AB(aq) ⇌ A+(aq) + B-(aq). The endpoint of a titration is the point at which the color of the indicator changes.
A titration curve is a mathematical graph that plots the pH of a solution (y-axis) against the volume of a titrant (x-axis) added during a titration. These curves are used to monitor the progress of a reaction, determine the concentration of an unknown solution (analyte), and identify key points like the equivalence point, where the moles of the acid and base are equal and the acid and base have neutralized each other. The shape of the curve depends on the strength of the acid and base involved. However, the shape is usually an S-shaped where the middle vertical section of the S-curve contains the equivalence point. The buffer region is the region of the curve before the equivalence point where the solution contains a weak acid/base and its conjugate, creating a buffer that resists large pH changes. The end point is the point at which the visual indicator changes color, signaling that the equivalence point has been reached.
Solubility Equilibria: Solubility Product Expression and Ksp: Quantitative predictions about how much of a given ionic compound will dissolve in water is possible with the solubility product constant, Ksp.
AB(s) ⇌ A+(aq) + B-(aq)
Ksp = [A+][B-]
Calculations Involving Ksp and Solubility: Molar solubility is the number of moles of solute in 1 L of a saturated solution (mol/L). Solubility is the number of grams of solute in 1 L of a saturated solution (g/L). To calculate a compound's molar solubility, construct an equilibrium table with initial concentration values for reactants and products and equilibrium concentrations. Substitute x for unknown variables of change in concentration and solve for x. Convert concentration units from M to mol/L to grams per L.
Predicting Precipitation Reactions: For the dissociation of an ionic solid in water: possible conditions are: (1) solution is unsaturated (2) solution is saturated (3) solution is supersaturated.
Predicting when a precipitate might form: Q< Ksp = no precipitate forms. Q = Ksp = no precipitate forms. Q > Ksp = a precipitate forms. Use Ksp values to determine the concentration of each compound's ions, and use them to determine the value of the reaction quotient, Qsp using the equation. Then compare each reaction quotient with the value of the corresponding Ksp. If the reaction quotient is greater than Ksp, a precipitate will form. If less, a precipitate will not form.
Factors that affect the solubility of ionic compounds: the common ion effect, pH, and the formation of complex ions. The common ion effect is the decrease in solubility of a salt that occurs when a second salt containing a common ion is added to the solution.
The solubilities of salts containing anions that do not hydrolyze are unaffected by pH.
Complex ion formation: A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions.
Fractional precipitation is the separation of mixture based upon the component's solubilities. Some compounds or ions can be separated based on fractional precipitation.
Qualitative analysis involves the principle of selective precipitation and can be used to identify the types of ions present in a solution.
A system at equilibrium will shift in response to being stressed. The addition of a reactant or product can be an applied stress.
After the initial concentration in M is given for the reactant, the change in x for reactants and products can be calculated using an equilibrium table. When there is an addition in the products, equilibrium is directed toward the reactant.
AB (aq) ⇌ A(aq) + B(aq)
Initial concentration (M)
Change in concentration (x)
Mr=concentration of reactant
Mp=concentration of product
Equilibrium concentration of reactant AB = Mr-x
Equilibrium concentration of product A = x
Equilibrium concentration of product B = Mp + x
Therefore, these equilibrium concentrations are substituted into the equilibrium expression to give:
[(x)(Mp + x)]/[(Mr - x)]
If the x is expected to be very small, it can be considered negligible and the expression becomes:
[(x)(Mp)]/[(Mr)]
A buffer is a solution that contains a weak acid and its conjugate base (or a weak base and its conjugate acid).
Acid (aq)(reacts with added base) ⇌ H+(aq) + conjugate base (aq)(reacts with added acid)
Buffer solutions resist changes in pH. To calculate the pH of a buffer knowing the initial concentration of reactant acid and conjugate base, and knowing the equilibrium constant,
Ka = (x)(Mp + x)/(Mr - x)
When additional acid is added to the same buffer, the added acid reacts with the conjugate base.
so Ka = (x)(Mp + x)/(Mr - x)
If ionization is suppressed by the presence of a common ion, x can be neglected. The pH of a buffer solution can often be calculated by using a derived expression, the Henderson-Hasselbalch equation: pH = pKa + log[conjugate base]/[weak acid]
pKa is the Ka of an acid (strength)
Preparing a buffer solution with a specific pH: Buffers must have conjugate acid and base concentrations within a factor of 10.
10>=[conjugate base]/[weak acid]>=0.1
The pH of a buffer cannot be more than one pH unit different than the pKa of the weak acid it contains.
To make a buffer with a specific pH: (1) Pick a weak acid whose pKa is close to the desired pH. (2) Substitute the pH and pKa into the equation to obtain the necessary [conjugate base]/[weak acid] ratio.
pH = pKa + log[conjugate base]/[weak acid]
Strong Acid-Strong Base Titrations: The reaction between a strong acid and strong base: Net ionic equation: OH-(aq) + H+(aq) => H2O(l)
A strong acid-strong base titration involves neutralizing a strong acid with a strong base (or vice-versa) to determine an unknown concentration. The reaction forms a neutral salt and water, and the equivalence point occurs at a pH of 7. A pH indicator, such as phenolphthalein, is used to signal the endpoint, where the indicator changes color.
Weak Acid-Strong Base Titrations: A weak acid and strong base titration is a process to determine the concentration of a weak acid by reacting it with a strong base. At the equivalence point, the moles of acid and base are equal, and the resulting solution is basic, or pH greater than 7 because the weak acid has been converted into its conjugate base, which hydrolyzes water to produce hydroxide ions. Common indicators for this type of titration are phenolphthalein, which changes from colorless to pink in the basic pH range of the equivalence point.
Strong Acid-Weak Base Titrations: involves adding a strong acid to a weak base to determine the base's concentration. The pH starts high and decreases as the acid is added, with an acidic equivalence point less than pH 7 because the reaction produces a conjugate acid that creates a weakly acidic solution. A distinct titration curve is formed, including a buffer region where the pH changes slowly until the equivalence point is reached.
Acid-Base Indicators can be used to visualize the equivalence point in a titration: AB(aq) ⇌ A+(aq) + B-(aq). The endpoint of a titration is the point at which the color of the indicator changes.
A titration curve is a mathematical graph that plots the pH of a solution (y-axis) against the volume of a titrant (x-axis) added during a titration. These curves are used to monitor the progress of a reaction, determine the concentration of an unknown solution (analyte), and identify key points like the equivalence point, where the moles of the acid and base are equal and the acid and base have neutralized each other. The shape of the curve depends on the strength of the acid and base involved. However, the shape is usually an S-shaped where the middle vertical section of the S-curve contains the equivalence point. The buffer region is the region of the curve before the equivalence point where the solution contains a weak acid/base and its conjugate, creating a buffer that resists large pH changes. The end point is the point at which the visual indicator changes color, signaling that the equivalence point has been reached.
Solubility Equilibria: Solubility Product Expression and Ksp: Quantitative predictions about how much of a given ionic compound will dissolve in water is possible with the solubility product constant, Ksp.
AB(s) ⇌ A+(aq) + B-(aq)
Ksp = [A+][B-]
Calculations Involving Ksp and Solubility: Molar solubility is the number of moles of solute in 1 L of a saturated solution (mol/L). Solubility is the number of grams of solute in 1 L of a saturated solution (g/L). To calculate a compound's molar solubility, construct an equilibrium table with initial concentration values for reactants and products and equilibrium concentrations. Substitute x for unknown variables of change in concentration and solve for x. Convert concentration units from M to mol/L to grams per L.
Predicting Precipitation Reactions: For the dissociation of an ionic solid in water: possible conditions are: (1) solution is unsaturated (2) solution is saturated (3) solution is supersaturated.
Predicting when a precipitate might form: Q< Ksp = no precipitate forms. Q = Ksp = no precipitate forms. Q > Ksp = a precipitate forms. Use Ksp values to determine the concentration of each compound's ions, and use them to determine the value of the reaction quotient, Qsp using the equation. Then compare each reaction quotient with the value of the corresponding Ksp. If the reaction quotient is greater than Ksp, a precipitate will form. If less, a precipitate will not form.
Factors that affect the solubility of ionic compounds: the common ion effect, pH, and the formation of complex ions. The common ion effect is the decrease in solubility of a salt that occurs when a second salt containing a common ion is added to the solution.
The solubilities of salts containing anions that do not hydrolyze are unaffected by pH.
Complex ion formation: A complex ion is an ion containing a central metal cation bonded to one or more molecules or ions.
Fractional precipitation is the separation of mixture based upon the component's solubilities. Some compounds or ions can be separated based on fractional precipitation.
Qualitative analysis involves the principle of selective precipitation and can be used to identify the types of ions present in a solution.