BIOLOGY Volume 3 - A Guide to General Biology - 2004

APPENDICES

Appendix 1. Biological Chemistry

P.1.1. Fundamentals of Chemistry

An atom is the smallest particle of any element capable of taking part in Chemical Reactions. An element is defined as a substance that cannot be broken down by chemical Methods into simpler substances; Examples of elements include carbon, oxygen, and nitrogen. A substance consisting of two or more chemically bonded elements is called a compound, for example:

A molecule is the smallest particle of a given element or compound capable of independent existence under normal conditions. Examples of molecules include the following: Н₂, О₂, СО₂, and Н₂О.

P.1.1.1. Atomic Structure

All elements are composed of atoms. The word "atom" comes from the Greek *atomos*, meaning "indivisible."

An atom consists of protons, neutrons, and electrons (Table P.1.1). The masses of a proton and a neutron are equal and together make up the mass of the atomic Nucleus. The mass of electrons is many times smaller than that of protons and neutrons; therefore, the mass of an atom is considered to be essentially the mass of its nucleus alone.

A neutron consists of one proton and one electron, making it electrically neutral, i.e., its charge is zero.

Class="center">Table P.1.1. Position within the atom, mass, and charge of protons, neutrons, and electrons

An atom is electrically neutral because the number of protons in its nucleus equals the number of electrons orbiting the nucleus.

The number of protons in an atomic nucleus is called the atomic number of the given element. It is also equal to the number of electrons in that atom. The sum of the number of protons and neutrons in a given atom is called its mass number.

Atoms of the same element can exist in different forms called isotopes, which differ from one another in their mass number (Section P.1.3). *Atomic mass* refers to the average mass of a given atom; this usually means the average value for a naturally occurring mixture of its isotopes. Chlorine, for example, is a mixture of isotopes with mass numbers 35 and 37; The ratio of these isotopes is such that naturally occurring chlorine has an atomic mass of 35.5.

Over 100 elements are currently known. They can be arranged in order of increasing atomic numbers, as shown in Table P.1.2. As can be seen from this table, electrons are arranged around the nucleus, filling a series of electron shells. The first electron shell (closest to the nucleus and therefore the smallest) can hold up to two electrons, the second up to eight, the third up to 18, and the fourth up to 32.

Table P.1.2. The first 20 elements (in order of increasing atomic number)

Larger atoms have more than four electron shells, but we will not discuss them here. Fig. P.1.1 illustrates the electron shell configuration of the first 12 elements.

Fig. P.1.1. Electron shell configuration of the first 12 elements. (The nucleus is shown only for the hydrogen atom.)

Any element with a completely filled outer electron shell is virtually unreactive. For this reason, helium and neon (Table P.1.2) extremely rarely form compounds with other atoms. Consequently, they are referred to as inert gases or noble gases.

All other elements tend to fill their electron shells by reacting with other elements. When two atoms react with each other, either an ionic or a covalent bond may form between them.

P.1.1.2. Ionic Bond

When ionic bonds are formed, electrons are transferred from one atom to another. Let us consider the reaction between sodium and chlorine (Fig. P.1.2, A). A sodium atom loses an electron, acquiring a positive charge of +1 (its nucleus, which contains 11 positively charged protons, is now surrounded by 10 negatively charged electrons). A chlorine atom gains an electron, acquiring a negative charge of -1. In both atoms, the outer shells become filled and, consequently, stable.

Fig. P.1.2. A. Formation of sodium chloride. B. Formation of calcium chloride. C. Formation of a chlorine molecule. D. Formula of methane. E. Formula of ethene. For clarity, electrons belonging to different atoms are marked with different symbols (crosses, black or white circles). In reality, all electrons are identical. Figures C, D, and E show only the outer electron shells.

Such charged particles are no longer true atoms; they are called ions. The sodium ion is denoted as Na⁺, and the chlorine ion (chloride ion) as Cl^-. Positively charged ions are called cations, and negatively charged ions are called anions. The compound formed As a result of the reaction between sodium and chlorine is called sodium chloride (its formula is NaCl), but discrete molecules of NaCl do not exist. Instead of molecules, the crystal lattice of sodium chloride consists of sodium ions and chloride ions in equal proportions (the ionic formula is Na⁺Cl^-). Such compounds, formed as a result of electron transfer, are called ionic compounds. They are formed primarily when metals react with non-metals. The metal forms a cation, and the non-metal forms an anion. All salts are ionic compounds.

Another typical example of an ionic compound is calcium chloride, CaCl₂ (Fig. P.1.2, B). In this case, the calcium atom loses two electrons, and each chlorine atom gains one. The calcium ion is therefore designated as Ca⁺⁺ or Ca²⁺.

The number of transferred electrons (gained or lost) characterizes the valence of an atom, i.e., its ability to form chemical bonds. From the foregoing, it can be seen that sodium and chlorine are monovalent, while calcium is divalent. Thus, the number preceding the plus or minus sign in an ion symbol indicates its valence; for example, the valence of the potassium ion and the hydroxyl ion (K⁺ and OH^-) is 1, the valence of the magnesium ion and the sulfate ion (Mg²⁺ and SO₄^2-) is 2, and the valence of the aluminum ion (Al³⁺) is 3.

Ionic Formulas

Ionic compounds exist not in molecular form, but in ionic form. An ionic formula shows the ratio in which elements enter into a given compound; for example, the ionic formula of aluminum oxide, Al₂O₃, means that Al³⁺ and O^2- ions are present in this compound in a 2:3 ratio. If an ion consists of several atoms, such as the sulfate ion (SO₄^2-), and there are multiple such ions in the compound, parentheses are used in the ionic formula, e.g., Al₂(SO₄)₃; parentheses are not needed in the ionic formula Na₂SO₄ because only one sulfate ion is present in this compound.

P.1.1.3. Covalent Bonds

When bonds of this type are formed, the participating atoms neither give away nor accept electrons; instead, the electrons are shared between two atoms. Let us imagine two chlorine atoms. Each of them has seven electrons in its outer electron shell (electron configuration 2.8.7). In covalent bonding, each of the two chlorine atoms contributes one electron to the shared electron pair, forming a chlorine molecule — Cl₂. In this way, both atoms acquire an electron configuration approaching that of a noble gas, and molecules are formed instead of ions (Fig. P.1.2, C). A shared pair of electrons is conventionally denoted as a single bond: C—C. Chlorine is monovalent (i.e., it shares one of its electrons). Another example is methane, CH₄. Carbon has an atomic number and its outer electron shell contains 4 electrons (2.4); hydrogen (atomic number 1) has only one electron in its outer electron shell (Fig. P.1.2, D).

In the ethene (Ethylene) molecule, C₂H₄, two pairs of electrons are shared between the carbon atoms, and these two pairs are represented by a double bond (Fig. P.1.2, E). In some compounds, notably ethyne (acetylene), C₂H₂, There is a triple bond, indicating three pairs of shared electrons.

Covalent compounds are much more common in biological systems than ionic ones.

The valences of some widespread elements and the charges of frequently encountered ions are listed in Table P.1.3.

Table P.1.3. Valences of some elements and charges of some ions

A. Valences of some elements

B. Charges of some monoatomic ions

C. Charges of some polyatomic ions

Formulas of Covalent Compounds

The formulas of simple covalent compounds indicate the number of atoms of each type contained in a given molecule; for example, the formula CO₂ (carbon dioxide) shows that in this molecule one carbon atom is bonded to two oxygen atoms.

P.1.1.4. Chemical Equations

In a chemical equation, not only must the formulas of all substances participating in the reaction be written correctly, but the equation must also be balanced, i.e., the number of atoms of each element must be the same on both the left-hand and right-hand sides. Let us explain how this is done.

1. Write the word equation, for example:

Methane + Oxygen → Carbon dioxide + Water

2. Replace the words with the formulas of these compounds:

СН₄ + О₂→ СО₂ + Н₂О.

3. Check whether the equation is balanced. The equation in step 2 is not balanced, because the right side contains three oxygen atoms while the left side contains only two; the hydrogen atoms are also unbalanced, with four on the left and two on the right.

4. Balance the equation by placing the required numbers in front of the corresponding formulas (recall that the formulas themselves cannot be changed):

СН₄ + 2O₂ → СO₂ + 2Н₂O

2O₂ means two oxygen molecules (4 oxygen atoms); 2Н₂O means two water molecules (4 hydrogen atoms, 2 oxygen atoms)

Ionic Equations

Reactions involving ionic compounds can be written simply as ionic equations. Consider the following reaction:

(All three compounds are in aqueous solution.)

The equation can be written in such a way as to show all the ions present in the system:

By cancelling out all the ions shown on both the left and right sides (those not participating in the reaction), we obtain:

Only this reaction actually took place in this case.

P.1.1.5. Acids, Bases, Salts, pH, and Buffers

A hydrogen atom consists of a single electron and a single proton. If the electron is lost, a single proton remains, so a proton can be regarded as a hydrogen ion, commonly denoted as Н⁺. An acid is defined as a substance capable of acting as a proton donor; upon dissociation (dissociation being the Separation into the ions of which a substance is composed), an acid yields Н⁺ as the cation. In this book, we will use the definition according to which an acid is a substance that, upon dissociation in water, yields Н⁺ ions as the cation. A strong acid (such as Hydrochloric acid, НСl) is an acid that dissociates almost completely. A strong acid is therefore a more effective proton donor than a weak acid, such as acetic or carbonic acid, in which only a small fraction of the molecules dissociate to form Н⁺ ions:

The Characteristic Properties of acids are listed below.

1. Many acids react with reactive metals such as zinc or magnesium, accompanied by the evolution of hydrogen gas.

2. Acids are neutralized by bases, resulting in The formation of salts of these acids and water.

3. Almost all acids react with carbonates, accompanied by the release of carbon dioxide.

4. In dilute solutions, acids have a sour taste, such as ethanoic (acetic) acid.

5. Acid solutions produce a characteristic color change with indicators; for example, blue litmus paper turns red under the action of an acid.

A base is a substance that reacts with an acid to form a salt and water (alternatively, a base can be defined as a substance capable of acting as a proton acceptor). Most bases are insoluble in water. Soluble bases yield solutions known as alkalis. Examples include sodium hydroxide, calcium hydroxide, and ammonium hydroxide. Other characteristic properties of bases are listed below.

1. Bases generally show little reactivity with metals.

2. Bases react with aqueous solutions of most metal salts, precipitating an insoluble hydroxide.

3. The reaction of bases with ammonium salts releases ammonia.

4. Base solutions produce a characteristic color change with indicators; for instance, red litmus paper turns blue in the presence of bases.

A salt is a compound formed by the complete or partial replacement of the hydrogen atoms of an acid by a metal. A classic example is sodium chloride, in which the hydrogen atom of hydrochloric acid is replaced by a sodium atom. When a salt dissolves in water, its constituent ions dissociate—that is, they separate into free ions surrounded by water molecules.

The pH scale

The acidity or alkalinity of a solution is determined by its hydrogen ion concentration. Hydrogen ion concentration is expressed as the pH of the solution (where 'p' denotes a mathematical operation and 'H' is the chemical symbol for hydrogen); pH is the negative logarithm of the hydrogen ion concentration. One cubic decimeter (1 L) of pure water contains 1 · 10^-7 moles of hydrogen ions.

Consequently, for water, the pH value equals log(1/10⁷) = 7.

This value (pH 7.0) characterizes a neutral solution (at room Temperature). Values below 7.0 indicate an acidic solution, while those above 7.0 indicate an alkaline one.

The pH scale spans roughly from -1 to 15 (mostly 0–14). Because it is a logarithmic scale, a one-unit change in pH corresponds to a tenfold change in hydrogen ion concentration.

Cells and Tissues require a pH of around 7.0, and deviations of more than one or two units from this value are fatal to them. Therefore, specific mechanisms exist to maintain body fluid pH within a relatively constant range. This is achieved, in part, by buffers.

Buffers

A buffer solution is a mixture of a weak acid and its soluble salt designed to resist changes in pH. Such changes can occur as a result of dilution or The addition of an acid or base.

When acidity (hydrogen ion concentration) increases, the free anions provided by the salt readily bind with free hydrogen ions, removing them from the solution. When acidity decreases, the tendency to release hydrogen ions increases. Thus, a buffer solution maintains a constant, balanced concentration of hydrogen ions. This can be illustrated by the following example:

Certain Organic compounds, particularly Proteins, can act as buffers; this property is especially vital for Blood.

P.1.2. Oxidation and reduction

All biological processes require energy, which is why biologists study various reactions that can serve as an energy source. Reactions that release energy are called exothermic or exergonic, whereas reactions that consume energy are called endothermic or endergonic. Synthetic processes (anabolic processes) are endergonic (such as Photosynthesis), while breakdown processes (catabolic processes) are exergonic (such as Respiration). The sum of all Catabolic and anabolic reactions taking place within a Cell at any given moment constitutes its METABOLISM.

The Cell obtains most of its required energy through The oxidation of nutrients during respiration. Oxidation is defined as the loss of electrons. The reverse process—the gain of electrons—is called reduction. These two processes always occur simultaneously: electrons are transferred from an electron donor, which is thereby oxidized, to an electron acceptor, which is thereby reduced. Reactions of this type are termed oxidation-reduction or redox reactions and are widespread in biological systems. There are several different mechanisms of oxidation and reduction, which we will examine in the following sections.

P.1.2.1. Oxidation

The oxidation of a given substance can occur directly through the addition of oxygen to that substance, in which case the reaction is simply called oxidation:

However, the most common form of Biological Oxidation involves the removal of hydrogen (dehydrogenation):

In this case, A is oxidized and B is reduced.

The cell contains A number of substances called hydrogen carriers, which act like B in the example above. Every dehydrogenation is catalyzed by a specific dehydrogenase. Hydrogen carriers are arranged in a specific sequence such that their potential energy level (Sec. P.1.6.2) decreases from one end of the chain (where hydrogen enters the process) to the other. This means that whenever hydrogen atoms are transferred from one carrier to another with a lower potential energy, the transition is accompanied by the release of a small amount of energy. In certain cases, this energy can be stored in the form of ATP.

In some reactions, each hydrogen atom (which can be viewed as a hydrogen ion, or proton, H⁺, plus a negatively charged electron, e^-) is not transferred as a single whole. Such a process involves only electron transfer. For example:

Iron(II) ions are oxidized to iron(III) as a result of losing electrons, one per each ion:

Electrons are transferred to a chlorine molecule, which is thereby reduced, forming two chloride ions. Thus, the overall ionic equation is

Cytochromes, which contain iron, function within Mitochondria by transferring electrons (derived from hydrogen atoms split into hydrogen ions and electrons) along the Electron Transport Chain. Here, electrons are passed from less electronegative atoms to more electronegative ones. The products of such reactions have a lower potential energy than the reactants, and The amount of energy corresponding to this difference is released and utilized in one form or another. At the end of The electron transport chain lies a cytochrome containing both Iron and copper, which directly transfers electrons to atmospheric oxygen, thereby becoming oxidized:

P.1.2.2. Reduction

Reduction occurs when molecular oxide is split off from a substance, when hydrogen atoms are added to a substance, or when a substance gains electrons.

P.1.3. Isotopes

The atoms of certain elements exist in multiple forms. These different forms of an atom are called its isotopes (from the Greek isos meaning equal, and topos meaning place, i.e., occupying the same place in the periodic table of elements). All isotopes of a given element have the same number of protons and electrons (the same atomic number) and, consequently, identical chemical properties; they differ, however, in the number of neutrons in their nuclei, meaning they have different atomic masses. To distinguish between isotopes, the mass number is added to the element symbol; for example, three isotopes of oxygen occur in nature: ¹⁶O, ¹⁷O, and ¹⁸O. One of the isotopes is usually present in a larger quantity than the others; for oxygen in particular, the ratio of its isotopes ¹⁶O : ¹⁷O : ¹⁸O is 99.759% : 0.037% : 0.204%.

Certain combinations of protons and neutrons are capable of existing indefinitely without change. Such atomic nuclei are referred to as stable. In other cases, nuclei are unstable, meaning they tend to disintegrate with the emission of particles or radiation. These nuclei are called radioactive. They are easily detected using various instruments such as Geiger-Müller counters, scintillation counters, and the like. As the atomic number of a nucleus increases, so does the proportion of neutrons in the nucleus required for it to remain stable. A uranium nucleus, for example, contains 92 protons and requires 138 neutrons for stability. Uranium isotopes with a larger number of neutrons are radioactive and have unstable nuclei.

The rate of decay is often expressed in terms of the half-life, i.e., the time during which, on average, half of all the atoms present undergo decay. For example, the half-life of the carbon isotope ¹⁴C is 5,570 years.

Radioactive isotopes can emit three Different types of "rays" (particles and radiation).

1. α-Particles. These are nothing other than helium atom nuclei, i.e., they consist of two protons and two neutrons. These particles carry a double positive charge.

An example of α-particle emission is given below (see also Fig. P.4.1). (The superscript next to the symbol of each element denotes its mass number, and the subscript denotes its atomic number.)

By emitting an α-particle, a ²³⁸U nucleus loses four units of mass and two units of charge, thereby transforming into a thorium isotope.

2. β-Particles. These are fast-moving electrons emitted by the nucleus when one of its neutrons converts into a proton. β-Particles carry a single negative charge (see also Fig. P.4.1).

The decay of thorium, for example, is accompanied by the emission of β-particles:

A thorium nucleus emits a single electron; thus, one of its neutrons is converted into a proton. The atomic mass remains unchanged, but the atomic number (number of protons) increases by one, resulting in the formation of a protactinium isotope.

3. Gamma rays. This is electromagnetic radiation with a very short wavelength, produced during α- and β-decay. Gamma rays possess very high energy and are extremely difficult to block; for instance, they are capable of penetrating a thick layer of lead.

α-Particles are easily stopped even by a layer of air or a thin sheet of paper. β-Particles have a higher penetrating power, but they can be blocked by a thick layer of aluminum or a thin layer of lead. These particles and radiation have harmful effects on living organisms.

P.1.4. Solutions and the colloidal state

Solutions consist of at least two parts, or phases: a continuous (dispersion) phase, or solvent, and a dispersed phase distributed within it, or solute.

In 1861, Graham identified Two Types of solutes, which he called crystalloids and colloids. Graham distinguished between them based on the ability of solute molecules to pass through a parchment (semipermeable) membrane. In biological systems, however, there is actually no clear-cut distinction between the two, because water always acts as the biological solvent, and The properties of any aqueous solution depend on the size of the solute molecules and the effects of gravity. There are Three types of solutions.

1) True solutions. In true solutions, the solute particles are small and comparable in size to the solvent molecules; that is, the system is homogeneous and its particles do not separate under The Influence of gravity. Examples include salt solutions and a sucrose solution. Chemists view such solutions as systems consisting of a single phase.

2) Colloidal solutions. The solute particles are large compared to the solvent molecules; that is, the system is heterogeneous, but its particles still do not separate under the influence of gravity. An example is clay in water.

3) Suspensions or emulsions. The solute particles are so large that they can remain in a dispersed state without settling out under gravity only with continuous agitation. If solid particles are suspended in the solvent, the system is called a suspension, and if liquid droplets are suspended, it is an emulsion. Silt is an example of a suspension.

All three of these systems can be considered disperse systems, since the particles are distributed within a dispersion medium. Disperse systems encompass all three states of matter: solid, liquid, and gas—for example, a gas in water (soda water), a solid in water (table salt solution), and a solid in a solid (zinc in copper, i.e., brass). In all these cases, we can speak of solutions, but solutions are usually defined as systems in which a liquid plays The Role of the solvent.

Many biological systems exist as colloidal solutions, which can be hydrophobic or hydrophilic; a hydrophobic sol (such as clay or charcoal in water) repels water, whereas a hydrophilic sol (starch paste, jelly, gelatin, and Agar) attracts it. Most colloidal solutions found in living organisms, protein solutions in particular, are hydrophobic sols. The viscosity of a hydrophobic sol, such as jelly, can be increased by raising its concentration or lowering the temperature. Ultimately, with an increase in viscosity, the sol may set. Such a set sol is called a gel. A gel is a more or less dense colloidal system, although, strictly speaking, there is no sharp boundary between a sol and a gel. Sol–gel transitions are also influenced by factors such as ionic composition, pH, and pressure. Under certain circumstances, all of these can play a vital role in living cells.

Table P.1.4 lists various CHARACTERISTICS OF THE colloidal state.

Table P.1.4. Characteristics of the colloidal state

P.1.5. Diffusion and osmosis

Molecules and ions in solution can move passively and spontaneously in a specific direction as a result of diffusion. (Osmosis is a special type of diffusion, described in detail in Section 5.9.8.) Unlike Active Transport, such movement in living organisms requires no Energy Expenditure. Another type of movement, namely bulk flow, is discussed in Chapter 13.

P.1.5.1. Diffusion

Diffusion is associated with the random and spontaneous movement of individual molecules and ions. If, for instance, a bottle of concentrated aqueous ammonia is left open, the smell of ammonia will very quickly spread throughout the room. This process of ammonia molecule dispersal represents diffusion, and although any individual molecule can move in any direction, the net flow of molecules is outward from the bottle—that is, from the source where their concentration is high to regions where their concentration is lower. Diffusion can therefore be defined as the movement of molecules or ions from a region of high concentration to a region of lower concentration, or in other words, movement down a concentration gradient. Unlike what is typical for bulk flow, net diffusion of different types of molecules or ions can occur simultaneously in opposite directions, with each type of molecule moving along its own concentration gradient. In the Lungs, for example, oxygen diffuses into the blood while carbon dioxide simultaneously diffuses from the blood into the alveoli, whereas bulk flow of blood in the lungs can only have a single direction. Given equal concentration gradients, small molecules and ions diffuse faster than large ones. There is a special type of diffusion known as Facilitated Diffusion, which is described in Section 5.9.8.

P.1.6. Laws of Thermodynamics

All chemical transformations obey the laws of thermodynamics. The First Law, known as the law of conservation of energy, states that for any chemical process, the total energy of the System and Its surroundings always remains constant. This means that energy cannot be created or destroyed; thus, if a chemical system gains energy, an equal amount of energy must be drawn from its surroundings, and vice versa. Energy can therefore be redistributed, converted from one form to another, or undergo both of these changes, but it cannot vanish.

It follows from The Second Law of thermodynamics that a system and its surroundings, left to themselves, generally approach a state of maximum disorder (Entropy). This means that highly ordered systems are easily degraded unless energy is expended to maintain their order. All biological processes obey and are governed by these two laws of thermodynamics.

P.1.6.1. Energy relationships in living systems

Let us consider the decomposition of hydrogen peroxide into oxygen and water:

In general, pure hydrogen peroxide can persist for a long time without undergoing noticeable decomposition. For decomposition to occur, colliding molecules must possess energy exceeding a certain threshold, known as the activation energy, E_a. Once this activation barrier is reached, The Nature of the bonds within the molecules changes, and the reaction generates sufficient energy to proceed spontaneously. The magnitude of the activation energy varies for different reactants.

Heating is the simplest way to reach the activation energy; most reactants require far greater amounts of thermal energy than they possess at ordinary temperatures. For instance, the decomposition of hydrogen peroxide at 150 °C proceeds so rapidly that the reaction becomes explosive. Water and oxygen are formed in this reaction, and energy is released. The overall energy change resulting from the reaction is called the Free energy change (ΔG). Since this reaction proceeds very rapidly, and its products—water and oxygen—do not recombine under normal conditions (i.e., the reverse reaction does not occur), the released energy is effectively lost, passing from the chemical system into the environment. The value of ΔG is therefore negative (Fig. P.1.3).

Fig. P.1.3. Activation energy.

High temperatures are lethal to biological systems, and here their effect is replaced by the action of Enzymes. Acting as catalysts, enzymes lower the activation energy required by the reactants, thereby ensuring higher reaction rates without adding external energy—specifically, without raising the temperature. In living systems, the rapid decomposition of hydrogen peroxide occurs through the action of the enzyme catalase.

P.1.6.2. Potential Energy

Potential energy is the energy possessed by a system by virtue of its position and existing conditions. Consider a sphere resting motionless at the edge of an inclined plane (Fig. P.1.4). This sphere possesses gravitational potential energy equivalent to the work that had to be done to place it in that position. If the sphere rolls down, part of its potential energy is converted into kinetic energy. When it subsequently comes to a rest at the bottom, its potential energy will be lower than it was at the top. To restore the sphere's potential energy to its initial value, it must be lifted back up by utilizing energy drawn from the surroundings.

Fig. P.1.4. Potential and kinetic energy.

Potential energy in biological systems is stored by green plants during photosynthesis when they synthesize sugars (Fig. P.1.5). During this process, certain electrons are raised to a higher energy level by solar energy, thereby acquiring potential energy. Later, when sugars are oxidized during respiration, the potential energy of these electrons is utilized in various forms by living systems.

Fig. P.1.5. Energy Flow in biological systems.

P.1.7. Chromatography

Chromatography is a method used to separate various mixtures into their constituent components. The method is based on the principle that in a stationary medium through which a solvent flows, each component carried along by the solvent moves at its own characteristic speed, independently of the others. If, for example, a mixture of plant pigments responsible for green coloration is dissolved in an appropriate solvent and passed through a stationary medium, such as powdered chalk, the mixture will separate into several distinctly colored pigments. This type of separation is described in Experiment P.1.3.

Depending on the Nature of the stationary medium used, three MAIN TYPES OF chromatography are distinguished: paper, Column, and thin-layer. Paper chromatography is described in Experiments P.1.1–P.1.3 and in Section 7.6.3. Various chromatographic techniques are now widely used in chemistry, biology, biochemistry, and specialized fields such as forensics.

Electrophoresis

Electrophoresis is a modification of chromatography used to separate charged molecules. In a chromatographic medium under the influence of an applied electric field, one side becomes positively charged and the other negatively charged. Individual molecules in the mixture being separated move toward one side or the other depending on their relative masses and charges. Electrophoresis is widely used for the isolation and identification of Amino Acids; further refinement of the technique in this case is achieved by adjusting the pH of the medium.

P.1.7.1. R_f Values

In chromatography, the mobility of a dissolved substance relative to the solvent front is constant for a given substance. This can be expressed by the R_f value as shown below:

If the solvent front moves beyond the edge of the paper, the mobility of the given solute can be expressed in comparison with the mobility of another standard substance. Then

See Fig. P.1.6, B.

P.1.7.2. Two-Dimensional Paper Chromatography

Chromatography proceeding in only one direction is not always sufficient to effectively separate a complex mixture of substances. In this case, for better spot resolution, additional separation must be carried out in a perpendicular direction using a second solvent (Fig. P.1.6, D).

Fig. P.1.6. Chromatographic methods.

A square sheet of paper is used for this Procedure. The sample is applied near one of the corners along the starting line, and development is carried out in the first direction. The paper is then removed from the chamber, dried, rotated by 90°, and subjected to a second chromatographic separation using a solvent with different properties. As a result, substances that were only partially resolved During the first run are fully separated In the second solvent. The paper is removed, dried, and treated with an appropriate reagent to visualize the separated compounds. Identification can be achieved by comparing THE POSITION OF the substance with that of known reference standards. This technique was famously employed by Calvin in his experiments to identify the primary products of photosynthesis (Section 7.6.3).

Before attempting to chromatograph unknown substances for identification, it is recommended to practice using colored inks or indicators. By carrying out the Procedures described below, you will find that the more concentrated the starting spot, the better the resolution. You will also observe that separation improves with a longer solvent run.

Practical P.1.1. Separation of indicators

Materials and equipment

Whatman No. 1 or No. 3 chromatographic paper

Screened methyl orange

Bottle of ammonia 0.880

Petri dish

Pipette

Procedure

Apply a drop of screened methyl orange to the center of the chromatographic paper. Dry the paper by waving it in the air, hold it briefly over an open bottle of 0.880 ammonia, and then place it over a Petri dish (Fig. P.1.6, A). Place a single drop of water onto the indicator spot.

Observed effect

The two indicators present in the screened methyl orange solution migrate outward from the center at different rates, with the blue ring moving faster than the yellow ring. The blue ring consists of bromothymol blue, while the yellow ring is methyl orange.

Practical P.1.2. Separation of colored inks into their component parts

Materials and equipment

Test tube

Bung (stopper)

Chromatographic paper

Drawing pin (push-pin)

Felt-tip pens

Distilled water

Procedure

1. Take a test tube and bung (see Fig. P.1.6, B) and attach a rectangular strip of chromatographic paper to the bung using the pin. Draw a pencil line parallel to the free end of the paper, approximately 1 cm from the edge.

2. Mark the starting line with crosses at regular intervals, allocating one cross for each ink color to be tested.

3. Using water-soluble felt-tip pens of various colors, apply a spot of a different ink to each cross and label its identity underneath in pencil. The spot should not exceed 2 mm in diameter. Allow the spots to dry completely.

4. Suspend the paper in the test tube so that the starting line is just above the solvent surface and the bottom end of the paper is immersed in the solvent. Distilled water is used as the solvent.

5. Continue chromatography until the solvent front is 1 cm from the top edge of the paper. Remove the chromatogram and dry it, having previously marked the position of the solvent front with a pencil.

6. If larger chromatography chambers are available, they can be used for both ascending and descending chromatography.

Experiment P.1.3. Separation of plant pigments by paper chromatography

Materials and Equipment

Nettle or spinach leaves

Homogenizer or knife, 90% propanone (acetone)

Mortar and pestle

Short piece of capillary tubing

Buchner funnel

Separatory funnel

Petroleum ether (boiling point 37.8 – 48.9 °C)

Chromatography paper

Laboratory test tube

Stopper

Drawing pin

Procedure

Chop several leaves of stinging nettle (Urtica dioica) or spinach using a homogenizer (or simply cut them into small pieces with a knife). Grind the leaves with 90% acetone in a mortar. Filter the extract through a Buchner funnel and a separatory funnel. Add an equal volume of petroleum ether. Shake the mixture vigorously. Wash the mixture with water several times, discarding the aqueous phase and its contents each time. The chromatography solvent consists of 100 parts petroleum ether and 12 parts 90% acetone. Suspend the chromatography paper in the test tube as described in the procedure above, then, in the same manner as before, draw a pencil line 1 cm from the bottom edge of the paper. Using a short piece of capillary tubing, apply the pigment mixture to the middle of the pencil line. Pour solvent into the test tube to a depth of 2 cm, then secure the stopper and the paper in the test tube. Allow the solvent to ascend the paper almost to the stopper. This should take about 1–2 hours. Keep the test tube in a dimly lit place during this time.

Results

The following colored spots should be visible: