Acid/Base

1. What is pH, and how do you calculate it?

👉 pH = −log₁₀[H⁺]. It measures the concentration of free hydrogen ions in a solution.

Example:

• 1 mM HCl = 0.001 M = 10⁻³ M H⁺ (assuming 100% dissociation).

• pH = −log(10⁻³) = 3.

2. What is an acid, and what happens when it dissociates?

👉 An acid (HA) releases a hydrogen ion (H⁺), leaving behind its conjugate base (A⁻).

• HA = acid.

• H⁺ = hydrogen ion.

• A⁻ = conjugate base.

3. What is pK, and what does it tell you?

👉 pK is the pH at which an acid is 50% dissociated and 50% remains in its HA form.

• 50% HA.

• 50% A⁻.

• At pH = pK, the concentrations of HA and A⁻ are equal.

• Memory trick: pK = 50/50 point.

4. What is the difference between a strong acid and a weak acid?

👉 Strong acids dissociate almost 100%; weak acids only partially dissociate.

• Strong acid → releases nearly all its H⁺.

• Weak acid → some remains as HA, while some becomes H⁺ and A⁻.

• Lower pK = stronger acid.

5. What is a buffer, and how does it work?

👉 A buffer is a weak acid/conjugate base system that resists changes in pH by absorbing added H⁺ or base.

• When acid (H⁺) is added → A⁻ binds H⁺ to form HA.

• When base is added → HA can release H⁺ to neutralize some of the added base.

6. When does a buffer work most effectively?

👉 When the solution pH is close to the buffer's pK, generally within ±1 pH unit.

At this point, both HA and A⁻ are available to handle added acid or base.

7. What is the Henderson-Hasselbalch equation?

It calculates the pH of a buffer solution using the ratio of conjugate base to acid.

• A⁻ = conjugate base.

• HA = weak acid.

8. What happens to pH when an acid is 50% dissociated?

👉 pH = pK because the acid and conjugate base concentrations are equal.

Example: 25 mM lactic acid with 50% dissociation.

• HA = 12.5 mM.

• A⁻ (lactate) = 12.5 mM.

\[ pH=3.86+\log(12.5/12.5) \]

\[ pH=3.86 \]

9. Which phosphate conjugate acid/base pair is used to make a buffer at pH 7.40, and why?

👉 H₂PO₄⁻ / HPO₄²⁻ because its pK is approximately 7.20, close to the desired pH.

• H₂PO₄⁻ = acid (HA).

• HPO₄²⁻ = conjugate base (A⁻).

10. How do you calculate the phosphate ratio needed to make a pH 7.40 buffer?

👉 Use Henderson-Hasselbalch:

\[ 7.40=7.20+\log\frac{[HPO_4^{2-}]}{[H_2PO_4^-]} \]

\[ \frac{[HPO_4^{2-}]}{[H_2PO_4^-]}=1.585 \]

Therefore, you need approximately 1.585 parts base for every 1 part acid.

Example:

• 1.585 mM HPO₄²⁻.

• 1 mM H₂PO₄⁻.

11. What are the three pK values of phosphoric acid, and which phosphate ion predominates at pH 7.30?

Dissociation	pK
H₃PO₄ → H⁺ + H₂PO₄⁻	2.12
H₂PO₄⁻ → H⁺ + HPO₄²⁻	7.21
HPO₄²⁻ → H⁺ + PO₄³⁻	12.3

At pH 7.30, HPO₄²⁻ predominates because the pH is slightly above the second pK of 7.21.

12. What is the normal blood pH, and why must it be tightly regulated?

👉 Normal blood pH = 7.40 (range 7.35–7.45).

Enzymes, transporters, and cellular processes function best within this narrow range.

13. What are the two types of acids produced by the body?

Type	Example	How produced/handled
Volatile	CO₂	Produced by metabolism and eliminated through the lungs
Nonvolatile	Sulfuric acid, phosphoric acid, HCl	Produced through cellular metabolism and eliminated by the kidneys

14. What is a nonvolatile acid?

👉 An acid that dissociates to release H⁺ and cannot be eliminated through the lungs as a gas.

Examples:

• Sulfuric acid.

• Phosphoric acid.

• Hydrochloric acid.

Produced through metabolism of substances such as amino acids and nucleic acids.

15. Why is CO₂ considered a volatile acid even though it does not contain H⁺?

👉 CO₂ combines with water to form carbonic acid, which dissociates and releases H⁺.

\[ CO_2+H_2O\rightleftharpoons H_2CO_3 \]

\[ H_2CO_3\rightleftharpoons H^++HCO_3^- \]

CO₂ is called volatile because it is a gas that can be eliminated through breathing.

16. What is the primary buffer system in the body, and why is it effective despite its pK being far from blood pH?

👉 The bicarbonate/CO₂ buffer system.

• Its pK is 6.1.

• Normal blood pH is 7.40.

• Normally, buffers work best near their pK.

However, the bicarbonate system is an open system: the lungs independently regulate CO₂, and the kidneys regulate HCO₃⁻. This allows it to effectively regulate blood pH.

17. How do the lungs regulate blood pH?

👉 The lungs regulate blood CO₂ by changing ventilation.

• Hyperventilation → increases CO₂ removal → decreases blood CO₂ → decreases H⁺ → pH increases.

• Hypoventilation → decreases CO₂ removal → increases blood CO₂ → increases H⁺ → pH decreases.

Memory trick: CO₂ acts like acid. More CO₂ = more acidic.

18. How do the kidneys regulate blood pH?

👉 The kidneys regulate bicarbonate (HCO₃⁻) and eliminate H⁺.

• Reabsorb HCO₃⁻ → retain base.

• Secrete/excrete H⁺ → eliminate acid.

Important: The kidneys are the only organ in this lecture's framework that actually eliminates H⁺ from the body. Buffers and lungs primarily buffer or regulate acid rather than directly eliminate the nonvolatile acid load.

19. What are the three mechanisms of acid-base regulation, and how quickly do they work?

Mechanism	Main function	Speed
Body buffers	Bind added H⁺ or base	Seconds–minutes
Pulmonary (lungs)	Regulate CO₂ (acid)	Minutes–hours
Renal (kidneys)	Regulate HCO₃⁻ and eliminate H⁺	Hours–days

Memory trick: Buffers are fastest, lungs are intermediate, kidneys are slowest.

20. A patient has tuberculosis causing lung scarring and impaired CO₂ exchange. What happens to blood pH, and why?

👉 Blood pH decreases (becomes more acidic).

Reasoning:

1. Damaged lungs cannot remove CO₂ efficiently.

2. CO₂ accumulates in the blood.

3. More CO₂ produces more carbonic acid.

4. More H⁺ is released.

5. Blood pH decreases.

This causes respiratory acidosis.

21. What are the four primary acid-base disorders, and what causes each?

Disorder	Primary change	Effect
Respiratory acidosis	↑ pCO₂	More acid
Respiratory alkalosis	↓ pCO₂	Less acid
Metabolic acidosis	↓ HCO₃⁻	Less base
Metabolic alkalosis	↑ HCO₃⁻	More base

Normal values to memorize:

• Blood pH = 7.40.

• pCO₂ = 40 mmHg.

• HCO₃⁻ = 24 mM.

22. How do you identify an acid-base disorder using blood pH, pCO₂, and HCO₃⁻?

Step 1: Look at pH.

• pH below 7.40 → acidosis.

• pH above 7.40 → alkalosis.

Step 2: Look at pCO₂.

• Increased pCO₂ → respiratory acidosis.

• Decreased pCO₂ → respiratory alkalosis.

Step 3: Look at HCO₃⁻.

• Decreased HCO₃⁻ → metabolic acidosis.

• Increased HCO₃⁻ → metabolic alkalosis.

Step 4: Identify the primary problem and compensation.

Example:

Measurement	Patient	Normal
pH	7.30	7.40
pCO₂	54 mmHg	40 mmHg
HCO₃⁻	28 mM	24 mM

• pH is low → acidosis.

• pCO₂ is high → respiratory acidosis.

• HCO₃⁻ is high → metabolic compensation.

Answer: Respiratory acidosis with metabolic compensation.

Quick exam memorization sheet

Concept	Must remember
pH	−log[H⁺]
pK	50% dissociation
Lower pK	Stronger acid
Best buffering	pH close to pK
Henderson-Hasselbalch	pH = pK + log(A⁻/HA)
Normal blood pH	7.35–7.45
Bicarbonate buffer pK	6.1
Normal pCO₂	40 mmHg
Normal HCO₃⁻	24 mM
Hyperventilation	↓ CO₂, ↑ pH
Hypoventilation	↑ CO₂, ↓ pH
Kidneys	Reabsorb HCO₃⁻, excrete H⁺
Fastest regulation	Buffers
Slowest regulation	Kidneys

One important exam distinction:

The lecture uses 7.40 as the reference point for deciding acidosis versus alkalosis. Clinically, a pH below 7.35 is acidemia and above 7.45 is alkalemia; the 7.35–7.45 interval is the normal range. Also, the elevated bicarbonate in the respiratory acidosis example represents compensation, not a second primary metabolic alkalosis.