pH & pOH Masters
A solution has [OH⁻] = 4.0 × 10⁻⁴ M. Without directly calculating pH, explain whether the solution is acidic or basic and why.
Basic, because [OH⁻] > 1.0 × 10⁻⁷ M.
Which acid is stronger and why: Ka = 1.0 × 10⁻³ or Ka = 1.0 × 10⁻⁶?
1.0 × 10⁻³; larger Ka → more ionization.
Predict whether KCN forms an acidic or basic solution and justify your reasoning.
Basic; CN⁻ is the conjugate base of a weak acid (HCN) and reacts with water to form OH⁻.
Strong base + weak acid → Basic solution
At the half-equivalence point, why does pH = pKa?
Because [A⁻] = [HA], making log(1) = 0 in the Henderson–Hasselbalch equation.
A 0.10 M acid has [H₃O⁺] = 1.0 × 10⁻³ at equilibrium. Calculate percent ionization.
(1.0×10⁻³ / 0.10) × 100 = 1.0%
A species reacts with H⁺ by forming a bond. Explain why this automatically makes it both a Brønsted–Lowry base and a Lewis base.
Accepting H⁺ requires donating an electron pair. The bond formation is the donation of electrons to a proton. → Lewis base; accepting H⁺ → Brønsted–Lowry base.
Given Ka = 1.8 × 10⁻⁵, calculate Kb for its conjugate base.
Kb = (1.0 × 10⁻¹⁴)/(1.8 × 10⁻⁵) = 5.6 × 10⁻¹⁰
Why does NaCl not affect pH, but NH₄Cl does?
Na⁺ and Cl⁻ come from strong base/acid (no hydrolysis), while NH₄⁺ is a weak acid and reacts with water.
FROM LAST WEEK
•Cl⁻ comes from HCl (strong acid)
•Strong-acid anions do nothing in water.
•NH₄⁺ comes from NH₃ (a weak base)
•NH₄⁺ can donate a proton:
•NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺
•This produces H₃O⁺, which makes the solution acidic.
Why is Ka₁ always larger than Ka₂ in a polyprotic acid?
For a polyprotic acid (an acid that can donate more than one proton, such as H2SO4 , the first ionization constant is always larger than the second because it is electrostatically harder to remove a positively charged proton from a negatively charged ion than from a neutral molecule.
Calculate Ka for a solution where [HA] = 0.090 M and [H₃O⁺] = [A⁻] = 0.010 M.
Ka = (0.010×0.010)/0.090 = 1.1 × 10⁻³
Two solutions have pH values of 3 and 5. How many times more acidic is the first solution, and why?
100× more acidic; each pH unit represents a 10× change in [H₃O⁺]. 2 pH units = 10 * 10 = 100x more acidic
Two acids have identical concentrations, but one produces a lower pH. What can you conclude about Ka and why?
Lower pH → higher [H₃O⁺] → larger Ka → stronger acid.
Why does a buffer resist pH change when small amounts of acid are added?
The conjugate base neutralizes added H⁺, minimizing pH change.
Why is the pH at equivalence not always 7 in titrations?
The pH at the equivalence point is not always 7 because it depends on the strength of the acid and base involved and the subsequent hydrolysis of the salt produced.
In a strong acid + strong base, the salt (NaCl for example), has ions that do not react with water, keeping the pH at 7.0.
In a weak acid + strong base, the conjugate base of the weak acid produces hydroxide ions, resulting in a basic solution
•Example: NaC₂H₃O₂ (sodium acetate)
•Na⁺ comes from NaOH (strong base)
•Does nothing in water.
•C₂H₃O₂⁻ (acetate) comes from acetic acid (weak acid).
•It can react with water:
•This produces OH⁻, which makes the solution basic.
Strong acid + weak base: The conjugate acid of the weak base forms, which reacts with water to produce hydronium ions
•Example: NH₄Cl
•NH₄⁺ comes from NH₃ (a weak base)
•NH₄⁺ can donate a proton:
•NH₄⁺ + H₂O ⇌ NH₃ + H₃O⁺
•This produces H₃O⁺, which makes the solution acidic.
A solution has pH = 3.00. After adding enough water to double the volume, what is the new pH? Show reasoning.
pH= 3.30
1) Determine the initial concentration (10-3)
2) Divide concentration of H+ by 2
3) Calculate new pH by taking negative logarithm of new concentration
If you dilute an acidic solution by a factor of 10, how does the pH change?
pH increases by 1 unit because [H₃O⁺] decreases by a factor of 10.
Why does a weaker acid always have a stronger conjugate base?
Because a weak acid does not donate protons easily, so its conjugate base has a higher tendency to accept them, and therefore the conjugate base is stronger.
A buffer has equal concentrations of HA and A⁻. What does this imply about the pH relative to pKa?
pH = pKa.
Henderson-Hasselbach Equation: pH= pka + log (A-/HA)
Log (1) = 0
Why do we notice a buffer region in a weak acid-strong base / weak base-strong acid titration, but not in a strong acid - strong base titration?
A "buffer" is termed as a solution that compensates for increases or decreases in [H+] due to the reversibility of the acid<->base reaction. If an acid/base is strong, the conj. base/acid (respectively) doesn't convert back to the original compound. Therefore, there is no region where the pH changes slightly due to the effects of the conj. base/acid since it's essentially inactive after it donated or accepted H+.
At a temperature where Kw = 2.4 × 10⁻¹³, calculate the pH of neutral water.
[H₃O⁺] = √Kw = 1.55 × 10⁻⁷ → pH ≈ 6.31
Why can pH be negative, and what does that physically mean about the solution?
If [H₃O⁺] > 1 M, the -log of a number greater than 1 will be negative. The -log of a number less than 1 is positive; it indicates a very high concentration of acid.
A molecule stabilizes its conjugate base through resonance. Predict how this affects its acid strength and justify.
Increases acid strength; resonance stabilizes conjugate base → favors deprotonation.
Why does buffer capacity depend on concentration, not just ratio?
Buffer capacity depends on concentration because it measures the total amount of weak acid & conjugate base available to neutralize added strong base or acid. While the ratio of components determines the pH, the absolute concentration determines how many moles of H+ or OH- can be added before the buffer is depleted and pH changes significantly
A 50.0 mL sample of 0.100 M HCl is titrated with 0.100 M NaOH. What is pH after 25.0 mL NaOH is added?
pH= 1.48
1) Initial moles of HCl= 0.0500 L x 0.100 M = 0.00500 moles
2) Moles of NaOH added = 0.0250 L x 0.100 M = 0.00250 moles
3) NaOH neutralizes an equal amount of HCl --> 0.00500 moles- 0.00250 moles= 0.00250 moles of HCl remaining
4) Concentration = moles/ liters. Finding liters= total volume after mixing = 50 mL + 25 mL= 75 mL = 0.0750 L
0.00250 moles/ 0.0750 L= 0.0333 M
5) Calculate pH = -log (0.0333)= 1.48
Mix 50.0 mL of pH 2.00 solution with 50.0 mL of pH 4.00 solution. What is the final pH?
Final pH= 2.2967
For pH 2.00, H+ = 0.01 M
For pH 4.00, H+ = 0.0001 M
Determine moles of H+ in each solution
Solution 1: 0.01 mol/L * 0.0500 L= 0.0005 moles
Solution 2: 0.0001 mol/L * 0.0500L = 0.000005 moles
Total moles = 0.000505 moles
Total volume: 100 mL
Final H+ concentration = 0.00505 M
Final pH= 2.2967