Physiologic pH and buffers - acid-base physiology

Summary

This script explains the concept of physiologic pH, which is crucial for maintaining the optimal functioning of cells and enzymes. It details how the body uses logarithmic pH to manage hydrogen ion concentrations, with a balanced pH of around 7.4 being essential. The video discusses the role of buffers, such as the bicarbonate system, phosphate buffers, and plasma proteins, in stabilizing pH levels against fluctuations caused by acids and bases. It also touches on the importance of intracellular buffers like hemoglobin and organic phosphates in maintaining cellular pH balance.

Takeaways

• 🔍 Physiologic pH is a measure of the balance between acids and bases in the body, determined by the concentration of hydrogen ions.
• ⚗️ The pH scale is logarithmic, with a pH of 7.4 representing an optimal hydrogen ion concentration of 40 nanoequivalents per liter.
• 📉 As hydrogen ion concentration increases, pH decreases due to the negative sign in the pH formula.
• 🔄 The relationship between pH and hydrogen ion concentration is not linear, with changes in pH reflecting exponential changes in ion concentration.
• 🌡️ The body maintains a narrow pH range between 7.37 and 7.42, which is crucial for cellular function.
• 🛡️ Buffers act as protective mechanisms to prevent rapid changes in pH, using a weak acid and its conjugate base or a weak base and its conjugate acid.
• 💧 The bicarbonate buffering system, involving carbonic acid (H2CO3) and bicarbonate ions (HCO3-), is the most important in the body.
• 🌬️ Excess carbon dioxide can be eliminated through the lungs, helping to regulate the bicarbonate buffer system.
• 🧬 Proteins, such as hemoglobin in red blood cells, serve as intracellular buffers, binding or releasing hydrogen ions based on the pH environment.
• 🧪 The phosphate buffer system involves dihydrogen phosphate (H2PO4-) and monohydrogen phosphate (HPO4^2-), contributing to pH balance.
• 🏗️ Organic phosphates like ATP and glucose 6-phosphate also act as intracellular buffers, managing excess hydrogen ions within cells.

Q & A

What is the definition of physiologic pH?

Physiologic pH is a measure of the balance between acids and bases in the body, defined by the concentration of hydrogen ions and expressed with the equation pH = -log base 10 of the hydrogen ion concentration.

What is the optimal hydrogen ion concentration for cells and enzymes in our tissues and organs?

The optimal hydrogen ion concentration for cells and enzymes is 40 times 10 to the minus 9 equivalents per liter, or 40 nanoequivalents per liter.

Why is it important to express the hydrogen ion concentration as pH rather than using the actual concentration?

Expressing the hydrogen ion concentration as pH is important because the actual concentrations are very small and difficult to work with. The pH scale is a logarithmic function that simplifies these tiny numbers into a more manageable scale.

How does the pH scale relate to the concentration of hydrogen ions?

The pH scale is a logarithmic inverse relationship to the hydrogen ion concentration. As hydrogen ion concentrations increase, the pH decreases, and vice versa, due to the negative sign in the pH equation.

What is the normal pH range for the human body?

The normal pH range for the human body is between 7.37 and 7.42.

What happens when the body's pH drops below 7.4?

When the body's pH drops below 7.4, it is considered acidemia, indicating a more acidic state.

What is the role of buffers in maintaining the body's pH?

Buffers act like a protective cushion or shield to prevent the pH from rising or falling too quickly. They handle excess hydrogen ions without causing a major shift in the overall pH.

What is the most important buffer in the body and how does it work?

The most important buffer in the body is the bicarbonate buffer system, consisting of the weak acid carbonic acid (H2CO3) and its conjugate base bicarbonate ion (HCO3-). It works by taking on extra hydrogen ions to form carbonic acid, which can then be converted to carbon dioxide and water, helping to maintain pH balance.

How does the body handle excess carbon dioxide or bicarbonate ions?

Excess carbon dioxide is eliminated through the lungs by breathing, while excess bicarbonate ions are removed by the kidneys in the urine.

What are the other buffering systems in the body besides the bicarbonate buffer system?

Besides the bicarbonate buffer system, the body also has the phosphate buffer system and plasma proteins, which serve as important buffering systems to maintain pH balance.

How do intracellular buffers, such as hemoglobin, help maintain the pH inside cells?

Intracellular buffers like hemoglobin can reversibly bind to hydrogen ions or oxygen. When there is a buildup of hydrogen ions, deoxygenated hemoglobin binds to them, preventing the pH from falling too quickly. This process also involves the chloride shift, which helps maintain charge balance.

What is the significance of organic phosphates like ATP in intracellular buffering?

Organic phosphates like ATP are significant intracellular buffers because their phosphate groups can serve as a source or sink for excess hydrogen ions, helping to maintain the pH balance inside cells.

Outlines

00:00

🧪 Understanding pH Balance and Buffers

This paragraph explains the concept of pH as a measure of acidity and alkalinity in the body, based on hydrogen ion concentration. It details the logarithmic relationship between pH and hydrogen ion concentration, emphasizing the importance of maintaining a pH of around 7.4 for optimal cellular and enzymatic function. The paragraph further describes how changes in pH levels affect hydrogen ion concentrations non-linearly and introduces the concept of acidemia and alkalemia. It also explains the role of buffers in stabilizing pH levels, using the bicarbonate buffer system as a primary example, illustrating how it prevents drastic pH changes through the reversible reactions involving carbonic acid and bicarbonate ions. The summary concludes with the explanation of how the body manages excess carbon dioxide and bicarbonate ions through respiration and urinary excretion.

05:02

🌡 Buffer Systems in Maintaining Physiological pH

The second paragraph delves into the mechanisms by which the body maintains its physiological pH using various buffer systems. It starts by describing the response of the bicarbonate buffer system to the introduction of a strong base (NaOH) and a strong acid (HCL), showing how these buffers neutralize the impact of added substances to keep pH stable. The paragraph also introduces the phosphate buffer system and the role of plasma proteins as buffers, highlighting their capacity to act as both acids and bases depending on pH conditions. Additionally, it discusses the limitations of protein buffering due to their limited quantities in plasma. The paragraph further explains the importance of intracellular pH balance, using hemoglobin in red blood cells as an example of an intracellular buffer, and describes the process of the chloride shift that helps maintain charge balance. It concludes by summarizing the role of organic phosphates like ATP as intracellular buffers and recaps the overall function of buffers in resisting pH changes.

10:03

🛡️ Key Buffers in Extracellular and Intracellular Fluids

The final paragraph succinctly outlines the most important buffers in both extracellular and intracellular environments. It identifies bicarbonate, phosphate, and plasma proteins as the key extracellular buffers, while hemoglobin and organic phosphates like ATP are highlighted as the primary intracellular buffers. This paragraph serves as a concise summary of the primary buffering agents in the body, emphasizing their crucial role in maintaining the pH balance necessary for life.

Keywords

💡Physiologic pH

Physiologic pH refers to the measure of the balance between acids and bases in the body. It is crucial for understanding the body's internal environment, as it directly relates to the concentration of hydrogen ions. In the video, it is defined by the equation pH = -log10 of the hydrogen ion concentration, and it is essential for the optimal functioning of cells and enzymes, which work best around a pH of 7.4.

💡Hydrogen Ion Concentration

Hydrogen ion concentration is the amount of hydrogen ions (H+) in a solution and is inversely related to pH. The video explains that a concentration of 40 nanoequivalents per liter (40 nEq/L) corresponds to a pH of 7.4, which is the ideal for human cells and enzymes. Changes in this concentration, even small ones, can significantly affect the body's pH balance.

💡Logarithmic Function

A logarithmic function is used to express the relationship between hydrogen ion concentration and pH. The video emphasizes that because of the negative sign in the pH equation, as hydrogen ion concentration increases, pH decreases. This non-linear relationship is fundamental to understanding how small changes in pH can correspond to larger changes in hydrogen ion concentration.

💡Acidemia

Acidemia is a condition where the body's pH is below the normal range of 7.37 to 7.42. The video describes it as a state of acidity, which can occur due to an increase in hydrogen ion concentration. It is important to note that the body has mechanisms to prevent such pH drops, maintaining homeostasis.

💡Alkalemia

Alkalemia is the opposite of acidemia, where the body's pH rises above the normal range. The video explains that this condition occurs when there is a decrease in hydrogen ion concentration, leading to a higher pH. Like acidemia, the body has buffers to prevent such pH increases.

💡Buffers

Buffers are substances that resist changes in pH by neutralizing added acids or bases. In the context of the video, they are crucial for maintaining the body's pH within a narrow range. They act like a protective shield against rapid pH changes, preventing the body from becoming too acidic or alkaline.

💡Carbonic Acid (H2CO3)

Carbonic acid is a weak acid formed when carbon dioxide combines with water. It plays a central role in the body's buffering system, as it can easily dissociate into hydrogen ions and bicarbonate ions. The video explains how this reversible reaction helps maintain the body's pH by adjusting the concentration of hydrogen ions.

💡Bicarbonate Ion (HCO3-)

The bicarbonate ion is the conjugate base of carbonic acid and is part of the body's primary buffering system. The video describes how it can bind with hydrogen ions to form carbonic acid, which can then be exhaled as carbon dioxide, thus helping to regulate pH.

💡Chloride Shift

The chloride shift is a process mentioned in the video that helps maintain the balance of charges in red blood cells. It involves the exchange of bicarbonate ions moving into the plasma for chloride ions, which is particularly important in the context of hemoglobin binding with hydrogen ions and releasing oxygen in the lungs.

💡Hemoglobin

Hemoglobin is a protein in red blood cells that can bind with both oxygen and hydrogen ions. The video explains its role in intracellular buffering, where it can bind hydrogen ions in tissues with high carbon dioxide levels, helping to prevent a rapid drop in pH.

💡Organic Phosphates

Organic phosphates, such as ATP and glucose 6-phosphate, are molecules that can serve as intracellular buffers. The video mentions that their phosphate groups can either donate or accept hydrogen ions, thus helping to maintain the pH balance within cells.

Highlights

Physiological pH is a measure of the balance between acids and bases in the body.

pH is calculated using the equation pH = -log10 of hydrogen ion concentration.

Cells and enzymes function optimally at a hydrogen ion concentration of 40 nEq/L.

A logarithmic function is used to express hydrogen ion concentration as pH.

An increase in hydrogen ion concentration leads to a decrease in pH.

pH and hydrogen ion concentration have a non-linear relationship.

Acidemia is defined as a body pH below 7.4, and alkalemia as above 7.4.

Buffers act as protective shields to prevent rapid pH changes.

The body constantly generates acids that can shift pH into the acidic zone.

Buffers are typically a weak acid and its conjugate base or a weak base and its conjugate acid.

Carbonic acid (H2CO3) and bicarbonate ion (HCO3-) are the most important body buffers.

The body maintains pH balance through reversible reactions of carbonic acid and bicarbonate.

Excess carbon dioxide is eliminated by the lungs, and bicarbonate by the kidneys.

The phosphate buffer system involves dihydrogen phosphate and monohydrogen phosphate.

Proteins like albumin in plasma serve as important buffer systems.

Proteins can function as both acids and bases, depending on pH conditions.

Intracellular pH balance is maintained by proteins like hemoglobin and organic phosphates.

Hemoglobin binds hydrogen ions when oxygen levels are low and releases them when oxygen levels are high.

Intracellular buffers like ATP help maintain pH balance by binding or releasing hydrogen ions.

Transcripts

foreign

physiologic pH is a way of quantifying

the balance between acids and bases in

the body

in fact the pH depends on the

concentration of hydrogen ions and can

be described with this equation pH

equals negative log base 10 of the

hydrogen ion concentration

now the cells and enzymes in our tissues

and organs work best when the

concentration of hydrogen ions is 40

times 10 to the minus 9 equivalents per

liter

otherwise known as 40 Nano equivalents

per liter

small changes to that number matter a

ton and because it can get annoying

working with such tiny numbers

scientists converted this concentration

into a logarithmic function and

expressed it as pH

in this case a hydrogen ion

concentration of 40 times 10 to the

minus 9 equivalents per liter works out

to a pH of about 7.4

with this logarithmic function there are

two important aspects to remember

first as hydrogen concentrations

increase the pH decreases this is

because of the negative sign in front of

the log

second since it's a logarithmic function

pH and the hydrogen ion concentration

don't have a linear relationship

for example an increase in PH from 7.4

to 7.6 means a decrease in the hydrogen

concentration of 15 Nano equivalents per

liter

taking the same change in PH but going

the other way from 7.4 to 7.2 means an

increase in the hydrogen concentration

of 23 Nano equivalents per liter

this is why the graph of hydrogen ion

concentration versus pH has a curve to

it rather than being a straight line

for simplicity's sake when the body's pH

drops below 7.4 it's considered acidemia

and when it goes above 7.4 it's

considered alkalemia

so due to this logarithmic relationship

a change in PH in the acidic range or

below 7.4 will show a larger change in

hydrogen concentrations than if the same

change happened in BH in the alkaline

range or pH above 7.4

now maintaining a pH between about 7.37

and 7.42 is essential for the human body

and this is accomplished with buffers

in everyday language a buffer is

something that acts like a protective

cushion or Shield

and the same is true of physiologic

buffers they Shield or prevent the pH

from rising or falling too fast

the reason the body needs buffers is

that acids or molecules that readily

give up their hydrogen ion are being

generated by the body all the time

these extra hydrogen ions would shift

the pH into the acidic Zone so the body

needs a way to handle them without a

major shift in the overall pH

to accomplish this buffers can

essentially take on some of these extra

hydrogen ions and therefore keep the pH

from dropping too much

buffers are usually a weak acid with its

conjugate base or a weak base with its

conjugate acid

the most important buffer in the body is

the weak acid carbonic acid or h2co3 and

its conjugate base bicarbonate ion hco3

minus

carbonic acid forms when carbon dioxide

combines with water with the help of the

enzyme Carbonic anhydrase

as a weak acid carbonic acid easily

dissociates into hydrogen ions and

bicarbonate ions

of course these reactions are reversible

and can happen in the opposite direction

as well

in fact because carbonic acid is such a

weak acid when hydrogen ion

concentrations Get Low it'll drop off

its hydrogen ion and the equation moves

to the right producing more bicarbonate

and a hydrogen ion and when there are

lots of hydrogen ions around the

bicarbonate will bind to one and form

carbonic acid which can go the other way

and split into water and carbon dioxide

at that point the extra carbon dioxide

can be breathed out through the lungs

all right so imagine you've got some

extracellular fluid it has a normal

amount of hydrogen ions in it which puts

it in the normal physiologic PH range

now we drop some NaOH in or sodium

hydroxide which is a strong base

that means that in water it completely

dissociates into sodium or na plus ions

and hydroxide or oh minus ions

the hydroxide ions bind to hydrogen ions

and form water which makes the fluid

more basic since there's a decrease in

hydrogen ions which increases the pH a

lot

when there's plenty of carbon dioxide

around though it reacts with the water

to form more carbonic acid which splits

into bicarbonate ions and hydrogen ions

which quickly replaces the hydrogen ions

and therefore buffers the ph and keeps

it in a normal range

on the flip side imagine tossing HCL

into the fluid or hydrochloric acid

which this time is a strong acid

it would fully dissociate into hydrogen

ions and a bunch of chloride ions

without a buffer this would cause it to

become acidic since there's more

hydrogen ions and therefore the pH would

go down

the new hydrogen ions though get

instantly picked up by bicarbonate ions

and converted to carbonic acid which

would then dissociate into carbon

dioxide in water once again this buffer

normalizes the pH

now the great thing here is that there's

a huge supply of carbonic acid in the

body

because it's formed from carbon dioxide

in water which are an abundant Supply in

the body

finally as a last point if there's too

much carbon dioxide it gets blown off by

the lungs and if there are too many

bicarbonate ions then the kidneys

eliminate them in the urine

apart from the bicarbonate buffer system

there's also the phosphate buffer system

this is a mixture of the weak acid

dihydrogen phosphate or h2po4 minus and

its conjugate base monohydrogen

phosphate

or hpo4 2 minus

dihydrogen phosphate has two hydrogens

and it's always ready to give off one of

them and get converted to monohydrogen

phosphate which has a single hydrogen

now the extracellular fluid and

especially the plasma is also full of

proteins like albumin which serve as

another extremely important buffer

system

this is because some of the proteins

amino acids have exposed carboxyl groups

or cooh that act as weak acids meaning

that they're ready to release hydrogen

ions when the pH starts to rise

at the same time other amino acids might

have exposed nh2 or amine groups that

act as weak bases meaning that they can

bind a hydrogen ion and prevent it from

decreasing the pH in other words a

single protein molecule can function

both as an acid as well as a base

depending on the pH it has to deal with

but unlike the bicarbonate buffering

system there are only a limited number

of proteins in our plasma which limits

how much they can buffer

okay until now we've explored the

extracellular fluid but intracellular pH

has to stay balanced as well and this is

the pH inside cells

all cells are full of proteins and those

proteins are the most significant

intracellular buffer

an example is hemoglobin inside red

blood cells

hemoglobin can reversibly bind either

hydrogen ions which bind to the protein

itself or oxygen which binds to the iron

of the heme group

and the interesting thing is that when

one of these is bound the other is

released

so when red blood cells are in the

capillaries of various tissues the

oxygen levels are low and the carbon

dioxide levels are high

in this case carbon dioxide diffuses

into red blood cells where Carbonic

anhydrous enzyme combines water and

carbon dioxide to form carbonic acid

which then dissociates the hydrogen ion

in bicarbonate

because there's so much carbon dioxide

around the reaction continues in the

direction of generating more hydrogen

ion and bicarbonate and over time the

buildup of hydrogen ions can cause the

pH to fall

to prevent the hydrogen ion

concentration from rising too quickly

each deoxygenated hemoglobin binds

hydrogen ions

the bicarbonate on the other hand moves

into the plasma in exchange for chloride

ions

this exchange called the chloride shift

keeps positive and negative charges

balanced

now once the red blood cell gets into

the capillaries of the lungs there's a

low level of carbon dioxide and a high

level of oxygen so the process reverses

oxygen binds to hemoglobin and hydrogen

ions get released

bicarbonate re-enters the red blood

cells and combines with hydrogen ions to

form carbonic acid which dissociates

into water and carbon dioxide then the

carbon dioxide gets breathed out by the

lungs

intracellular fluid also has large

amounts of organic phosphates like

adenosine triphosphate or ATP and

glucose 6-phosphate

which can both serve as intracellular

buffers

the phosphate group of these organic

molecules can serve as a source or a

sink for excess hydrogen ions

alright as a quick recap

between 7.37 and 7.42 is physiologic pH

and the body uses buffering systems to

maintain pH within this range

buffers are pairs of a weak acid in its

conjugate base or a weak base in its

conjugate acid and their physiologic

role in our body is to resist pH changes

the most important buffers in the

extracellular fluid are bicarbonate

phosphate and plasma proteins

and the most important buffers within

cells are proteins like hemoglobin and

organic phosphates like ATP