Acids and Bases.mp4


One chemical process that has important and
diverse biological implications is the dissociation of acids and bases into ions when in water.
Acids and bases are all around us – citrus fruits are acidic because of citric acid,
and ammonia, used in cleaning solutions, is basic. In this tutorial, we’ll use bonding
principles from the Foundation II tutorial to understand what acids and bases are. We’ll
then see how these principles apply to protein folding as well as our bodies’ regulation
of pH. When placed into water, some molecules and
compounds release H+ ions into the water. These are acids. For example, think about
hydrochloric acid, made of H+ and Cl-. These ions are formed through electron transfer
– hydrogen transfers one electron to chlorine, resulting in full valence electron shells
for both ions. (As a side note – when hydrogen loses an electron to form H+, it now only
has a proton, because hydrogen has no neutrons. So H+ is often referred to simply as a proton.)
In a solution of pure hydrochloric acid, these charged ions form ionic bonds. But when in
water, more favorable interactions can be formed if H+ associates with water molecules
instead of with chloride. These associations can form H3O+ ions, called hydronium ions.
Adding H+, also called a proton, to a molecule is called protonation. These hydronium ions make the water more acidic.
Acidity is a measure of hydronium ion concentration. Because hydronium is made of water and H+,
acidity is often referred to as the H+ concentration – but these H+ ions rarely exist alone and
instead form favorable interactions with partially negatively charged oxygen atoms of water.
Even pure water has some hydronium ions, because the covalent bonds between oxygen and hydrogen
can transiently be broken and reformed, creating OH- and H3O+ ions as H+ skips from one molecule
to another. These double arrows indicate that the process readily occurs in either direction.
In pure water, the concentration of hydronium ions is equal to the concentration of OH-
ions, and is 10-7 M – approximately 2 out of a billion water molecules exist as H3O+. Because these numbers are hard to use, we
use the pH scale as a measure of acidity instead. pH stands for “power of hydrogen” and
is equal to the negative log of the hydronium ion concentration (also written as the negative
log of the H+ concentration). So the pH of pure water is 7. As the hydronium ion concentration
increases, the pH decreases – so a more acidic solution has a lower pH. Think back to hydrochloric acid. When in water,
the favorability of interactions between H+ and water is much greater than the favorability
of interactions between H+ and Cl-; in other words, it’s much more likely that the H+
ions from HCl will exist as hydronium ions than as HCl. HCl is likely to ionize into
H+ and Cl-, which then interact separately with water molecules. Because of this, virtually
all H+ ions from HCl form hydronium ions, contributing to the acidity of water. This
property makes HCl a strong acid. Note that chloride doesn’t directly affect pH; in
other words, pH is not a measure of charge, it’s a measure of hydronium concentration. Many acids found in the body are weak acids,
which less readily give up H+ ions and thus less readily ionize. For example, carboxylic
acids, which are found on some amino acid side chains as well as the C-terminus of all
polypeptide chains, are weak acids. Some carboxylic acids give up H+ ions to water, but others
keep the H+ ion interacting with the carboxylic acid. This is because carboxylic acids interact
more strongly with H+ than chloride does in the case of hydrochloric acid. So the difference
in favorability between the H+ interacting with water and the H+ interacting with the
carboxylic acid is much lower than we saw in the case of hydrochloric acid. Some carboxylic
acids give up H+, while others do not. Only the H+ ions that are given up contribute to
the acidity of the water. The ratio of carboxylic acids that have and have not given up H+ remains
constant at a given pH. If we put the same amount of HCl and carboxylic
acid into water, HCl would donate more H+ than carboxylic acid would. HCl would make
the water more acidic than carboxylic acid would. Molecules or compounds that bind H+ from water,
and thus decrease the concentration of hydronium ions, are bases. Like we saw for acids, there
are two categories of bases: strong bases and weak bases. An example of a strong base
is sodium hydroxide, or NaOH. When in water, sodium hydroxide fully dissociates into Na+
and OH- to maximize favorable interactions. OH- is called hydroxide, and it can interact
with a hydronium ion to form two water molecules. This converts a hydronium ion into a water
molecule, decreasing the hydronium ion concentration – and thus decreasing acidity. Remember
that because the pH scale is the negative log of hydronium ion concentration, a decrease
in acidity is an increase in pH. Most bases in the body are weak bases. An
example of a weak base is an amino group, or -NH2, which is found on some amino acid
side chains, the N-terminus of polypeptide chains, and nitrogenous bases. (In fact, the
basicity of these amino groups is what gives nitrogenous bases their name, although they
are weak bases.) The nitrogen of amino groups can interact with a hydrogen ion to form -NH3+.
This hydrogen ion could have come from hydronium, leaving water. Or the hydrogen ion could have
come from water, leaving OH- – which can then take a hydrogen ion from hydronium, leaving
water. Either way, the effect is to lower the hydronium ion concentration and thus increase
pH. Like weak acids, weak bases only partially ionize – only some amino groups pick up
a hydrogen ion to form NH3+, while others stay as NH2. How do acidity and basicity relate to protein
folding? Remember from the Amino Acids and Protein Structure tutorial that one type of
interaction in the tertiary structure of proteins is ionic bonding between charged atoms of
the side chain. For example, the positively charged nitrogen of lysine can interact with
the negatively charged oxygen of aspartic acid. But these charges only exist in a certain
pH range. If the cellular environment is very acidic, then the concentration of hydronium
ions is high, and it’s likely that a hydrogen ion from a hydronium ion can skip over to
the oxygen of aspartic acid, protonating it. Now, aspartic acid is no longer negatively
charged, so it no longer interacts favorably with the positively charged lysine. Without
these types of interactions, the folded form of the protein is no longer as stable. In
other words, the change in enthalpy going from the unfolded form to the folded form
is less negative. This can tip the
delta G of protein folding from negative – spontaneous – to positive – non-spontaneous. As a
result, proteins can denature in an acidic environment. Our digestive system takes advantage
of this – the acidic environment of the stomach causes proteins we eat to unfold,
increasing access for enzymes like pepsin to cleave peptide bonds and thus break down
proteins. While the pH of the stomach very low pH, the
pH of the rest of the blood is tightly maintained between 7.3 and 7.4 – slightly more basic
than pure water. Even in the presence of external influences that have the effect of increasing
or decreasing pH, the body adjusts to remain at homeostasis. This occurs at three levels:
chemical buffers, control of breathing rate, and control of bicarbonate and acid secretion
in the kidneys. Buffers are weak acids or bases that can either
donate or accept hydrogen ions, depending on the pH. An example of a buffer in our blood
is bicarbonate, or HCO3-. Bicarbonate comes form carbonic acid, or H2CO3, which can dissociate
into bicarbonate and H+. This H+ associates with water to form hydronium, but for this
demonstration, I’ll leave it as H+ for simplicity. Carbonic acid is a weak acid, so some carbonic
acid molecules give up H+ while others do not. Under normal cell conditions, the ratio
of carbonic acid to bicarbonate is about 1:10 – for every carbonic acid molecule, there
are about 10 bicarbonate ions and 10 H+ ions that came from that dissociation. But if the pH of the blood decreases, in other
words, if the hydronium ion concentration increases, there’s more H+ around. As a
result, more bicarbonate ions combine with H+ from hydronium to form carbonic acid. And
the hydronium ion concentration decreases – and thus the pH increases. So buffers
respond to a decrease in pH by increasing pH. Likewise, if the pH of the blood increases,
in other words, if the hydronium ion concentration decreases, there’s less H+ around and relatively
more OH- around. These hydroxide ions can take H+ from carbonic acid, forming bicarbonate.
As a result, the OH- concentration decreases – and remember from above that a decrease
in OH- is functionally equivalent to an increase in hydronium ions. So buffers respond to an
increase in pH by decreasing pH. How does the respiratory system regulate blood
pH? Because of the enzyme carbonic anhydrase, which converts carbon dioxide and water into
carbonic acid, carbon dioxide produced by our cells travels through the blood as carbonic
acid. The more carbonic acid there is, the more bicarbonate there is – because for
every 10 carbonic acid molecules produced, all but one of them will dissociate into bicarbonate
and H+. The H+ formed from dissociation combines with water to form hydronium, decreasing pH.
By controlling how much we breathe, we can control the amount of carbon dioxide in our
body, which means controlling carbonic acid, bicarbonate and H+, and pH. Breathing more
means exhaling more carbon dioxide, which decreases carbonic acid. Because the ratio
of carbonic acid to bicarbonate stays around 1:10, a decrease in carbonic acid causes bicarbonate
and H+ to come back together to return to that ratio. This decreases H+ and increases
pH. The reverse – breathing less – has the opposite effect, increasing H+ and decreasing
pH. The kidneys control the secretion of bicarbonate.
If the blood is very acidic, then the kidneys reabsorb more bicarbonate, moving it back
into the blood where it can act as a buffer for those H+ ions. If the blood isn’t acidic
enough, the opposite happens – bicarbonate is secreted and subsequently excreted in the
urine. This secretion of bicarbonate disrupts the ratio of carbonic acid to bicarbonate.
To compensate, more carbonic acid dissociates into bicarbonate and H+. This dissociation
increases H+ concentration, which increases hydronium concentration and thus decreases
pH, compensating for the initial basicity. As you can see, the ability of our bodies
to control pH is crucial. Protein structure and function is totally dependent on pH. We
have three mechanisms to maintain pH homeostasis, all of which act on different time scales
to ensure optimal physiological pH.

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