Step by Step
Acid
Acids — donating H+, lowering pH
pH measures hydrogen ion (H+) concentration — the lower the pH, the more acidic and the more H+ present. The scale is logarithmic, meaning pH 6 has 10 times more H+ than pH 7. Acids donate H+ (examples include HCl, lactic acid, and carbonic acid).
Base
Bases — accepting H+, raising pH
Bases accept H+ (examples include NaOH and bicarbonate, HCO3-).
7.4
Blood pH — a narrow, tightly controlled range
Normal blood pH sits between 7.35 and 7.45 — slightly alkaline. Acidosis (pH below 7.35) can cause proteins to denature and enzymes to stop functioning, leading to death if severe. Alkalosis (pH above 7.45) can cause muscle tetany and cardiac arrhythmias.
Buff
Buffer systems — preventing drastic pH swings
Buffer systems resist drastic pH changes: the bicarbonate buffer acts fastest, alongside protein and phosphate buffers. For longer-term control, the respiratory system (adjusting CO2 levels) and the renal system (adjusting HCO3- levels) both play a role.
A patient in diabetic ketoacidosis develops a dangerously low blood pH as acidic ketone bodies accumulate — the body's buffer systems and rapid, deep breathing (blowing off CO2) both work to compensate, but severe, uncorrected acidosis can ultimately cause proteins to denature and enzymes to stop functioning properly.
Applied Walkthrough
1
A patient in diabetic ketoacidosis is breathing unusually fast and deep, and their blood pH is found to be below the normal 7.35-7.45 range.
2
Ask: why would rapid, deep breathing occur in this situation, and how does it relate to blood pH? The body is trying to compensate for the acidosis by blowing off CO2 through increased breathing — since CO2 combines with water to form carbonic acid in the blood, removing CO2 faster helps reduce the acid load and push pH back toward normal.
3
This illustrates the respiratory system's role as one of the body's long-term pH control mechanisms, working alongside the faster-acting bicarbonate buffer system to resist the acidosis caused by the excess ketone bodies.
4
If this compensation isn't sufficient and pH continues to fall, the consequences become serious — proteins denaturing and enzymes losing function, which is why severe, uncorrected acidosis is a medical emergency.
Exam Application
Exams test the definition of pH as a logarithmic measure of H+ concentration, the normal blood pH range (7.35-7.45) and what acidosis versus alkalosis means clinically, and the buffer systems (bicarbonate, protein, phosphate) plus the respiratory and renal systems' roles in long-term pH regulation.
⚠ Common Trap
The most common trap is forgetting that the pH scale is logarithmic, not linear — a difference of just one pH unit represents a 10-fold difference in H+ concentration, which is why even a seemingly small pH change (like from 7.4 to 7.0) represents a genuinely dramatic shift in acidity.
✓ Quick Self-Check
1. What does pH measure, and how does the scale behave (linear or logarithmic)?
It measures hydrogen ion (H+) concentration; the scale is logarithmic, so each one-unit change represents a 10-fold change in H+ concentration.
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2. What is the normal blood pH range, and is it acidic or alkaline?
7.35-7.45, which is slightly alkaline.
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3. What is acidosis, and what can it cause if severe?
A blood pH below 7.35; if severe, it can cause proteins to denature and enzymes to stop functioning, potentially leading to death.
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4. What is alkalosis, and what can it cause?
A blood pH above 7.45; it can cause muscle tetany and cardiac arrhythmias.
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5. What is the fastest-acting buffer system in the blood, and what two organ systems provide longer-term pH control?
The bicarbonate buffer system acts fastest; the respiratory system (via CO2) and renal system (via HCO3-) provide longer-term control.
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