Acid Base Intro Flashcards
normal pH, viable range, why is it necessary to regulate pH
Normal = 7.40 (7.36-7.44)
(Barely) Viable range = 6.90 - 7.80
Why is it necessary to regulate pH?
Optimum pH for protein structure – enzymes that regulate virtually every biological process require a narrow pH range. If outside normal range, reactions proceed too slowly to permit physiologic reactions sufficient to sustain life.
strong and weak acids and bases; how they ionize and what they release, examples
- Strong acids (HCl, H 2SO4) ionize more completely and release more H+
- Weak acids (H 2CO3, HAc) ionize less completely and release fewer H+
- Strong bases (NaOH) ionize more completely and release more OH -
- Weak bases (HCO3-) ionize less completely and release fewer OH -
factors causing blood pH changes; volative vs non volatile
-Changes in Levels of Volatile Acids (leads to disturbances in the body’s natural buffer systems):
CO2 + H2O <-> H2CO3 <-> HCO3- + H+
If CO2 accumulates ® acid load on body. If it is lost excessively ® base load on body
Changes in non-Volatile (fixed acids/bases):
* Byproducts of protein, lipid, carbohydrate metabolism
* Sulfuric and phosphoric acid, antacids, ketone bodies, lactic acid
how does body defend body fluid pH (maintain within a small range) (3)
- Buffers
– Instantaneous - Respiratory system
– Rapid – changes in breathing - Kidneys
– Long term regulation of pH: Conservation or excretion of HCO3-
– Excess H + ions excreted in urine
general concept for buffer systems; combination, general equations, what they prevent, consist of what
Strong acid + Buffer salt ® Neutral salt + Weak acid
- General equation: H+ + A- –> HA
- Buffers prevent drastic changes in pH
- Typically consist of mixtures of weak acids and their salts
physiological buffers; most important, other details
- Most important extracellular buffer system is HCO3-/CO2
- Hemoglobin most important intracellular buffer
- Phosphates supplement HCO3-/CO2 system in ECF
- Proteins are important supplementary buffers, especially within cells.
- Carbonate in bone – buffering in chronic states of acidosis. An extra source for HCO3-
physiological buffers (2 systems and their details)
Bicarbonate (CO2/HCO3- ) System
* Most important extracellular buffer. Bicarbonate ion converts a strong acid to a weak acid (“binds” the hydrogen ion so it doesn’t contribute to H+ level)
* Carbonic acid converts a strong base to a weak base (HCO3- ).
Phosphate System
* The monohydrogen phosphate ion converts a strong acid to a weak acid
* The dihydrogen phosphate ion converts a strong base to a weak base
physiologic buffers; protein buffer systems, importance and details
Protein Buffer Systems – Also an important part of the fast buffering systems.
* NH 3+ group on protein releases a H+ ion in the presence of excess base
* COO- group on protein accepts a H+ ion in the presence of excess
acid
Through longer term equilibrium of these “fixed” buffers with the CO2/HCO 3- buffer system, H + ions can then be excreted or retained by the lung and kidneys.
HCO3-/CO2 system; equation, specific reasons
CO2 + H2O <-> H2CO3 <-> H+ + HCO3-
- Special features
– Large amount of HCO3- and CO2 are present in blood
– “Open” system – regulated by both kidneys (HCO3- ) and lungs (CO2)
– Presence of CA allows rapid equilibrium between CO2 and H 2CO3
– Components can be measured to diagnose acid-base disturbances
– Works effectively with Hemoglobin which plays an important role by
binding both CO2 and H+ as required (based on pH conditions).
– Optimum buffering capacity is near physiological pH (pKa is 6.1)
Titration Curve for the Bicarbonate Buffer System Defines a pKa Value (where buffering activity is optimal)
The pH at which 50% of the buffer is dissociated. It is the point at which the system can buffer maximally against pH change. Since increased H+ is the most common pH change physiologically the HCO3-/CO2 system is an optimal buffer in this range
what do we need to know about the henderson hasselbach equation
You do not need to know how to do this, or apply it to numbers. You do need to know what the equation allows you to calculate and the principle of how it can be applied to calculate blood pH, CO2 or HCO3- levels. (tells you pH)
For any buffer system: H+ + A- <-> HA
hemoglobin and HCO3-/CO2 cooperate as buffer systems
blood transport of CO2
HCO3- = 65%,
dissolved CO2 = 8%,
hemoglobin associated = 27%
what is body fluid pH regulated by, how does the body use CO2/HCO3-, what happens when saturated
Body fluid pH is regulated by ventilation and renal function
The body uses CO2/HCO3- to keep pH in the normal range then regulates HCO3- and CO2 concentrations with the lungs and kidneys to make sure they are available as to buffer against further pH change.
When the body saturates these systems (by generating too much acid or base) or when the systems can’t be regenerated (loss of HCO3- by the kidneys, diarrhea etc., change in CO2 concentrations due to respiratory problems) then acid/base
disturbances result.
how does the body deal with an acid/base load; 1st line of defense
H+ Ion concentration rises due to more H+ generation:
First Line of Defense - Rapid Chemical Buffering: Basically, the body uses Bicarbonate to rapidly “absorb” the
excess (or deficiency) of H+ ions.
This results in the formation of a Weak Acid (H2CO3 ) that has a pKa is
much closer to physiologic pH than the dissolved H+ ions in the blood
acid/base load second line of defense; fast respiratory component
The Carbonic Acid (H 2 CO3 ) generated by the previous buffering reaction dissociates into CO2 and water.
Alveolar ventilation normally maintains pCO2 at 40 mm Hg. This rapidly removes the extra CO2 that was generated in the “Buffering Step”.
12 H2CO3 –> 12 CO2 + 12 H2O
In addition, the change in the blood pH stimulates increased alveolar ventilation, decreasing the pCO2 even further (25 mm Hg) and increasing pH even more. (This is known as Respiratory COMPENSATION)
