Water: Physical, Chemical, and Biological Properties PDF

**WATER: PHYSICAL, CHEMICAL AND BIOLOGICAL PROPERTIES** *Water* is a transparent and virtually colorless chemical substance that is the main constituent of the earth's surface. It is indispensable to life. Its chemical formula is H2O, meaning that each of its molecules contains one oxygen and two h...

**WATER: PHYSICAL, CHEMICAL AND BIOLOGICAL PROPERTIES** *Water* is a transparent and virtually colorless chemical substance that is the main constituent of the earth's surface. It is indispensable to life. Its chemical formula is H2O, meaning that each of its molecules contains one oxygen and two hydrogen atoms that are connected by covalent bonds. **BIOLOGICAL PROPERTIES** **From the biological viewpoint,** water has many distinct properties that are critical for the proliferation of life that set it apart from other substances. It carries out this role by allowing organic compounds to react in ways that ultimately allow replication. All known forms of life depend on water. Water is vital both as a solvent in which many of the body's solutes dissolve and as an essential part of many metabolic processes within the body. Water is fundamental to photosynthesis and respiration. Photosynthetic cells use the suns energy to split off water's hydrogen from oxygen. Hydrogen is combined with CO2 (absorbed from air or water) to form glucose and release oxygen. All living cells use such fuels and oxidize the hydrogen and carbon to capture the suns energy and reform water and CO2 in the process (cellular respiration). **Structure of Water** The two hydrogen atoms of water are linked covalently to oxygen, each sharing an electron pair, to give a nonlinear arrangement. This "bent" structure of the H2O molecule is of enormous significance to its properties. If H2O were linear, it would be a non-polar substance. In the bent configuration, however, the electronegative O atom and the two H atoms form a dipole that renders the molecule distinctly polar. Furthermore, this structure is ideally suited to Hydrogen (H) -bond formation. **Water can serve as both an H donor and an H acceptor in H-bond formation**. The potential to form four H bonds per water molecule is the source of the **strong intermolecular attractions** that endow this substance with its **anomalously** **high boiling point, melting point, heat of vaporization, and surface tension.** Hydrogen bonding in water is cooperative. That is, an H-bonded water molecule serving as an acceptor is a better H-bond donor than an unbounded molecule (and an H2O molecule serving as an H-bond donor becomes a better H-bond acceptor). Thus, participation in H bonding by H2O molecules is a phenomenon of mutual reinforcement. The H bonds between neighboring molecules are weak (23 kJ/mol each) relative to the H-O covalent bonds (420 kJ/mol). As a consequence, the hydrogen atoms are situated asymmetrically https://upload.wikimedia.org/wikipedia/commons/thumb/c/c6/3D\_model\_hydrogen\_bonds\_in\_water.svg/220px-3D\_model\_hydrogen\_bonds\_in\_water.svg.png **Fig. Model of hydrogen bond between molecules of water** ***Water Forms H Bonds with Polar Solutes*** In the case of nonionic but polar compounds such as sugars, the excellent solvent properties of water stem from its ability to readily form hydrogen bonds with the polar functional groups on these compounds, such as hydroxyls, amines, and carbonyls. These polar interactions between solvent and solute are stronger than the intermolecular attractions between solute molecules caused by van der Waals forces and weaker hydrogen bonding. Thus, the solute molecules readily dissolve in water. ***Hydrophobic Interactions*** The behavior of water toward nonpolar solutes is different from the interactions just discussed. Nonpolar solutes (or nonpolar functional groups on biological macromolecules) do not readily H bond to H2O, and, as a result, such compounds tend to be only sparingly soluble in water. The process of dissolving such substances is accompanied by significant reorganization of the water surrounding the solute so that the response of the solvent water to such solutes can be equated to "structure making." Because nonpolar solutes must occupy space, the random H-bond network of water must reorganize to accommodate them. At the same time, the water molecules participate in as many H-bonded interactions with one another as the temperature permits. Consequently, the H-bonded water network rearranges toward formation of a local cagelike **(clathrate)** structure surrounding each solute molecule. Under these conditions, nonpolar solute molecules experience a net attraction for one another that is called **hydrophobic interaction.**  ***Amphiphilic Molecules*** Compounds containing both strongly polar and strongly nonpolar groups are called **amphiphilic molecules** (from the Greek *amphi* meaning "both," and *philos* meaning "loving"), also referred to as **amphipathic molecules** (from the Greek *pathos* meaning "passion"). Salts of fatty acids are a typical example that has biological relevance. They have a long nonpolar hydrocarbon tail and a strongly polar carboxyl head group, The ionic carboxylate function hydrates readily, whereas the long hydrophobic tail is intrinsically insoluble. Sodium palmitate and other amphiphilic molecules readily disperse in water because the hydrocarbon tails of these substances are joined together in hydrophobic interactions as their polar carboxylate functions are hydrated in typical hydrophilic fashion. Such clusters of amphipathic molecules are termed **micelles**. Of enormous biological significance is the contrasting solute behavior of the two ends of amphipathic molecules upon introduction into aqueous solutions. Image result for amphiphilic molecules **PHYSICAL AND CHEMICAL PROPERTIES** Indeed, normal metabolic activity can occur only when cells are at least 65% H2O. This dependency of life on water is not a simple matter, but it can be grasped through a consideration of the unusual chemical and physical properties of H2O. The Major Chemical and Physical Properties of Water Are: 1. Water is a liquid at standard temperature and pressure. 2. It's a good solvent, often called 'the universal solvent'. All the major components in cells (proteins, DNA and polysaccharides) are also dissolved in water. 3. The boiling point of water (and all other liquids) is dependent on the barometric pressure. 4. At 4181.3 J/(kg·K), water has a **high specific heat capacity**, as well as a **high heat of vaporization** (40.65 kJ·mol−1), both of which are a result of the extensive hydrogen bonding between its molecules. 5. The **maximum density of water occurs at 3.98 °C (39.16 °F)**. 6. Water is **miscible with many liquids**, such as ethanol, in all proportions, forming a single homogeneous liquid. On the other hand, water and most oils are immiscible, usually forming layers according to increasing density from the top. As a gas, water vapor is completely miscible with air. 7. Water can be split by electrolysis into hydrogen and oxygen. 8. As an **oxide of hydrogen, water is formed** when hydrogen or hydrogen-containing compounds burn or react with oxygen or oxygen-containing compounds. 9. Elements which are more electropositive than H e.g lithium, sodium, calcium, potassium and caesium displace hydrogen from water, forming **hydroxides**. Being a flammable gas, the hydrogen given off is dangerous and the reaction of water with the more electropositive of these elements may be violently explosive. 10. **Taste and odor**: Water can dissolve many different substances, giving it varying tastes and odors. However, pure H2O is tasteless and odorless. **ACIDITY, ALKALINITY, pH and BUFFER SYSTEMS** For the purpose of enzyme activity and to maintain the shape and structure of the proteins, the body must maintain the right state of acidity and alkalinity. If there are more H^+^ (hydrogen ions), a solution is acidic. If there are more OH^-^ (hydroxyl) ions, the solution is alkaline. The acidity and alkalinity of a solution is measured in terms of **pH** (hydrogen ion concentration). As the quantity of hydrogen ion is so small, it is cumbersome to express in actual numbers. If needed, the number would be something like 0.0000001. To make it easier, this is expressed by pH. The pH is actually a measure of hydrogen ion concentration in the body fluid; the pH scale extends from 0 to 14. Water is considered to be a pH of 7.0; a neutral pH. This means that water contains 0.0000001, or 1x 10^-7^ of a mole of hydrogen ions per liter. If the pH is lower than 7.0, it denotes that the fluid has more hydrogen ions or that it is acidic. For example, if a solution has a pH of 5.0, it contains 0.00001 or 1 x10^-5^ of a mole of hydrogen ions per liter (i.e., more hydrogen ions than a solution of pH 7.0). If a solution has a pH above 7.0, it has less hydrogen ions than water and is alkaline. **The pH of the body is 7.4 (range, 7.35--7.45)** (i.e., slightly alkaline). For body enzymes to be active and for chemical reactions to proceed optimally, it is vital that pH be maintained at this level. This implies that the body needs regulatory mechanisms that monitor the hydrogen ion levels carefully and get rid of them as and when they form above normal levels. One of the body's compensatory mechanisms is the presence of many **BUFFERS**. Buffers are compounds that **prevent the hydrogen ion concentration from fluctuating too much and too rapidly to alter the pH**. The body uses buffers to convert strong acids (that dissociate easily into hydrogen ions) to weak acids (that dissociate less easily). Examples of buffers that are present in the body fluid include: 1. Proteins, 2. Hemoglobin, and 3. A combination of **bicarbonate and carbonic acid compounds**. The later is an important buffer. The following chemical reaction indicates how a combination of bicarbonate and carbonic acid compounds work as buffers. **HCO~3~^-^  +  H → H~2~CO~3~ → H~2~O   +  CO~2~** In this chemical reaction, HCO~3~ (bicarbonate), a weak base, combines with the hydrogen ions to form H~2~CO~3~ (carbonic acid), a weak acid. This weak acid can be further broken down to CO~2~ (carbon dioxide), which can be breathed out, and H~2~O (water), which can be used for other reactions or excreted by the kidneys. Alternately, if the pH becomes acidic, the weak carbonic acid H~2~CO~3~ can break down to form HCO~3~^-^(a weak base) and H^+^ (hydrogen ions). **[Ionization of Water]** Water shows a small but finite tendency to form ions; **H~2~O -- H^+^ + OH^-^** Free protons (H+\] are immediately hydrated to form **hydronium ions,** H~3~O^+^.Indeed, because most hydrogen atoms in liquid water are hydrogen-bonded to a neighboring water molecule, this protonic hydration is an instantaneous process and the ion products of water are H~3~O^+^ and OH^-^: **H^+^ + H~2~O - H~3~O^+^** The amount of **H~3~O^+^** or **OH^-^** in 1 L (liter) of pure water at 25°C is **1 X 10^-7^ mol** ***[Kw, the Ion Product of Water]*** The dissociation of water into hydrogen ions and hydroxyl ions occurs to the extent that **10^-7^ mol** of **H^+^** and **10^-7^ mol** of **OH^-^** are present at equilibrium in 1 L of water at 25°C. **H~2~O -- H^+^ + OH^-^** The equilibrium constant for this process is  where brackets denote concentrations in moles per liter. Because the concentration of H2O in pure water is essentially constant, a new constant, ***K***w**,** the **ion product of water (or water equilibrium constant),** can be written as ***K*w = \[H^+^\]\[OH^-^\]** Experimentally, \[H^+^\] has been found to be same as \[OH^-^\] which is equal to 10^-7^ , Kw is therefore equal to 10^-7^ x 10^-7^ =10^-14^ The equation shows that there is a reciprocal relationship between H^+^ and OH^-^ concentrations of aqueous solutions. If a solution is acidic, that is, of significant \[H^+^\], then the ion product of water dictates that the OH^-^ concentration is correspondingly less. **For example, if \[H^+^\] is 10^-2^ *M,* \[OH^-^\] must be 10^-12^ *M*** ***i.e: K*w = 10^-14^ = \[10^-2^\]\[OH\_\];** **\[OH\_\] = 10^-12^ *M*).** Similarly, in an alkaline, or basic, solution in which \[OH^-^\] is great, \[H^+^\] is low. ***Questions*** *1.* What is the concentration of H^+^ in a solution of 0.1 M NaOH? *2.* What is the concentration of OH^-^ in a solution with an H^+^ concentration of 1.3x10^-4^ ? **pH** To avoid the cumbersome use of negative exponents to express concentrations that range over 14 orders of magnitude, Sørensen, a Danish biochemist, devised the **pH scale** by defining **pH** as *the negative logarithm of the hydrogen ion concentration*. **pH = -log~10~ \[H^+^\]** Note also that from water equilibrium constant, we can have: **p*Kw = pH + pOH = 14*** Note that the pH scale is logarithmic, not arithmetic. To say that two solutions differ in pH by 1 pH unit means that one solution has ten times the H^+^ concentration of the other, but it does not tell us the absolute magnitude of the difference. For example, a cola drink (pH 3.0) or red wine (pH 3.7) has an H^+^ concentration approximately 10,000 times that of blood (pH 7.4). Measurement of pH is one of the most important and frequently used procedures in biochemistry. The pH affects the structure and activity of biological macromolecules; for example, the catalytic activity of enzymes is strongly dependent on pH. Measurements of the pH of blood and urine are also commonly used in medical diagnoses. The pH of the blood plasma of people with severe, uncontrolled diabetes, for example, is often below the normal value of 7.4; this condition is called **acidosis**. In certain other disease states the pH of the blood is higher than normal, the condition of **alkalosis**. **[Dissociation of Electrolytes (Bases and Acids)]** Hydrochloric, sulfuric, and nitric acids, commonly called strong acids, are completely ionized in dilute aqueous solutions; the strong bases NaOH and KOH are also completely ionized. Of more interest to biochemists is the behavior of weak acids and bases---those not completely ionized when dissolved in water. These are common in biological systems and play important roles in metabolism and its regulation. The behavior of aqueous solutions of weak acids and bases is best understood if we first define some terms: 1. ACID AND BASES: Acids may be defined as proton donors and bases as proton acceptors. A proton donor and its corresponding proton acceptor make up a **conjugate acid-base pair**. Image result for equilibrium constant E.g, Acetic acid (CH3COOH), a proton donor, and the acetate anion (CH3COO^-^), the corresponding proton acceptor, constitute a conjugate acid-base pair, related by the reversible reaction CH~3~COOH = H^+^ + CH~3~COO^-^ Each acid has a characteristic tendency to lose its proton in an aqueous solution: the stronger the acid, the greater its tendency to lose its proton. 2. ACID DISSOCIATION CONSTANT: The tendency of any acid (HA) to lose a proton and form its conjugate base (A^-^) is defined by the equilibrium constant (*K*eq) for the reversible reaction HA ⇆ H^+^ + A^-^, Which is  Equilibrium constants for ionization reactions involving acids are usually called acid ionization or acid **dissociation constants,** often designated ***K*a**. Stronger acids, such as **hydrochloric acids**, **phosphoric and carbonic acids**, have **larger dissociation constants**; weaker acids, such as **acetic acid** (CH3COOH) and **monohydrogen phosphate (HPO~4~^2-^)**, have smaller dissociation constants. Reason: For strong acids, there are more \[H+\] in solution compared to \[HA\] (because they dissociate completely), hence the value for Ka from the equation becomes large. Reverse is true for weak acids 3. **pKa: To avoid the cumbersome use of negative exponents associated with Ka, we can take a negative logarithm of Ka to give pKa (like the case for pH).** **For weak acids, [Titration Curves] can be used to calculate p*K*a** 4. TITRATION CURVE: Titration is used to determine the amount of an acid in a given solution. A measured volume of the acid (V1) is titrated with a solution of a strong base, usually sodium hydroxide (NaOH), of known concentration (C1. The NaOH is added in small increments (V2) until the acid is consumed (neutralized), as determined with an indicator dye or a pH meter. The [concentration of the acid] in the original solution can be calculated from the volume and concentration of NaOH added (Recollect: C1V1 = C2V2). **A plot of pH against the amount of NaOH added (a titration curve) reveals the p*K*a of the weak acid**. 5. **HENDERSON HASSELBACH EQUATION:** Once the concentration of the acid in the solution is determined using C1V1=C2V2, and the pH is taken using a pH meter/indicator, as described above, the pKa of the acid can be calculated using HH equation. How is this equation derived? It is as shown below: Image result for henderson hasselbalch equation **METABOLIC ACIDOSIS AND ALKALOSIS** **Acidosis** refers to an excess of acid in the blood that causes the pH to fall below 7.35, and **alkalosis** refers to an excess of base in the blood that causes the pH to rise above 7.45. Many conditions and diseases can interfere with pH control in the body and cause a person\'s blood pH to fall outside of healthy limits. **Causes of Metabolic acidosis** Causes of primary metabolic acidosis are commonly classified by the [anion gap](https://www.acid-base.com/clinical.php#Gap): Anion Gap. Some causes of metabolic acidosis, e.g., lactic acidosis, release anions into the ECF which are not normally measured. When this occurs there will be an unexpected discrepancy between the sums of the principal cations and anions. The usual sum is: #### Gap   =   Na^+^   +   K^+^   -   Cl^-^   -   HCO~3~^-^ ####   15    =   140   +   5     -  105  -     25     mMol/L In addition to Cl^-^ + HCO~3~^-^ there are extra unmeasured anions, e.g., lactate, phosphate, sulphate, which increase the \"gap\". **A gap greater than 30** indicates a significant concentration of unmeasured anions. If information is required about the unmeasured anions, it is probably more appropriate to measure their concentration, i.e., lactate in tissue hypoxia, 3-hydroxybutyrate in diabetic ketosis, and phosphate or sulfate in renal failure. #### Metabolic Acidosis with a Normal Anion Gap: - Longstanding diarrhea (bicarbonate loss) - Uretero-sigmoidostomy - Pancreatic fistula - Renal Tubular Acidosis - Intoxication, e.g., ammonium chloride, acetazolamide, bile acid sequestrants - Renal failure #### Metabolic Acidosis with an Elevated Anion Gap: - lactic acidosis - ketoacidosis - chronic renal failure (accumulation of sulfates, phosphates, uric acid) - intoxication, e.g., salicylates, ethanol, methanol, formaldehyde, ethylene glycol, paraldehyde, INH, toluene, sulfates, metformin. - rhabdomyolysis Treatment: involves treatment of the underlying cause ### Metabolic Alkalosis **Causes:** - Loss of acid via the urine, stools, or vomiting - Transfer of hydrogen ions into the cells - Excessive bicarbonate administration, e.g. alkali given to patients with renal failure. - Contraction of the extracellular space due to excessive diuretic treatment **Prolonged Metabolic Alkalosis** may be caused by a number of different mechanisms: - **Decrease in renal perfusion:** occurs in dehydration, cardiac failure, or cirrhosis, stimulates the renin-angiotensis system which increases sodium reabsorption in the nephron. - **Chloride Depletion:** may occur via vomiting or through the use of loop diiuretics and this enhances bicarbonate reabsorption with associated hydrogen ion loss. - **Hypokalemia:** Metabolic alkalosis may be associated with hypokalemia which can then maintain metabolic alkalosis by various mechanisms such as - Shift of hydrogen ions intracellularly which enhances bicarbonate reabsorption in the collecting duct. - Stimulation of the H^+^/K^+^ ATPase in the collecting duct: this leads to potassium ion reabsorption and hydrogen ion secretion. The net gain of bicarbonate maintains the metabolic alkalosis. - Renal ammonia genesis: Ammonium ions (NH4+) are produced in the proximal tubule from glutamine metabolism. Alpha-ketoglutarate is produced The metabolism generates bicarbonate. - Impaired chloride ion reabsorption in the distal nephron increases luminal electronegativity with enhanced hydrogen ion secretion. - Lowered glomerular filtration rate (GFR). Hypokalemia may decrease GFR, which in turn decreases the filtered load of bicarbonate. In volume depletion this impairs excretion of the excess bicarbonate. ### Treating Severe Metabolic Alkalosis **A**dequate hydration normally allows the kidneys to correct the problem. However, in severe cases accompanied by hypokalemia, correction of the hypokalemia may be necessary first. **As with metabolic acidosis**, ideal treatment is the correction of the underlying abnormality. **Other therapies:** Intravenous dilute hydrochloric acid is occasionally used but carries the risk of hemolysis. Potassium chloride may also be used unless there is kidney failure. In severe cases which are unresponsive to other measures ammonium chloride may be given What is respiratory alkalosis? ------------------------------ Respiratory alkalosis occurs when the levels of carbon dioxide and oxygen in the blood are not balanced. Your body needs oxygen to function properly. When you inhale, you introduce oxygen into the lungs. When you exhale, you release carbon dioxide, which is a waste product. Normally, the respiratory system keeps these two gases in balance. Respiratory alkalosis occurs **when you breathe too fast or too deep and carbon dioxide levels drop too low. This causes the pH of the blood to rise and become too alkaline. When the blood becomes too acidic, respiratory acidosis occurs.** [Hyperventilation](https://www.healthline.com/symptom/hyperventilation) is typically the underlying cause of respiratory alkalosis. Hyperventilation is also known as overbreathing. Someone who is hyperventilating breathes very deeply or rapidly. ### Causes of hyperventilation Panic attacks and anxiety are the most common causes of hyperventilation. However, they're not the only possible causes. Others include: - [heart attack](https://www.healthline.com/health/heart-attack) - pain - drug use - [asthma](https://www.healthline.com/health/asthma) - fever - [chronic obstructive pulmonary disease](https://www.healthline.com/health/copd) - infection - [pulmonary embolism](https://www.healthline.com/health/pulmonary-embolus) - [pregnancy](https://www.healthline.com/health/pregnancy) Symptoms of respiratory alkalosis --------------------------------- Overbreathing is a sign that respiratory alkalosis is likely to develop. However, low carbon dioxide levels in the blood also have a number of physical effects, including: - Dizziness. bloating. feeling lightheaded. numbness or muscle spasms in the hands and feet. discomfort in the chest area.confusion - dry mouth - tingling in the arms - [heart palpitations](https://www.healthline.com/symptom/palpitations) - feeling short of breath What is respiratory acidosis? ----------------------------- Respiratory acidosis is a condition that occurs when the lungs can't remove enough of the carbon dioxide (CO~2~) produced by the body. Excess CO~2~ causes the pH of blood and other bodily fluids to decrease, making them too acidic. Normally, the body is able to balance the ions that control acidity. This balance is measured on a pH scale from 0 to 14. Acidosis occurs when the pH of the blood falls below 7.35 (normal blood pH is between 7.35 and 7.45). Respiratory acidosis is typically caused by an underlying disease or condition. This is also called respiratory failure or ventilatory failure. Normally, the lungs take in oxygen and exhale CO2. Oxygen passes from the lungs into the blood. CO2 passes from the blood into the lungs. However, sometimes the lungs can't remove enough CO2. This may be due to a decrease in respiratory rate or decrease in air movement due to an underlying condition such as: - [asthma](https://www.healthline.com/health/asthma) - [COPD](https://www.healthline.com/health/copd) - [pneumonia](https://www.healthline.com/health/pneumonia) - [sleep apnea](https://www.healthline.com/health/sleep/obstructive-sleep-apnea) Types ----- Forms of respiratory acidosis ----------------------------- There are two forms of respiratory acidosis: acute and chronic. **Acute respiratory acidosis** occurs quickly. It's a medical emergency. Left untreated, symptoms will get progressively worse. It can become life-threatening. **Chronic respiratory acidosis** develops over time. It doesn't cause symptoms. Instead, the body adapts to the increased acidity. For example, the kidneys produce more bicarbonate to help maintain balance. Chronic respiratory acidosis may not cause symptoms. Developing another illness may cause chronic respiratory acidosis to worsen and become acute respiratory acidosis. Symptoms -------- Symptoms of respiratory acidosis -------------------------------- Initial signs of acute respiratory acidosis include: - headache - anxiety - blurred vision - restlessness - confusion Without treatment, other symptoms may occur. These include: - sleepiness or fatigue - lethargy - delirium or confusion - shortness of breath - coma The chronic form of respiratory acidosis doesn't typically cause any noticeable symptoms. Signs are subtle and nonspecific and may include: - memory loss - sleep disturbances - personality changes