BCM 201 Cell Biology PDF

**BCM 201 CELL BIOLOGY** What is pH A figure expressing the [acidity](https://www.google.com/search?sa=X&sca_esv=8686b2b8d75a1600&biw=1366&bih=599&sxsrf=AHTn8zq4_MtyGyyfaLeprWswP_qL1Uz97g:1737935592493&q=acidity&si=APYL9btR06w9iCpfOKePDdxJ7dfzjnB_SBmKgp3jGaHM4W7CSg2p4acBpChao4FVpz8bjcEOa3s82gXzR9k...

**BCM 201 CELL BIOLOGY** What is pH A figure expressing the [acidity](https://www.google.com/search?sa=X&sca_esv=8686b2b8d75a1600&biw=1366&bih=599&sxsrf=AHTn8zq4_MtyGyyfaLeprWswP_qL1Uz97g:1737935592493&q=acidity&si=APYL9btR06w9iCpfOKePDdxJ7dfzjnB_SBmKgp3jGaHM4W7CSg2p4acBpChao4FVpz8bjcEOa3s82gXzR9kOwbUb0yTIk-n2Vw%3D%3D&expnd=1&ved=2ahUKEwjagYjNypSLAxUgSkEAHSpeDI0QyecJegQIJRAO) or alkalinity of a solution on a [logarithmic](https://www.google.com/search?sa=X&sca_esv=8686b2b8d75a1600&biw=1366&bih=599&sxsrf=AHTn8zq4_MtyGyyfaLeprWswP_qL1Uz97g:1737935592493&q=logarithmic&si=APYL9btTB54oNzRD0c75DM-v-cL-IA2wXVj46BEUEZ3J9fnel9rgToZvuraqBigtSJPz4Wi3t11QV9_cFYLoX7C5OJCUHK46d-mlIo1Z_ZLGPuRxnxZ8pT8%3D&expnd=1&ved=2ahUKEwjagYjNypSLAxUgSkEAHSpeDI0QyecJegQIJRAP) scale on which 7 is neutral, lower values are more acid and higher values more alkaline. The pH is equal to −log~10~ *c*, where *c* is the hydrogen ion concentration in [moles](https://www.google.com/search?sa=X&sca_esv=8686b2b8d75a1600&biw=1366&bih=599&sxsrf=AHTn8zq4_MtyGyyfaLeprWswP_qL1Uz97g:1737935592493&q=moles&si=APYL9bseOkRtPplmiFwUxbDEwT9fJavHSuGky_39hxbB8skhnDWQzPUzdYvD1Wgcw8BLz9ZzUfLYWe_KocHn-umqNMUl4sJZig%3D%3D&expnd=1&ved=2ahUKEwjagYjNypSLAxUgSkEAHSpeDI0QyecJegQIJRAQ) per litre. **BIOLOGICAL IMPORTANCE OF pH** In biology, pH is critically important as it directly influences the structure and function of most biological molecules, including enzymes, proteins, and DNA, meaning that maintaining a stable pH within a narrow range is essential for proper cellular processes and organism homeostasis; significant fluctuations in pH can disrupt these processes and lead to cellular damage or malfunction.  Key points about the biological importance of pH: - **Enzyme activity:** Enzymes have optimal pH ranges where their structure and function are maximized; even slight changes in pH can significantly alter their catalytic activity by affecting the ionization state of amino acid residues in their active site.  - **Protein structure:** The charge on protein molecules is influenced by pH, which affects their folding pattern and ability to interact with other molecules.  - **Membrane function:** Cell membranes are selectively permeable to ions, and the pH of the surrounding environment can influence the movement of ions across the membrane.  - **Blood pH regulation:** The pH of blood is tightly regulated by bicarbonate buffer systems, which are crucial for maintaining proper oxygen transport and cellular function.  - **Cellular processes:** Many cellular processes, including DNA replication, transcription, and translation, are sensitive to pH changes.  Factors affecting pH in biological systems: - **Carbon dioxide levels:** The presence of carbon dioxide in the body can contribute to acidity due to the formation of carbonic acid.  - **Metabolic activity:** Cellular metabolism produces acidic byproducts which can alter pH if not properly regulated.  - **Diet:** The food we consume can impact our body\'s pH balance.  **Consequences of pH disruption:** - **Denaturation of proteins:** Extreme pH changes can cause proteins to unfold and lose their functional shape.  - **Enzyme inhibition:** Improper pH can significantly reduce enzyme activity, disrupting metabolic pathways.  - **Cellular damage:** Significant pH fluctuations can disrupt cellular processes and lead to cell death.  Important concepts related to pH in biology: - **Buffers:** Molecules or systems that resist large changes in pH by accepting or donating hydrogen ions.  - **pKa value:** A measure of the acidity of a molecule, indicating its tendency to donate a hydrogen ion.  - **pH gradient:** A difference in pH across a membrane that can be used to drive cellular processes.  In biochemistry, pH is crucial because it significantly impacts the structure and function of biological molecules like proteins and enzymes, directly affecting the biochemical reactions that can occur within an organism, essentially dictating which enzymes are active and thus controlling vital biological processes by maintaining a stable internal environment (homeostasis) within the body; a slight change in pH can disrupt these functions drastically.  **Key points about pH in biochemistry:** - **Enzyme activity:** Most enzymes have an optimal pH range where they function best, and any deviation from this range can significantly reduce their activity.  - **Protein structure:** The 3D structure of proteins is heavily influenced by pH due to the charge of amino acid side chains, which can change their interactions with each other and affect protein function.  - **DNA stability:** The hydrogen bonds holding DNA strands together are sensitive to pH, and significant changes can disrupt DNA structure and function.  - **Membrane transport:** The movement of ions across cell membranes is influenced by pH gradients, which are crucial for various cellular processes.  - **Blood pH regulation:** Maintaining a stable blood pH is vital for proper physiological function, and the body has sophisticated mechanisms to buffer against pH changes.  Potential consequences of pH imbalance: - **Acidosis:** When the body becomes too acidic, it can lead to symptoms like confusion, fatigue, and impaired breathing. - **Alkalosis:** When the body becomes too alkaline, it can cause muscle cramps, tremors, and irregular heartbeat.  **What are the main uses of pH in different industries?** Any industry that uses water, a lot of them---needs to test pH levels, at least on a periodic basis, to ensure water composition falls within desirable ranges. Otherwise, when pH levels fall outside preferred ranges, food may be unsafe to eat, beverages may be too acidic to drink, spa water may be too irritating to the eyes and skin, and research experiments may reached skewed conclusions, among other bad outcomes. **Here are seven industries that commonly test pH.\ ** **1. Wastewater treatment** During wastewater treatment, heavy metals, organic compounds, and other toxic substances are removed from water; pH levels need to be adjusted during this process by adding chemicals to the water to separate dissolved waste from the liquid. Managers at facilities that process sewage or [recycle water used](https://blog.jencoi.com/4-instruments-used-in-industrial-wastewater-treatment) in manufacturing need to monitor pH levels to ensure that water can safely pass to the next phase of treatment so that they end up with pure, safe water at the end of the cleaning process. **2. Aquaculture** Aquaculture is a multibillion-dollar industry that is responsible for supplying nearly half of the world's seafood. To keep aquatic creatures alive and ensure they're healthy enough to be consumed, companies in this space need to monitor changes in pH levels on a regular basis. Generally speaking, fish thrive in pH levels between 6.5 and 9.0. **3. Food and beverage** It is important to monitor pH levels in the production of any food or beverage product. For example, when water used in beverage manufacturing is too acidic, consumers can potentially damage their dentition. Companies also need to monitor pH levels in [food production](https://ourdailybrine.com/how-to-test-the-ph-of-food-and-drink/#role-of-ph-in-food) to ensure their products are safe to eat and of high quality. For example, pork with a pH of 5.6 to 6.0 indicates that a pig was raised well, while a lower pH reading indicates the pig might have lived a stressful life. What's more, a meat product like salami should have a pH under 5.3 to protect against bacteria growth. **4. Pool and spa water** Whether you're managing a large community swimming pool or a private spa, you need to measure pH levels to determine how many sterilizing agents (e.g., chlorine) you need to put in the water to maintain a safe swimming environment. When pH levels are too high, water gets cloudier, and swimmers' skin and eyes may get irritated. When pH levels are too low, swimmers' skin and eyes may get irritated, too, and your pool or spa's plaster can get damaged as well. The ideal pH reading for a spa is between 7.2 and 7.8. **5. Aquariums and fish tanks** To keep fish and other aquatic creatures alive---and to keep fish tanks and aquariums clean---many water composition characteristics need to be proactively monitored, including pH. Otherwise, at best, the water in the aquarium or fish tank will look cloudy and, at worst, fish may struggle to survive.\   **6. Research** Water plays an integral role in many research projects. While the requirements of each experiment are unique, lab science demands highly accurate results. In experiments, researchers need to monitor pH levels to ensure high-quality findings that can be repeated. For example, different types of bacteria thrive in environments with different pH levels. To reach accurate conclusions, it is, therefore, critical to monitor pH levels in any experiment involving bacteria. **7. Hydroponics** The practice of growing plants in a nutrient-enriched water-based solution instead of soil is called[ hydroponics](https://blog.jencoi.com/the-best-water-quality-instruments-for-commercial-hydroponic-systems). If the nutrient balance or pH level of the water swings too much in one direction, plants can die rather quickly. To this end, managers at hydroponic facilities keep a close eye on pH levels at all times. When water is involved in an industrial application, it's likely that water quality characteristics---including pH levels---need to be measured to ensure safety and compliance. The most common methods to determine the pH of a medium are using a pH meter (which measures the electrical potential between a glass electrode and a reference electrode), pH indicator paper (which changes color depending on the pH level), and indicator solutions (which exhibit distinct color changes at specific pH ranges) - all of these methods rely on the principle of comparing the solution\'s hydrogen ion concentration to a known standard; with pH meters providing the most accurate readings, while indicator methods offer a quicker, less precise assessment.  Key points about pH determination methods: - **pH meter:** - Considered the most accurate method.  - Uses a glass electrode to measure the potential difference between the solution and a reference electrode.  - Can measure a wide range of pH values with good precision.  - **pH indicator paper:** - Simple and portable method.  - Provides a rough estimate of pH based on color change when dipped in the solution.  - Can only indicate whether a solution is acidic, neutral, or basic.  - **Indicator solutions:** - Similar to pH paper, but can provide a more refined pH range based on the specific indicator used.  - Examples of indicators include phenolphthalein, methyl orange, and bromothymol blue.  Other methods (less commonly used): - **Conductivity meter:** Measures the electrical conductivity of a solution, which can be related to the concentration of ions (including hydrogen ions).  - **Hydrogen electrode:** Considered the standard method for pH measurement but is not often used due to its complexity.  - **Antimony electrode:** It can be used to measure pH in certain situations, but is less accurate than a glass electrode.  **BUFFER SYSTEM** **Acid-Base Balance** *Learning Objectives* By the end of this section, you will be able to: - Identify the most powerful buffer system in the body - Identify the most rapid buffer system in the body - Describe the protein buffer systems. - Explain the way in which the respiratory system affects blood pH - Describe how the kidney affects acid-base balance Proper physiological functioning depends on a very tight balance between the concentrations of acids and bases in the blood. Acid-balance is measured using the pH scale, as shown in Figure 1. A variety of buffering systems permits blood and other bodily fluids to maintain a narrow pH range, even in the face of perturbations. A buffer is a chemical system that prevents a radical change in fluid pH by dampening the change in hydrogen ion concentrations in the case of excess acid or base. Most commonly, the substance that absorbs the ions is either a weak acid, which takes up hydroxyl ions, or a weak base, which takes up hydrogen ions. **Figure 1 -- The pH Scale:**  This chart shows where many common substances fall on the pH scale. **Buffer Systems in the Body** The buffer systems in the human body are extremely efficient, and different systems work at different rates. It takes only seconds for the chemical buffers in the blood to make adjustments to pH. The respiratory tract can adjust the blood pH upward in minutes by exhaling CO~2~ from the body. The renal system can also adjust blood pH through the excretion of hydrogen ions (H^+^) and the conservation of bicarbonate, but this process takes hours to days to have an effect. The buffer systems functioning in blood plasma include **plasma proteins, phosphate, and bicarbonate and carbonic acid buffers**. **The kidneys help control acid-base balance by excreting hydrogen ions and generating bicarbonate that helps maintain blood plasma pH within a normal range**. Protein buffer systems work predominantly inside cells. **Protein Buffers in Blood Plasma and Cells** Nearly all proteins can function as buffers. Proteins are made up of amino acids, which contain positively charged amino groups and negatively charged carboxyl groups. The charged regions of these molecules can bind hydrogen and hydroxyl ions, and thus function as buffers. Buffering by proteins accounts for two-thirds of the buffering power of the blood and most of the buffering within cells. **Hemoglobin as a Buffer** Hemoglobin is the principal protein inside of red blood cells and accounts for one-third of the mass of the cell. During the conversion of CO~2~ into bicarbonate, hydrogen ions liberated in the reaction are buffered by hemoglobin, which is reduced by the dissociation of oxygen. This buffering helps maintain normal pH. The process is reversed in the pulmonary capillaries to re-form CO~2~, which then can diffuse into the air sacs to be exhaled into the atmosphere. This process is discussed in detail in the chapter on the respiratory system. **Phosphate Buffer** Phosphates are found in the blood in two forms: sodium dihydrogen phosphate (Na~2~H~2~PO~4~^−^), which is a weak acid, and sodium monohydrogen phosphate (Na~2~HPO4^2-^), which is a weak base. When Na~2~HPO4^2-^ comes into contact with a strong acid, such as HCl, the base picks up a second hydrogen ion to form the weak acid Na~2~H~2~PO~4~^−^ and sodium chloride, NaCl. When Na~2~HPO4^2−^ (the weak acid) comes into contact with a strong base, such as sodium hydroxide (NaOH), the weak acid reverts back to the weak base and produces water. Acids and bases are still present, but they hold onto the ions. HCl + Na~2~HPO~4~→NaH~2~PO~4~ + NaCl (strong acid) + (weak base) → (weak acid) + (salt) NaOH + NaH~2~PO~4~→Na~2~HPO~4~ + H~2~O (strong base) + (weak acid) → (weak base) + (water) **Bicarbonate-Carbonic Acid Buffer** The bicarbonate-carbonic acid buffer works in a fashion similar to phosphate buffers. The bicarbonate is regulated in the blood by sodium, as are the phosphate ions. When sodium bicarbonate (NaHCO~3~), comes into contact with a strong acid, such as HCl, carbonic acid (H~2~CO~3~), which is a weak acid, and NaCl are formed. When carbonic acid comes into contact with a strong base, such as NaOH, bicarbonate and water are formed. NaHCO~3~ + HCl →  H~2~CO~3~+NaCl (sodium bicarbonate) + (strong acid) → (weak acid) + (salt) H~2~CO~3~ + NaOH→HCO~3-~ + H~2~O (weak acid) + (strong base)→(bicarbonate) + (water) As with the phosphate buffer, a weak acid or weak base captures the free ions, and a significant change in pH is prevented. Bicarbonate ions and carbonic acid are present in the blood in a 20:1 ratio if the blood pH is within the normal range. With 20 times more bicarbonate than carbonic acid, this capture system is most efficient at buffering changes that would make the blood more acidic. This is useful because most of the body's metabolic wastes, such as lactic acid and ketones, are acids. Carbonic acid levels in the blood are controlled by the expiration of CO~2~ through the lungs. In red blood cells, carbonic anhydrase forces the dissociation of the acid, rendering the blood less acidic. Because of this acid dissociation, CO~2~ is exhaled (see equations above). **The level of bicarbonate in the blood is controlled through the renal system, where bicarbonate ions in the renal filtrate are conserved and passed back into the blood**. However, the bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. CO~2~ + H~2~O ↔ H~2~CO~3~ ↔ H^+^ + HCO~3~^--^ **Respiratory Regulation of Acid-Base Balance** The respiratory system contributes to the balance of acids and bases in the body by regulating the blood levels of carbonic acid (figure 2). CO~2 ~in the blood readily reacts with water to form carbonic acid, and the levels of CO~2 ~and carbonic acid in the blood are in equilibrium. When the CO~2 ~level in the blood rises (as it does when you hold your breath), the excess CO~2~ reacts with water to form additional carbonic acid, lowering blood pH. Increasing the rate and/or depth of respiration (which you might feel the "urge" to do after holding your breath) allows you to exhale more CO~2~. The loss of CO~2~ from the body reduces blood levels of carbonic acid and thereby adjusts the pH upward, toward normal levels. In summary, this process also works in the opposite direction. Excessive deep and rapid breathing (as in hyperventilation) rids the blood of CO~2~ and reduces the level of carbonic acid, making the blood too alkaline. This brief alkalosis can be remedied by rebreathing air that has been exhaled into a paper bag. Rebreathing exhaled air will rapidly bring blood pH down toward normal. **Figure 2 -- Respiratory Regulation of Blood pH:**  The respiratory system can reduce blood pH by removing CO~2~ from the blood. The chemical reactions that regulate the levels of CO~2~ and carbonic acid occur in the lungs when blood travels through the lung's pulmonary capillaries. Minor adjustments in breathing are usually sufficient to adjust the pH of the blood by changing how much CO~2~ is exhaled. In fact, doubling the respiratory rate for less than 1 minute, removing "extra" CO~2~, would increase the blood pH by 0.2. This situation is common if you are exercising strenuously over a period of time. To keep up the necessary energy production, you would produce excess CO~2~ (and lactic acid if exercising beyond your aerobic threshold). In order to balance the increased acid production, the respiration rate goes up to remove the CO~2~. This helps to keep you from developing acidosis. **The body regulates the respiratory rate by the use of chemoreceptors, which primarily use CO~2~ as a signal. Peripheral blood sensors are found in the walls of the aorta and carotid arteries. These sensors signal the brain to provide immediate adjustments to the respiratory rate if CO~2 ~levels rise or fall.** Yet other sensors are found in the brain itself. Changes in the pH of CSF affect the respiratory center in the medulla oblongata, which can directly modulate breathing rate to bring the pH back into the normal range. Hypercapnia, or abnormally elevated blood levels of CO~2~, occurs in any situation that impairs respiratory functions, including pneumonia and congestive heart failure. **Reduced breathing (hypoventilation) due to drugs such as morphine, barbiturates, or ethanol (or even just holding one's breath) can also result in hypercapni**a. Hypocapnia, or abnormally low blood levels of CO~2~, occurs with any cause of hyperventilation that drives off the CO~2~, such as salicylate toxicity, elevated room temperatures, fever, or hysteria. **RENAL REGULATION OF ACID-BASE BALANCE** The renal regulation of the body's acid-base balance addresses the metabolic component of the buffering system. W**hereas the respiratory system (together with breathing centers in the brain) controls the blood levels of carbonic acid by controlling the exhalation of CO~2~, the renal system controls the blood levels of bicarbonate**. A decrease of blood bicarbonate can result from the inhibition of carbonic anhydrase by certain diuretics or from excessive bicarbonate loss due to diarrhea. Blood bicarbonate levels are also typically lower in people who have Addison's disease (chronic adrenal insufficiency), in which aldosterone levels are reduced, and in people who have renal damage, such as chronic nephritis. Finally, low bicarbonate blood levels can result from elevated levels of ketones (common in unmanaged diabetes mellitus), which bind bicarbonate in the filtrate and prevent its conservation. Bicarbonate ions, HCO~3~^--^, found in the filtrate, are essential to the bicarbonate buffer system, yet the cells of the tubule are not permeable to bicarbonate ions. The steps involved in supplying bicarbonate ions to the system are seen in figure 3 and are summarized below: - Step 1: Sodium ions are reabsorbed from the filtrate in exchange for H^+^ by an antiport mechanism in the apical membranes of cells lining the renal tubule. - Step 2: The cells produce bicarbonate ions that can be shunted to peritubular capillaries. - Step 3: When CO~2~ is available, the reaction is driven to the formation of carbonic acid, which dissociates to form a bicarbonate ion and a hydrogen ion. - Step 4: The bicarbonate ion passes into the peritubular capillaries and returns to the blood. The hydrogen ion is secreted into the filtrate, where it can become part of new water molecules and be reabsorbed as such, or removed in the urine. This diagram depicts a cross section of the left wall of a kidney proximal tubule. The wall is composed of two block-shaped cells arranged vertically one on top of each other. The lumen of the proximal tubule is to the left of the two cells. Yellow-colored urine is flowing through the lumen. There is a small strip of blue interstitial fluid to the right of the two cells. To the right of the interstitial fluid is a cross section of a blood vessel. A loop of chemical reactions is occurring in the diagram. Within the lumen of the proximal tubule, HCO three minus is combining with an H plus ion that enters the lumen from a proximal tubule cell. This reaction forms H two CO three. H two CO three then breaks into H two O and CO two, a reaction catalyzed by the enzyme carbonic anhydrase. The CO two then moves from the lumen of the proximal tubule into one of the proximal tubule cells. There, the reaction runs in reverse, with CO two combining with H two O to form H two CO three. The H two CO three then splits into H plus and HCO three minus. The H plus moves into the lumen, reinitiating the first step of the loop. The HCO three minus leaves the proximal tubule cell and enters the blood stream. **Figure 3 Conservation of Bicarbonate in the Kidney.**  Tubular cells are not permeable to bicarbonate; thus, bicarbonate is conserved rather than reabsorbed. Steps 1 and 2 of bicarbonate conservation are indicated. It is also possible that salts in the filtrate, such as sulfates, phosphates, or ammonia, will capture hydrogen ions. If this occurs, the hydrogen ions will not be available to combine with bicarbonate ions and produce CO~2~. In such cases, bicarbonate ions are not conserved from the filtrate to the blood, which will also contribute to a pH imbalance and acidosis. The hydrogen ions also compete with potassium to exchange with sodium in the renal tubules. If more potassium is present than normal, potassium, rather than the hydrogen ions, will be exchanged, and increased potassium enters the filtrate. When this occurs, fewer hydrogen ions in the filtrate participate in the conversion of bicarbonate into CO~2~ and less bicarbonate is conserved. If there is less potassium, more hydrogen ions enter the filtrate to be exchanged with sodium and more bicarbonate is conserved. Chloride ions are important in neutralizing positive ion charges in the body. If chloride is lost, the body uses bicarbonate ions in place of the lost chloride ions. Thus, lost chloride results in an increased reabsorption of bicarbonate by the renal system. **ROLE OF PHYSIOLOGICAL BUFFER IN A LIVING SYSTEM** In a living system, physiological buffers play a critical role in maintaining a stable pH level within the body fluids by readily absorbing excess hydrogen ions (H+) or hydroxide ions (OH-) when introduced, thus preventing drastic changes in pH that could disrupt essential biological processes; the primary physiological buffers include the bicarbonate buffer system (most important in blood), phosphate buffer system (intracellular), and protein buffer system (both intracellular and extracellular), with each contributing to maintaining homeostasis across different body compartments.  **Key points about physiological buffers**: - **Function:** Buffers resist pH changes by acting as a reservoir for H+ ions, allowing them to be released or absorbed as needed to maintain a stable pH range.  - **Components:** A buffer system typically consists of a weak acid and its conjugate base, which can readily donate or accept H+ ions depending on the pH of the solution.  - **Major Buffer Systems in the Body:** - **Bicarbonate Buffer System:** The most important extracellular buffer, primarily found in blood, where it consists of carbonic acid (H2CO3) and bicarbonate ions (HCO3-).  - **Phosphate Buffer System:** Primarily intracellular, composed of dihydrogen phosphate (H2PO4-) and monohydrogen phosphate (HPO42-) ions.  - **Protein Buffer System:** Proteins in the blood and cells have amino acid side chains with ionizable groups that can act as buffers, especially the imidazole group of histidine residues.  How these buffers work: - **Bicarbonate Buffer:** - When H+ ions are added to the blood (becoming more acidic), bicarbonate ions (HCO3-) readily react with them to form carbonic acid (H2CO3).  - Carbonic acid then dissociates into carbon dioxide (CO2) and water, which can be eliminated by the lungs through exhalation.  - When OH- ions are added (becoming more alkaline), carbonic acid can donate H+ ions to neutralize the excess base.  - **Phosphate Buffer:** - Functions similarly to the bicarbonate buffer, with the phosphate ions (HPO42-) acting as a base to accept H+ ions when the pH drops.  - **Protein Buffer:** - The amino acid side chains of proteins, particularly histidine residues, can readily accept or donate H+ ions depending on the pH of the environment.  **The Importance of Maintaining pH Balance**: - **Enzyme Function:** Most enzymes have optimal activity within a narrow pH range, and significant pH changes can disrupt their function.  - **Cell Membrane Integrity:** Extreme pH levels can damage cell membranes and disrupt cellular processes.  - **Nervous System Function:** Proper neural transmission depends on a stable pH environment.  **Disorders of the Fluid Balance: Acid-Base Balance: Ketoacidosis** Diabetic acidosis, or ketoacidosis, occurs most frequently in people with poorly controlled diabetes mellitus. When certain tissues in the body cannot get adequate amounts of glucose, they depend on the breakdown of fatty acids for energy. When acetyl groups break off the fatty acid chains, the acetyl groups then non-enzymatically combine to form ketone bodies, acetoacetic acid, beta-hydroxybutyric acid, and acetone, all of which increase the acidity of the blood. In this condition, the brain isn't supplied with enough of its fuel---glucose---to produce all of the ATP it requires to function. Ketoacidosis can be severe and, if not detected and treated properly, can lead to diabetic coma, which can be fatal. A common early symptom of ketoacidosis is deep, rapid breathing as the body attempts to drive off CO~2~ and compensate for the acidosis. Another common symptom is fruity-smelling breath, due to the exhalation of acetone. Other symptoms include dry skin and mouth, a flushed face, nausea, vomiting, and stomach pain. Treatment for diabetic coma is ingestion or injection of sugar; its prevention is the proper daily administration of insulin. A person who is diabetic and uses insulin can initiate ketoacidosis if a dose of insulin is missed. ***Chapter Review*** A variety of buffering systems exist in the body that helps maintain the pH of the blood and other fluids within a narrow range---between pH 7.35 and 7.45. A buffer is a substance that prevents a radical change in fluid pH by absorbing excess hydrogen or hydroxyl ions. Most commonly, the substance that absorbs the ion is either a weak acid, which takes up a hydroxyl ion (OH^--^), or a weak base, which takes up a hydrogen ion (H^+^). Several substances serve as buffers in the body, including cell and plasma proteins, hemoglobin, phosphates, bicarbonate ions, and carbonic acid. The bicarbonate buffer is the primary buffering system of the IF surrounding the cells in tissues throughout the body. The respiratory and renal systems also play major roles in acid-base homeostasis by removing CO~2~ and hydrogen ions, respectively, from the body. **What are the definitions of**: Acidemia Alkalemia Acidemia ◦ The condition of increased \[H+\] in blood ◦ Low blood pH \` Alkalemia ◦ The condition of decreased \[H+\] in blood ◦ High blood pH What are the definitions of: ◦ Acidosis ◦ Alkalosis Acidosis ◦ The disease process which increases \[H+\] \` Alkalosis ◦ The disease process which decreases \[H+\] **What are 5 effects of severe acidemia?** Impairs enzyme function Interferes with electrophysiology Disturbs electrolyte balance Blocks calcium influx into cells Inhibits catecholamine action