Acids and Bases
Acids and Bases
Definition: Acids are substances that donate hydrogen ions (H+) in solution, while bases accept them or donate hydroxide ions (OH-); the pH scale measures how acidic or basic a solution is on a range typically from 0 to 14.
How It Works
Three definitions layer on top of each other, each broader than the last:
- Arrhenius: acid increases [H+] in water, base increases [OH-]. Only applies to aqueous solutions.
- Brønsted-Lowry: acid is a proton (H+) donor, base is a proton acceptor. Explains why ammonia (NH3) is a base with no OH- to donate, it accepts a proton to form NH4+.
- Lewis: acid accepts an electron pair, base donates one. Broadest definition, covers reactions with no protons at all, like BF3 reacting with NH3.
Conjugate pairs:
- Every Brønsted-Lowry acid-base reaction produces a conjugate pair: when HA donates a proton, it becomes A-, its conjugate base.
- The base that accepted the proton becomes its conjugate acid.
- Strong acids dissociate essentially completely in water: HCl, HNO3, H2SO4, HBr, HI, HClO4.
- Weak acids only partially dissociate, establishing an equilibrium: acetic acid, carbonic acid, HF.
Water autoionization:
- Water is amphoteric, it can act as both acid and base.
2H2O ⇌ H3O+ + OH-- At 25°C, the autoionization equilibrium constant is Kw:
Kw = [H+][OH-] = 1.0 × 10⁻¹⁴- This is why pH + pOH = 14 in any aqueous solution at 25°C.
Polyprotic acids:
- Some acids donate more than one proton, in sequential, increasingly difficult steps.
- Sulfuric acid (H2SO4) loses its first proton almost completely, a strong acid in that step.
- The resulting HSO4- is only moderately weak for the second proton.
- Each successive Ka is smaller because pulling a proton from an increasingly negative species takes more energy.
- Phosphoric acid (H3PO4) has three distinct Ka values:
Ka1 ≈ 7.5 × 10⁻³
Ka2 ≈ 6.2 × 10⁻⁸
Ka3 ≈ 4.2 × 10⁻¹³- Each is about five orders of magnitude smaller than the last.
- This stepwise behavior is what makes phosphate useful as a biological buffer across a wide pH range.
Under the Hood
Core pH definitions:
pH = -log[H+]
pOH = -log[OH-]
pH + pOH = 14 (at 25°C)- Because the scale is logarithmic, each unit change means a tenfold change in H+ concentration.
Weak acid dissociation:
Ka = [H+][A-]/[HA]
pKa = -log(Ka)
Ka × Kb = Kw = 1.0 × 10⁻¹⁴- A lower pKa means a stronger acid.
Worked example 1: strong acid.
- Given: 0.010 M HCl, dissociates completely
- Step 1: [H+] = 0.010 M
pH = -log(0.010)- Answer: pH = 2.00
Worked example 2: weak acid.
- Given: 0.100 M acetic acid, Ka = 1.8 × 10⁻⁵
- Step 1: set up the equilibrium expression
CH3COOH ⇌ H+ + CH3COO-
Ka = x²/(0.100 - x) ≈ x²/0.100- Step 2: solve for x
x² = (1.8 × 10⁻⁵)(0.100) = 1.8 × 10⁻⁶
x = [H+] = 1.34 × 10⁻³ M- Step 3: convert to pH
pH = -log(1.34 × 10⁻³)- Answer: pH = 2.87
- Note: only about 1.3% of the acetic acid molecules ionize, far less acidic than the strong acid at the same concentration.
Worked example 3: strong base.
- Given: 0.0050 M NaOH
- Step 1: [OH-] = 0.0050 M
pOH = -log(0.0050) = 2.30
pH = 14.00 - 2.30- Answer: pH = 11.70
Worked example 4: weak base.
- Given: 0.050 M ammonia, Kb = 1.8 × 10⁻⁵
- Step 1: set up the equilibrium expression
NH3 + H2O ⇌ NH4+ + OH-
Kb = x²/(0.050 - x) ≈ x²/0.050- Step 2: solve for x
x² = (1.8 × 10⁻⁵)(0.050) = 9.0 × 10⁻⁷
x = [OH-] = 9.49 × 10⁻⁴ M- Step 3: convert to pH
pOH = -log(9.49 × 10⁻⁴) = 3.02
pH = 14.00 - 3.02- Answer: pH = 10.98
Why It Matters
- pH governs whether biological and industrial processes can proceed at all.
- Enzymes have narrow pH ranges where they fold and function correctly.
- Pepsin works around pH 2 in the stomach; trypsin needs the near-neutral small intestine.
- Acid-base chemistry drives fertilizer production, water treatment, and food preservation.
- It also drives the etching processes used in semiconductor manufacturing.
- Ocean acidification, driven by CO2 dissolving into carbonic acid, is measurably shifting seawater pH.
- This threatens calcium-carbonate-shelled marine organisms.
pH measurement in practice:
- Indicators are weak acids or bases whose ionized and neutral forms have different colors.
- Litmus turns red below about pH 4.5 and blue above about pH 8.3.
- Phenolphthalein is colorless below pH 8.2 and pink above pH 10.
- A pH meter measures the voltage across a glass membrane, generated by the [H+] difference between sample and internal reference.
- That voltage converts to a pH reading via the Nernst equation, giving a continuous numeric value.
Common Pitfalls
- Confusing “strong” with “concentrated”: a strong acid dissociates completely, a concentrated acid just has a lot of it dissolved. Dilute HCl is still a strong acid.
- Forgetting pH is logarithmic: pH 3 is ten times more acidic than pH 4, not three-quarters as acidic.
- Applying the “x is small” approximation in weak acid ICE tables when Ka is too large relative to concentration. Rough validity check: if Ka/C < 10⁻⁴, it’s safe to use.
- Mixing up Ka and pKa, or forgetting the negative sign when converting between them.
- Assuming neutralization always gives pH 7: a strong acid + weak base gives an acidic salt solution, a weak acid + strong base gives a basic one, because the conjugate species hydrolyzes water.
- Treating polyprotic acids as if all protons dissociate equally; usually only the first step matters.
- Reporting pH to more decimal places than the underlying [H+] measurement justifies; because pH is a logarithm, only digits after the decimal point count as significant figures.
Comparison
| Definition | Acid | Base | Scope |
|---|---|---|---|
| Arrhenius | Increases [H+] in water | Increases [OH-] in water | Aqueous only |
| Brønsted-Lowry | Proton (H+) donor | Proton acceptor | Any proton-transfer reaction |
| Lewis | Electron pair acceptor | Electron pair donor | Broadest, includes non-proton reactions |
| Conjugate acid | Forms when a base accepts H+ | — | Brønsted-Lowry framework |
| Conjugate base | Forms when an acid donates H+ | — | Brønsted-Lowry framework |
| Property | Strong Acid | Weak Acid |
|---|---|---|
| Dissociation | ~100% | Partial, governed by Ka |
| Example | HCl, H2SO4 | CH3COOH, H2CO3 |
| Conducts electricity | Strongly | Weakly |
| pH at given concentration | Lower (more acidic) | Higher |
| Equilibrium expression needed | No | Yes |
| Reaction with active metals | Vigorous, fast | Slower, less vigorous |
| Everyday Substance | Approximate pH | Classification |
|---|---|---|
| Battery acid | 0-1 | Strong acid |
| Stomach acid | 1.5-2 | Strong acid (dilute HCl) |
| Lemon juice | 2 | Weak acid |
| Coffee | 5 | Weak acid |
| Pure water | 7 | Neutral |
| Blood | 7.35-7.45 | Weakly basic (buffered) |
| Baking soda solution | 9 | Weak base |
| Ammonia cleaner | 11-12 | Weak base |
| Drain cleaner (NaOH) | 13-14 | Strong base |
Example
Antacid tablets like Tums use calcium carbonate to neutralize excess stomach acid:
CaCO3(s) + 2HCl(aq) → CaCl2(aq) + H2O(l) + CO2(g)- The reaction consumes H+ ions directly, raising the stomach’s pH.
- This relieves the burning sensation of acid reflux.
- The fizzing some antacids produce is the CO2 gas escaping.
Real-World Application
Industrial wastewater treatment often requires neutralizing acidic effluent before it can legally be discharged.
- Given: 10,000 L of wastewater at pH 2.00 (treated as a strong acid), neutralize to pH 7 using Ca(OH)2, molar mass 74.09 g/mol
- Step 1: find [H+]
[H+] = 10⁻²·⁰⁰ = 0.0100 M- Step 2: find total moles of H+
mol H+ = 0.0100 mol/L × 10,000 L = 100 mol- Step 3: Ca(OH)2 supplies two OH- per formula unit
mol Ca(OH)2 needed = 100 mol H+ / 2 = 50 mol- Step 4: convert to mass
mass = 50 mol × 74.09 g/mol = 3,705 g- Answer: about 3.7 kg of Ca(OH)2 needed
- Scaled to the millions of liters a real facility processes daily, this is why neutralizing agents are a significant ongoing line item in a treatment plant’s budget.
FAQ
Can a substance be both an acid and a base?
- Yes, this is called being amphoteric, or amphiprotic for proton-transfer specifically.
- Water is the classic example.
- Bicarbonate (HCO3-) is another: it can donate a proton to become CO3²- or accept one to become H2CO3.
Does a higher concentration always mean a lower pH?
- Only for a given acid.
- Comparing across different acids, a dilute strong acid can have a lower pH than a concentrated weak acid.
- Dissociation, not just concentration, controls [H+].
Why do pKa values matter for drug design?
- A drug’s pKa determines what fraction is ionized at a given pH, via the Henderson-Hasselbalch equation.
- Ionized versus neutral forms cross cell membranes at very different rates.
- This directly affects absorption and how the drug distributes through the body.
Related Terms
Referenced by