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11/1/2011
Fundamentals of pH
Is diet really the issue?
Is diet an issue?
Fundamental Terms and Concepts
Mole (mol): The mole is the amount of substance which contains as
many elementary entities as there are atoms in 0.012 kilogram of
carbon 12; its symbol is "mol." (6.022 x 1023)
Equivalent: 1 mole of ionic charges
Example: HCO3- has a single charge, thus one mole of HCO3- is
one mole of charges, or one “equivalent”.
Ca2+, however, has two charges, thus one mole of this
ion yields two equivalents.
Moles x Valence = Equivalents
When do I use “moles”, when do I use “equivalents”?
Molecules must be quantified by “moles” because they have no charge.
Ions may be quantified by either “moles” or “equivalents”
Typical usage: I liter of blood plasma contains about 0.004 moles of
K+. This = 4 mmol/l or 4 meq/l
Test question: If 0.004 moles/l of Ca2+, ____ meq/l?
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The Hydrogen Ion and pH
The hydrogen ion consists of a single positively charged particle
(the proton) that is not orbited by any electrons. The hydrogen ion
is, therefore, the smallest ionic particle and is extremely reactive.
It is this fact that accounts for its profound effect on the functioning
of biological systems at very low concentrations.
In the environment hydrogen ion concentrations vary over an
enormous scale (from less than 10-14 mol/l to more than 1mol/l).
The pH scale was developed in order to simplify (or perhaps
further complicate!) the mathematics of handling such a large
range of numbers. The pH is calculated by taking the negative
logarithm of the hydrogen ion concentration, as shown below.
pH = -log10[H+]
where [H+] is the hydrogen ion concentration.
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Acids: An acid is defined as any compound,
which forms hydrogen ions in solution. For this
reason acids are sometimes referred to as
"proton donors". To aid understanding of these
concepts consider an imaginary acid with the
chemical formula HA. In the first example in
Figure 2, the acid dissociates (separates) into
hydrogen ions and the conjugate base when in
solution.
Bases: A base is a compound that combines
with hydrogen ions in solution. Therefore,
bases can be referred to as "proton
acceptors".
Strong Acids: A strong acid is a compound
that ionizes completely in solution to form
hydrogen ions and a base. Example 2
illustrates a strong acid in solution, where this
dissociation is complete.
Weak Acids and Bases: these are
compounds that are only partially ionized in
solution.
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The Importance of Hydrogen Ion Concentration
•Hydrogen ion concentration has a widespread effect on the
function of the body's enzyme systems. The hydrogen ion is
highly reactive and will combine with bases or negatively charged
ions at very low concentrations.
•Proteins contain many negatively charged and basic groups
within their structure. Thus, a change in pH will alter the degree
ionization of a protein, which may in turn affect its functioning. At
more extreme hydrogen ion concentrations a protein's structure
may be completely disrupted (the protein is then said to be
denatured).
•Enzymes function optimally over a very narrow range of
hydrogen ion concentrations. For most enzymes this optimum pH
is close to the physiological range for plasma (pH= 7.35 to 7.45,
or [H+]= 35 to 45nmol/l). Figure 5 shows a typical graph obtained
when enzyme activity is plotted against pH. Notice that the curve
is a narrow bell shape centered around physiological pH.
WHY is pH important?
Regulation of pH is critical to the maintenance of homeostasis
because almost all enzyme systems in the body are directly
influenced by [H+]
…but why does
this happen?
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•Proton binding (H+) is reversible and depends on the concentration
of protons in solution.
•When proton concentration rises (pH falls), more anionic sites in the
solution are occupied by protons. These occupied sites are
neutralized, lose their negative charge.
•Conversely, when proton concentration falls (pH rises) fewer anionic
sites are occupied and they retain the negative charge.
•Proteins are especially vulnerable to these changes because the
three-dimensional structure is greatly influenced by the charges on
constituent amino acids, and these structural changes affect
function. Enzyme activity may be decreased, the shape of the
structural proteins may be altered, and the activity of transmembrane channels and pumps may be reduced.
•Example: Calcium. When blood pH rises, more anionic sites on
albumin and other plasma proteins are exposed. These sites bind
with Ca+2 and may reduce plasma concentration so low as to induce
tetany, paresthesia, cardiac arrythmia!
pH and H+ Concentration of Body Fluids
Extracellular fluid
Arterial blood
Venous blood
Interstitial fluid
Intracellular fluid
Urine
Gastric HCl
H+ (mEq/L)
pH
4.0 x 10-5
4.5 x 10-5
4.5 x 10-5
7.40
7.35
7.35
1 x 10-3 to 4.0 x 10-5
3 x 10-2 to 1 x 10-5
160
6.0 to 7.4
4.5 to 8.0
0.8
The greatest body burden is the very acid it makes for digestion!
How does the body “protect” itself?
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Production of Hydrogen Ions
•The processes of metabolism generate hydrogen ions. Small amounts
(40-80mmol/24h) are formed from the oxidation of amino acids and the
anaerobic metabolism of glucose to lactic and pyruvic acid.
•Far more acid is produced as a result of carbon dioxide (CO2) release
from oxidative (aerobic) metabolism - 15,000mmol/24h (1.5x103
mmol/24h).
•Although CO2 does not contain hydrogen ions it rapidly reacts with
water to form carbonic acid (H2CO3), which further dissociates into
hydrogen and bicarbonate ions (HCO3-). This reaction is shown below:
CO2 + H20 <= H2CO3 => HCO3- + H+
•This reaction occurs throughout the body and in certain circumstances is
speeded up by the enzyme carbonic anhydrase. Carbonic acid is a weak
acid and with bicarbonate, its conjugate base, forms the most important
buffering system in the body.
•Acids or bases may also be ingested, however, it is uncommon for these
to make a significant contribution to the body's hydrogen ion
concentration, other than in deliberate overdose.
Buffers: A buffer is a compound that limits the change in hydrogen ion
concentration (and so pH) when hydrogen ions are added or removed from
the solution. It may be useful to think of the buffer as being like a sponge. When
hydrogen ions are in excess, the sponge mops up the extra ions. When in short
supply the sponge can be squeezed out to release more hydrogen ions!
All buffers are weak acids or bases.
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“…like a sponge.”
The effects of buffers can also
be illustrated graphically. If a
strong acid is added slowly to
a buffer solution and the
hydrogen ion concentration
[H+] is measured then a plot
similar to the one in figure 4
will be generated. Notice that
during the highlighted portion
of the curve a large volume of
acid is added with little
change in [H+] or pH.
As we shall see later buffers
are crucial in maintaining
hydrogen ions within a narrow
range concentrations in the
body.
Normal physiology is
well protected against
pH changes by several
buffer systems that
react over different
timeframes.
There are essentially 3
rapidly acting chemical
buffering systems in
the body, shown here.
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Defending against pH changes: Buffers, Lungs, Kidneys
There are three primary systems that regulate pH in the body
fluids to prevent acidosis or alkalosis:
1. Chemical Acid-Base Buffer Systems: (see previous slide)
immediately combine with acid or base to prevent excessive
changes in pH. Acts within a fraction of a second
2. Respiratory Center: regulates removal of CO2 (and,
therefore, H2CO3) from the ECF. Acts within a few minutes.
3. Kidneys: can excrete either acid or base, thus readjusting
the ECH pH to normal. Acts within hours to days.
•The first two act as the “sponge” to minimize the effects of
changes in [H+]. Buffers do not remove or add anything!
•It is the third line of defense, the kidneys, that actually eliminate
excess acid or base as needed. This is, by far, the most powerful
of the three systems.
Bicarbonate Buffer System
…discussion next…
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The primary role of the carbonic acid–bicarbonate buffer system is to prevent
pH changes caused by organic acids and fixed acids in the ECF. This buffer
system has three important limitations:
•It cannot protect the ECF from pH changes that result from elevated or
depressed levels of CO2. A buffer system cannot protect against changes
in the concentration of its own weak acid. In the case above, the addition
of excess H+ drove the reaction to the left. But if we had added excess
CO2 instead of excess H+, the elevated CO2 would have driven the
reaction to the right. Additional H2CO3 would have formed and dissociated
into H+ and HCO3–. This reaction would have reduced the pH of the
plasma.
•It can function only when the respiratory system and the respiratory
control centers are working normally. Normally, the elevation in PCO2 that
occurs when fixed or organic acids are buffered will stimulate an increase
in the respiratory rate. This increase accelerates the removal of CO2 at the
lungs. If the respiratory passageways are blocked, blood flow to the lungs
will be impaired, or if the respiratory centers do not respond normally, the
efficiency of the buffer system will be reduced. (normal lung activity!)
•The ability to buffer acids is limited by the availability of bicarbonate
ions. Every time a hydrogen ion is removed from plasma, a bicarbonate
ion goes with it. When all the bicarbonate ions have been tied up, buffering
capabilities are lost. (normal kidney activity!)
Proteins as buffers
Protein buffer systems depend on the ability of amino acids to respond to pH changes
by accepting or releasing H+. At the normal pH of body fluids (7.35–7.45), the carboxyl
groups of most amino acids have already given up their hydrogen ions. (Proteins carry
negative charges primarily for that reason.) However, some amino acids, notably
histidine and cysteine, have R groups (side chains) that will donate hydrogen ions if the
pH climbs outside the normal range. Their buffering effects are very important in both the
ECF and ICF.
If the pH drops, the amino group (—NH2) can act as a weak base and accept
an additional hydrogen ion, forming an amino ion (—NH3+). This effect is primarily
limited to free amino acids and the last amino acid in a polypeptide chain, because the
amino groups in peptide bonds cannot function as buffers.
Plasma proteins contribute to the buffering capabilities of blood. Interstitial fluid
contains extracellular protein fibers and dissolved amino acids that also assist in the
regulation of pH. In the ICF of active cells, structural and other proteins provide an
extensive buffering capability that prevents destructive pH changes when organic acids,
such as lactic acid or pyruvic acid, are produced by cellular metabolism.
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Phosphate Buffer system
The phosphate buffer system consists of the anion H2PO4–, which is a weak
acid. The operation of the phosphate buffer system resembles that of the
carbonic acid–bicarbonate buffer system. The reversible reaction involved is
The weak acid is dihydrogen phosphate (H2PO4–), and the anion released is
monohydrogen phosphate (HPO42–). In the ECF, the phosphate buffer system
plays only a supporting role in the regulation of pH, primarily because the
concentration of HCO3– far exceeds that of HPO42–. However, the phosphate
buffer system is quite important in buffering the pH of the ICF. In addition,
cells contain a phosphate reserve in the form of the weak base sodium
monohydrogen phosphate (Na2HPO4). The phosphate buffer system is also
important in stabilizing the pH of urine. The dissociation of Na2HPO4 provides
additional HPO42– for use by this buffer system:
Kidneys: The three major buffers involved are
(1) the carbonic acid–bicarbonate buffer system,
(2) the phosphate buffer system, and
(3) ammonia.
Glomerular filtration puts components of the carbonic acid–bicarbonate
and phosphate buffer systems into the filtrate. The ammonia is
generated by tubule cells (primarily those of the PCT).
Reversal in alkalosis
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The Bicarbonate Buffer System:
CO2 + H2O = H2CO3 = H+ + HCO3K = 6.1 (dissociation constant) see Guyton, p. 385
Kidneys
pH = 6.1 + log
HCO 3
0.03 X PCO
2
Lungs
What is the truth about the
impact of diet on ph?
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Acid/Alkaline Theory of Disease Is Nonsense
Gabe Mirkin, M.D.
Have you seen advertisements for products such as coral
calcium or alkaline water that are supposed to neutralize acid in
your bloodstream? Taking calcium or drinking alkaline water does
not affect blood acidity. Anyone who tells you that certain foods or
supplements make your stomach or blood acidic does not
understand nutrition.
You should not believe that it matters whether foods are
acidic or alkaline, because no foods change the acidity of anything
in your body except your urine. Your stomach is so acidic that no
food can change its acidity. Citrus fruits, vinegar, and vitamins
such as ascorbic acid or folic acid do not change the acidity of
your stomach or your bloodstream. An entire bottle of calcium pills
or antacids would not change the acidity of your stomach for more
than a few minutes.
All foods that leave your stomach are acidic. Then they
enter your intestines where secretions from your pancreas
neutralize the stomach acids. So no matter what you eat, the food
in stomach is acidic and the food in the intestines is alkaline. You
cannot change the acidity of any part of your body except your
urine.
Your bloodstream and organs control acidity in a very
narrow range. Anything that changed acidity in your body would
make you very sick and could even kill you. Promoters of these
products claim that cancer cells cannot live in an alkaline
environment and that is true, but neither can any of the other cells
in your body.
All chemical reactions in your body are started by
chemicals called enzymes. For example, if you convert chemical A
to chemical B and release energy, enzymes must start these
reactions.
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All enzymes function in a very narrow range of acidity. (The
degree of acidity or alkalinity is expressed as "pH."). If your blood
changes its acidity or alkalinity for any reason, it is quickly
changed back to the normal pH or these enzymes would not
function and the necessary chemical reactions would not proceed
in your body.
For example, when you hold your breath, carbon dioxide
accumulates in your bloodstream very rapidly and your blood
turns acidic, and you will become uncomfortable or even pass out.
This forces you to start breathing again immediately, and the pH
returns to normal. If your kidneys are damaged and cannot
regulate the acidity of your bloodstream, chemical reactions stop,
poisons accumulate in your bloodstream, and you can die.
Certain foods can leave end-products called ash that can
make your urine acid or alkaline, but urine is the only body fluid
that can have its acidity changed by food or supplements.
ALKALINE-ASH FOODS include fresh fruit and raw
vegetables.
ACID-ASH FOODS include ALL ANIMAL PRODUCTS, whole
grains, beans and other seeds. These foods can change the
acidity of your urine, but that's irrelevant since your urine is
contained in your bladder and does not affect the pH of any other
part of your body.
When you take in more protein than your body needs, your
body cannot store it, so the excess amino acids are converted to
organic acids that would acidify your blood. But your blood never
becomes acidic because as soon as the proteins are converted to
organic acids, calcium leaves your bones [WRONG!!] to neutralize
the acid and prevent any change in pH. Because of this, many
scientists think that taking in too much protein may weaken bones
to cause osteoporosis.
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Cranberries have been shown to help prevent recurrent
urinary tract infections, but not because of their acidity. They
contain chemicals that prevent bacteria from sticking to urinary
tract cells.
Taking calcium supplements or drinking alkaline water will
not change the pH of your blood. If you hear someone say that
your body is too acidic and you should use their product to make it
more alkaline, you would be wise not to believe anything else the
person tells you. ________________
Dr. Mirkin is an associate clinical professor of pediatrics at Georgetown
University School of Medicine and is board-certified in four specialties: allergy
and immunology; sports medicine; pediatrics; and pediatric immunology. He
practices medicine in Kensington, Maryland; produces and hosts a syndicated
radio that can be heard online; publishes a monthly newsletter (The Mirkin
Report); and has written books on sports medicine, weight control, and low-fat
eating. His Web site contains reports on hundreds of topics (www.drmirkin.com)
Acid-Base Balance of Diets Which Produce Immunity to Dental Caries
Among the South Sea Islanders and Other Primitive Races
by Weston A. Price, DDS, MS, FACD
Read before the New York Dental Centennial Meeting, New York, N.Y., December 4, 1934;
reprinted from the Dental Cosmos for September 1935
http://www.price-pottenger.org/Articles/Acid_base_bal.htm
It is very important that dependable data be accumulated as rapidly as
possible which bear upon this problem of acid-base balance of foods, since
many enthusiasts are advocating strongly the elimination or reduction of
potentially acid foods such as cereals, meats and fish. Indeed, a great deal of
propaganda is reaching the profession and laity which places great stress upon
the importance of keeping the diet potentially alkaline.
It is my personal belief, based on the extensive data that I am accumulating,
from a study of these various primitive groups and their breakdown at the point
of contact with civilization and its foods, that several constitutional factors may
be involved besides tooth decay, and which are very important. My
investigations are showing that primitive groups have practically complete
freedom from deformity of the dental arches and irregularities of the teeth in the
arches and that various phases of these disturbances develop at the point of
contact with foods of modern civilization.
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It is not my belief that this is related to potential acidity or
potential alkalinity of the food but to the mineral and activator
content of the nutrition during the developmental periods, namely,
prenatal, postnatal and childhood growth. It is important that the
very foods that are potentially acid have as an important part of
the source of that acidity the phosphoric acid content, and an
effort to eliminate acidity often means seriously reducing the
available phosphorus, an indispensable soft and hard tissue
component.
It is my belief that much harm has been done through the
misconception that acidity and alkalinity were something apart
from minerals and other elements. Many food faddists have
undertaken to list foods on the basis of their acidity and alkalinity
without the apparent understanding of the disturbances that are
produced by, for example, condemning a food because it contains
phosphoric acid, not appreciating that phosphorus can only be
acid until it is neutralized by combining with a base.
In my clinical practice, in which I am endeavoring to put into practice the lessons
I am learning from the primitive people, I do not require that the foods of the
primitive races be adopted but that our modern foods be reinforced in body building
materials to make them equivalent in mineral and activator content to the efficient
foods of the primitive people. This usually is accomplished by displacing white-flour
products with whole-wheat products, together with eliminating or reducing the high
caloric foods such as sugars and other sweets, and adding foods that are good
providers of the fat-soluble activators, such as the butter of milk as produced by
cows that are eating liberally of fresh or cured rapidly growing green wheat or rye,
together with the organs of animals and the use of sea foods such as these
primitive people have used so successfully in providing not only high immunity to
dental caries but excellent bodies, with high defense for the degenerative diseases.
We are learning Nature’s methods and undertaking to utilize them. The chemical
content of all of these primitive foods is comparably high in minerals and activators,
especially the fat-soluble activators, while being relatively low in calories. In no
instance have I found the change from a high immunity to dental caries to a high
susceptibility among these primitive racial stocks to be associated with a change
from a diet with a high potential alkalinity to a high potential acidity, as would seem
to have been the case had the high alkalinity balance theory been the correct
explanation. If the requisite is so simple as a potential alkalinity, why has not the
addition of sodium bicarbonate to a deficient diet controlled dental caries?
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It’s all
about
balance!
Let’s
examine
the pH of
foods and
see if there
is a
“problem”
with
popular
concepts!!
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The American Journal of CLINICAL NUTRITION
Volume 68 Number 64 October 1998 pp859-865
Dietary Protein Affects Intestinal Calcium Absorption
Background: Changes in dietary protein in adults are associated
with changes in urinary calcium excretion. The mechanisms
underlying this effect are not completely understood, but
alterations in intestinal absorption of calcium are not thought to be
involved
Method: The effect of two weeks of a well balanced diet followed
by five days of either a low-protein or a high protein diet on
urinary excretion of calcium was determined.
Results: Subjects developed hypocalcuria and secondary
hyperparathyroidism on day 4 of the low-protein diet. Urinary
excretion of calcium and the glomerular filtration rate were
elevated significantly by day 4 of the high-protein diet as
compared to the low-protein diet. (Controversy exists regarding
the effect of dietary protein on bone and intestine.)
American Journal of Clinical Nutrition, Vol. 78, No. 3, 584S-592S, September 2003
Dietary protein, calcium metabolism, and skeletal homeostasis, revisited
Jane E Kerstetter, Kimberly O O’Brien and Karl L Insogna 1 From the School of Allied Health,
University of Connecticut, Storrs (JEK); the Johns Hopkins Bloomberg School of Public Health,
Center for Human Nutrition, Baltimore (KOO); and the Yale University School of Internal
Medicine, New Haven, CT (KLI).
• High dietary protein intakes are known to increase urinary calcium
excretion and, if maintained, will result in sustained hypercalciuria.
• To date, the majority of calcium balance studies in humans have not
detected an effect of dietary protein on intestinal calcium absorption
or serum parathyroid hormone. Therefore, it is commonly concluded
that the source of the excess urinary calcium is increased bone
resorption.
• Recent studies from our laboratory indicate that alterations in dietary
protein can, in fact, profoundly affect intestinal calcium absorption. In
short-term dietary trials in healthy adults, we fixed calcium intake at
20 mmol/d while dietary protein was increased from 0.7 to 2.1 g/kg.
Increasing dietary protein induced hypercalciuria in 20 women [from
3.4 ± 0.3 ( ± SE) during the low-protein to 5.4 ± 0.4 mmol/d during the
high-protein diet]. ………………
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American Journal of Clinical Nutrition, Vol. 78, No. 3, 584S-592S, September 2003
• The increased dietary protein was accompanied by a significant
increase in intestinal calcium absorption from 18.4 ± 1.3% to
26.3 ± 1.5% (as determined by dual stable isotopic
methodology). Dietary protein intakes at and below 0.8 g/kg
were associated with a probable reduction in intestinal calcium
absorption sufficient to cause secondary hyperparathyroidism.
• The long-term consequences of these low-protein diet–induced
changes in mineral metabolism are not known, but the diet
could be detrimental to skeletal health. Of concern are several
recent epidemiologic studies that demonstrate reduced bone
density and increased rates of bone loss in individuals habitually
consuming low-protein diets. Studies are needed to determine
whether low protein intakes directly affect rates of bone
resorption, bone formation, or both.
American Journal of Clinical Nutrition, Vol. 75, No. 4, 773-779, April 2002
Calcium intake influences the association of protein intake with rates of bone
loss in elderly men and women
Bess Dawson-Hughes and Susan S Harris 1 From the Calcium and Bone Metabolism Laboratory at the Jean
Mayer US Department of Agriculture Human Nutrition Research Center on Aging at Tufts University, Boston.
• Background: There is currently no consensus on the effect of dietary protein intake on
the skeleton, but there is some indication that low calcium intakes adversely influence
the effect of dietary protein on fracture risk.
• Objective: The objective of the present study was to determine whether supplemental
calcium citrate malate and vitamin D influence any associations between protein
intake and change in bone mineral density (BMD).
• Design: Associations between protein intake and change in BMD were examined in
342 healthy men and women (aged 65 y) who had completed a 3-y, randomized,
placebo-controlled trial of calcium and vitamin D supplementation. Protein intake was
assessed at the midpoint of the study with the use of a food-frequency questionnaire
and BMD was assessed every 6 mo by dual-energy X-ray absorptiometry.
• Results: The mean (±SD) protein intake of all subjects was 79.1 ± 25.6 g/d and the
mean total calcium intakes of the supplemented and placebo groups were 1346 ± 358
and 871 ± 413 mg/d, respectively. Higher protein intake was significantly
associated with a favorable 3-y change in total-body BMD in the supplemented
group (in a model containing terms for age, sex, weight, total energy intake, and
dietary calcium intake) but not in the placebo group. The pattern of change in femoral
neck BMD with increasing protein intake in the supplemented group was similar to
that for the total body.
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American Journal of Clinical Nutrition, Vol. 77, No. 6, 1517-1525, June 2003
Protein intake: effects on bone mineral density and the rate of
bone loss in elderly women
Prema B Rapuri, J Christopher Gallagher and Vera Haynatzka 1 From the Bone
Metabolism Unit (PBR and JCG), Creighton University, School of Medicine (VH),
Omaha.
• Background: The role of dietary protein in bone metabolism is
controversial.
• Objective: We investigated the associations of dietary protein intake
with baseline bone mineral density (BMD) and the rate of bone loss
over 3 y in postmenopausal elderly women.
• Design: Women aged 65–77 y (n = 489) were enrolled in an
osteoporosis intervention trial. We studied the associations of protein
intake as a percentage of energy with baseline BMD and the rate of
bone loss in 96 women in the placebo group (n = 96). We also
examined the effect of the interaction of dietary calcium intake with
protein intake on BMD.
………………………………
• Results: In the cross-sectional study, a higher intake of
protein was associated with higher BMD. BMD was
significantly higher (P < 0.05) in the spine (7%), midradius (6%),
and total body (5%) in subjects in the highest quartile of protein
intake than in those in the lower 2 quartiles. This positive
association was seen in women with calcium intakes > 408
mg/d. There was no significant effect of protein intake on hip
BMD. In the longitudinal study of the placebo group, there was
no association between protein intake and the rate of bone loss.
• Conclusions: The highest quartile of protein intake ( : 72 g/d)
was associated with higher BMD in elderly women at baseline
only when the calcium intake exceeded 408 mg/d. In the
longitudinal study, no association was seen between protein
intake and the rate of bone loss, perhaps because the sample
size was too small or the follow-up period of 3 y was not long
enough to detect changes.
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• American Journal of Clinical Nutrition, Vol 37, 924-929,
Further studies of the effect of a high protein diet as meat on calcium
metabolism
• H Spencer, L Kramer, M DeBartolo, C Norris and D Osis
• Previous studies in this Unit have shown that a high protein intake, given as
meat, did not induce hypercalciuria, except for the initial and temporary
increase in two subjects.
• In the present investigation the long-term effect of a high meat diet on calcium
metabolism was studied for 78 to 132 days in four adult males and the shortterm effect for 18 to 30 days in three subjects.
• Calcium and phosphorus balances and calcium absorption studies, using
47Ca as the tracer, were carried out.
• During the long-term high meat intake and during the short- term high meat
studies, there was no significant change of the urinary or fecal calcium nor of
the calcium balance.
• There was also no significant change of the intestinal absorption of calcium
during the high meat intake.
• These long- and short-term studies have confirmed our previous results that a
high protein intake, given as meat, does not lead to hypercalciuria and does
not induce calcium loss.
Published online ahead of print November 16, 2004, 10.1210/jc.2004-0179
The Journal of Clinical Endocrinology & Metabolism Vol. 90, No. 1 26-31
The Impact of Dietary Protein on Calcium
Absorption and Kinetic Measures of Bone Turnover
in Women
Jane E. Kerstetter, Kimberly O. O’Brien, Donna M. Caseria, Diane E. Wall and Karl L. Insogna
Although high-protein diets induce hypercalciuria in humans, the source of the additional
urinary calcium remains unclear. One hypothesis is that the high endogenous acid load
of a high-protein diet is partially buffered by bone, leading to increased skeletal
resorption and hypercalciuria. We used dual stable calcium isotopes to quantify the
effect of a high-protein diet on calcium kinetics in women. The study consisted of 2 wk of
a lead-in, well-balanced diet followed by 10 d of an experimental diet containing either
moderate (1.0 g/kg) or high (2.1 g/kg) protein. Thirteen healthy women received both
levels of protein in random order. Intestinal calcium absorption increased during the
high-protein diet in comparison with the moderate (26.2 ± 1.9% vs. 18.5 ± 1.6%, P <
0.0001, mean ± SEM) as did urinary calcium (5.23 ± 0.37 vs. 3.57 ± 0.35 mmol/d, P <
0.0001, mean ± SEM). The high-protein diet caused a significant reduction in the fraction
of urinary calcium of bone origin and a nonsignificant trend toward a reduction in the rate
of bone turnover. There were no protein-induced effects on net bone balance. These
data directly demonstrate that, at least in the short term, high-protein diets are not
detrimental to bone.
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BONE LOSS - Protein - There is a tight connection between protein intake and bone metabolism.
Some of the studies show benefits of protein supplementation, while excess protein intake has
been hypothesized to increase bone loss. ….. There has been a positive correlation found between
protein intake and bone mass in premenopausal women. In women on low-calorie diets,
insufficient protein intake can have an adverse effect on bone mass integrity. …... The hypothesis
that high protein diets may be harmful for bone comes from studies indicating that urinary
calcium is positively associated with protein intake, which would suggest a negative calcium
balance and subsequent increase in bone loss. Other studies have shown that a reduction in
dietary protein leads to a decline in calcium absorption and secondary hyperparathyroidism.
There is evidence that increasing protein intake improves bone mineral mass due to an adequate
supply of both calcium and vitamin D. The hypothesis that animal protein generates more sulfuric
acid from sulfur-containing amino acids than a vegetarian diet does not seem to be valid. Protein
derived from grains and legumes delivers as many millimoles of sulfur per gram of protein as
would a purely meat-based diet. The net release of proton buffers from bone mineral does not
appear to contribute significantly to blood acid-base equilibrium. …. In the Nurses’ Health Study,
which had a follow-up of 12 years, the trend toward hip fracture incidence was inversely related
to protein intake. This epidemiologic observation was not seen in other studies. …. In 30,000
women from Iowa who were studied, higher protein intake was associated with a reduced risk of
hip fracture. The association was particularly evident with protein of animal vs vegetable origin.
…. Dietary protein contributes to maintaining bone integrity from early childhood to old age, and
should be recommended for the prevention and treatment of postmenopausal and age-dependent
osteoporosis. “Protein Intake and Bone Health,” Bonjour J-P, Ammann P, et al, Nutrition and
Bone Health, 2004;17:261-277, edited by M. F. Holick and B. Dawson-Hughes, Humana Press
Inc., Totowa, NJ.
Influence of High Protein Diets on Calcium
A variety of studies have shown that increasing dietary protein can induce
excessive losses of calcium in the urine and a negative calcium balance. This effect
is called the calciuric effect of protein. The effect may occur in some persons but
not in others. It is most marked with low intakes of calcium and high intakes of
protein. Many persons in the United States consume such diets. The calciuric effect
of proteins has been demonstrated in a number of controlled studies with human
subjects. The effect can be reduced, and perhaps minimized, by phosphate. A
simultaneous increase in phosphate intake with protein intake may result in only a
small increase in urinary calcium and maintenance of calcium balance. Foods that
are high in protein, such as meat and eggs, also contain high levels of phosphate.
Thus, the potential hypercalcemic effect of the meat or egg protein may be largely
reduced by the phosphate in the same food.
The mechanism by which dietary protein induces an increase in urinary calcium is
not clear. The effect has been attributed, in part, to the catabolism of sulfur
containing amino acids to yield sulfate. Elevated levels of plasma sulfate can form
a complex with calcium. The complex passes into the renal tubule, where it is
poorly reabsorbed, resulting in its excretion in the urine. The mechanism by which
phosphate reverses the hypercalciuric effect of protein is also not clear.
Nutritional Biochemistry, 2nd Edition, Tom Brody, Academic Press, 1999, p772
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Metabolic Acidosis
May result from either an excess of acid or reduced buffering capacity due
to a low concentration of bicarbonate. Excess acid may occur due
increased production of organic acids or, more rarely, ingestion of acidic
compounds.
a) Excess H+ Production: this is perhaps the commonest cause of
metabolic acidosis and results from the excessive production of organic
acids (usually lactic or pyruvic acid) as a result of anaerobic metabolism.
b) Ingestion of Acids: this is an uncommon cause of metabolic acidosis and
is usually the result of poisoning with agents such as ethylene glycol
(antifreeze) or ammonium chloride.
c) Inadequate Excretion of +: this results from renal tubular dysfunction
and usually occurs in conjunction with inadequate reabsorption of
bicarbonate. Any form of renal failure may result in metabolic acidosis.
d) Excessive Loss of Bicarbonate: gastrointestinal secretions are high in
sodium bicarbonate. The loss of small bowel contents or excessive
diarrhea results in the loss of large amounts of bicarbonate resulting in
metabolic acidosis. This may be seen in such conditions as Cholera or
Crohn's disease.
Clinical Manifestations: Headache, lethargy in early stages.
Progresses to coma with severe acidosis. Deep, rapid respirations
(Kussmaul respirations) are indicative of respiratory compensation.
Anorexia, nausea, vomiting, diarrhea, and abdominal discomfort are
common.
CO2 + H20 <= H2CO3 => HCO3- + H+
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Metabolic Alkalosis
May result from the excessive loss of hydrogen ions, the excessive reabsorption
of bicarbonate or the ingestion of alkalis.
a) Excess H+ loss: gastric secretions contain large quantities of hydrogen ions.
Loss of gastric secretions, therefore, results in a metabolic alkalosis. This
occurs in prolonged vomiting for example, pyloric stenosis or anorexia
nervosa.
b) Excessive Reabsorption of Bicarbonate: Bicarbonate and chloride
concentrations are linked. Therefore, if chloride concentration falls or
chloride losses are excessive then bicarbonate will be reabsorbed to maintain
electrical neutrality. Chloride may be lost from the gastro-intestinal tract,
therefore, in prolonged vomiting it is not only the loss of hydrogen ions that
results in the alkalosis but also chloride losses resulting bicarbonate
reabsorption. Chloride losses may also occur in the kidney usually as a result
of diuretic drugs. The thiazide and loop diuretics a common cause of a
metabolic alkalosis. These drugs cause increased loss of chloride in the urine
resulting in excessive bicarbonate reabsorption.
c) Ingestion of Alkalis: alkaline antacids when taken in excess may result in
mild metabolic alkalosis. This is an uncommon cause of metabolic alkalosis.
Clinical Manifestations: dizziness, confusion, paresthesias (tingling),
convulsions, and coma. Carpopedal spasm and other symptoms of
hypocalcemia are similar to those of metabolic acidosis.
Causes of
Alkalosis
Renal and
Respiratory
Buffering
Compensation
to return pH to
normal range
CO2 + H20 <= H2CO3 => HCO3- + H+
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Respiratory Acidosis
•This results when the PaCO2 is above the upper limit of normal, >6kPa
(45mmHg). The relationship between hydrogen ion concentration and
CO2 was discussed earlier (Production of Hydrogen Ions).
•Respiratory acidosis is most commonly due to decreased alveolar
ventilation causing decreased excretion of CO2.
•Less commonly it is due to excessive production of CO2 by aerobic
metabolism.
a) Inadequate CO2 Excretion: the causes of decreased alveolar
ventilation are numerous, they are summarized on the next slide.
b) Excess CO2 Production: respiratory acidosis is rarely caused by
excess production of CO2. This may occur in syndromes such
as malignant hyperpyrexia, though a metabolic acidosis usually
predominates.
•More commonly, modest overproduction of CO2 in the face of
moderately depressed ventilation may result in acidosis. For example,
in patients with severe lung disease a pyrexia or high carbohydrate diet
may result in respiratory acidosis.
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Clinical Manifestations: headache, restlessness, blurred vision,
apprehension followed by lethargy, muscle twitching, tremors,
convulsions, coma. Neurologic symptoms are caused by a decrease in
pH of the CSF and vasodilation because CO2 readily crosses the
blood-brain barrier
Respiratory Alkalosis
•Results from the excessive excretion of CO2, and occurs when
the PaCO2 is less than 4.5kPa (34mmHg).
•This is commonly seen in hyperventilation due to anxiety states.
In more serious disease states, such as severe asthma or
moderate pulmonary embolism, respiratory alkalosis may occur.
•Here hypoxia, due to ventilation perfusion (V/Q) abnormalities,
causes hyperventilation (in the spontaneously breathing
patient).
• As V/Q abnormalities have little effect on the excretion of CO2
the patients tend to have a low arterial partial pressure of
oxygen (PaO2) and low PaCO2.
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Clinical Manifestations: dizziness confusion, paresthesias, convulsions,
and coma. Carpopedal spasms and other symptoms of hypocalcemia
are similar to those of metabolic acidosis.
Dr. Royal Lee on Acid-Base Regulatory Factors:
The following factors tend to promote normalization of pH through
physiological mechanisms. They do not supply mineral elements which
directly influence body chemistry.
Catalyn
3
Most important organ in acid-base balance regulation.
Renatrophin
1-3
Drenatrophin
1-3
Regulation of sodium-potassium-chloride balance.
Thytrophin
1-3
Regulation of calcium-phosphorus ratio.
Note: Normally the endocrine glands, along with the kidneys, regulate the pH of
the blood, just as the pancreas with its insulin regulates the blood sugar levels. The
sex glands, along with the adrenals, seem to be the main endocrines involved.
Particularly after menopause or male climacteric there may be serious changes in
the body’s economy whereby acid-base disorders develop, which may be
aggravated, if not caused by dietary circumstances-just as diabetes is aggravated by
a high sugar intake. Ringer’s experiments help to explain this phenomenon,
“calcium rigor” being the significant observation here. (See “Calcium Therapy in
Diseases of the Cardiovascular System” by Edward Podolsky, M.D., Lee
Foundation Reprint No. 68.)
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Synergist Products:
Pituitrophin
1-3 Trophic control of endocrine system.
Ovatrophin
1-3⎬
Utrophin
1-3⎬ ⎨Specific cell activators
Prostate PMG
1-3⎬
Orchic PMG
1-3⎬
Normalization of the function of the glands and kidneys which
regulate the pH balance should be the primary consideration
(as outlined above). However, in most cases specific support
of either acid or alkaline mineral therapy is necessary where
immediate results are to be obtained, particularly in acute
situations. We therefore list below acidosis and alkalosis as
separate entities.
ALKALOSIS
Physiological Considerations:
The calcium may precipitate out of body fluids (with excess tissue
calcium) where alkalosis exists. Increased contractility of muscles is
noted (“calcium rigor”). The patient may have calcium deficiency
symptoms, but cannot assimilate calcium in any form. The neuritic
pains which may develop are no doubt due to calcium crystals (calcium
carbonate) that form at nerve endings which accumulate and disperse
according to the variations of pH of the blood that follow dietary levels of
acid and alkaline food ingestion, unable to be compensated for by a
depleted endocrine system. See “Acid-Alkaline Diet Control chart”
PREDISPOSING FACTORS
Endocrine insufficiency (kidneys, thyroid, adrenals, gonads, etc.)
Excess carbohydrate intake (citrus fruits, sugar, etc.)
Excessive alkaline ash foods
Mental stress (Loss from gastric hyperacidity, acid urine, etc.)
Fluid loss (perspiration, diarrhea, vomiting, etc.)
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The following products tend to combat the effects of alkalosis:
PRIMARY:
Cal-Amo
Chloride ion source
Phosfood
Ortho-phosphoric acid source
Lecithin Perles
Phosphorus and choline source
Betaine Hydrochloride Hydrochloric acid source
SECONDARY:
Cataplex G
Bio-Dent
Calcifood
Biost
Prostex
Acetic acid metabolism
Bone-source of phosphorus
Bone source of phosphorus
Phosphatase source (enzyme)
Phosphatase source (enzyme)
ACIDOSIS
Physiological Considerations:
When a depletion of the bicarbonates occurs the carbon dioxide
accumulates in the tissues and oxygen brought to the tissues by the
arterial blood cannot be utilized and is carried away by the venous
blood. The effect is equivalent to depriving the individual of oxygen.
Such persons suffer from symptoms of suffocation, dehydration and
are hyperirritable.
PREDISPOSING FACTORS
•Kidney overload, “possible” failure of adrenal mechanisms caused by
excessive mental stress or shock.
•Liver insufficiency (unable to synthesize urea or detoxify waste acids)
•Deficient intake of alkaline-ash foods
•Excessive intake of acid-ash foods
•Inability to metabolize (or excessive intake of) carbohydrates
•Starvation (ketosis tends to develop as stored fats are utilized from
body reserves in liver dysfunction in the absence of carbohydrates)
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The following products tend to combat the effects of acidosis:
PRIMARY:
AC Carbamide
Blood buffer salt
Potassium Bicarbonate Blood buffer salt
Organic Minerals
Source alkaline minerals
SECONDARY:
Cataplex B
Oxidation of lactic and pyruvic acids
Cataplex A
Combats acidosis via kidney function
Cataplex C
Increases O2 carrying capacity of blood
Cataplex F
Calcium diffuser
Arginex
Kidney overload
Note: ARGINEX may be indispensable in chronic cases of acidosis, particularly
where edema is present as in liver cirrhosis, congestive heart failure, nephrosis, ascites,
etc.
CONCLUSION:
Practically the whole array of chronic and acute diseases may in some degree be related
to acid-base disorders. The most significant mineral may be potassium. Recent tests
with radioactive isotopes show that potassium (as a trace mineral may be deficient in
both acidosis and alkalosis and that it functions both as a base and as a trace mineral. In
the Addison’s Syndrome (adrenal insufficiency) the serum potassium is excessive and
side reactions upon potassium administration indicate nutritional support of the
adrenals.
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