Electrolytes are substances whose molecules dissociate into ions
when placed in a solution. An ion is an atom or group of atoms with an
electrical charge. Those electrolytes with a positive charge are cations and
those with negative charge are anions. The total number of cation is balanced
by an equal number of anions.
Cations mg% MEq/L
Sodium 326 142
Potassium 20 5
Calcium 10 5
Magnesium 2.4 2
Total 358.4 154
Anions mg% MEq/L
Bicarbonates 60.5 27
Chlorides 365.7 108
Phosphates 3.4 2
Sulfates 1.6 1
Proteinates 6500 16
Total 6948.7 154
The electrolytes in the body in general are extremely small so
that they are expressed in terms of milliequivalent weight. You will also note
that the total cation measured in milliequivalent weight per liter is equal to
the total anion but not when they are measured in milligram percent.
Functions of electrolytes
1. Maintenance of osmotic pressure and hydration
2. Maintenance of proper body pH.
3. Regulation of the proper function of the
heart and muscles.
4. Involvement in oxidation–reduction (electron
transfer)
5. Participation as an essential part of
co–factor of enzymes.
Distribution of electrolytes
1. Interacellular fluid (ICF) – that which is contained within
the cells, represents about 66% of total body water.
2. Extracellular fluid (ECF) – that which is outside the
cells, represents about 34% of total body water.
a. Interstitial fluid – immediately surrounds the
cells and is separated from ICF by the cell membrane.
b. Intravascular fluid – separated from interstitial
fluid by capillary wall.
************ SODIUM
************
Sodium is an important electrolyte involved in the distribution of
water between extra and intracellular fluid. It is also involved in the
maintenance of acid–base equilibrium as it makes up 90% of the total
extracellular ions.
The total amount of sodium in the body is 80 grams. About 44% is
extracellular and 35% is found in the bones. About 82% of the total sodium is
inexchangeable sodium and 18% of the total sodium is inert and found mostly in
the bones.
Sodium in interstitial fluid is in equilibrium with that of serum.
The pH of serum can alter the distribution of interstitial sodium and
intercellular potassium. At a pH 7.5, sodium and hydrogen ions move into the
cells while potassium leaves the cells.
The adrenal glands control the metabolism of sodium. Adrenal
steroids (aldosterone) cause a decrease in the urinary output of sodium and
facilitate renal excretion of water. Excess steroids increase the intracellular
sodium and decreases intracellular potassium.
A decrease in adrenal steroids is usually followed by an increase
loss of NaCl in the urine and maybe associated with a shift of sodium to other
sites. Sodium is excreted through the kidneys. Small amounts of sodium appear
in the feces and sweats.
Other hormones which regulate sodium are: anti–diuretic hormone
(ADH) and atrial natriuretic peptide (ANP)
Methods of determination
1. Precipitation as pyroantimonate and
quantitation by gravimetric method.
2. Precipitation as sodium manganous uranyl
acetate followed by oxidation with periodate and photometric quantitation of
the permanganate ion produced.
3. Precipitation of sodium magnesium uranyl
acetate which is then dissolved in alkali and photometrically measured.
4. Precipitation as sodium uranyl zinc acetate
followed by quantitation of the precipitate.
5. Photometric measurement of sodium violurate
which requires ashing of the precipitate.
6. Isolation of the sodium into a cationic
exchange resin followed by removal with barium chloride, precipitation of the
barium ion in the eluate, conversion of the NaC to NaOH on the anion resin and
finally titration of NaOH.
7. Precipitation of the sodium salt of
L–methoxyphenyl–acetic acid which is washed and titrated with alkali.
8. Determination of the change in erythrocytic
volume after equilibrium with 2 known concentration of NaCl solution.
9. Neutron activation analysis
10. Flame
photometry
11. Ion
selective electrode
Trindler method
Sodium is added to an alcoholic solution of magnesium uranyl
acetate. Sodium is precipitated as sodium magnesium uranyl acetate and the
proteins are precipitated by ethyl alcohol. The precipitated protein and
precipitated sodium salts are removed by centrifugation and the excess uranyl
in the supernatant fluid are read against a blank that contains a known
concentration of uranyl
Albanese and Lein method
Sodium is precipitated as sodium uranyl zinc acetate which is then
dissolved in water and determined photometrically by its yellow color.
Flame photometry
Serum is diluted 1:200. The solution is subjected to a
non–luminous flame emitting light with a characteristic wavelength for sodium
and the intensity of light emitted is measured, which is directly proportional
to the concentration of sodium in the sample.
Precaution in sodium determination
1. Sodium excretion decrease during exercise
2. Anticoagulants lowers sodium levels like
oxalates which tend to increase the plasma volume.
3. Unpurified distilled water contains traces of
sodium ions which will erroneously elevate sodium levels.
4. Mixing tube using thumbs as cover causes an
elevation of sodium levels since NaCl is present in the skin.
5. Glasswares maybe contaminated with oxide of
sodium
Normal values: 135
– 145 mEq/L
Clinical significance
A. Hyponatremia – means abnormally low plasma Na+.
It is a reflection of the ratio of Na+ to volume in plasma and tells
nothing directly about the total Na+ content.
Classification of hyponatremia
1. Depletional hyponatremia
a. Renal losses – come from the use of diuretics
and hypoaldosteronism
(1) Diuretics
are useful in the treatment of edema because they block Na+ or Cl
reabsorption in the nephron, resulting in an increase urinary excretion of Na+
and water.
(2) Hypoaldosteronism refers to impairment in the
renin–angiotensin– aldosterone axis due to either deficient aldosterone or
renin secretion.
(a) In primary
hypoaldosteronism, there is
deficiency of aldosterone because of some defects in its synthesis within the
adrenal cortex. With a deficiency of aldosterone, sodium and its anions are
lost in urine.
(b) Secondary
hypoaldosteronism refers to
deficient aldosterone production due to renin deficiency resulting from damage
to the JGA. Without adequate renin secretion, aldosterone production is not
stimulated and Na+ is not adequately reabsorbed.
(c) Addison’s disease is an adrenocortical
insufficiency due to destruction of adrenal tissue, most commonly from
autoimmune destruction. There is a deficiency of both mineralocorticoid and
glucocorticoid hormones. As a result, maximal sodium reabsorption cannot occur.
b. Non–renal losses
(1) Gastrointestinal
diarrhea
(2) Skin burns
and trauma
2. Dilutional hyponatremia – results from conditions in
which there is a greater proportion of water to sodium than normal so that it
appears as if sodium concentration is low.
a. Syndrome of Inappropriate Anti–Diuretic
Hormone (SIADH) – an
excessive amount of ADH is produced regardless of the plasma osmolality and
water needs of the body. The increased ADH leads to increased water retention,
resulting in a mild hypo osmolality and hyponatremia. SIADH may be due to many
clinical disorders, especially malignant tumors with ectopic production of ADH,
CNS disturbances and head injuries that damage the osmoreceptors in the
hypothalamus.
b. Generalized edema occurs in patients who have
a markedly increased level of total body sodium. These patients also have a
severe defect in water excretion so that water retention is proportionately
greater than sodium retention with resultant hyponatremia and edema. This
occurs in congestive heart failure, cirrhosis, and nephrotic syndrome.
c. Glucose, when present in high concentration
in plasma, exerts a significant osmotic force, causing a movement of water out
of the cells until osmotic equilibrium is restored. It is estimated that the
plasma Na+ decreases by 1.6 mmol /L for each 100 mg / dl increment
in blood glucose.
3. Artificial or pseudohyponatremia occurs in condition in which the
volume of plasma contains a significant amount of non–aqueous components such
as triglycerides or proteins.
(1) Pseudohyponatremia
due to elevated triglycerides can occur in such diseases as diabetes mellitus,
nephrotic syndrome and some types of cirrhosis of the liver. However, the
triglyceride level must be 1500 mg/dl or higher to cause significant
displacement of body water.
(2) Hypernatremia
as seen in multiple myeloma and other dysproteinemias may rarely cause
pseudohyponatremia.
Signs and symptoms of
hyponatremia
(a) >125 mmol
/ L – nausea and malaise occur.
(b) 110–120 mmol
/L – headache, lethargy and obtundation appear.
(c)
<110 mmol / L – seizure and coma.
B. Hypernatremia – occurs when the plasma Na+
concentration is greater than 145 mmol / L.
1. Water loss
a. Gastrointestinal losses – vomiting, diarrhea
b. Excessive sweating: fever, exercise
c. Diabetes insipidus
2.
Sodium gain
a. Ingestion / Infusion
b. Hypoaldosteronism (Conn’s disease)
************ POTASSIUM ************
Potassium is the major intracellular cation of the body. The
normal daily intake of K+ is 80 to 100 mmol; 98% of which is
intracellular and 2% extracellular, 95% of the total potassium is exchangeable.
Potassium is absorbed into the blood stream from the small
intestines. It is filtered in the glomeruli and completely reabsorbed in the
proximal renal tubular epithelium but it is then re–excreted by the distal
renal tubular epithelium.
Accumulation of potassium in the blood above 7.5 mEq / L results
in disturbance in the muscle irritability (inhibitory effect), respiration and
myocardial infarction. There is a movement of Na+ into the cells and
shift of K+ from the cells into the extracellular fluid during the
state of alkalosis.
Functions of potassium
1. Catalytic action in various enzymatic
reactions in the cells.
2. Helps maintain normal pH of the body fluid.
3. Helps maintain normal movements of
intracellular fluid.
Methods of determination
1. Precipitation as cholorplatinate followed by
a. Gravimetric measurement
b. Determination of chloride released by
reduction of platinum.
2. Precipitation as the dipricylaminate followed
by photometric measurement
3. Precipitation of potassium silver
cobaltinitrite (introduced by Breh and Gaebler) followed by determination of
cobalt either photometrically and thiocyanate or dimethyl glaxamine and benzidine
or photometrically by diazotization or by adaptation of the Bratton–Marshall
method for sulfanilamide.
4. Neuron activation analysis
5. Photometric measurement of the blue color of
undissociated (NH4)2Co(SCN)3 complex in
alkaline solution.
6. Flame photometry with 1:50 dilution
7. Ion selective electrode
Colorimetric method of potassium
determination
1. Lockhead and Purcell method
Potassium is precipitated
directly from the serum or plasma as potassium sodium cobaltinitrite. The
alkaline solution of cobalt in the presence of a trace of amino acid like
glycine reduces the Folin–Ciocalteau reagent to a blue color. The intensity of
the color is then measured against a standard.
2. Jacob and Rowland
The potassium is precipitated
directly from the serum as potassium sodium cobaltinitrite salt. The
precipitate is washed as the amount of cobalt in the precipitate is determined
photometrically using the green color produced by the addition of chloride by
the action of sodium ferrocyanide.
3. Flame Photometry
A 1:50 dilution of serum is
prepared. This is subjected to a non–luminous flame which emits light of a
characteristic wavelength for potassium. The intensity of the light is directly
proportional to the concentration of potassium present in the serum.
Precautions in potassium determination
1. The patient must be fasting.
2. Specimen desired is serum or heparinized
blood.
3. Serum or plasma must be obtained 20 minutes
after blood is withdrawn from the red blood cells into the serum or plasma.
4. Flame photometry is preferred for potassium
determination and if not available, the chemical or colorimetric method is
used.
· Regardless of the method used, no hemolysis
of serum should be allowed. The blood must be allowed to clot, centrifuged;
serum is removed from the clot.
Clinical significance
1. Hypokalemia
a. Potassium deficiency
b. Prolonged diarrhea
c. Hyperadrenalism
d. Prolonged vomiting
e. Hyperinsulinism
f. Diabetes
g. Chronic nephritis
h. Hereditary periodic paralysis
2. Hyperkalemia
a. Addison’s Disease
b. Anuria or urinary obstruction
c. Renal tubular acidosis
d. Pneumonia
e. Lack of steroid
************ CALCIUM ************
Ninety–five percent of calcium in the body is found in bones and
teeth. The metabolism of calcium, phosphorous and sometimes Mg are related to
one another. Calcium and phosphorous are absorbed into the blood stream from
the intestine.
Functions of
calcium
1.
Mineralization of the skeletal system
2.
Involved in the process of blood coagulation
3.
Plays a role in the transmission of nerve impulses
4.
Is essential in the preservation of skeletal
muscle
5.
Takes part in the activation of several
enzymes
Increased absorption of calcium is brought
about by the following:
1.
High calcium concentration in the diet
2.
Low pH in the intestine
3.
Presence of Vitamin D
4.
Parathormone
Decreased absorption of calcium occurs when:
1. Dietary concentration ratio (Ca:P) is above
2:1
2. Presence of phytic acid in diet
3. Presence of steatorrhea. Fatty acids from
insoluble soap with calcium
All the calcium of the blood is present in serum, while
phosphorous is present mainly in the cell as organic phosphate, with a small
amount in the serum as inorganic phosphate.
Forms of calcium in serum in general:
1. Non–diffusible, protein bound calcium – 50%;
4.6 mg% or 2.3 mEq/L
2. Diffusible form
a. Ionized calcium – 45%; 4.2–4.5 mg% or 2.1–2.2
mEq/L
b. Complexed calcium (with citrate and
phosphorous) – 0.4 – 0.6 mg% or 0.2 – 0.3 mEq /L
Tetany will result if there is a significant decrease of the
ionized calcium regardless of the total calcium. The normal CSF calcium levels
are 4.2 – 5.8 mg%. Some authors recommend this value as an index of the ionized
calcium in the serum. The approximate concentration of ionized calcium may also
be calculated with the formula based on the monogram by McLean and Hastings:
Mg. Ca++/100
ml = (6 x total serum Ca in mg%) – (gram serum protein / 100 ml)
(gram
protein / 100 ml) ÷ 6
The formula is based on the relationship between calcium and
protein of serum at pH 7.55 at 25oC. The following factor can alter
the formula:
1.
Difference in pH
2.
Proportions of proteins
3.
Increased phospholipids
Factors influencing serum calcium levels are:
1.
Parathyroid hormone
2.
Plasma proteins
3.
Plasma phosphate
4.
Vitamin D
Parathormone regulates the level of serum calcium. It mobilizes
calcium from the bones; it increases calcium absorption in the intestine and
increases the excretion of phosphates by the kidneys.
Thyrocalcitonin (hypocalcemic factor) is produced by the
parafollicular “C” cells of the thyroid glands in response to the presence of
hypercalcemia. This hormone negates the action of parathormone and vitamin D on
bone resorption including osteoclastic osteolysis. It also causes a decreased
urinary excretion of calcium, magnesium and hydroxyproline.
Gonadal hormones (androgens and estrogen) depress bone
reabsorption. Growth hormones increases both intestinal absorption and renal
excretion of calcium.
Conditions with hypercalcemia accompanied by
calcium loss:
1.
Hyperparathyroidism
2.
Sarcoidosis
3.
Metastatic neoplasm to bone
4.
Vitamin D intoxication
5.
Milk alkali syndrome
6.
Addison’s disease
7.
Disuse atrophy
Conditions with hypocalcemia
1.
Hypoparathyroidism
2.
Rickets
3.
Renal failure
4.
Thyrocalcitonin excess
5.
Hyperparathyroidism
6.
Advanced and sustained vitamin D deficiency
Methods of calcium determination
1. The classical method for calcium
determination depends on the precipitation of calcium to insoluble oxalate and
the measurement of the oxalate in the precipitate.
a. Precipitated oxalate is transferred to a
platinum crucible, dried, converted to calcium oxide dissolved in acid and
titrated with a standard base.
b. Oxalated precipitate is washed, dissolved in
sulfuric acid and the oxalic acid formed with standardized potassium
permanganate solution
(1) Calcium is
precipitated as calcium oxalate
Ca++ + (NH4)2C2O4
-------------> CaC2O4 + 2H4+
(2) The
precipitate is converted
CaC2O4 + H2SO4
-------------> H2C2O4
+ CaSO4
(3) The oxalic
acid is titrated with potassium permanganate
2 KMnO4 + 5H2C2O4
+ 3 H2SO4 ----------> K2SO4 + 2MnSO4
+ 10CO2 + 8H2O
2. Tricholoroacetic acid is added to precipitate
the serum proteins. Sodium hydroxide and trisodium phosphate are added to the
filtrate to precipitate calcium as tricalcium phosphate. The mixture is
centrifuged and the precipitate is washed with an alkaline alcohol solution to
remove impurities and excess phosphates. Acid molybdate and aminonapthosulfonic
acid reagent are added to form a color complex with tricalcium phosphate and
the depth of color is measured and compared with a standard.
3. Determination of calcium by EDTA titration
(Bachra, Dauer and Sobel)
Certain dye solution like cal–red, calcein, ammonium purate and
Eriochrome Black T have a characteristic color in the presence of ionized
calcium. The addition of chelating agent (EDTA) binds the calcium and a change
in color results.
4. Chloranilic acid method (Ferro–Ham)
The calcium in the sample is precipitated as calcium chloranilate
by saturated solution of sodium chloranilate. The excess chloranilic acid is
removed by washing the precipitate with isopropyl alcohol and the precipitate
is then treated with EDTA, which chelates calcium and releases chloranilic
acid.
Ca++ + chloranilate
---------------------> Ca–chloranilate
Ca–chloranilate + EDTA
--------------> Ca–EDTA + chloranilic acid (purple color)
5. Flame photometry
This method is not accurate.
Difficulties encountered in this method are the following:
a. Positive interference by sodium and potassium
b. Inhibition of calcium omission by phosphates
and sulfates
c. Difficulty of exciting calcium
6. Atomic absorption spectroscopy
Calcium compound dissociate into free calcium atoms when
introduced into a flame. The calcium atoms absorb light of a characteristic
wavelength produced by a hollow cathode lamp.
7. Calcein–fluorometric method
Calcium ions in combination with 3,6–dehydroxy–2–4–bis fluoran forming a fluorescent complex which is measured fluorometrically.
8. Ion selective electrode
************ INORGANIC PHOSPHOROUS ************
Phosphorous metabolite is directly related with that of calcium.
About 80–90% of phosphorous is absorbed into the blood stream from the small
intestine. Normal values of phosphorous are from 3.0 – 4.5 mg% with slightly
higher values in children.
Functions of phosphorous are:
1. Phosphorous forms a major intermediate and
high energy phosphate bonds in carbohydrate metabolism.
2. It is an important constituent of nucleic
acids, phospholipids, nucleotides as well as bone.
3. Plays a role in the regulation of the pH of
the body.
In the blood, phosphorous occurs as monovalent and bivalent
phosphate. At pH 7.4 for each part of inorganic phosphate, 0.8 parts is
bivalent phosphate and 0.2 parts is monovalent phosphate.
The determination of inorganic phosphorous requires the conversion
of phosphorous in a protein free filtrate to form phosphomolybdic acid and the
subsequent reduction of the acid to produce color.
Method of determination
1. Fiske – Subarrow method
Ammonium molybdate in an acid
medium of serum protein free filtrate forms phospholybdic acid. The addition of
para–aminonapthol sulfonic acid reagent reduces the phosphomolybdic acid to
phosphomolybdous acid with blue molybdenum for photometric measurement.
Points to consider
1. Patient must be fasting
2. Serum must be obtained 30 minutes after the
blood is withdrawn. There must be no hemolysis of serum.
3. Fresh serum, less than 24 hours is necessary
to get accurate and precise phosphorous value.
4. Addition of a suitable mild reducing agent
that reacts with phosphomolybdate complex causes the formation of a blue color
of molybdenum. Aminonapthol sulfonic acid reagent is a mild reducing agent.
Other mild reducing substances like stannous chloride, ascorbic acid,
elonpicotol have been used.
Strong reducing agents are undesirable
as they will also reduce the excess molybdate and will yield a false elevated
phosphorous value.
5. The phosphorous should not be kept standing
more than after color development because reoxidation and fading of the color
may occur. Phosphorous in urine can also be determined using the method
described.
Organic phosphorous like
phosphorous bound to protein and lipids can also be determined using the method
described. However, ashing or digestion of the sample must first be done to
liberate the bound phosphorous and the analysis is done by the method described
before.
************ MAGNESIUM ************
Magnesium is an important intracellular cation second to potassium
in quantity. It is absorbed in the intestines. About 85% of the magnesium in
blood is diffusible and the remainder is bound to protein. Magnesium ions serve
as activators for a number of important enzyme systems engaged in hydrolysis
and transfer of phosphate groups, e.g. alkaline phosphatase, prostatic acid
phosphatase, hexokinase and creatinine kinase.
A reciprocal relationship exists between serum calcium and serum
magnesium.
Decrease Magnesium is found in
1.
Malabsorption syndrome
2.
Acute pancreatitis
3.
Chronic alcoholism
4.
Delirium tremens
5.
Aldosteronism
Increased Magnesium is found in
1.
Dehydration
2.
Severe diabetic acidosis
3.
Addison’s disease
4.
Uremia
A Magnesium deficiency, tetany, has been described in which the
laboratory findings is a low magnesium and normal calcium level.
Methods of magnesium determination
1. Chemical method
Calcium is removed from the serum
and the magnesium is precipitated as ammonium phosphate salt (MgNH4PO4).
The precipitate is washed and redissolved and the phosphate determined by the
inorganic phosphorous method of Fisk and Subarrow.
2. Titan yellow method
A TCA filtrate of serum is treated
with the dye, titan yellow, in an alkaline solution. The red lake that forms is
thought to be dye absorbed on the surface of colloidal particles of magnesium
present and is then measured photometrically and compared with a standard.
3. Complexiometric method
Titration with EDTA chelates the
magnesium removing it from ionic form and destruction of the dye complex,
resulting in a change in color. During the titration process, calcium is also
chelated so that a second titration with murexide is needed which will
represent the calcium concentration. The value of magnesium is determined by
subtracting the value from the use of Eriochrome dye.
4. Fluorometric method
Magnesium ions and
gamma–hydroxyl–5–quinolone sulfonic acid form a chelate compound that floresce
when excited at wavelength 380 – 410 nm. EGTA increases the specificity by
preventing calcium.
5. Atomic absorption spectroscopy
Lanthanum and strontium are
contained in the diluent to bind with phosphate, preventing the formation of
magnesium–phosphate compounds that are not measured.
********** IRON AND IRON–BINDING CAPACITY **********
Total amount of iron in an adult is about 4 – 5 grams, of which
66% is found in hemoglobin, 4% in myoglobin and the cytochrome system, 30% is
found in various storage sites of the spleen, liver and bone.
Regulation
Normally, the body only absorbs 5 – 10% supply of the daily
requirement of 1 – 2 mg from newly absorbed dietary iron. Recycling of iron
provides most of the 20 mg required daily for normal erythropoiesis.
Iron absorption occurs primarily in the duodenum and can be
increased up to 20% in states of iron deficiency and during growth and
pregnancy when iron needs are greater.
Dietary iron exists as both heme and non–heme iron pools. Heme
iron is found primarily in hemoglobin and myoglobin and can be absorbed
directly by the intestinal mucosal cells.
Non–heme irons are found in foods such as vegetables and eggs,
where it exists in the form of ferric hydroxide. Dietary iron can be absorbed
only in the ferrous (Fe++) form. Thus, all dietary iron in the
ferric (Fe+++) form must first be reduced. Certain factors such as
ascorbic acid and the acidic conditions of the gastric contents facilitated
this reduction and, hence, absorption. Other factors acts as blocking agents
and reduce absorption on non–heme iron. Phytates and phosphates found in
certain foodstuff, including cereals, inhibit iron reduction as do antacids and
antibiotics.
Once reduced to the ferrous form, approximately 5 – 10% of dietary
iron enters the mucosal cells. The remainder is lost in the feces. The exact
mechanism by which iron is absorbed is unknown but appears to be carrier
mediated. The control of iron absorption resides at the intestinal mucosal level;
iron homeostasis is not regulated by excretion.
Upon entering the mucosal cells, iron is oxidized back to the
ferric form (Fe3+) and is stored with a protein, apoferritin, to
form storage iron. The main storage complex is known as ferritin. Ferritin is
the major iron storage protein found in all cells of the body. However, the
major storage sites are cells of the reticuloendothelial system and liver.
Although most ferritin is found in the tissues, a small percentage
exists in the plasma, where its concentration is proportional to tissue
concentration. Thus, measurement of serum ferritin levels is a direct
indication of the amount of storage iron.
When iron is needed by the body for incorporation into the iron–containing
molecules, it is released from ferritin, enters the plasma, and is bound to the
transport protein, transferrin. Once transferrin delivers it’s bound iron, the
red cell precursors actively making hemoglobin, it recirculates to the
transport more iron.
Senescent red blood cells are engulfed by macrophages in the
spleen and other organs, where iron is liberated from catabolism of the
hemoglobin molecule. The amount of iron released is approximately 20 mg which
is the amount required for daily hemoglobin synthesis. Iron released from
hemoglobin remains temporarily in the cells of the RE system. It then slowly
leaves these cells and is bound to transferrin, which recirculates it for
incorporation into developing cells.
Methods of iron determination
1. Serum iron
a. Splitting of Fe+++ from
transferrin complex by exposure to acid.
(1) HCl
(2) H2SO4
(3) Trichloroacetic
acid
b. Separation of Fe+++ and protein
(1) Protein
precipitation
(2) Protein
remains in solution without interfering with analysis
c. Reduction of Fe+++ to Fe++
with
(1) Ascorbic
acid
(2) Hydrazine
(3) Thioglycollic
acid
(4) Hydroxylamine
d. Reaction of Fe++ with a chromogen
(1) Bathopenanthroline
(2) Diphenylpananthroline
(3) Ferrozine
(4) Tripyridylriazine
Iron must be separated from its
protein complex and reduced to Fe++ because the chromogen reacts
only with iron in its reduced state.
2. Total iron binding capacity
The following steps outline the
general procedure for TIBC measurement
a. Saturation of transferrin with excess Fe+++
(1) Ferric
ammonium citrate
(2) Ferric
chloride
b. Removal of excess unbound Fe+++
(1) Ion exchange
resin
(2) Iron chelator
(MgCO3)
c. Performance of conventional iron
determination of supernatant
Precautions in iron determination
1. Serum is the specimen of choice, preferably
collected in amber–stoppered tube that has been specifically treated to remove
any trace of iron contamination.
2. A fasting, morning sample allows the most
accurate assessment of iron status since iron levels are subject to diurnal
variation, which causes afternoon samples to be decreased by as much as 30%.
3. Hemolyzed samples are unsuitable because of the
iron content of erythrocytes.
4. Glasswares used for iron analysis should be
acid washed to remove trace contamination or iron and double distilled water
should be used in reagent preparation.
5. Separated serum is stable for 1 week at
refrigerated temperatures for both iron and TIBC analyses.
Clinical significance
1. Iron deficiency
Iron deficiency occurs when the
amount of iron absorbed is inadequate to meet the needs of the body. The serum
iron concentration is <40 mg/dl.
Causes of iron deficiency
a. Iron deficiency anemia
b. Chronic diarrhea and malabsorption syndrome
c. Chronic blood loss due to menstruation,
peptic ulcer, hemorrhoids, esophageal varices and gastritis due to salicylate
ingestion.
d. Long standing infection
Signs and symptoms
a. Fatigue, headache and pallor.
b. A distorted appetite, called pica, with
cravings for such substances as earth, clay or ice may also be present
c. Sore tongue or mouth
d. Koilonychia – thinning or spooning of the
fingernails
e. Plummer–Vinson syndrome – opening of the
esophagus is partially occluded, leading to a sensation of food sticking in the
throat.
2. Iron overload
a. Increased absorption
(1) Primary
hemochromatosis
(2) Hemosiderosis
(3) Iron
poisoning: dietary, medicinal, transfusional
b. Increased red cell destruction
(1) Hemolytic
anemia
c. Ineffective erythropoiesis
(1) Thalassemia
(2) Sideroblastic
anemia
************ CHLORIDES
************
This is an essential and important anion of extracellular fluid.
It is closely associated with sodium and potassium in body tissue and their
excretion.
Chlorides are absorbed into the blood from the small intestine and
its distribution is the same as sodium.
Functions of chlorides
1. Plays a role in acid–base equilibrium
2. Acts indirectly as a factor in the
maintenance of body water
3. Maintains osmotic pressure
There is a marked difference between the chloride concentration of
the intracellular and extracellular components. About 65% of the total
extracellular anions are made up of chlorides.
Whole blood has a lower value of chlorides than serum or plasma as
red cells make up almost half of the whole volume. There is only one half as
much plasma chloride in the red cells because of a relatively high
concentration of protein and low water content within the red blood cells.
About 4% of the chlorides are turned in daily.
Concept of a chloride shift
CO2 accumulates in tissue cells as a product of normal
cellular metabolism. It diffuses out of tissue cells into the plasma, where a
small amount is dissolved. Most CO2, however, diffuses down a concentration
gradient into the red blood cells, where it combines with H2O to
form H2CO3. The reaction is catalyzed by the enzyme
carbonic anhydrase. H2CO3 dissociate into H+,
which is buffered by hemoglobin and HCO3. As the HCO3
concentration builds up in the cell, its concentration becomes greater than the
extracellular concentration and it diffuses out of the cell. To maintain
electronegativity, Cl– flows into the cell in exchange for HCO3–.
This process is known as chloride shift.
Method of determination
1. Fantus in 1735 introduced a method in which
chloride in urine was estimated by adding potassium chromate to the sample and
titrated with silver nitrate to a red brown color
2. Volhard (1874) described a method in which
chloride is precipitated with standard thiocyanate using ferric ions as
indicators. However, the presence of proteins, an organic material interfered
with the results so that two variations were made.
a. Destruction of the organic material by wet
digestion with nitric acid – Open Carius technique
b. Application of the Volhard technique to
protein free filtrate as suggested by Whitehorn, Osterberg and Schmidt
3. Introduction of the iodometric method by
Sendroy in 1937. The chloride ions react with soluble AgIO3 which is
added in excess to form insoluble AgCl and IO3. The insoluble salts
are filtered off and the IO3 in the filtrate is determined:
a. Gasometrically
b. Titrimetrically by a thiosulfate – starch
titration of the I2 evolved from the reaction of IO3 with
acidified KI.
c. Photometrically – measurement of either the
(1) Yellow color
of the I2 itself
(2) Blue color
formed with the starch
4. Schales and Schales in 1940 adapted the
mercuric method. Chloride ions and mercury forming undissociated but soluble
HgCl2. The end point is obtained when a violet blue (purple) color
is seen resulting from the combination of the excess Hg with the indicator
(diphenyl carbazone or diphenyl carbazide).
Bromides present in the sample
will also combine with mercury and will be calculated as chloride. Normally, the
amount of bromide in the blood is not detectable but in bromide poisoning, the
results will be falsely elevated.
5. Automated electrometric titration method
(Cotlove–Buchler chloridometer)
The serum is diluted in an acid
solution (HNO3–CH3COOH) mixture containing small amount
of gelatin (25 mg / 100 ml)
HNO3 – provides good
electrolytic conductivity
Acetic acid – renders the
solution less polar thus reducing the solubility of silver chloride and thus
providing a sharper end point.
Gelatin – provides a smoother and
more reproducible titration curve by being absorbed preferentially to high
spots of the electrode and thus equalizing the reaction rate over the entire
electrode surface.
Acid solutions also aid in
preventing reduction of precipitated silver chloride at the indicator cathode.
Excess protein may introduce some
error due to the reaction of silver ions with sulfhydryl group of proteins.
When titration is started, a
silver generator electrode is feed by a constant current that oxidizes silver
to Ag+ at a constant rate proportional to Q (Coulomb). Silver ions produced
combine with chloride to form a precipitate of AgCl3. After the
equivalence point is reached (sufficient Ag+ has been generated to
react with all chloride present) additional generation of Ag+ will
result in an increase in electroactivity of the titrant which is measured
amperometrically by a set of silver indicator electrode. The increase in
current activates a relay which in turn will stop an automatic timer and also
stops the generation of additional Ag+. Since the current feeding,
the silver generator electrode is constant, the rate of generation of Ag+
is also constant, thus the time necessary to reach the titration end point can
be taken as a measure of the chloride concentration. Titration of a blank
solution as well as standard solution should be done.
Other halogens as well as CN–,SCN,
and –SH interferes with the determination
Greatest accuracy is obtained if
titration is held between 70–160 seconds.
Colorimetric measurement with
mercuric thiocyanate – Zoll, Fisher, Gawer, adapted to the Auto Analyzer by
Skeggs is based on the following principle:
The sample is mixed with a
solution of Hg(SCN)2. As a result of the higher affinity of Cl to
Hg, there is formation of undissociated HgCl2 resulting in the
release of free SCN–.
The SCN– reacts with
Fe+++ of ferric chloride reagent to form an intense reddish color
complex of Fe(SCN)3 with an absorption peak at 480 nm.
Hg(SCN)2 + 2Cl
-------------> HgCl2 + 2(SCN)–
3 (SCN)– + Fe+++
-------------> Fe (SCN)3
6. Potentionmetric method
7. Conductimetric method
8. Polarographic method
Clinical significance
1. Hyperchloremia
a.
Nephritis
b.
Eclampsia
c.
Prostatic obstruction
d.
Anemia
e.
Hyperventilation
f.
Hypoproteinemia
g.
Serum sickness
h.
Urinary obstruction
i.
Increase chloride intake
j.
Dehydration
k.
Decrease renal blood
2. Hypochloremia
a.
Addison’s disease
b.
Burns
c.
Fever
d.
Intestinal obstruction
e.
Metallic poisoning
f.
Pneumonia
g.
Heat cramps
h.
Diarrhea
i.
Vomiting
j.
Uremia
k.
Polycythemia vera
l.
Hypercortico–adrenalism
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