Chapter 2 / The Chemical Level of Organization (8e+)
I. INTRODUCTION
[460=title]
A. Since chemicals compose your body and all body activities are chemical in nature, it is important to become familiar with the language and fundamental ideas of chemistry.B. To understand the nature of the matter that composes your body and the changes it undergoes in health and in disease, you will need to know how it is organized and how different elements of matter interact with one another.
II. MATTER AND ENERGY
A. Basic Principles 1. Matter is anything that occupies space and has mass. a. Matter exists as a solid, liquid, or gas. b. Mass is the amount of matter a substance contains; weight is the force of gravity acting on a mass.
2. Energy is the capacity to do work (put mass into motion).
a. Potential energy is inactive or stored energy. 1) Examples include the energy stored in a battery, a compressed spring, water behind a dam, or an object that can drop some distance. b. Kinetic energy is the energy of matter in motion.
1) The matter can be of any size or mass, as long as it is moving relative to some other matter, like an atom or a planet. c. Potential and kinetic energy exist in several forms, including radiant, electrical, heat, chemical, and mechanical.
[470-475=clinical,laser,medical]
1) Radiant energy, such as visable light, has the properties of moving waves. a) The more waves that move past a point in a second (higher frequency) is associated with higher energy. (1) This means the waves must be closer together and so they have a shorter wave length, since all light moves at the same speed. b) Several forms of radiant energy may be illustrated using examples from long wave length (low frequency, low energy) to short wave length (high frequency, high energy).
(1) radio waves (2) microwaves
(3) infrared waves (heat)
(4) visable light waves
(5) ultrviolet waves (sunburn)
(6) x-rays
(7) gamma rays
2) Electrical energy is due to the movement of charged particles, like electrons in a computer or ions (atoms with a positive or negative charge) in the body.
a) An impulse moving along a nerve is an example of eletrical energy in the body. 3) Heat energy is in the form of the vibrational motion of atoms that can be transferred by either contact or by infrared light.
a) Heat will be transferred from warmer to cooler areas. 4) Chemical energy is potential energy that is held by the electrons in bonds that hold atoms together.
a) This energy can be released when a bond is broken, as when the body breaks down food. b) Chemical energy may be transferred from one bond and stored in another chemical bond, as occurs when energy from food is moved to special chemicals in the body to hold it for later use.
(1) This also happens when chemical structures of the body are made by attaching atoms together with chemical bonds. 5) Mechanical energy is the energy due to either the position or the movement of a mass.
3. Mass and energy can neither be created nor destroyed, but one can be converted into the other.
B. Chemical Elements
1. All forms of matter are composed of chemical elements. a. Elements (atoms) are the small stable units of matter, and cannot be broken down by chemical reactions. b. 92 of the 109 known elements are found naturally, the others have been produced briefly by humans.
2. Elements are given letter abbreviations called chemical symbols
3. Oxygen (O), carbon (C), hydrogen (H), and nitrogen (N) make up 96% of body mass.
a. These elements together with calcium (Ca) and phosphorus (P) make up 98.5% of total body weight. b. There is also a significant amount potessium, sulfur, sodium, chlorine, magnesium, iodine, and iron.
4. The roughly 13 trace elements are present in very small amounts (0.1%) but are also essential to the human body.
(Exhibit2.1 )
C. Structure of Atoms
1. Units of matter of all chemical elements are called atoms. a. An element is simply a quantity of matter composed of atoms of the same type. b. Atoms are effectively the smallest units of matter, with the smallest atom measuring only 10-8 cm in diameter, and the lagest being only five times bigger.
2. Atoms consist of a nucleus, which contains protons and neutrons (nucleons), and electrons that move about the nucleus in energy levels.
(Fig.2.1[461])
a. Protons are defined as having one atomic mass unit, and a single unit positive electrical charge. b. Neutrons also have one atomic mass unit, but no electrical charge.
c. Electrons have a mass about 1/2000 of a mass unit (proton), and a single unit negative electrical charge.
3. The number of protons in an atom's nucleus is called the atom's atomic number and so each element has a different atomic number.
a. The number of protons in the nucleus makes the atoms of one element chemically different from those of another element. (Fig.2.2[462-469])
1) For instance hygrogen has one proton, whereas the element helium has 2 , carbon 6, and sodium 11, which is the same as their atomic number. b. Since all atoms are electrically neutral (same number of positive and negative electrical charges), the atomic number also equals the number of electrons in each atom.
4. The mass number of an atom is the sum of the total number of protons and neutrons.
5. Atoms of an element that have the same number of protons but different numbers of neutrons are called isotopes, and so their chemistry is the same, but their only difference is mass number (weight).
a. An element is defined by the number of protons becuase their positive charges (in the nucleus) interact with the negatively charged electrons within the atom, and with electrons of other atoms to form bonds, in accordance with the chemical nature of each element. b. Certain isotopes called radioisotopes are unstable and may emit radiation that can be monitored with radiation detectors.
1) The amount of radiation declines to half the original amonut emitted as half the isotopes are converted to a different form, at a rate determined to be a half-life. 2) Examples of normal / radioactive forms of familiar elements are; 16O/19O, and 12C/14C, where the superscripts indicate the atomic mass number of each element.
c. Radioisotopes can be used to study both the structure and function of particular tissues as described in the Clinical Application on "Medical Imaging Using Radioactive Isotopes."
6. The atomic mass (atomic weight) of an atom is the sum of the masses of its protons, neutrons, and electrons.
a. The atomic weight is the average sum of protons, and neutrons (average mass number) for all the isotopes of an element, given in grams. 1) For example; most carbon atoms have six protons, and six neutrons and so a mass number of 12, and so are called carbon-12 (12C). a) However, a small percentage of carbon atoms have six protons and seven or eight neutrons, and so are called carbon-13 (13C) and carbon-14 (14C) isotopes. b) If all carbons were 12C (mass number =12) then the atomic weight would be 12.000 grams.
c) But when the small number of heavier isotopes are averaged-in the atomic weight becomes 12.010 grams.
d) The atomic weight of each element contains exactly the same number of atoms as the atomic weight of all other elements.
(1) It is such a large number (6.023X1023 atoms) that it has its own name, mole. D. Electrons and Chemical Reactions
1. When atoms combine or break apart from other atoms, a chemical reaction occurs producing other chemicals with different properties than the original combination of atoms. a. Chemical reactions are the foundation of all life processes, and electrons interacting with each other and the protons in the nuclei of nearby atoms are the basis of all chemical reactions. 2. Electrons revolve around the nucleus of an atom, tending to spend most of the time in speciflc atomic regions, called shells.
(Fig.2.2[462-469])
a. Each shell can hold a certain maximum number of electrons. 1) The inner most shell can hold 2 electrons, the second 8, and the third 18. 2) The outermost shell is called the valence shell, and its electrons are the only ones involved in bond formation.
b. To achieve stability, atoms tend to either empty their valence shell or fill it to the maximum extent.
(PERIODIC TABLE OF THE ELEMENTS)
1) This may be achieved by giving up, accepting or sharing electrons with other atoms in the simplest manner possible. a) Chlorine atoms (Cl) have 7 electrons in their outer shell and can obtain a stable configuration of 8 electrons in that shell by receiving one from another atom. (1) Even though the third shell can take 18 electrons, one stable state is reached with the first 8 electrons. (2) This extra electron gives the chloride atom 18 total negatively charged electrons with only 17 positively charged protons in the nucleus, making a chloride ion with a single negative electrical charge (Cl-).
b) Sodium atoms have only one electron in the outer, third, shell and so may go to a stable state with the maximum of 8 in the second shell by giving up the single electron that is in the third shell.
(1) When sodium (Na) gives up an electron there is a total of only 10 negatively charged electrons and 11 positively charged protons, and so it becomes a sodium ion with a single positive charge (Na+). 2) Atoms that naturally have completely filled outer shells are called inert elements and do not often participate in chemical reactions.
3) Atoms with incompletely filled outer shells tend to combine with each other in chemical reactions that provides each with the number of electrons they need to obtain a stable state (8 electrons in the outer shell).
a) Sodium and chlorine atoms would make the proper combination to form a stable state for each, with sodium donating an electron to chlorine. 3. When two or more atoms combine in a chemical reaction, a molecule is formed, by either exchanging or sharing electrons to establish a stable state for all atoms involved.
a. The term molecule formally defined as two atoms of the same type forming a bond with each other (e.g. H2 or O2). 4. A compound is a substance that can be broken down into two or more different elements by chemical means (e.g. HCl or NaCl).
a. Compounds are usually referred to as molecules also. E. Chemical Bonds
1. The atoms of a molecule are held together by forces of attraction called chemical bonds. a. A bond is formed when electrons form an optimum distribution in the space between the atoms that balances the attraction of the electrons for the protons of the atoms at each end of the bond, against the repulsion between all the electrons within the bond. b. In chemical reactions that occur in the body, breaking bonds usually releases energy, and forming bonds usually requires energy.
2. Ionic Bonds
a. In an ionic bond, valance-shell electrons are transferred from one atom to another, creating electrically charged ions, whose unlike charges attract each other and this holds them together as an ionic bonds.
(Fig2.3[476-481])
1) Cations are positively charged ions that have given up one or more electrons (i.e., they are electron donors). a) Examples are; Na+, K+, and Ca++. 2) Anions are negatively charged ions that have picked up one or more electrons from another atom (i.e., they are electron acceptors).
a) Examples are; Cl-, I-, H2PO43-.
3) Compounds/molecules that dissociate (break up) into ions in solution are called electrolytes because they are charged and so can conduct an electrical current.
4) Ionic bonds form when there is a big difference in attraction for electrons (electronegativity) between the two atoms forming a bond.
a) Electronegativity is related to how easy it is for an atom to either let electrons go, or take on more to complete the outer shell. b) Atoms with 4 or less electons in the valence shell (left side of periodic table, like Na, K, and Ca) tend to have low electronegativity, and so tend to let electrons go in order to have only the completed shell below.
c) Atoms with 5 or more electons in the valence shell (right side of periodic table, like Cl, and O) tend to have high electronegativity, and so tend to take on more electrons in order to have a completed shell.
b. In a covalent bond, there is sharing of pairs of valance-shell electrons.
(Fig.2.4[482-486])
1) Single, double, or triple covalent bonds are formed by sharing one, two, or three pairs of electrons, respectively.
a) One electron of each shared pair comes from each atom involved in forming the bond, or bonds. b) Hydrogen atoms need one more electron in their valance shell to be stable, and so two hydrogen atoms may share their electrons with each other to effectively have two in each of their outer shells (H2).
(1) The subscript 2 indicates that there are two hydrogen atoms bonded together in the hydrogen molecule. c) Oxygen atoms in an oxygen molecule (O2) share two electrons from each atom (4 in all) and so form two covalent bonds, a double bond, with two electrons in each bond.
d) Nitrogen atoms in an nitrogen molecule (N2) share three electrons from each atom (6 in all) and so form three covalent bonds, a thriple bond, with two electrons in each bond.
e) Different types of atoms can form covalent bonds with one or more other atoms as long as they all obtain the proper number of electrons they need to fill their outer shells.
(1) Carbon atoms require 4 additional electrons to have 8 in its outer shell, and so may obtain them by sharing electrons with 4 hydrogen atoms, which need to share one electron each (=CH4 = mehtane). 2) Covalent bonds may be nonpolar or polar.
a) If there is some small difference in attraction for the electrons (negative) between the nuclei (positive protons) then the electrons spend more time around the nucelus with the greatest attraction. (1) So that end of the bond tends to have more electrons than protons, thus has a slightly negative electrical charge. (2) The nucleus with the least attraction for the electrons tends to have fewer electrons around it than the number of protons, and so tends to have a slightly positive electrical charge.
(3) Since the two ends of the bond have slightly opposite charges (like the poles of a magnet or battery) the bond is called polar.
(4) The bonds between hydrogen and oxygen in a water molecule (H2O) is an example of polar bonds.
b) If the two atoms at either end of a bond have similar attraction (electronegativity) for the electrons then the electrons (negative charges) spend their time more equally around both nuclei, and so negative and positive charges are about equal for each nucleus and the bond is non-polar.
(1) Bonds between identical atoms (H-H, N-N) or where there is a geometrical balance between bonds on the opposite sides of an atom (H-C-H) are non-polar bonds. c. In a hydrogen bond, two other atoms (usually oxygen or nitrogen) associate with a hydrogen atom.
(Fig.2.6b)
(Fig.2.16[560])
1) This type of bond is not very strong, but when there is a large number of them in a big molecule then they become important for things like setting the shape of a protein. F. Chemical Reactions
1. A number of factors determine the rate (number of reactions per second) of a chemical reaction. a. Each reaction involves the breaking and formation of bonds. 1) When atoms have sufficient kinetic energy (activation energy) upon collision to overcome the repulsion between electrons in the old bonds, then the nuclei get close enough to the electrons to form new, stronger attractions and hence a new electron distribution that make the new bonds. b. A higher concentration of molecules (greater number per volume) increases the number of collisions with the proper energy and so increases the number of chemical reactions that occur.
d. The speed at which atoms and molecules are traveling, and so their kinetic energy, is influenced by temperature and size of the particles.
1) A higher temperature increases the average kinetic energy of all particles and so the number of collisions that are higher than the activation energy, producing completed chemical reactions. 2) Atoms and molecules with a smaller mass have a higher speed, at any particular temperture, and greater kinetic energy, and so have a greater number of collisions that are higher than the activation energy, producing completed chemical reactions.
e. The proper orientation of the colliding particles is necessary so the atoms that are to react come into direct contact with each other at the angle in which the new bond will exist.
f. Metabolism is defined as all the chemical reactions that occur in the body.
1) The total number of atoms is the same before and after every chemical reaction. 2. Energy balance is the amount of energy released or absorbed by a chemical reaction in which one bond is being broken and another is formed and depends on the difference in the amount of energy in the two bonds.
a. An exergonic reaction is one in which the bond being broken has more energy than the one formed so that there is extra energy released, usually as heat. 1) This occurs during catabolism of food molecules. b. An endergonic reaction is just the opposite and thus requires energy, usually from a molecule called ATP (adenosine triphosphate) to form a bond, as in covalently binding amino acids molecules together to form proteins.
3. Synthesis reactions involve the combination of reactants to produce a new molecule, called a product, which are anabolic reactions (anabolism), meaning that bonds are formed with an input of energy (endergonic).
[487=diagram,reaction,chemical,synthetic]
a. (reactants) A + B --> AB (product) b. (reactants) C + O2 --> CO2 (product)
4. In decomposition reactions, a substance breaks down into other substances through catabolic reactions (catabolism), which means that chemical bonds are broken and energy is usually release (exergonic) in the process.
[488=diagram,reaction,chemical,decomposition]
a. AB --> A + B b. H2O --> H+ + OH-
5. Exchange reactions involve the replacement of one atom or atoms by another atom or atoms.
[489=diagram,reaction,chemical,exchange]
a. AB + CD --> AC + BD 6. In reversible reactions, end products can revert to the original combining molecules.
[490=diagram,reaction,chemical,reversible]
a. AB <--> A + B b. H2O <--> H+ + OH-
7. In oxidation-reduction reactions electrons are taken from the atom being oxidized by the atom being reduced.
a. In cellular oxidations two electrons are usually removed from a molecule along with two hydrogens [H+ + H-= 2(H+ + e-)], along with energy held by the electrons. b. CH3-(HC-OH)-COOH -(oxidation)-> CH3-(C=O)-COOH + 2(H++e-)
lactic acid pyruvic acid
1) The 2(H++e-) usually becomes bound to a special molecule (NAD), which is consequently reduced (gains electrons and energy). a) NAD + 2(H++e-) -(reduction)-> NADH2
III. CHEMICAL COMPOUNDS AND LIFE PROCESSES
A. Inorganic substances usually lack carbon and are small molecules. B. Organic substances always contain carbon and hydrogen and most organic substances contain covalent bonds.
C. Organic materials make the bulk of living things, however inorganic substances are also essential to life.
IV. INORGANIC COMPOUNDS
A. Inorganic Acids, Bases, and Salts 1. When molecules of inorganic acids, bases, or salts dissolve in water, they undergo ionization or dissociation, that is, they separate into ions. 2. Acids ionize into one or more hydrogen ions (H+) and one or more anions (negative ions).
(Fig.2.5a[492])
a. An acid molecule, or just acid, may be defined as a compound that can release some or all of its hydrogen in the form of an ion (H+= acid). b. HCl --> H+ + Cl-
hydrochloric acid --> hydrogen ion (acid) + chloride ion
c. Hydrogen ion (H+) is the acid component of an acid molecule, and the amount of H+ in a volume of solution (its concentration) is the definition of the acidity of that solution.
3. Bases dissociate into one or more hydroxide ions (OH-) and one or more cations (positive ions).
(Fig.2.5b[492])
a. A base molecule, or just base, may be defined as a compound that produces a component that can bond a H+, with the hydroxide ion being only one example. b. NaOH --> Na+ + OH-
OH- + H+ --> H2O
c. By bonding to H+ a base decreases the concentration of H+ in a solution and so decreases the acidity, or to say it another way it makes the solution more basic, or alkaline.
4. A salt, when dissolved in water, dissociates into cations and anions, neither of which is H+ or OH-.
(Fig.2.5c[492])
a. Many salts are present in the body and are formed when acids and bases react with one another. b. Salts are often made of substances essential to the operation of cells in the body.
B. Water and Solutions
1. Water is the most abundant substance in the body. a. A solution consists of a solute, like NaCl (table salt), dissloved in a solvent, like H2O (water). 2. The concentration of a molecule is a way of stating the amount of that molecule dissolved in a solution.
(Exhibit2.2)
a. A standard unit that expresses the number of molecules of solute in a liter of solution is called Molarity = moles/liter. b. A mole is just the name for the number of atoms there is in an atomic weight (=atomic mass number in grams) of that element, or the number of molecules there are in a molecular weight of that type molecule.
1) The molecular weight is the addition of all the atomic weights (in grams) of the atoms that make the molecule. 2) The molecular weight of NaCl = 23 g/mole for Na + 35 g/mole for Cl = 58 g/mole of NaCl molecule.
c. As an example, if there is 58 g NaCl in a liter of solution then the concentration of NaCl is 1.0 Molar (=1.0 mole NaCl in one liter of solution).
1) If there 1/1000 of a mole of NaCl (=0.058g) in a liter of solution then the concentration is 0.001 Molar = 1.0 millimolar (the usual units of concetrations seen in body fluids). 3. Water has many properties that make it a vital compound in living systems.
a. It participates in chemical reactions. 1) This occurs with digestion of proteins in the diet where H+ and OH- (=H2O) are added to the atoms on either end of a bond, causing the bond to break, called hydrolysis (=water cleavage).
(Fig.2.13[540])
1) The components of water (H+ and OH-) can be taken from adjacent atoms to help the formation of a bond, referred to as dehydration synthesis.
(Fig.2.8a[505])
b. Water can absorb or release a lot of heat energy, with little change in the temperature of the water, which is the definition of a high heat capacity.
1) Since the body contains a large proportion of water, this helps keep the body temperature stable. c. Water requires a large amount of heat to change from a liquid to a gas therefore its heat of vaporization is considered to be high.
1) When water (sweat) evaporates from the surface of the body it takes the heat from the body to change from a liquid to a vapor, and so cools the body. d. It serves as a lubricant, between moving organs and joints, and as part of mucus help food transit the digestive tract.
e. It is an excellent solvent and suspending medium.
(Fig.2.6c[491])
1) Water molecules contain polar bonds that are attracted to other molecules with polar bonds, which allows them to interact and so dissolve in the water, as a solute. a) NaCl dissolves in water because there is a greater attraction between the polar water and the Na+ and Cl- ions than there is between the ions for for each other. 2) Water also suspends, floats, materials that do not dissolve in it, like red blood cells, and allows them to move around in the water and body as the water in blood plasma is transported.
3) It is absolutely neccessary for chemicals to move around in order to come in contact with each other and react.
a) So water acts as an indispensable reaction medium. C. Acid-Base Balance: The Concept of pH
1. Body fluids must constantly contain balanced quantities of acids and bases. 2. Biochemical reactions are very sensitive to even small changes in acidity or alkalinity.
3. A solution's acidity or alkalinity is expressed on the pH scale, which runs from 0 (=100=1.0 moles H+/l) to 14 (=10-14 =0.00000000000001 moles H+/l).
(Fig.2.7[493])
a. pH 7.0 = 10-7 moles H+/l = 0.0000001 moles H+/l 1) H2O <---> H+ + OH- => [H+] = [OH-] 2) This is considered a neutral solution since there is equal amounts of acid (H+) and base (OH-).
b. Values below 7 indicate acid solutions ([H+] > [OH-]).
c. Values above 7 indicate alkaline solutions ([H+] < [OH-]).
[494-95=measurement,pH,indicator,color]
d. The definition of pH is just a mathematical description of the concentration of H+ (=[H+]), which can also be transformed to give the [H+] from the pH number.
1) pH = -log[H+] 2) [H+] = 10-pH
3) So roughly speaking pH is the number of places to the right of the decimal point after where the first number for the [H+] occurs.
4) These definitions also mean that, starting from pH 7.0, as the [H+] increases, the pH number decreases, or as [H+] decreases, the pH number increases.
D. Maintaining pH: Buffer Systems
1. The pH values of different parts of the body are maintained fairly constant by buffer systems, which usually consist of a weak acid and a weak base. a. An example of a weak acid is carbonic acid that forms from a reaction of CO2 with water and then release only a set proportion of its H+ and its weak base bicarbonate (HCO3-). H2O + CO2 <---> H2CO3 <---> H+ + HCO3-
2. Buffer systems bind excess H+ and excess OH- ions in order to maintain pH homeostasis, as described in Chapter 27.
a. If H+ is added to this solution then bicarbonate will bind it so that it does not increase the [H+]. H+ + HCO3- ---> H2CO3
b. If a base, such as OH-, is added to the solution then carbonic acid will release a H+ that will bind with the base and so the [H+] does not change.
OH- + H2CO3 ---> OH- + H+ + HCO3- ---> H2O + HCO3-
3. Show the pH values for certain body fluids compared with common substances.
(Exhibit2.3[495-504])
V. ORGANIC COMPOUNDS
A. The carbon that organic compounds always contain has several properties which make it particularly useful to living organisms. 1. It can react with up to four elements and combine with several hundred other carbon atoms to form large molecules of many different shapes. 2. Many carbon compounds do not dissolve easily in water, making them useful materials for building body structures.
3. Carbon compounds are mostly or entirely held together by covalent bonds and can be broken easily by special molecules; this means that organic compounds are a good source of energy.
4. Small organic molecules can combine to form very large molecules (macromolecules, or polymers, when composed of repeating subunits called monomers).
a. Dehydration synthesis occurs when two monomers join together, eliminating a molecule of water in the process. b. Hydrolysis (digestion) breaks macromolecules down into monomers by adding a molecule of water.
B. Carbohydrates
1. Carbohydrates, including sugars, starches, glycogen, and cellulose, provide most of the energy needed for life. a. Some carbohydrates are converted to other substances, which are used to build structures and to generate an energy holding molecule ATP. 1) Two such structures include branched chains of sugars attached to glycoproteins, and a sugar deoxyribose that is part of the molecules that make DNA (=deoxyribonucleic acid). b. Other carbohydrates function as food reserves called glycogen, which are stored in the liver and skeletal muscles.
c. Carbohydrates are given this name because the contain only C, H, and O in a proportion of 1C:2H:1O (=C H2O), which effectively is a hydrated (watered) carbon.
d. The elements C and H form polar colvalent bonds with O in carbohydrates and so are very soluble in water, unless they are part of extremely large molecules, like glycogen.
e. Carbohydrates constitute only about 2-3% of body weight.
2. Carbohydrates are divided into three major groups based on size.
a. Monosaccharides contain from three to seven carbon atoms and include glucose, a hexose that is the main energy-supplying sugar of the body. b. Disaccharides are formed from two monosaccharides by dehydration synthesis and they can be split back into simpler sugars by hydrolysis.
(Fig.2.8[505-06])
1) Glucose and fructose combine to produce sucrose. c. Polysaccharides are the largest carbohydrates and may contain hundreds of monosaccharides.
[507-515=carbohydrates,sources]
1)The princlpal polysaccharide in the human body is glycogen.
C. Lipids
1. Lipids, like carbohydrates, contain carbon, hydrogen, and oxygen, but have few polar covalent bonds and thus are mostly insoluble in polar solvents such as water (i.e., they are hydrophobic).
(Fig.2.9[523])
a. Hydrophobic/non-polar, or lipid/oil-soluble molecules are attracted to each other, and so separate from water, forming structures that isolate aqueous (water) solutions into distinct compartments where different types of chemical reactions occur, called cells and organelles. 2. Lipids are a diverse group of compounds that include triglycerides (neutral fats), phospholipids, steroids, carotenes, vitamins A, D, E, and K, and eicosanoids.
(Exhibit2.4)
a. For efficient transport in blood, lipids combine with proteins to form water-soluble lipoproteins. b. Triglycerides are the most plentiful lipids in the body and siet, and provide protection, insulation, and energy (both immediately and stored as fat).
1) Triglycerides are composed of glycerol and fatty acids.
(Fig.2.9[523])
a) These can be made from excess sugars and proteins in the diet. b) The name triglyceride comes from the fact that a three-carbon glycerol molecule has three fatty acid molecules attached, by dehydration synthesis.
c) Fatty acid structures consist of a carboxylic acid molecule (-COOH), bonded to a long chain of hydrogenated carbons (n-CH2, where n = 15 to 17 units).
2) The type of covalent bonds (and by inference, number of hydrogen atoms) found in the fatty acids determines whether a triglyceride is saturated, monounsaturated, or polyunsaturated.
a) A monounsaturated fatty acid is formed when two carbons in a saturated hydrocarbon chain (all Cs have two Hs attached) form a double bond with each other and so there are two less hydrogens on the molecule, such that it is no longer "saturated" with hydrogens. b) There can be more than one double bond in the hydrocarbon chain, making polyunsaturated fatty acids, which tend to be more fluid at body temperatures.
3) A Clinical Application on "Saturated Fats and Atherosclerosis" describes the formation of fatty plaques in the walls of arteries and cites the role of dietary saturated fats in increasing the risk of plaque development.
a) Sources of lipids are from any fat containing food.
[516-22=fats,sources]
(1) Animal products contain saturated fatty acids, whereas almost all plant products contain unsaturated fatty acids. d. Phospholipids are important membrane components.
(Fig.2.10[523-25])
1) These are just triglycerides with one of the fatty acids replaced on the glycerol molecule by a phosphate containing molecule. 2) The glycerol portion with the phosphate group attached is very polar ("polar head") and so will interact and dissolve with water.
3) The hydrocarbon chains portion ("tails") are very non-polar and so are attracted to other similar fatty acids (lipids) and will not mix (dissolve) with water.
4) Because of these chemical properties phospholipids will form a double layer (bilayer), since the nonpolar tails will line up side-by-side and end-to-end, with the polar heads facing out toward the water.
a) This ability of phospholipids to partially dissolve (interact) with polar water and nonpolar lipids is called amphipathic. b) The "world-shaking" result of these chemical properties is that the phospholipid bilayer will wrap around a volume of water to complete the lipid-to-lipid contact and thus automatically form an enclosed compartment.
(1) This is the basic membrane structure of all cells and organelles from which all life is made. e. Steroids include sex hormones, cortisol, vitamin D, and cholesterol.
(Fig.2.11[526-29])
[520-533=steroids,abuse,effects]
1) Cholesterol serves as an important component of cell membranes and as starting material for synthesis of other steroids. 2) The basic structure of steroids has four rings of carbon with various chemical groups as appendages.
f. Eicosanoids include prostaglandins and leukotrienes, and are derived from a 20-carbon fatty acid called arachidonic acid.
1) Prostaglandins; a) modify responses to hormones, b) contribute to the inflammatory response,
c) prevent stomach ulcers,
d) dilate airways to the lungs,
e) regulate body temperature, and
f) influence formation of blood clots, among other things.
2. Leukotrienes participate in allergic and inflammatory responses.
D. Proteins
1. Proteins are constructed from combinations of amino acids. 2. Proteins provide a wide range of essential properties of an living organism.
(Exhibit2.5)
[547-50=protein,structures,body]
a. Many connective tissue structures of the body are make of proteins, like collagen in bones and tendons. b. Certain forms of proteins act to regulate processes, like blood glucose levels with the hormone insulin.
c. Some white blood cells provide protection (immunity) by producing protein antibodies.
d. Specialized proteins allow muscles to contract.
e. Certain types of proteins transport substances like oxygen, lipids, or hormones.
f. A complex group of proteins serve as enzymes, which perform and control most of the chemical reactions that occur in the body.
3. Amino Acids and Polypeptides
a. Chemically, proteins always contain carbon, hydrogen, oxygen, and nitrogen as the components of amino acids.
(Fig.2.12[534-39])
[542-46=protein,sources]
b. All 20 different amino acids have a three-atom atomic structure base that is the same, and corresponds to the name. 1) One end of the three-atom core is a base, a nitrogen with two hydrogens, called an amino group. 2) The other end of the three-atom core is an acid, a carbon with a double bond to an oxygen, and a single bond to an OH (hydroxyl) group, all together called a carboxylic acid group (O=C-OH).
3) The carbon in the middle has a bond to a hydrogen and one to a R-group molecule that is different for each of the 20 different amino acids.
c. Amino acids are joined together in a stepwise fashion with each covalent peptide bond joining the amino nitrogen of one amino acid to the carboxylic carbon of the next amino acid, through dehydration synthesis.
(Fig.2.13[540-41]) 1) Two amino acids coupled together is called a dipeptide, three a tripeptide, 4-10 a peptide, and longer chains polypeptides.
d. Resulting polypeptide chains, generally referred to as proteins, may contain 10 to more than 2000 amino acids.
1) Examples of polypeptides are the molecule hemoglobin that transports oxygen in the blood, and enzymes. 4. Levels of structural organization include primary, secondary, tertiary, and quatemary structures.
(Fig.2.14[551-55]
a. The primary level is the genetically determined order of specific amino acids in the chain and is the principle factor defining all other properties of the protein. 1) A change of only two amino acids in the hemoglobin protein, that contains 350, is responsible for the dysfunction called sicle cell anemia. b. The secondary level is produced by hydrogen bond formation between different parts of the protein's core chain.
1) When hydrogen bonds develop between the amino nitrogen and carboxylic oxygen of every fifth amino acid of the core chain this forms a spiral called an alpha helix. 2) When two protein chains are parallel they form hydrogen bonds that make pleted sheets.
c. The tertiary structure is the overall three-dimensional shape of a protein that is formed by attractions between the r-groups in different parts of the protein.
1) The interactions include hydrogen bonds, ionic bonds, covalent bonds between sulfur atoms, polar, and non-polar (hydrophobic) attractions. 2) The shape of a protein is very important to its function.
d. The quaternary structure is formed by the attraction of separate proteins for each other, that holds them in a set orientation.
e. The resulting shape of the protein greatly influences its ability to recognize and bind to other molecules.
[3544PLAY=protein,hemoglobin,3-D]
1) The shape is strongly altered by the temperature, pH, and electrolyte concentration of the solution around a protein.
2) The function of enzymes, receptors, and antibodies are particularly sensitive.
f. Denaturation of a protein by a hostile environment causes loss of its characteristic shape and function.
5. Enzymes, which are proteins, speed up chemical reactions by increasing frequency of collisions, lowering the activation energy, and properly orienting the colliding molecules.
a. They do these things without themselves beinq altered in the reaction and are therefore classified as catalysts. 1) Enzymes contain a protein component called an apoenzyme, and often a cofactor. 2) The cofactor may be a metal ion, like iron or calcium, or a coenzyme, derivatives of vitamins.
3) Together the apoenzyme and coenzyme make a holoenzyme.
b. Although enzymes catalyze selected reactions, they do so with great efficiency and with many built-in controls.
1) Enzymes are highly specifc in terms of the substrates with which they react. (Fig.2.15[556-59])
a) Substrates are molecules that fit the unique three-dimensional shape and chemistry of the active site of an enzyme.
b) The active site may be static or have an induced fit when the site changes its shape to accomodate the shape of the substrates.
c) Enzymes bind only specific molecules, in such a way that the electrons of certain atoms are redistributed to either form or breaks bonds.
2) Enzymes are extremely efficient in terms of the number of substrate molecules converted from reactant to product per second, called turnover number.
a) This is generally between 1 to 10,000 but sometimes as many as 600,000 per second). 3) Enzymes are subject to a great variety of cellular controls.
a) The number of enzymes present is genetically controlled, and is a balance between the number produced and degraded per minute. b) The number of active enzymes is often regulated by special signal molecules and enzymes in a cell.
c) The availability of cofactors can be controlled which impacts the level of enzyme activity.
4) Enzymes are thought to operate through several common steps.
(Fig.2.15[556-559])
a) An enzyme-substrate complex is formed when a substrate molecule(s), reactant(s), binds to the active site of an enzyme. b) Products are made as bonds within or between molecules are formed or broken.
c) Product molecules are released from the active site, which then can bind new reactant molecules to perform the process again.
[3617PLAY=enzyme,superoxide-dismutase, operation]
[5567PLAY=enzyme,adenylate-cyclase, operation]
c. The names of enzymes usually end in the suffix -ase; e.g., oxidase, kinase, and lipase.
d. The lack of an enzyme that converts the sugar galactose into glucose is found in the inherited disorder, galactosemia, which can lead to loss of appatite, vomiting, and diarrhea resulting in small size, and retardation.
1) The immediate treatment is elimination of lactose and galactose from the diet. E. Nucleic Acids: Deoxyribonucleic Acid (DNA) and Ribonucleic Acid (RNA)
1. Nucleic acids are huge organic molecules which contain carbon, hydrogen, oxygen, nitrogen, and phosphorus. a. Deoxyribonucleic acid (DNA) forms the genetic code inside each cell for the production of proteins and thereby regulates most of the activities that take place in our cells throughout a lifetime. b. Ribonucleic acid (RNA) relays instructions from the genes in the cell's nucleus to guide each cell's assembly of amino acids into proteins.
[563=RNA,in-oocyte]
2. The basic units of nucleic acids are nucleotides, composed of a nitrogenous base, a pentose sugar, and a phosphate group.
(Fig.2.16[560])
[593PLAY=DNA,packing]
a. There are four types of nucleotides, classified into two chemical groups due to the only portion of the molecule that is different, the nitrogenous base. 1) The purine group has two nitrogen/carbon rings in the nitrogenous base molecule. a) Adenine (A) and Guanine (G) differ only by the position and type of atoms attached to the outer part of their second ring. 2) The pyrimidine group has only one nitrogen/carbon ring in the nitrogenous base molecule.
a) Thymine (T) and Cytocine (C) differ only by the type of atoms attached to the same position on the outer part of their single ring. 3) The purine nucleotides are said to "complement" the pyrimedine nucleotides.
a) T and A form two hydrogen bonds between the outer atoms of their rings. b) C and G form three hydrogen bonds between the outer atoms of their rings.
c) These complementary hydrogen bonds turn out to be very important in the basic DNA structure.
b. A deoxyribose pentose sugar holds together the base and phosphate molecules of each DNA nucleotide.
1) Deoxyribose is the "D" part of DNA. c. The phosphate group (PO43-) attached to the pentose sugar of each nucleotide also forms a bond to the pentose sugar of another nucleotide to make the backbone of the entire DNA molecule.
d. The structure of the complete DNA molecule is a double helix, that is two DNA strands that spiral around each other.
1) The pentose-phosphate molecules constitute the continuous strand of the backbone of each spiral. 2) The base molecules extend into the center of the double spiral where they form hydrogen bonds with the complementary nucleotide (A to T and C to G) on the opposing strand.
3) When a cell divides to make another cell it needs to make a second copy of all DNA instructions to go with the new cell.
a) This is accomplished when the double helix separates and enzymes put together a new strand of complementary nucleotides to match each of the original strands b) This makes two complete DNA molecules, each with an old and new strand of nucleotides.
3. DNA fingerprinting is used in research and in legal situations to determine the genetic identity of an individual.
a. This technique relies on the fact that certain DNA segments contain base sequences that are repeated several time, and the number of repeat copies in one region and the number of regions subject to repeat are different for each individual.
[5867PLAY=DNA,cystic-fibrosis,detection]
4. A second variety of nucleic acid is Ribose Nucleic Acid (RNA) that is used to transfer the information on DNA from the nucleus to the cytoplasm where it is used in the process to make proteins.
a. The difference between DNA and RNA is that RNA contains a ribose sugar in all nucleotide molecules and the nitrogenous base Uricil replaces Thymine. b. There are three types of RNA.
1) Messenger RNA (mRNA) transfers the code for the sequence of amino acids that will be in a protein, from DNA to the cytoplasm. 2) Ribosomal RNA (rRNA) acts as the site of protein synthesis.
3) Transfer RNA (tRNA) bring specific amino acids to the rRNA site to be bonded together into proteins.
F. Adenosine Triphosphate (ATP)
1. Adenosine triphosphate (ATP) is the principal energy-storing molecule in the body. 2. Among the cellular activities for which ATP provides energy are muscular contractions, chromosome movement during cell division, cytoplasmic movement within cells, membrane transport processes, and synthesis reactions.
3. ATP consists of three phosphate groups attached to an adenosine unit composed of adenine and the five-carbon sugar ribose.
(Fig.2.17[561])
a. When energy is liberated from ATP, it is decomposed to adenosine diphosphate (ADP) and inorganic phosphate (Pi). b. ATP is manufactured from ADP and P using the energy supplied by various decomposition reactions, particularly that of glucose.
1) In anaerobic metabolism glucose can be partially used to make only four ATP. 2) Whereas in aerobic metabolism glucose is used completely to make 36 ATP.
c. Another form, called cyclic-AMP is used as an intracellular signal.
[562=RNA,cyclic-AMP]