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    Biology Super Review, 2nd. Ed.
    Biology Super Review, 2nd. Ed.
    Biology Super Review, 2nd. Ed.
    Ebook426 pages3 hours

    Biology Super Review, 2nd. Ed.

    By The Editors of REA

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    About this ebook

    Need help with Biology? Want a quick review or refresher for class? This is the book for you!
    REA’s Biology Super Review gives you everything you need to know!

    This Super Review can be used as a supplement to your high school or college textbook, or as a handy guide for anyone who needs a fast review of the subject.

    • Comprehensive, yet concise coverage – review covers the material that students must know about biology. Each topic is presented in a clear and easy-to-understand format that makes learning easier.

    • Questions and answers for each topic – let you practice what you’ve learned and build your biology skills.

    • End-of-chapter quizzes – gauge your understanding of the important information you need to know, so you’ll be ready for any homework assignment, quiz, or test.

    Whether you need a quick refresher on the subject, or are prepping for your next exam, we think you’ll agree that REA’s Super Review provides all you need to know!
    LanguageEnglish
    PublisherResearch & Education Association
    Release dateJun 15, 2013
    ISBN9780738684055
    Biology Super Review, 2nd. Ed.

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      Biology Super Review, 2nd. Ed. - The Editors of REA

      Ecology

      CHAPTER 1

      The Chemical and Molecular

      Basis of Life

      1.1 The Elements

      Element – An element is a substance which cannot be decomposed into simpler or less complex substances by ordinary chemical means.

      Compound – Compounds are a combination of elements present in definite proportions by mass. These are substances which can be decomposed by chemical means.

      Mixtures – Mixtures contain two or more substances, each of which retains its original properties and can be separated from the others by relatively simple means. They do not have a definite composition.

      Atoms – Each element is made up of one kind of atom. An atom is the smallest part of an element which can combine with other elements. Each atom consists of:

      A)  Atomic Nucleus – small, dense center of an atom.

      B)  Proton – positively charged particle of the nucleus.

      C)  Neutron – electrically neutral particle of the nucleus.

      D)  Electron – negatively charged particle which moves around the nucleus.

      In normal, neutral atoms, the number of electrons is equal to the number of protons.

      Figure 1.1 Atomic Structure of carbon and nitrogen

      Atomic Weight – The total number of protons and neutrons in a nucleus is the atomic weight (mass number). This number approximates the total mass of the nucleus.

      Atomic Number – The atomic number is equal to the number of protons in the nucleus of an element.

      Isotope – Atoms of the same element that have a different number of neutrons are known as isotopes. All isotopes of the same element have essentially the same chemical properties but their physical properties may be affected.

      Ions – Atoms or groups of atoms which have lost or gained electrons are called ions. One of the ions formed is always electrically positive and the other electrically negative.

      Figure 1.2 Representation of oxygen from

      Periodic Table of the Elements

      Problem Solving Examples:

         Define the following terms: atom, isotope, ion. Could a single particle of matter be all three simultaneously?

         An atom is the smallest particle of an element that can retain the chemical properties of that element. It is composed of a nucleus, which contains positively charged protons and neutral neutrons. Negatively charged electrons revolve in orbits around the nucleus. For example, a helium atom contains 2 protons, 2 neutrons, and 2 electrons.

      An ion is an electrically charged atom or group of atoms. The electrical charge results when a neutral atom or group of atoms loses or gains one or more electrons during a chemical reaction. An ion which is negatively charged is called an anion, and a positively charged ion is called a cation.

      Isotopes are alternate forms of the same element. An element is defined in terms of its atomic number, which is the number of protons in its nucleus. Isotopes of an element have the same number of protons, but a different number of neutrons. Since atomic mass is determined by the number of protons plus neutrons, isotopes of the same element have varying atomic masses. For example, deuterium (²H) is an isotope of hydrogen, and has one neutron and one proton in its nucleus. Hydrogen has only a proton and no neutrons in its nucleus.

      A single particle can be an atom, an ion, and an isotope simultaneously. The simplest example is the hydrogen ion (H+). It is an atom which has lost one electron and thus developed a positive charge. Since it is charged, it is therefore an ion. A cation is a positively charged ion (i.e., H+), and an anion is a negatively charged ion (i.e., Cl−). If one compares the atomic number of H+ (1) with that of deuterium (1), it is seen that although they have different atomic masses, since their atomic numbers are the same, they must be isotopes of one another.

         Describe the differences between an element and a compound.

         All substances are composed of matter, which means that they have mass and occupy space. Elements and compounds constitute two general classes of matter. Elements are substances that consist of identical atoms (i.e., atoms with identical atomic numbers). This definition of an element includes all isotopes of that element. Hence O¹⁸ and O¹⁶ are both considered to be elemental oxygen. A compound is a substance that is composed of two or more different kinds of atoms (two or more different elements) combined in a definite weight ratio. This fixed composition of various elements, according to law of definite proportions, differentiates a compound from a mixture. Elements are the substituents of compounds. For example, water is a compound composed of the two elements hydrogen and oxygen in the ratio 2:1, respectively. This compound may be written as H2O, which is the molecular formula of water. The subscript 2 that appears after the hydrogen (H) indicates that every molecule of water has two hydrogen atoms. There is no subscript after the oxygen (O) in the molecular formula of water, which indicates that there is only one oxygen atom per molecule of water. Hence, water is a compound whose molecules are each made up of two hydrogen atoms and one oxygen atom.

      1.2 Chemical Bonds

      Ionic Bond – This involves the complete transfer of an electron from one atom to another. Ionic bonds form between strong electron donors and strong electron acceptors. Ionic compounds are more stable than the individual elements.

      Covalent Bond – This involves the sharing of pairs of electrons between atoms. Covalent bonds may be single, double, or triple.

      Polar Covalent Bond – A polar covalent bond is a bond in which the charge is distributed asymmetrically within the bond.

      Non-Polar Covalent Bond – A non-polar covalent bond is a bond where the electrons are distributed equally between two atoms.

      Hydrogen Bond – A hydrogen bond is the attraction of a hydrogen atom, already covalently bonded to one electronegative atom, to a second electronegative atom of the same molecule or adjacent molecule. Usually these bonds are found in compounds that have strong electronegative atoms such as oxygen, fluorine, or nitrogen.

      Van der Waals Forces – Van der Waals forces are weak linkages which occur between electrically neutral molecules or parts of molecules which are very close to each other.

      Hydrophobic Interactions – Hydrophobic interactions occur between groups that are insoluble in water. These groups, which are non-polar, tend to clump together in the presence of water.

      Problem Solving Examples:

         Distinguish between covalent and ionic bonds.

         A covalent bond is a bond in which two atoms are held together by a shared pair of electrons. An ionic bond is a bond in which oppositely charged ions are held together by electrical attraction.

      In general, the electronegativity difference between two elements influences the character of their bond (see the table on the following page).

      Electronegativity measures the relative ability of an atom to attract electrons in a covalent bond. Using Pauling’s scale, where fluorine is arbitrarily given the value 4.0 units and other elements are assigned values relative to fluorine, an electronegativity difference of greater than 1.7 units is mostly ionic in character. Therefore, a bond between two atoms with an electronegativity difference of greater than 1.7 units is mostly ionic in character. If the difference is less than 1.7, the bond is predominantly covalent.

      Electronegativities of main groups of elements

         What are hydrogen bonds? Describe fully the importance of hydrogen bonds in the biological world.

         A hydrogen bond is a molecular force in which a hydrogen atom is shared between two atoms. Hydrogen bonds occur as a result of the uneven distribution of electrons in a polar bond. Here, the bonding electrons are more attracted to the highly electronegative oxygen atom, resulting in a slightly positive charge (δ+) on the hydrogen and a slightly negative charge (δ-) on the oxygen. A hydrogen bond is formed when the relatively positive hydrogen is attracted to a relatively negative atom of some other polarized bond. For example:

      The atom with which it forms the polar bond is called the hydrogen donor, while the other atom is the hydrogen acceptor. Note, however, that the bond is an electrostatic one – no electrons are shared or exchanged between the hydrogen and the negative dipole of the other molecule of the bond.

      Hydrogen bonds are highly directional (note the arrows in the figure), and are strongest when all three atoms are colinear (when the bond angles between the atoms are 180°).

      Bond energies of hydrogen bonds are in the range of about 3 to 7 kcal/mole. This is intermediate between the energy of a covalent bond and a van der Waals bond. However, only when the electronegative atoms are either F, O, or N is the energy of the bond enough to make it important.

      Hydrogen bonds are responsible for the structure of water and its special properties as a biological solvent. There is extensive hydrogen bonding between water molecules, forming what is called the water matrix. The formation of the water matrix is what causes frozen H2O (water) to float and be less dense than liquid form, sustaining life even in frozen ponds or rivers. This structure has profound effects on the freezing and boiling points of water and its solubility properties. Any molecule capable of forming a hydrogen bond can do so with water, which results in dissociation, or solubility of the molecule.

      Hydrogen bonds are also most responsible for the maintenance of the shape of proteins. Since shape is crucial to protein function, this bonding is extremely important. For example, hydrogen bonds maintain the helical shape of keratin and collagen molecules and give them their characteristic strength and flexibility.

      DNA helices are held together by hydrogen bonds. Bonding occurs between the base pairs. The intermediate bond strength of the hydrogen bond is ideal for the function of DNA – it is strong enough to give the molecule stability, yet weak enough to be broken with sufficient ease for replication and RNA synthesis.

         What are van der Waals forces? What is their significance in biological systems?

         Van der Waals forces are the weak attractive forces that molecules of non-polar compounds have for one another. These are the forces that allow non-polar compounds to liquefy and/or solidify. These forces are based on the existence of momentary dipoles within molecules of non-polar compounds. A dipole is the separation of opposite charges (positive and negative). A non-polar compound’s average distribution of charge is symmetrical, so there is no net dipole. But, electrons are not static; they are constantly moving about. Thus, at any instant in time a small dipole will exist. This momentary dipole will affect the distribution of charge in nearby non-polar molecules, inducing charges. This induction happens because the negative end of the temporary dipole will repel electrons and the positive end attracts electrons. Thus, the neighboring non-polar molecules will have oppositely oriented dipoles:

      These momentary, induced dipoles are constantly changing, short range forces. But, their net result is attraction between molecules.

      The attraction due to van der Waals forces steadily increases when two non-bonded atoms are brought closer together, reaching its maximum when they are just touching. Every atom has a van der Waals radius. If the two atoms are forced closer together than the minimum radius, van der Waals attraction is replaced by van der Waals repulsion (because of the positively charged nuclei). The atoms then try to return to equilibrium of the maximum radius.

      Both attractive and repulsive van der Waals forces play important roles in many biological systems. These forces acting between non-polar chains of phospholipids serve to hold the membranes of living cells together.

      1.3 Acids and Bases

      Acid – An acid is a compound which dissociates in water and yields hydrogen ions [H+]. It is referred to as a proton donor.

      Base – A base is a compound which dissociates in water and yields hydroxyl ions [OH-]. Bases are proton acceptors.

      Figure 1.3. Reaction between hydrochloric acid (proton donor)

      and ammonia (proton acceptor)

      pH – The degree of acidity or alkalinity is measured by pH.

      Problem Solving Examples:

         Differentiate between acids and bases. Give examples of each. How is water defined?

         There are essentially 2 widely used definitions of acids and bases: the Lowry-Bronsted definition and the Lewis definition. In the Lowry-Bronsted definition, an acid is a compound with the capacity to donate a proton, and a base is a compound with the capacity to accept a proton. In the Lewis definition, an acid has the ability to accept an electron pair and a base the ability to donate an electron pair.

      Some common acids important to the biological system are acetic acid (CH3COOH), carbonic acid (H2CO3), phosphoric acid (H3PO4), and water. Amino acids, the building blocks of protein, are compounds that contain an acidic group (−COOH). Some common bases are ammonia (NH3), pyridine (C5H5N), and water. The nitrogenous bases important in the structure of DNA and RNA carry the purine or pyridine functional group. Water has the ability to act both as an acid (H2O → H+ + OH-) and as a base (H2O → H+ + H3O+) depending on the conditions of the reaction, and is thus said to exhibit amphiprotic behavior.

         What does the pH of a solution mean?

         The pH (an abbreviation for potential of hydrogen) of a solution is a measure of the hydrogen ion (H+) concentration. Specifically, pH is defined as the negative log of the hydrogen ion concentration. A pH scale is used to quantify the relative acid or base strength. It is based upon the dissociation reaction of water: H2O → H+ + OH-. The dissociation constant (K) of this reaction is 1.0 × 10-14. [H2O] is the concentration of water (which is equal to one). The pH of water can be calculated from its dissociation constant K. Since one H+ and one −OH are formed for every dissociated H2O molecule, [H+] = [−OH].

      A pH of 7 is considered to be neutral since there are equal concentrations of hydrogen and hydroxide ions (OH-). The pH scale ranges from 0 to 14. Acidic compounds have a pH range of 0 to 7 and basic compounds have a range of 7 to 14.

      1.4 Chemical Changes

      Chemical Reaction – A chemical reaction refers to any process in which at least one chemical bond is either broken or formed. The outcome of a chemical reaction is a rearrangement of atoms and bonding patterns.

      Laws of Thermodynamics

      A)  First Law of Thermodynamics (Conservation of Energy) – In any process, the sum of all energy changes must be zero.

      B)  Second Law of Thermodynamics – Any system tends toward a state of greater entropy, meaning randomness or disorder.

      C)  The Third Law of Thermodynamics – A perfect crystal, which is a completely ordered system, at absolute zero (0° Kelvin) would have perfect order, and therefore its entropy would be zero.

      Stability of chemical system depends on:

      A)  Enthalpy – total energy content of a system.

      B)  Entropy – energy distribution.

      Exergonic Reaction – Exergonic reactions release free energy; all spontaneous reactions are exergonic.

      Endergonic Reaction – Endergonic reactions require the addition of free energy from an external source.

      Figure 1.4 Exergonic and Endergonic Reactions

      Problem Solving Examples:

         What are the three laws of thermodynamics? Discuss the biological significance of the first two.

         The first law thermodynamics states that energy can be converted from one form into another, but it cannot be created or destroyed. In other words, the energy of a closed system is constant. Thus, the first law is simply a statement of the law of conservation of energy; the sum of all energy changes must be zero.

      The second law of thermodynamics states that the total entropy (a measure of the disorder or randomness of a system) of the universe is increasing. This is characterized by a decrease in the free energy, which is the energy available to do work. Thus, any spontaneous change that occurs (chemical, physical, or biological) will tend to increase the entropy of the universe.

      The third law of thermodynamics refers to a completely ordered system, particularly, a perfect crystal. It states that a perfect crystal at absolute zero (0 Kelvin) would have perfect order, and therefore its entropy would be zero.

      These three laws affect the biological as well as the chemical and physical worlds. Living cells do their work by using the energy stored in chemical bonds. The first law of thermodynamics states that every chemical bond in a given molecule contains an amount of energy equal to the amount that was necessary to link the atoms together. Thus, living cells turn other forms of energy into chemical bond energy and use it to do work. A living organism is a storehouse of potential chemical energy due to the many millions of atoms bonded together in each cell, so it might appear that the same energy could be passed continuously from organism to organism with no required extracellular energy source. However, the second law of thermodynamics tells us that every energy transformation results in a reduction in the usable or free energy of the system. Consequently, there is a steady increase in the amount of energy that is unavailable to do work (an increase in entropy). In

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