pH of Salt Solutions

pH of Salt Solutions

• A salt can change the pH of a solution when dissolved in water.

Definition of a Salt

• A salt is any water-soluble ionic compound, not just sodium chloride (table salt).
• A salt is the product of acid-base neutralization, other than water.
• Example: Sulfuric acid ($H<em>2SO</em>4$) plus sodium hydroxide (NaOH) yields sodium sulfate ($Na<em>2SO</em>4$) and water.

Weak Acids and Conjugate Bases

• As the strength of a weak acid decreases, the strength of its conjugate base increases.
• As the strength of a weak base decreases, the strength of its conjugate acid increases.

Example

• Nitrate is a neutral anion because it's the conjugate base of a strong acid.
• Fluoride is more basic than nitrate because hydrofluoric acid is less acidic than nitric acid.

$K_a$ Values and Basicity

• For a generic acid: $HA \rightleftharpoons H^+ + A^-$; $K_a$ is usually much less than 1.
• For any weak acid, $A^-$ is more basic than $HA$ is acidic.
• $K<em>a = \frac{[H^+][A^-]}{[HA]}$; if Ka << 1, then the reaction lies strongly in the reverse direction.
  • This means the concentration of products is much lower than reactants at equilibrium.
  • Weak acids have an equilibrium that lies strongly in the reverse direction, indicated by a very small $K_a$ value.

Impact of Very Weak Acids

• For hydrocyanic acid, $K_a$ is around $10^{-10}$, meaning the equilibrium lies even more in the reverse direction.
• Cyanide ($CN^-$) is a better base than $F^-$ because it pushes the equilibrium further in the reverse direction.

Hydrolysis of Salts

• Salts with ions that are conjugate bases of weak acids (e.g., sodium cyanide, potassium fluoride) undergo hydrolysis to form basic solutions.
• Salts with ions that are cation conjugate acids of weak bases (e.g., ammonium) form acidic solutions via hydrolysis.
• Example: Ammonium nitrate will create an acidic solution because ammonium is the conjugate of the weak base ammonia.

Ions Conjugate to Strong Acids/Bases

• Ions conjugate to strong acids or bases have a negligible effect on the pH of a solution.
• Example: Dissolving sodium chloride (NaCl) in water.
  • Sodium ($Na^+$) is the conjugate of a strong base (NaOH).
  • Chloride ($Cl^-$) is the conjugate base of a strong acid (HCl).
  • Neither ion is strong enough to react with water, so the pH remains neutral (approximately 7).

Explanation

• Chloride is too weak a base to take an $H^+$ back from hydronium ($H_3O^+$) or to break a covalent bond between H and O in water.
• It takes more energy to break the H-O bond in water than to take back a proton.

Hydrolysis Reactions Explained

Acidic Cations

• Ammonium chloride ($NH<em>4Cl$) is not neutral because ammonium ($NH</em>4^+$) is the conjugate acid of the weak base ammonia ($NH_3$).
• Therefore, ammonium is a fairly strong acid and undergoes hydrolysis:
  • $NH<em>4^+ + H</em>2O \rightleftharpoons H<em>3O^+ + NH</em>3$
• Excess hydronium ions ($H_3O^+$) make the solution acidic (pH < 7).

Competing Reactions

• The $NH_3$ formed can react with water to create more ammonium and hydroxide ($OH^-$), but this is negligible.
• $NH<em>3 + H</em>2O \rightleftharpoons NH<em>4^+ + OH^-$; $K</em>b \approx 10^{-5}$ for this reaction.
• $K_h$ for the hydrolysis reaction is much greater, so the hydrolysis dominates.

Basic Anions

• Sodium acetate ($NaC<em>2H</em>3O<em>2$) is a basic solution because acetate ($C</em>2H<em>3O</em>2^-$) is the conjugate base of the weak acid acetic acid ($HC<em>2H</em>3O_2$).
• Acetic acid is a weak acid (small $K_a$), so its conjugate base is more basic than the acid is acidic.
• Acetate undergoes hydrolysis:
  • $C<em>2H</em>3O<em>2^- + H</em>2O \rightleftharpoons HC<em>2H</em>3O_2 + OH^-$
• The reaction breaks a covalent bond in water, ripping a proton away, to make excess hydroxide ions ($OH^-$) in solution, so the pH is greater than 7.

Role of Molecular Acetic Acid

• The molecular acetic acid ($HC<em>2H</em>3O_2$) formed will also react with water to generate a little hydronium, but it's negligible compared to the hydroxide generated by the hydrolysis reaction.

Hydrolysis Revisited

• Cations donate $H^+$ to water to make excess hydronium and pH less than 7 if they are acidic (conjugate acid of a weak base).
• Example: Ammonium is acidic because it's conjugate to the weak base ammonia.
• If sodium can't even react with free hydroxide, it won't break a covalent bond in water to hydrolyze it.

Anions

• If the anion ($A^-$) is basic (conjugate base to a weak acid), it will break an H-O bond in water to make excess hydroxide ($OH^-$), resulting in a pH greater than 7.

Example: Fluoride vs. Chloride

• Chloride is not basic because it's conjugate to the strong acid hydrochloric acid (HCl).
• Fluoride is basic because it's the conjugate base to the weak acid hydrofluoric acid (HF).
• $HF \rightleftharpoons H^+ + F^-$; $K_a$ is much less than 1, so $F^-$ is very basic.
• Fluoride breaks the H-O bond in water (hydrolysis) to make HF plus hydroxide.
• The excess hydroxide in solution causes the pH to be greater than 7.

Neutral Salts

• Ions are conjugates to strong acids and bases.
• When sodium chloride (NaCl) is dissolved in water, it dissociates into its ions, does not hydrolyze with water, and the pH remains neutral.

Sample Test Question

• Predict whether each of the following salt solutions will be acidic, basic, or neutral.
  • You do not have to memorize which acids are strong or weak; this information is provided.
  • Example: Is the pH going to be less than seven (acidic), greater than seven (basic), or approximately equal to seven (neutral)?

Example A: Potassium Iodide (KI)

• $KI \rightarrow K^+ + I^-$
• $K^+$ is the conjugate acid to a strong base, so it is not acidic and does not affect pH.
• $I^-$ is the conjugate base to the strong acid HI, so it is also neutral.
• Therefore, letter A is neutral.

Example B: Potassium Fluoride (KF)

• $KF \rightarrow K^+ + F^-$
• $K^+$ is neutral.
• $F^-$ is the conjugate base to the weak acid hydrofluoric acid, so it is a basic salt.
• Therefore, letter B is basic, and the pH will be greater than 7.

Hydrolysis Reaction for $F^-$

• $F^- + H_2O \rightleftharpoons HF + OH^-$
• Excess hydroxide makes it basic.

Example C: Ammonium Bromide ($NH_4Br$)

• Ammonium is the conjugate acid to ammonia; ammonia + H+ = ammonium.
• Ammonia is a weak base, so ammonium is an acidic cation.
• Bromide is the conjugate base to the strong acid HBr, so it is a neutral anion.
• The pH will be less than 7 (acidic).

Example D: Methyl Ammonium Chloride ($CH<em>3NH</em>3Cl$)

• $CH<em>3NH</em>3^+$ is the conjugate acid to $CH<em>3NH</em>2$.
• $CH<em>3NH</em>2$ is a nitrogenous base, so $CH<em>3NH</em>3^+$ is an acidic cation.
• Chloride is conjugate to HCl which is strong, so we have another acidic salt two in a row.

Example E: Lithium Formate ($LiCHO_2$)

• Lithium is the conjugate to a strong base, so it is not acidic.
• Formate ($CHO_2^-$) is the conjugate base to this weak acid HCHO2.
• Remember how these organic acids work: the hydrogen is always written out front.
• A basic anion with a neutral cation = a base: pH greater than seven or basic.

Example F: Sodium Benzoate ($C<em>6H</em>5COONa$)

• Sodium does nothing; that's the conjugate to a strong base.
• Benzoate is the conjugate to $C<em>6H</em>5COOH$. This is an organic acid (also seen written as COOH to the right).
• This is the conjugate base of a weak acid, which means that it's basic and the pH will be greater than seven.

What If…?

• If given ammonium fluoride, ammonium is acidic (conjugate to a weak base), and fluoride is basic (conjugate base to a weak acid).
  • They essentially neutralize each other, but it will lean one way or the other based on the relative value of $K_a$ of their conjugates.
  • The higher the $K_a$ of the conjugate, the weaker the acidity or basidity of this conjugate.

More Examples

Other Sections

The Stronger the Acid, the Weaker the Conjugate Base

• $K<em>w = K</em>a \times K_b$
• $K<em>w$ is equal to $K</em>a$ for a weak acid times $K_b$ for the conjugate base of that acid.
• Example: If you had $K<em>a$ for hydrofluoric acid and $K</em>b$ for the hydrolysis of fluoride anion, those two Ks when multiplied together equal Kw which is a constant.
• Kb for ammonia times K hydrolysis for ammonium, when multiplied together, give us Kw.
• As one direction is stronger, the other is weaker.

Molecular Structure vs. Acid Strength

• Skip this. This would be just something you'd have to memorize.

Acid Base Properties of Salts

• There are some odd exceptions that the only way you can know them would be to memorize them.
• Some hydrates of small cations are acidic.
• Some ions can form acid or base upon hydrolysis.
• The way you will determine whether a salt is acidic or basic is based on the strength of its conjugate acid or base as we just went over.

Acid-Base Properties of Hydroxides and Oxides

• Most oxides, if you dissolve them with water, they'll react to make either an acid or a hydroxide (a base).
• Example: Carbonated beverages. If you suffer from acid reflux, the doctor tells you, well don't drink soda water, drink still water.
• And that's because the carbon dioxide reacts with the water to make carbonic acid.
• Carbonic acid then goes into this equilibrium. It's a diprotic acid.
• As the concentration of carbon dioxide increases in the atmosphere, since I've been teaching about this in the coral reef ecology class, the amount of CO2 that's being absorbed by the oceans has been greater than scientists predicted maybe, like, twenty years ago.
• So that's good news for the environment because the air the, I'm sorry. Good news for the atmosphere.
• The air may not warm so much, but it's really bad news for the ocean and its creatures because the acidification of the oceans affect anything that has limestone or calcium carbonate, things like coral reefs, shellfish that have hard shells.
• Their formation is being stunted by acidification.

Lewis Acids and Bases

• Another definition of acidity and basicity. Instead of it has to do with donating and accepting, but instead of the h plus ion, it's the pair of electrons.
• Lewis acids accept a pair of electrons; Lewis bases donate a pair.
• Standard reaction: Ammonia + H+.
• Ammonia is a Bronsted base because it accepts the H plus cation from an acid or from water for that matter.
• Lewis acids accept a pair of electrons. Lewis bases donate a pair.
• Why does it accept H plus? Because it has this lone pair of electrons that can form what is called a dative bond, d a t i v e, which is a covalent bond in which one atom provides both electrons.
• Well, H plus of course is just a free proton. It doesn't have any electrons. So it forms the covalent bond right here.
• I need to add my plus sign outside of my dot structure before you guys call me on it, right?
• So, if you see the blue, the pair of electrons that were here as a lone pair now form this covalent bond.
• The green hydrogen is the one from the H plus So H plus we know is acidic, but what we say is the H plus accepted this pair of electrons to become a bond. So it's a Lewis acid.
• Ammonia donated this pair of electrons to make this bond, and therefore it is a Lewis base.