2: Small Molecules: Structure and Behavior

Introduction

· Living things are composed of the same chemical elements as the nonliving world and obey the same physical and chemical laws.

· This chapter provides essential information on the characteristics and properties of atoms and molecules, which will be useful in the study of biology—the study of life.

Atoms: The Constituents of Matter

· All matter is composed of atoms.

· Atoms vary in size, but all are amazingly small.

· Each atom consists of at least one proton and one electron (as does a hydrogen atom, the simplest atom).

· The proton is found in the center of the electron's orbit, in a region called the nucleus.

· The nucleus possesses a positive electrical charge relative to the electron, which is negatively charged.

· Atoms have mass.

· Mass is a measure of quantity.

· In an atom, the mass comes mostly from the proton and a neutrally charged body called a neutron.

· A neutron (if present) is found in the nucleus, the same region where the proton is found.

· The mass of a proton or a neutron is 1.7 ´ 10^–24 grams.

· Grams are abbreviated g and are a widely used unit of mass.

· The mass of an electron, 9 ´ 10^–28 g, is usually ignored because it is so much less than the mass of a proton or neutron.

· An atom is electrically neutral when its number of protons equals its number of neutrons.

· Atoms are mostly empty space. If an atom were expanded proportionately to the size of the largest dome in the world, the nucleus would still only be the size of a grain of salt, and the electrons would be too small to see. The remainder of this enormous atom would be empty space.

An element is made up of only one kind of atom

· Each type of element contains only one type of atom. Hydrogen, for instance, consists solely of hydrogen atoms.

· More than 100 different elements are found in the universe.

· Information on elements is arranged in logical order in a table called the periodic table. (See Figure 2.1.)

· The periodic table groups the elements according to their physical and chemical properties. It arranges elements left to right based on their atomic number, and in columns based on similarities in their properties.

· About 98% of the mass of living organisms is made of carbon, hydrogen, nitrogen, oxygen, and sulfur.

· Elements contain just one kind of atom; substances that are composed of more than one kind of atom are called compounds.

The number of protons identifies the element

· Each element has a unique atomic number.

· The atomic number is the number of protons found in an atom of an element.

· The mass number is the total number of protons plus the number of neutrons (See the figure in the upper left corner of page 18.)

· The mass number is used as the weight of the atom in units called daltons.

· Each element has either a one- or two-character symbol: H is hydrogen, C is carbon, Na is sodium, and Fe is iron.

· Each box of the periodic table has a two-number designation. In the box for carbon, for example, the number 12 is the mass of carbon and the number 6 is the atomic number.

Isotopes differ in number of neutrons

· All atoms of an element have the same number of protons, but not necessarily the same number of neutrons.

· Atoms of the same element that have different atomic weights are called isotopes. (See Figure 2.2.)

· The common form of hydrogen is ¹H. One of every 6,500 hydrogen atoms on Earth has a neutron, which makes it ²H.

· The ³H form has 2 neutrons and weighs three times as much as ¹H.

· ²H is called deuterium; ³H is called tritium.

· Carbon exists as ¹²C, ¹³C, and ¹⁴C in nature. Like the isotopes of most elements, carbon's isotopes do not have distinct names. They are referred to as carbon-12, carbon-13, and carbon-14.

· Each element's mass number is calculated based on the naturally occurring mix of its isotopes.

· Some isotopes are radioisotopes.

· Radioisotopes emit energy as alpha, beta, and gamma radiation from their nuclei.

· Radioactive decay transforms the original atom into another atom, usually of another element.

· This means that a change in the number of protons may occur.

Electron behavior determines chemical bonding

· The location of an electron cannot be determined, but the area it travels can. This region is called the electron's orbital.

· The orbitals are a series of electron shells or energy levels around the nucleus.

· Two electrons at most can occupy each orbital.

· Electron shells, which can be viewed as energy levels, are composed of orbitals.

· The first shell is the innermost shell and has just one orbital, called the s orbital.

· The s orbital fills first and its electrons have the lowest energy.

· Hydrogen has one electron in an s orbital, and helium has two.

· The second shell is next closest to the nucleus and has one s and three p orbitals. (See Figure 2.4.)

· The second shell can accommodate eight electrons, two per orbit.

· The subsequent shells are found progressively outward from the nucleus.

· The outermost shell of an atom determines how it reacts with other atoms.

· Generally, if eight electrons are in the outer shell, the atom is stable and does not tend to react.

· Inert elements include helium, neon, argon, and krypton. (See Figure 2.5.)

· Other atoms, which do not have eight electrons in the outermost shell, share, gain, or lose electrons to arrive at a stable state.

· Atoms that share electrons are bonded together.

· A molecule is two or more atoms bonded together.

· The tendency of atoms to be stable when they have eight electrons in their outermost shells is called the rule of eight or the octet rule.

· Of the elements important to life, hydrogen and phosphorus are exceptions to the octet rule: Hydrogen is stable with two electrons in its orbital, phosphorus, commonly with 10 electrons in its outermost orbitals.

Chemical Bonds: Linking Atoms Together

· Bonds vary in strength; the strongest bonds are covalent bonds, which involve sharing electrons. (See Table 2.1.)

· The next strongest bonds are ionic and hydrogen bonds, which have approximately one-tenth the strength of covalent bonds.

Covalent bonds consist of shared pairs of electrons

· A covalent bond is the sharing of a pair of electrons between two atoms.

· In hydrogen molecules, H₂, a pair of electrons share a common orbital and spend equal amounts of time around each of the two nuclei. (See Figure 2.6.)

· The nuclei stay some distance from each other due to mutually repelling positive charges. The balanced distance between the two nuclei in H₂ is 0.1 nm.

· Each covalent bond has a predictable length, angle, and direction. These features make it possible to predict the 3D structures of molecules. (See Figure 2.7.)

· A double covalent bond occurs when atoms share two pairs of electrons, and triple covalent bond is the sharing of three electron pairs.

· The gas ethylene, C₂H₄, has a double-bonded pair of carbon atoms.

· Nitrogen gas, N₂, the form of nitrogen found in the atmosphere, has a triple covalent bond (NºN).

· Electrons are not always shared equally between covalently bonded atoms.

· N₂, H₂, O₂, and other bonds between atoms of the same element share electrons equally because their nuclei are identical to each other. Equal sharing also occurs when different atoms have about the same attraction for electrons, called attractive force.

· When a molecule has nuclei with different attractive forces, an electron spends most of its time around the nucleus with the greater attractive force.

· The attractive force that an atom exerts on electrons is called electronegativity. (See Table 2.3.) Electronegativity is determined by the number of protons in an atom and the distance of its electrons from its nucleus.

· Examples of atoms that form covalent bonds with high electronegativity are oxygen and nitrogen.

· Oxygen has six electrons in its outermost shell and requires two more to fill it. When oxygen forms covalent bonds with atoms that have weaker electronegativities, such as carbon or hydrogen, the electrons are shared unequally.

· Unequal sharing of electrons causes a partial negative charge (symbolized d^–, “delta neg”) around the more electronegative atom, and a partial positive charge (symbolized d⁺, “delta pos”) around the less electronegative atom, resulting in a polar covalent bond. (See Figure 2.8.)

· Molecules that have polar covalent bonds are called polar molecules.

· Some molecules are so large that they are said to have polar regions and nonpolar regions.

Hydrogen bonds may form between molecules

· The d^– portion of a water molecule, the area around the oxygen, has a weak attraction to the d⁺ portion of another water molecule, the area around the hydrogen. (See Figure 2.9.)

· Each of these attractions is called a hydrogen bond. Hydrogen bonds differ from covalent bonds in one important way: Hydrogen bonds do not share electrons.

· Hydrogen bonds can occur between molecules other than water.

· For example, hydrogen bonds hold together the DNA molecules found within our cells.

· Although hydrogen bonds are weak, they tend to be additive, and they are of profound biological importance, as we shall learn later.

Ions form bonds by electrical attraction

· Unlike the sharing—or even the unequal sharing—that characterizes covalent bonds, ionic bonds involve a complete transfer of one or more electrons.

· Ions are formed when an atom completely loses or gains electrons. (See Figure 2.10.)

· Ions are always symbolized by ⁺ or ^– superscripts.

· The ⁺ symbol stands for the electrical charge of the ion. A single ⁺ means the ion has one more proton than it has electrons. A single ^– means the ion has one more electron than it has protons.

· Positively charged ions are called cations. Na⁺ is an example.

· Negatively charged ions are called anions. Cl^– is an example.

· Some ions have gained or lost more than one electron.

· Calcium ion is Ca²⁺ and has two more protons than electrons.

· The Ca²⁺ ion is a cation and is said to be divalent.

· The aluminum ion is Al³⁺.

· Iron can be either Fe²⁺ (ferrous ion) or Fe³⁺ (ferric ion). In solution Fe²⁺ sometimes changes to Fe³⁺ and vice versa.

· Cuprous ion is Cu⁺, and cupric ion is Cu²⁺.

· Complex ions commonly occur in biological systems. These are groups of covalently bonded atoms that together carry an electrical charge. Examples are NH⁴⁺ (ammonium ion), SO₄^2– (sulfate ion), and PO₄^3– (phosphate ion).

· Ionic bonds are formed by the attractions of opposite charges.

· Table salt has chloride and sodium ions, which are held together by opposite attraction.

· These attractions are strong, but when introduced into water, the partial charges of the water molecules can easily interfere with the ionic bonds. For example, Figure 2.11 shows how water molecules cluster around cations and anions to "dissolve" table salt.

Polar and nonpolar substances interact best among themselves

· Like attracts like among polar and nonpolar molecules.

· Polar molecules tend to be hydrophilic (hydro means water, philic means loving). Substances that are ionic or polar often dissolve in water due to weak attractions with the hydrogen bonds.

· Nonpolar molecules are called hydrophobic (phobic means hating) because they tend to aggregate with other nonpolar molecules rather than with polar water. (See Figure 2.12.)

· The term hydrophobic is an exaggeration because even nonpolar molecules may interact very weakly with water due to the differing electronegativities of their atoms.

· Nonpolar molecules are also attracted to each other via relatively weak attractions called van der Waals forces.

· Van der Waals forces are transient attractions caused by brief variations in electron distributions between closely adjacent molecules.

· Even though each interaction is weak, the summation of many such attractions can have important effects, making van der Waals forces biologically significant.

Chemical Reactions: Atoms Change Partners

· Chemical reactions occur when atoms combine or change partners. (See Figure 2.13.)

· In a chemical reaction, reactants are converted to products.

· A chemical reaction can be written as an equation. The equation must balance because matter is neither created nor destroyed. For example, CH₄ + 2O₂ ® CO₂ + 2H₂O.

· Heat is often released from reactions.

· Changes in energy usually accompany chemical reactions.

· Stored energy, such as that in chemical bonds, is called potential energy and is available for future use.

· For example, plants store surplus collected solar energy in molecules.

· We can measure the potential energy of molecules and express it in units of heat called calories.

· A calorie is the amount of heat required to raise the temperature of one gram of pure water from 14.5°C to 15.5°C.

· A newer unit called joules (J) is growing in usage.

· The following equivalents convert between Joules and calories: 1 J = 0.239 cal, and 1 cal = 4.184 J.

· The calorie unit that is used popularly is actually 1,000 of the calorie units used by scientists.

Water: Structure and Properties

· Life now has been found in layers of rock, far below the surface of Earth.

· In the absence of sunlight, organisms derive energy from chemical sources.

· Life can exist without the need for sunlight or oxygen, but not without water. Some scientists think that wherever liquid water exists for a long enough time, life might be found in our universe.

Water has a unique structure and special properties.

· Water is the most common molecule (45 to 90%) in all organisms and also participates in or is the medium for most of an organism’s chemical reactions.

· A water molecule is composed of one oxygen and two hydrogen atoms (H₂O).

· Due to its shape, polarity, and ability to form hydrogen bonds, water has some unusual properties. (See the in-text figure on page 26.)

· For example, water is an excellent solvent; it takes a lot of heat to change its temperature relative to its weight; it has high cohesion; and it expands when frozen. (The Instructor’s Resource CD-ROM includes photos of water in various states and photos illustrating surface tension.)

· Ice is held in a crystalline structure by the orientation of water molecules’ hydrogen bonds. (See Figure 2.15.)

· Each molecule forms hydrogen bonds with four other molecules.

· These four hydrogen bonds actually increase the space water molecules take up in their solid state, so water expands as it freezes, and ice is less dense than liquid water.

· For these reasons, ice floats in liquid water. Few other substances expand when they change from liquid to solid. Glass, or silicon, is an example of another substance that does.

· Because ice floats, it forms an insulating layer on lakes and helps keep them from freezing solid.

· This insulation protects organisms in the cool, lower liquid layer from subfreezing temperatures.

· Melting and freezing

· Compared to other nonmetallic substances, it takes a lot of heat to melt ice because hydrogen bonds must be broken.

· Melting 1 mole of water requires 5.9 kJ, or 1410.1 cal, of energy.

· Each mole of water weighs 18 grams, so it takes 328 J, or 78.34 cal, to melt just 1 g of water! Recall that by definition, it only takes 1 cal to heat 1 g of water from 14.5°C to 15.5°C.

· The opposite process, freezing, requires water to lose a great deal of heat.

· This property of water helps to moderate Earth's temperature.

· Heat and cooling

· It takes a lot of heat energy to change the temperature of liquid water because of hydrogen bonds.

· Specific heat is the number of calories needed to raise one gram of a substance 1^oC. The specific heat of liquid water is 1.

· Liquid water has a higher specific heat than most other small molecules in liquid form.

· This property also helps to moderate the fluctuation of Earth's temperature.

· Evaporation and cooling

· The heat of vaporization is the amount of heat needed to change a substance from its liquid state to its gaseous state.

· It requires a lot of heat to change water to a gaseous state because the hydrogen bonds of the liquid water must be broken.

· Evaporation has a cooling effect by absorbing calories.

· Condensing has the opposite effect, releasing heat.

· Evaporation is an important means for cooling large multicellular organisms.

· A person can survive a temperature well in excess of their body, yet maintain normal body temperature by sweating and evaporating the sweat.

· Evaporation occurs less readily in high humidity, which is why hot, humid areas seem hotter at the same temperature than those places that have low humidity.

· Cohesion and surface tension

· Water has a cohesive strength even though it is a liquid.

· The attraction of water molecules to each other results in the transport of water from the roots to the tops of trees.

· Water has high surface tension, which means that the surface of liquid water is relatively difficult to puncture. (See Figure 2.16.)

· Even though water is liquid, diving into water from a significant height can cause injury. This is due to water’s surface tension.

Most biological substances are dissolved in water.

· A solution is a substance dissolved in a liquid.

· Qualitative analysis is the study of substances dissolved in a solvent and their reactions.

· Quantitative analysis measures the amounts of substances and solvents. The following are some terms used in quantitative analysis:

· Molecular formula: a compound depicted by its chemical symbols

· Molecular weight (or mass): the sum of all the atomic weights in a molecule. The molecular weight of H₂ is 2.

· Mole: the amount of a substance in grams whose weight is equal to its molecular weight. One mole of H₂ weighs 2 g. One mole of any given compound contains approximately 6.03 ´ 10²³ molecules of that compound (Avogadro's number).

· A 1 molar (1 M) solution is one mole of a compound dissolved in water to make one liter.

· One mole of sodium chloride (table salt) is the atomic weight of sodium (23) plus the atomic weight of chlorine (35.5), or 58.5, in grams. When 58.5 grams of sodium chloride are dissolved some water, and then additional water is added to create a final volume of a liter, the solution is 1 molar.

· A molar solution is abbreviated M or is symbolized by placing brackets, [ ], around the other symbols.

· 1 mM is 1/1000 of a molar solution.

· 1 µM is 1/1000 of a mM. A 1 µM solution contains one-millionth the concentration of solute as a 1 M solution (i.e., approximately 6.03 ´ 10¹⁷ molecules per liter.)

· Many water-soluble molecules in living tissues are present in the micromolar to millimolar range. Hormones are even less concentrated.

Acids, Bases, and the pH Scale

Acids donate H⁺; bases accept H⁺

· Some substances dissolve in water and release hydrogen ions (H⁺); these are called acids. Their release is called ionization.

· Other substances dissolve in water and release (also by ionization) hydroxide ions (OH^–); these are called bases.

· Acids donate H⁺; bases accept H⁺.

· If a compound increases the H⁺ ion concentration when added to water, then the compound is acidic. If the reaction is complete, such as HCl ® H⁺ + Cl^–, it is a strong acid.

· The carboxyl group (—COOH) is common in biological compounds. It functions as an acid, because —COOH ® —COO^– + H⁺.

· Not all acids dissolve fully in water. Acetic acid, for instance, does not completely react and is therefore called a weak acid.

· If a compound increases the OH^– ion concentration when added to water, then the compound is basic. There are strong and weak bases. A strong base completely reacts: NaOH ® Na⁺ + OH^–.

· A weak base, such as bicarbonate, does not completely react, and accepts H⁺ ions in several ways, one being the formation of weak carbonic acid (see page 29).

· The amino group (—NH₂) is an important part of many biological compounds; it functions as a weak base by accepting H⁺: —NH₂ + H⁺ ® —(NH₃)⁺.

· Amino acids, the building blocks of proteins, contain both carboxyl groups and amino groups, so they are simultaneously acids and bases.

· Reversible chemical reactions, in principle, can proceed in either direction, but the extent of reversibility may vary.

· Ionization of strong acids is virtually irreversible.

· Many large molecules contain weak acid or base groups.

Water is a weak acid

· Water is really a weak acid and has a slight tendency to ionize (break apart) into H⁺ and OH^–.

· This ionization is very important for living creatures and the chemical reactions they must perform.

pH is the measure of hydrogen ion concentration

· The pH scale indicates the strength of a solution of an acid or base. The scale is arrayed as a set of values 1 through 14. These values may be measured by electronic instruments. (See Figure 2.18.)

· The pH value is defined as the negative logarithm of the hydrogen ion concentration in moles per liter (molar concentration).

· A pH 7 means the concentration of hydrogen ions (or hydronium ions, which are H⁺ attached to a water molecule) is [1 ´ 10^–7]. This is 1 ´ 10^–7 moles per liter of water.

· When water is at pH 7, both the H⁺ and OH^– are at a concentration of 10^–7 molar. (See Figure 2.18.)

· When water is pH 8, it is 10^–8 molar for H⁺ and 10^–6 molar for OH^–. The higher the number of the pH, the greater the OH^– concentration and the lower the H⁺ concentration.

· When an aqueous solution is at pH 6, the H⁺ concentration is at 10^–6 molar (10 times greater than at pH 7), and the OH^– concentration becomes 10^–8 molar (1/10 that of pH 7).

· Even strongly acidic solutions have mostly water molecules and not ions. A solution with pH 1 has just one H⁺ for every 556 water molecules.

· A solution with pH 1 can have a powerfully corrosive effect on a variety of materials including metals, polysaccharides, proteins, nucleic acids, and bone.

Buffers minimize pH change

· A Buffer is a mixture of a weak acid and its corresponding base.

· Because buffers can react with both added bases and acids, they make the overall solution more resistant to pH change. (See Figure 2.19.)

· Different buffers transition to and from ionic form at their particular characteristic pH ranges.

· Buffers are common in biology and extremely important in the regulation of the internal environments of organisms.

· The many important biological buffers transition around pH 7, which keeps pH near neutral.

· Buffers illustrate the law of mass action: The addition of components to one side of a reaction drives the reaction in the direction that uses that component. As acid or base is added to a solution, the buffer will change form, transitioning between ionic and non-ionic forms.

The Properties of Molecules

· Molecules range in size and molecular weight from the very smallest, H₂, to the massive, such as the DNA molecule that makes up the length of a chromosome and contains millions of atoms.

· Carbon-containing molecules are called organic molecules and are common in life forms.

· Most organic molecules also contain hydrogen and oxygen, and many contain nitrogen and phosphorus.

· Nitrogen, oxygen, and carbon are also all found in air.

· Hydrogen and oxygen are found in water.

· All molecules have three-dimensional shapes.

· CH₄ is tetrahedral.

· Larger molecules have complex shapes that are the result of the atoms present and the ways they are linked together.

· Some molecules are long and ropelike; some are ball-shaped.

· Shape influences the behavior or function of molecules.

· Chemists use the characteristics of composition, structure, reactivity, and solubility to help classify molecules.

Functional groups give specific properties to molecules

· Some biologically important functional groups are listed in Figure 2.20.

· Functional groups are covalently bonded to organic molecules.

· Amino acids are examples of biological molecules that contains functional groups. Both a carboxyl group and an amino group are attached to the same carbon atom, which is also bonded to a hydrogen atom and a variable side chain.

Isomers have different arrangements of the same atoms

· Isomers are molecules that have the same chemical formula but different arrangements of the atoms.

· Iso means same, and is encountered in many biological terms.

· Structural isomers differ in how atoms are joined together.

· Optical isomers are mirror images of each other.

· Optical isomers can occur whenever a carbon has four different atoms or groups attached to it. (See Figure 2.21.)

· A carbon such as this is called asymmetric.

· Your left and right hands are examples of this kind of symmetry.

· Amino acids (except glycine) exist in two optical isomeric forms called d- and l-amino acids. l-amino acids are those commonly found in most organisms.

Photographs on the Instructor’s Resource CD-ROM include a scintillation counter, properties of water, and crystals of various small compounds.