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Fluid and electrolyte balance in pregnancy
During the course of a normal pregnancy, extracellular fluid accumulates:
plasma volume rises typically from 3 to 4.5 litres
interstitial-fluid volume rises typically from 11 to 15 litres.
The intracellular fluid volume also rises minimally.
These changes account for about 6–8 kg of the weight gain experienced during pregnancy. There is an accompanying accumulation of Na+ and K+, so that there is no overall significant change in the concentrations of these ions.
Because of the accumulation of H2O in the plasma compartment during a normal pregnancy, there is a fall in the concentration of plasma albumin and therefore a fall in the colloid osmotic pressure of plasma. This results in an imbalance in the so-called Starling forces, which control the movement of fluid across the capillary membrane:
intravascular pressure pushes fluid out of the vessel
colloid osmotic pressure tends to keep it in.
As a result of the drop in colloid osmotic pressure during pregnancy, accumulation of interstitial fluid and development of oedema become more likely, especially during the third trimester when arterial pressure rises.
It should be noted that in certain disease states, most importantly pre-eclampsia, there may be a fall in plasma volume.
Acid–base balance
Some Definitions
Understanding of the acid–base balance is helped greatly by a clear understanding of the relevant terminology.
Acids and bases
An acid is a proton donor, meaning that it tends to give up H+ ions. A base is a proton acceptor, meaning that it tends to take up H+ ions. It follows that acids and bases may be anions, cations or electrically neutral.
For example, ammonium (NH4+) is an anionic acid in equilibrium with ammonia (NH3):
Dihydrogen phosphate (H2PO4−) is a cationic acid in equilibrium with hydrogen phosphate (HPO42−):
Hydrochloric acid (HCl) is a neutral acid in equilibrium with H+ and Cl−:
These are all acids because they tend to give up H+ ions.
pH value
The pH value of a solution is the negative logarithm to the base 10 of the H+ concentration and is a dimensionless quantity (it has no units). This means that the higher the H+ concentration (the more acidic the solution), the lower the pH value.
Acidaemia and acidosis
These terms are often used interchangeably in everyday clinical practice and this can be a source of confusion. Strictly speaking, acidaemia means that the blood is more acidic than normal, while acidosis refers to the underlying process (for example ketoacidosis, lactic acidosis) that may cause acidaemia. Alkalaemia and alkalosis have comparable meanings.
Acid Load
The normal pH value of the blood and extracellular fluid is about 7.4, with a range of about 7.35–7.45. Because of the logarithmic derivation of the pH value, this represents an H+ concentration in the order of nmol/l. Table 7.1 shows how physiological pH values in the normal range correspond to the concentration of H+.
| pH | H+ concentration (nmol/l) |
|---|---|
| 7.0 | 100 |
| 7.2 | 63 |
| 7.4 | 40 |
| 7.6 | 25 |
| 7.8 | 16 |
Clearly, the pH value is normally controlled within a narrow range. The importance of this must be appreciated in the context of the amount of H+ production that takes place in the body:
15 000 mmol/day from CO2 production
50–150 mmol/day from the production of phosphoric acid and sulphuric acid
1000 mmol/day from lactic acid production.
The impact of this on the overall pH value of the extra-cellular fluid is minimised by three mechanisms:
Buffers
A buffer system is a system that tends to minimise changes in the pH value of a solution. It normally consists of a weak acid with its conjugate base (this is what remains when H+ has left the acid molecule). An example would be:
Addition of further H+ ions will cause the reaction to shift towards the left, resulting in the formation of more undissociated acid because the conjugate base tends to combine with the extra H+ ions, while the reverse happens if H+ ions are removed from the system. The buffering power of a system is greatest when half the buffering groups are undissociated and half are dissociated; that is, when [R–COOH] equals [R–COO−]. The pH value at this point equals the pK value of the system, with the pK value being the negative logarithm to the base 10 of the dissociation constant, K.
The most important buffers in extracellular fluid are the bicarbonate (HCO3−) system and haemoglobin. Other proteins also contribute some buffering power.
Bicarbonate as a buffer
Carbonic acid (H2CO3) is a relatively strong acid, so the HCO3− system is a relatively weak buffer. Its importance lies in the link between H2CO3, HCO3− and dissolved CO2:
Hence the contribution of ventilation and CO2 excretion to the acid–base balance.
The relationship between the components of the HCO3− buffer system can be described by the following equation, where K is the dissociation constant of H2CO3, α is the solubility coefficient for CO2 in plasma and pCO2 is the partial pressure of CO2:
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