"Why Does Your Blood Stay at pH 7.4 — Even After Drinking Lemonade?"
Your blood is pH 7.4 right now. It was the same an hour ago. It will be the same after lunch. One chemical system is responsible. Today you meet it.
An acid is an H⁺ (proton) donor. When dissolved in water, acids release H⁺ ions, creating an excess of H₃O⁺ (hydronium).
A base is an H⁺ (proton) acceptor. In water, bases accept H⁺ from water molecules, producing OH⁻ (hydroxide).
Every acid–base reaction creates a conjugate pair. The acid loses H⁺ → conjugate base. The base gains H⁺ → conjugate acid.
A buffer is a mixture of a weak acid and its conjugate base. It resists pH changes when small amounts of acid or base are added.
Higher concentration of H⁺ ions.
Example: 0.1 M HCl → [H⁺] = 0.1 M → pH = 1
[H⁺] = [OH⁻] = 1.0 × 10⁻⁷ M
Pure water at 25°C.
Higher concentration of OH⁻ ions.
Example: 0.1 M NaOH → pOH = 1 → pH = 13
Click an acid, then click its conjugate base. Match all 5 pairs!
Hover over or click any pH value to explore common substances and their acidity/basicity.
These are the 10 solutions you'll test in Part 1. Colors approximate the pH indicator reading.
Dissociate 100% in water. Every molecule breaks apart.
For strong acids: [H⁺] = [acid]
For strong bases: [OH⁻] = [base]
| Strong Acids | Strong Bases |
|---|---|
| HCl — hydrochloric acid | NaOH — sodium hydroxide |
| HBr — hydrobromic acid | LiOH — lithium hydroxide |
| HI — hydroiodic acid | KOH — potassium hydroxide |
| H₂SO₄ — sulfuric acid | Ba(OH)₂ — barium hydroxide |
| HNO₃ — nitric acid | |
| HClO₄ — perchloric acid |
Only partially dissociate — an equilibrium exists between the undissociated molecule and its ions.
For weak acids: [H⁺] ≪ [acid]
Most acid molecules stay intact!
See the difference between strong and weak acid dissociation:
Click a button above to visualize dissociation
Acetic acid (CH₃COOH) is weak, so its conjugate base CH₃COO⁻ (acetate) is strong enough to accept an H⁺ from water:
This is why sodium acetate solution is basic (pH > 7).
Ammonia (NH₃) is a weak base, so its conjugate acid NH₄⁺ (ammonium) donates H⁺ to water:
This is why ammonium chloride solution is acidic (pH < 7).
A buffer is a solution that resists changes in pH when small amounts of acid or base are added. It consists of:
Acetate buffer: CH₃COOH + CH₃COO⁻ (from sodium acetate)
• Weak acid neutralizes added base
• Conjugate base neutralizes added acid
Ammonia buffer: NH₃ + NH₄⁺ (from ammonium chloride)
• Weak base neutralizes added acid
• Conjugate acid neutralizes added base
Added H⁺ reacts with the conjugate base (CH₃COO⁻), converting it back to weak acid. The H⁺ is consumed, so pH barely changes.
Added OH⁻ reacts with the weak acid (CH₃COOH), converting it to conjugate base + water. The OH⁻ is consumed, so pH barely changes.
Choose a solution, then add acid or base to see how pH responds. Compare buffered vs unbuffered!
The carbonate buffer keeps blood pH in a 0.1-unit window. Outside that window, organs stop working in a specific, predictable order.
| Blood pH | Condition | What the Body Does | If Untreated |
|---|---|---|---|
| > 7.8 | Severe Alkalosis | Muscle tetany, seizures | Cardiac arrest |
| 7.45–7.8 | Alkalosis | Muscle cramps, tingling, confusion | Worsening neurological symptoms |
| 7.35–7.45 | ✅ Normal | All systems functioning | — |
| 7.0–7.35 | Acidosis | Rapid breathing (exhaling CO₂ to compensate), confusion, fatigue | Organ failure |
| < 7.0 | Severe Acidosis | Cardiac arrhythmia, coma | Death within hours |
Compare what happens when you add 5 drops of HCl to each solution:
| Solution | Initial pH | After 5 drops HCl | ΔpH | Verdict |
|---|---|---|---|---|
| Distilled Water | 7.00 | ~3.0 | −4.0 | ⚠️ Huge swing! |
| Acetate Buffer | ~4.74 | ~4.56 | −0.18 | ✅ Barely moved |
| Ammonia Buffer | ~9.25 | ~9.07 | −0.18 | ✅ Barely moved |
You'll measure the pH of 10 solutions using pH paper (test strips).
Place them in a test tube rack. Pour about 3 mL of each solution into its own test tube.
Dip a clean glass rod into the solution. Touch the rod to the pH test strip. Use a fresh strip for each solution!
Compare the color on the strip to the chart on the pH paper box. Record the pH and classify as acidic, basic, or neutral.
You'll compare how water, an acetate buffer, and an ammonia buffer respond to added acid and base.
Tubes 1 & 2: Add 10 mL distilled water to each.
Tubes 3 & 4: Add 5 mL of 0.1 M CH₃COOH + 5 mL of 0.1 M CH₃COONa to each (acetate buffer).
Tubes 5 & 6: Add 5 mL of 0.1 M NH₄OH + 5 mL of 0.1 M NH₄Cl to each (ammonia buffer).
Use the pH meter to record the pH of all 6 test tubes. Rinse the probe with RO water between measurements!
Add 5 drops of 0.1 M HCl to tubes 1, 3, and 5. Mix well. Measure and record the new pH.
Add 5 drops of 0.1 M NaOH to tubes 2, 4, and 6. Mix well. Measure and record the new pH.
| Solution | pH | Classification | Strong/Weak/N/A |
|---|
| Test Tube | Solution | Initial pH | Addition | Final pH | ΔpH |
|---|---|---|---|---|---|
| 1 | Distilled Water | 5 drops HCl | — | ||
| 2 | Distilled Water | 5 drops NaOH | — | ||
| 3 | Acetate Buffer | 5 drops HCl | — | ||
| 4 | Acetate Buffer | 5 drops NaOH | — | ||
| 5 | Ammonia Buffer | 5 drops HCl | — | ||
| 6 | Ammonia Buffer | 5 drops NaOH | — |
| [H⁺] (M) | pH | pOH | [OH⁻] (M) | Acidic/Basic? |
|---|---|---|---|---|
| 1.0 × 10⁻¹ | 1.0 | 13.0 | 1.0 × 10⁻¹³ | Strongly Acidic |
| 1.0 × 10⁻³ | 3.0 | 11.0 | 1.0 × 10⁻¹¹ | Acidic |
| 1.0 × 10⁻⁵ | 5.0 | 9.0 | 1.0 × 10⁻⁹ | Weakly Acidic |
| 1.0 × 10⁻⁷ | 7.0 | 7.0 | 1.0 × 10⁻⁷ | Neutral |
| 1.0 × 10⁻⁹ | 9.0 | 5.0 | 1.0 × 10⁻⁵ | Weakly Basic |
| 1.0 × 10⁻¹¹ | 11.0 | 3.0 | 1.0 × 10⁻³ | Basic |
| 1.0 × 10⁻¹³ | 13.0 | 1.0 | 1.0 × 10⁻¹ | Strongly Basic |
Test your understanding before, during, or after the lab. Click an answer to check it — explanations included!