Ideas of pH and solubility connect directly to the common ion effect and Le Chatelier’s Principle. Remember that pH describes the concentration of hydrogen ions (H+) and, conversely, hydroxide ions (OH-) in a solution. By observing connections between solubility equilibria, Le Chatelier’s Principle, and ion concentrations, we will be able to achieve some insights as to how pH can impact solubility equilibria!
Effects of Acidic Solutions
Like most other conditions, pH has an important effect on solubility, specifically when dealing with compounds that decompose into the conjugate base of a weak acid. Remember that the conjugate base is the ion formed from the dissociation of an acid into H+ and A-, where A- is the conjugate base.
A necessary factor to consider when discussing different conjugate bases is that the conjugate base of a weak acid, one that does not fully dissociate, is basic. The same cannot be said for the conjugate bases of strong acids such as Cl-, Br-, and ClO4-. Why is this important? It means that the conjugate acid of a weak base A- can react with water in the following fashion:
A- + H2O ⇌ HA + OH-
If the dissolution took place in a strongly acidic solution, the H+ would react with A- to form HA. This will decrease [A-], decreasing Q and pushing the dissolution to the right. Therefore, an acidic solution will increase the solubility of a compound that forms the conjugate base of a weak acid. Generally, the more basic the anion (ie. the weaker the acid it derives from), the more soluble a compound will be in an acidic solution. Similarly, compounds that dissolve into hydroxides, such as strong bases, are more soluble in acidic solutions.
However, an acidic compound will be less soluble in an acidic solution due to the common-ion effect. Because of the already present H+ concentration, a soluble acid will be less soluble in an acidic solution.
As an example problem, consider the following reaction:
Fe(OH)3 ⇌ Fe3+ + 3OH-
Will the solubility increase or decrease in an acidic solution? Explain.
In an acidic solution, we will have a high concentration of H+ ions (also seen as H3O+ or hydronium). These H+ ions will react with the OH- ions to form H2O in the autoionization of water. Therefore, in an acidic solution, there will be a lower concentration of OH- ions. According to Le Chatelier’s Principle, decreasing the concentration of OH- will push equilibrium to the right because Q will become less than K. Therefore, in an acidic solution, Fe(OH)3 is more soluble than it is in a neutral solution.
Effects of Basic Solutions
The effects of a basic solution are essentially the opposite of those in an acidic solution. In a basic solution, compounds that dissolve into significant conjugate acids, such as NH4+, will be more soluble in a basic solution. This is because the following reaction will take place:
NH4+ ⇌ NH3 + H+
In a basic solution, the OH- will react with the NH4+ and decrease [NH4+], therefore decreasing Q and pushing the overall dissolution towards the products.
On the other hand, basic compounds such as strong bases like NaOH or compounds that dissolve into significant conjugate bases will be less soluble. For strong bases, such as Ba(OH)2, this is because of the common ion effect. The already present OH- concentration will serve as a common ion in the solution.
For weak conjugate bases, such as the case of CH3COO-, the reaction of CH3COO- and H2O forming OH- and CH3COOH will be shifted left because of Le Chatelier’s Principle. Thus, in a basic solution, there will be less solubility for compounds with conjugate bases as ionic components.
pH Neutral Compounds
An important note is that compounds that do not have any basicity or acidity are not impacted by pH. For example, NaCl dissociates into Na+ and Cl-, the conjugate acid of NaOH and the conjugate base of HCl, respectively. However, note that Na+ and Cl- are both insignificant. This is because they are the conjugates of strong acids/bases. Therefore, they will not be impacted by any acidity or basicity of the solution. The only time solubility may be affected by the presence of an acid or base is in the presence of a common ion, but this is not connected to pH itself.
Learning Summary
Although computations of solubility as functions of pH will not be tested on the AP Exam, the test expects a solid understanding of the qualitative effects that changes in pH have on the solubility of salts. In this topic, we learned how acidic conditions, basic conditions, and neutral conditions affect solubility in the context of Le Chatelier’s principle.
Frequently Asked Questions
What is pH and how does it affect solubility?
pH measures H+ concentration (pH = −log[H+]). It affects solubility whenever a dissolved ion is a weak acid, weak base, or OH− (use Le Châtelier). If an anion is the conjugate base of a weak acid (like CO3^2−), adding H+ (lower pH) protonates the anion to form the acid (HCO3−/H2CO3), removing that ion from solution and shifting the dissolution equilibrium right—solubility increases. Conversely, if a cation is a weak acid (metal hexaaqua complexes), raising pH deprotonates or forms insoluble hydroxides, decreasing solubility (unless the hydroxide is amphoteric, in which case it dissolves in both strong acid and strong base). Also remember the common-ion effect: adding a common ion lowers solubility. On the AP exam you’ll be asked to identify these qualitative shifts (no Ksp-pH calculations)—see the Topic 8.11 study guide for examples (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV). For more review and practice, check the unit overview (https://library.fiveable.me/ap-chemistry/unit-8) and 1000+ practice questions (https://library.fiveable.me/practice/ap-chemistry).
Why does changing pH make some salts more or less soluble?
Changing pH shifts equilibria that involve a salt’s ions. If a salt contains the conjugate base of a weak acid (like CO3^2– from Na2CO3), adding acid (↑[H+]) protonates that anion (CO3^2– + H+ → HCO3–), removing the anion from solution and driving more solid to dissolve (Le Châtelier). Conversely, if a salt contains a weak base or OH– (like a metal hydroxide), adding OH– (making solution basic) supplies a common ion and suppresses solubility (less dissolves). Amphoteric metal hydroxides (e.g., Zn(OH)2) can dissolve in both strong acid (protonation) and strong base (forming soluble complexes), so pH can either increase or decrease their solubility. Remember this qualitatively: ask which ion is protonated or has a common-ion shift. The AP CED expects you to identify these effects qualitatively (no Ksp/pH calculations on the exam). For a focused review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
I'm confused about how Le Chatelier's principle relates to pH and solubility - can someone explain?
Think of solubility equilibria (salt ⇌ ions) like any equilibrium: Le Châtelier says if you change concentration of one component, the system shifts to oppose that change. For pH effects, the “component” that changes is H+ or OH−. If one ion of the salt is a weak base (A−, conjugate base of HA), adding acid (↑H+) protonates A− to HA, removing A− from solution—equilibrium shifts right and more solid dissolves. Conversely, if a salt contains a weak acid cation (BH+), making solution basic (↑OH−) deprotonates BH+ to B, removing BH+ and increasing solubility. Hydroxides are pH-sensitive too: increasing [OH−] (high pH) drives M(OH)x precipitation (less soluble), while very low pH can dissolve amphoteric hydroxides by forming soluble complexes. Qualitative reasoning like this is all AP expects (CED allows qualitative effects; pH-dependent solubility calculations aren’t tested). For extra practice and a focused study guide on this topic, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and more unit review (https://library.fiveable.me/ap-chemistry/unit-8).
What's the difference between salts that are pH sensitive and those that aren't?
Short answer: a salt is pH-sensitive if one of its ions is a weak acid, a weak base, or OH−—otherwise it’s not. For pH-sensitive salts, changing H+ (pH) shifts equilibria (use Le Châtelier) by protonating/deprotonating the ion or removing/adding OH−, so solubility changes. Examples: salts with basic anions (CO3^2−, A− from weak acids) become more soluble in acid (they get protonated to HCO3− or HA); salts with acidic cations (NH4+, metal ions that hydrolyze) become more soluble in base if deprotonation occurs. Amphoteric hydroxides (Zn(OH)2, Al(OH)3) dissolve in both acid and base. Non-pH-sensitive salts contain only spectator ions from strong acids/bases (NaCl, KNO3) so pH changes don’t affect Ksp. Remember AP won’t ask you to compute solubility vs pH, but you should predict direction qualitatively (CED 8.11.A.1). Review the Topic 8.11 study guide here (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and more unit review at (https://library.fiveable.me/ap-chemistry/unit-8).
How do I know if a salt's solubility will change with pH?
Check whether one of the ions is a weak acid, a weak base, or OH−—if so, solubility is pH sensitive (CED 8.11.A.1). Use Le Châtelier qualitatively: - If the anion is the conjugate base of a weak acid (A−), adding acid protonates A− → HA, removing A− and shifting the Ksp equilibrium to dissolve more salt (solubility increases in acid). - If the cation is the conjugate acid of a weak base (e.g., NH4+), adding base deprotonates it → NH3, removing the cation and increasing solubility in basic conditions. - Metal hydroxides and OH−: they’re more soluble in acid (H+ consumes OH−) and some are amphoteric (e.g., Al(OH)3)—they dissolve in both strong acid and strong base. Don’t try to calculate pH-dependent solubility on the exam (excluded), but be ready to state direction qualitatively using Ksp/common-ion and Le Châtelier. For quick review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and the unit overview (https://library.fiveable.me/ap-chemistry/unit-8). For extra practice, try problems at https://library.fiveable.me/practice/ap-chemistry.
Why do salts with weak acids or weak bases dissolve differently at different pH levels?
Salts with a weak acid or weak base dissolve differently with pH because changing [H+] or [OH−] shifts equilibria (Le Châtelier). If the anion is the conjugate base of a weak acid (like A− from HA), adding acid (↑H+) protonates A− to form HA, removing A− from solution and driving more salt to dissolve to re-establish Ksp. Conversely, adding base (↑OH−) pushes deprotonation or adds a common ion that can suppress solubility. For salts containing a weakly basic cation (conjugate acid of a weak base), raising pH deprotonates the cation and can increase solubility; lowering pH protonates it and decreases solubility. Metal hydroxides and amphoteric hydroxides are especially pH-sensitive. Qualitatively, think: if added H+ reacts with one of the ions, the Ksp equilibrium shifts to dissolve more solid; if you add a common ion, solubility decreases. See the AP CED keywords (Le Châtelier, Ksp, protonation/deprotonation) and the Topic 8.11 study guide for examples (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV). For more practice, try problems at (https://library.fiveable.me/practice/ap-chemistry) and review Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8).
What happens to the solubility of calcium carbonate when you add acid?
Adding acid increases the solubility of CaCO3. Start with the dissolution equilibrium: CaCO3(s) ⇌ Ca2+(aq) + CO3^2−(aq). H+ from the acid reacts with carbonate (CO3^2− → HCO3− → H2CO3 → CO2(g)/H2O), removing CO3^2− from solution. By Le Châtelier’s principle the equilibrium shifts right to replace the removed carbonate, so more CaCO3 dissolves and [Ca2+] rises. In strong acid you often drive formation of CO2(g), which further pulls the equilibrium toward dissolution. This is a classic pH-dependent solubility effect (protonation of an anion → increased solubility) covered in Topic 8.11; you’d use qualitative reasoning, not Ksp calculations, on the AP exam. For a quick review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8). For extra practice, check the AP Chem practice problems (https://library.fiveable.me/practice/ap-chemistry).
Can someone explain in simple terms why hydroxide salts are more soluble in acidic solutions?
Hydroxide salts (like M(OH)n) are more soluble in acidic solutions because H+ removes OH− from the equilibrium, so Le Châtelier’s principle shifts the dissolution reaction to the right. For example, M(OH)2(s) ⇌ M2+(aq) + 2 OH−(aq). If you add acid, H+ + OH− → H2O, which lowers [OH−]. The Ksp relation still holds, so to restore equilibrium more solid dissolves and [M2+] increases—net: greater solubility. This is a qualitative, pH-dependent solubility effect covered in the CED (use Le Châtelier, common-ion concept, and Ksp ideas). Note some metal hydroxides are amphoteric (they also dissolve in strong base by forming complex ions), but the acid case is simply protonation of the hydroxide ion. For AP-style questions, you’ll usually explain qualitatively (no heavy Ksp math). For a short study guide on pH and solubility, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV). For extra practice, Fiveable has many AP Chem practice problems (https://library.fiveable.me/practice/ap-chemistry).
I don't understand why some ionic compounds care about pH and others don't - what makes them different?
Short answer: a salt’s solubility is pH-sensitive when one of its ions can be protonated or deprotonated (i.e., it’s a weak-acid/conjugate-base or weak-base/conjugate-acid) or when the anion is OH–. Why: changing [H+] shifts equilibria by Le Châtelier’s principle and changes the effective concentration of the ion that appears in the Ksp expression. Examples and how to think: - Basic anions (CO3^2–, F–, S^2–): adding acid protonates them (CO3^2– + 2H+ → H2CO3), removing the common anion and so increasing solubility of the salt (more dissolves to replace it). - Acidic cations (NH4+ is the conjugate acid of NH3): making solution basic can deprotonate NH4+ → NH3, reducing solubility for some salts or forming volatile NH3. - Metal hydroxides: adding acid consumes OH–, so many M(OH)n become more soluble. Amphoteric hydroxides (Al(OH)3, Zn(OH)2) dissolve in both strong acid and strong base. Remember Ksp and common-ion ideas but don’t worry about numerical solubility vs pH on the AP (computations are excluded). For a quick review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and the unit overview (https://library.fiveable.me/ap-chemistry/unit-8). For extra practice, check Fiveable’s AP Chem practice bank (https://library.fiveable.me/practice/ap-chemistry).
How does adding HCl affect the solubility of magnesium hydroxide?
Adding HCl increases the solubility of Mg(OH)2. HCl supplies H+ that reacts with OH− to make H2O (H+ + OH− → H2O). That removes OH− from the dissolution equilibrium Mg(OH)2(s) ⇌ Mg2+ + 2 OH−, so by Le Châtelier the solid dissolves to replace the lost OH−—more Mg2+ goes into solution. This is exactly the pH-dependent solubility behavior described in the CED (8.11.A.1): salts with the hydroxide ion become more soluble as you acidify. (By contrast, adding OH− would decrease solubility via the common-ion effect.) For AP prep, you won’t be asked to compute exact solubilities vs. pH on the exam, but you should be able to qualitatively explain this using Le Châtelier and Ksp ideas. See the Topic 8.11 study guide for a quick review (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and more practice problems at (https://library.fiveable.me/practice/ap-chemistry).
What are some real world examples where pH affects how much salt dissolves?
pH changes solubility any time a dissolved ion is a weak acid/base or involves OH−. Real-world examples: - Limestone/building stone (CaCO3): acidic rain (pH ~4–5) protonates CO3^2− → HCO3−, shifting equilibrium to dissolve more CaCO3 (think marble statues and cave formation). Le Châtelier explains this: removing CO3^2− as HCO3− increases dissolution. - Antacids (Mg(OH)2, CaCO3): in stomach acid (low pH) these solids dissolve more because H+ reacts with OH− or CO3^2−, increasing solubility; that's why antacids neutralize acid. - Metal hydroxides and amphoteric oxides (Al(OH)3, Zn(OH)2): they’re more soluble in strong acid (protonation) and also in strong base (forming soluble complex anions), so solubility depends nonlinearly on pH. - Carbonate scaling in pipes: raising pH (adding OH−) precipitates CaCO3; lowering pH dissolves scale. On the AP exam you should explain these qualitatively (Le Châtelier, protonation/deprotonation, common-ion/OH− effects)—no Ksp/pH computations are required. For a topical review see the pH & solubility study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and try practice problems (https://library.fiveable.me/practice/ap-chemistry).
Why do antacids work better in your acidic stomach than they would in neutral water?
Antacids work better in your acidic stomach because H+ reacts with the basic anions from the antacid, shifting solubility equilibria by Le Châtelier’s principle. Example: Mg(OH)2(s) ⇌ Mg2+(aq) + 2 OH−(aq). In stomach acid, H+ + OH− → H2O removes OH−, so the equilibrium shifts right and more solid dissolves (more neutralization occurs). For carbonate antacids (CaCO3), H+ protonates CO32− → HCO3− → CO2(g), again consuming the anion and driving more solid into solution. In neutral water there’s almost no extra H+ to consume OH− or CO32−, so the salt stays less soluble (common-ion / no protonation effect). This is exactly the “pH-sensitive solubility” idea in the CED (protonation of anions, deprotonation of cations, Le Châtelier). For more review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and practice problems (https://library.fiveable.me/practice/ap-chemistry). Note: AP won’t ask you to compute pH-dependent solubilities, just to explain the qualitative effect.
How do I predict whether increasing or decreasing pH will make a salt more soluble?
Think about whether one of the ions can react with H+ or OH–. Use Le Châtelier: if the ion is the conjugate base of a weak acid (A–), adding H+ (lowering pH) protonates A– to HA and pulls the dissolution equilibrium right → more soluble. If the ion is the conjugate acid of a weak base (BH+), removing H+ (raising pH) deprotonates BH+ to B and shifts equilibrium toward more dissolved salt → more soluble. For metal hydroxides (M(OH)n), lowering pH (adding H+) consumes OH– and increases solubility; amphoteric hydroxides (e.g., Al(OH)3) dissolve in both acid (protonation) and base (forming complex anions). Also watch common-ion effects: adding a spectator that’s already in solution reduces solubility. This is all qualitative—AP won’t ask detailed solubility computations (CED 8.11.A). For more review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV), Unit 8 overview (https://library.fiveable.me/ap-chemistry/unit-8), and extra practice problems (https://library.fiveable.me/practice/ap-chemistry).
What's the connection between weak acids/bases in salts and pH sensitivity?
If one ion of a salt is the conjugate base of a weak acid (A−) or the conjugate acid of a weak base (BH+), the salt’s solubility will change when pH shifts—Le Châtelier explains why. Protonating a basic anion (A− + H+ → HA) removes A−, so more solid dissolves to replace it (solubility increases in acid). Conversely, deprotonating a cationic acid (BH+ → B + H+) removes BH+, so solubility increases in base. Hydroxide salts are pH-sensitive too: most metal hydroxides dissolve in acid; amphoteric hydroxides (like Al(OH)3) also dissolve in strong base by forming complex anions. These are qualitative ideas only on the AP (you won’t be asked to compute pH-dependent Ksp values)—align answers with Ksp/common-ion and Le Châtelier reasoning (CED 8.11.A). For a quick refresher, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV) and more practice at Fiveable (https://library.fiveable.me/practice/ap-chemistry).
I missed the lab on pH and solubility - what should I know about how we test this?
In the lab you’d test pH effects on solubility by putting the same solid into solutions of different pH (acidic, neutral, basic) and watching whether it dissolves or precipitates. Qualitative tests you’d use: visual (precipitate forms/disappears), pH adjustment (add HCl or NaOH), and simple probes like conductivity or ion strips to see ion concentration change. Use Le Châtelier: if the anion is a weak base (e.g., CO3^2−), adding H+ protonates it and increases solubility; if the cation is the conjugate acid of a weak base, raising pH can deprotonate and increase solubility. For metal hydroxides, adding OH− generally lowers solubility except for amphoteric hydroxides (e.g., Al(OH)3) that dissolve in strong base. Remember AP focus: identify qualitative effects and explain with equilibria/Le Châtelier and common-ion ideas—quantitative pH-solubility calculations aren’t tested (CED exclusion). For a quick review, see the Topic 8.11 study guide (https://library.fiveable.me/ap-chemistry/unit-7/ph-solubility/study-guide/QD1VMGuBFQJ1Rcw5ifOV), unit overview (https://library.fiveable.me/ap-chemistry/unit-8) and practice problems (https://library.fiveable.me/practice/ap-chemistry).