Chapter 2: The Chemical Context of Life

1. Chapter 2: The Chemical Context of Life

2. Reminder from Chapter 1: Organisms are natural systems to which basic concepts of chemistry and physics apply. One of the main themes of biology is the organization of life on a hierarchy of structural levels, with additional properties emerging at each successive level. In this chapter, we will see how the theme of emergent properties applies to the lowest level of biological organization.

3. Matter consists of chemical elements in pure form and in combinations called compounds •All life is composed of matter. Matter = anything that takes up space and has mass. Matter exists in diverse forms, each with its own characteristics. •All matter is made up of chemical elements. •Element = a substance that cannot be broken down to other substances by chemical reactions. •Compound = combination of two or more elements in a fixed ratio. Compounds have emergent properties - i.e. characteristics different from those of its elements.

4. Life requires about 25 elements •Six elements make up 97.6% of living matter (carbon, oxygen, hydrogen, nitrogen, phosphorous, sulfur)( Table 2.1). These elements form stable covalent bonds. •Trace elements = those required by an organism in minute quantities. •Deficiencies in trace elements can cause illness. e.g. lack of iron and iodine cause anemia and goiter, respectively.

5. Atomic structure determines the behavior of an element •Atoms = fundamental unit of matter. Smallest possible amount of an element that retains that element’s properties. •Each element consists of a certain kind of atom, which is different from the atoms of other elements. •Atoms are composed of even smaller parts. Only three kinds of subatomic particles are relevant from a biological perspective: protons, neutrons, and electrons.

6. •Neutrons and protons are packed together tightly at the center of an atom to form a nucleus. The electrons move about this nucleus at almost the speed of light. •Electrons (-) and protons (+) are electrically charged, whereas the neutron is neutral. Protons give the nucleus a positive charge, and it is the attraction between opposite charges that keeps the rapidly moving electrons orbiting around the nucleus. •The unit of measurement for atomic particles is the dalton. Neutrons and protons have a mass of 1 dalton each. Electron mass is negligible.

7. •Atoms of the various elements vary in their number of subatomic particles. All atoms of a particular element have the same number of protons in their nuclei. This is their atomic number. •Mass number = sum of protons and neutrons in the nucleus of an atom. •Atomic mass (weight) = since neutrons and protons have a mass close to 1 dalton, the mass number tells us the approximate mass of the whole atom. •Isotopes = variant forms of elements. Have same number of protons and electrons, but different number of neutrons. •The nucleus of 14C is unstable and therefore radioactive. •Radioactive isotopes = nucleus decays spontaneously giving off particles and energy. Dangerous to life because it causes mutations in DNA. However, radioactive isotopes can also be useful in biological research and medicine as tracers. Living cells cannot distinguish radioactive isotopes from nonradioactive atoms of the same elements.

8. Energy levels •Atoms are mostly empty space. When two atoms approach each other during a chemical reaction, their nuclei do not come close enough together to interact. Only electrons are directly involved in the chemical reactions between atoms. •An atom’s electrons vary in the amount of energy they possess (Fig 2.7). •Energy = ability to do work. •Potential energy = energy that matter stores because of its position or location. •Matter has a natural tendency to move to the lowest possible state of potential energy. Electrons of an atom also have potential energy because of their position in relation to the nucleus. The negatively charged electrons are attracted to the positively charged nucleus. The more distant the electrons are from the nucleus, the greater their potential energy.

9. Energy levels (continued) •Changes in the potential energy of electrons can only occur in steps of fixed amounts. The different states of potential energy for electrons in an atom are called energy levels or electron shells. Electrons in first shell closest to nucleus have the lowest energy. Electrons in the second shell have more energy, electrons in third shell have more energy still, and so on. •An electron can change its shell, but only by absorbing or losing an amount of energy equal to the difference in potential energy between the old shell and the new shell.

10. Electron orbitals •We can never know the exact trajectory of an electron. Instead, we describe the volume of space in which an electron spends most of its time (Fig 2.9). •Orbital = three-dimensional space where an electron is found 90% of the time. •No more than two electrons can occupy the same orbital. •First shell has a single spherical orbital and can hold only 2 electrons. An atom with more electrons must use higher shells. •The second electron shell can hold 8 electrons, two in each of four orbitals (1 spherical and 3 dumbbell-shaped).

12. Electron configuration and chemical properties •The chemical properties of an atom depend mostly on the number of electrons in its outermost shell (Fig 2.8). •valence electrons = electrons in outer most shell •valence shell = outermost energy shell. •Valence = an atom’s bonding capacity (# of electrons needed to fill outer shell). •Atoms with a complete valence shell are not reactive. All other atoms are chemically reactive because they have incomplete valence shells with unpaired electrons.

13. Atoms combine by chemical bonding to form molecules •When atoms with incomplete outer shells react, each atom gives up or acquires electrons so that partners end up with completed outer shells. •Atoms do this by either sharing (covalent bonds) or transferring outer electrons (ionic bonds) resulting in chemical bonds. •The strongest chemical bonds are covalent bonds and ionic bonds.

14. Covalent bonds •Covalent bond = two atoms sharing one or more pairs of outer shell electrons ( Fig 2.10 and 11). •Molecule = two or more atoms held together by covalent bonds. •The number of single covalent bonds an atom can form is equal to the number of additional electrons needed to fill its outer shell (i.e. it's valence). •Double bond = sharing of 2 pairs of electrons. Stronger than single bonds. •Atoms in a covalently bonded molecule are constantly in a tug-of-war for the electrons of their covalent bonds. •Electronegativity = an atom’s attraction for the shared electrons of the bond. The more electronegative an atom, the more strongly it pulls electrons towards itself. •Nonpolar covalent bonds = electrons shared equally between the atoms of equal electronegativity (H2, O2, CH4 ). •Water is made up of 2 kinds of atoms with differing electronegativity (O>H). Oxygen attracts electrons more strongly than hydrogen. •Polar covalent bond = chemical bond in which shared electrons are pulled closer to the more electronegative atom, making it partially negative and the other atom partially positive (Refer to Fig 2.12). •H2O, even though is neutral overall, has a slightly negative pole and two slightly positive poles, making it a polar molecule.

15. Ionic bonds •Refer to Fig 2.13 Ionic bonds = attractions between ions of opposite charge (e.g. table salt, NaCl). Much weaker than covalent bonds. •When atoms of chlorine and sodium collide, chlorine atom strips sodium’s outer electron away. This results in sodium having a positive charge and chlorine having a negative charge. Two ions of opposite charge attract each other; when the attraction holds them together, it is called an ionic bond. Ion = atom or molecule with an electrical charge resulting from a gain or loss of one or more electrons. anion = ion with negative charge; cation = ion with a positive charge •NaCl is a type of salt. Salts are ionic compounds that often exist as crystals in nature.

17. Weak chemical bonds play important roles in the chemistry of life •Weak bonds, unlike covalent bonds, allow interactions between molecules to be brief; molecules may come together, change in some way and then separate. •The most important weak bond in living matter is the hydrogen bond. Hydrogen bond = occurs when a hydrogen atom covalently bonded to one electronegative atom is also attracted by another electronegative atom. In living cells, the electronegative partner involved are usually oxygen and nitrogen atoms. (refer to Fig 2.15) •Hydrogen bonds, ionic bonds, and other weak bonds, form between and within molecules. Although these bonds are individually weak, their cumulative effect can re-enforce the 3-D shape of a large molecule.

18. A molecules biological function is related to its shape •Molecular shape is important in biology because it is the basis for how most molecules of life recognize and respond to one another. •Recognition and binding of neurotransmitters to cell surface receptors in synapses of brain cells is basis on intercellular communication in the nervous system (Fig 2.16 and 17).

19. Chemical reactions make and break chemical bonds •Living matter is not static. There is constant flux, as new molecules are being built and others are being broken down. The goal of biochemistry is not simply to catalogue the molecules that make up the living world, but to understand how these molecules are transformed into others in biochemical pathways. These transformations always involve chemical reactions. •In a chemical reaction, reactants interact, atoms rearrange, and products result. •Matter is conserved in a chemical reaction. Reactions cannot create nor destroy matter but can only rearrange it. •Living cells carry out thousands of chemical reactions that rearrange matter in significant ways. •Some chemical reactions go to completion, others are reversible. •Chemical equilibrium = point at which rate of forward reaction equals that of reverse reaction

20. Chapter 3: Water and the Fitness of the Environment

21. The polarity of water molecules results in hydrogen bonding •Oxygen is more electronegative than hydrogen. Consequently, the electrons of the polar bonds spend more time near the oxygen atom. This makes water a polar molecule. •The unique (emergent) properties of water arises from attractions among these polar molecules. •Each water molecule can hydrogen bond (H-bond) to a max of four neighbors. •H-bond = electrostatic attraction between a hydrogen in a polar bond to an electronegative atom of another molecule. •The charged regions of a polar molecule are attracted to opposite charges of neighboring polar or ionic molecules.

22. Organisms depend on the cohesion of water molecules •Water molecules stick together as a result of H-bonding. H-bonds form, break, and reform very frequently. At any given time, a substantial portion of all molecules are bonded to their neighbors, giving water more structure than most liquids. •Cohesion = tendency of molecules to stick together. Much stronger for water than for other liquids. Important in water transport in plants. •Adhesion = the clinging of one substance to another. Also important in water transport in plants. •Surface tension = a measure of how difficult it is to stretch or break the surface of a liquid. Higher for water than for most liquids.

23. Water contributes to earth's habitability by moderating temperatures •Water stabilizes air temperature by absorbing heat from air that is warmer and releasing the stored heat to the air that is cooler. •Water can store a lot of energy (heat) with only a slight increase in its own temperature. •Heat = measure of the total quantity of kinetic energy (energy of motion) due to molecular motion in a body of matter. •Temperature = measures intensity of heat due to the average kinetic energy of the molecules. •Calorie = amount of heat energy it takes to raise the temp of 1 gram of water by 1°C (Food calorie = 1000 calories)

24. Water's high specific heat •Specific heat= amount of heat that must be absorbed or lost for 1 gram of that substance to change temperature by 1 °C. •Compared to most substances, water has an unusually high specific heat (10x that of Fe) This is due to H-bonding. •A calorie of heat causes a relatively small change in temperature of water because much of that heat energy is used to disrupt H-bonds before water molecules can begin to move faster. •Conversely, when the temperature of water drops slightly, many additional H-bonds form, releasing a lot of heat energy. •Water buffers against extreme changes in temperature.

25. Evaporative cooling •Molecules in a liquid stay close together because they are attracted to one another. Molecules moving fast enough to overcome these attractions can depart from the liquid and enter into gas state. •Heat of Vaporization = quantity of heat a liquid must absorb for 1 gram of it to be converted from liquid to gaseous state. •Water's high heat of vaporization helps moderate earth's climate. A considerable amount of solar heat absorbed by tropical seas is consumed during evaporation of surface water. Thus, as moist tropical air circulates poleward, it releases heat as it condenses to form rain. •Evaporative cooling also helps moderate temperature in lakes and ponds, and prevents terrestrial organisms from overheating.

26. Ice floats •Water is one of the few substances that is less dense as a solid than as a liquid. (Fig 3.5) •If ice sank, eventually all bodies of water would freeze solid since floating ice insulates liquid water below. •Ice floats because as temperature decreases, there is less energy to break H-bonds, so eventually all water molecules are H-bonded to one another resulting in a crystal lattice structure in which water molecules are less densely packed.

27. Water as the solvent of life •Water dissolves more solutes than any other liquid - called “Universal Solvent.” • Solution = liquid that is a homogenous mix of 2 or more substances •Solvent = dissolving agent •Solute = substance that is dissolved •Aqueous solution = one in which water is solvent•The versatility of water as a solvent is based on its polarity. •Ions and polar water molecules have a mutual affinity through electrical attractions. E.g. Cl- attracted and surrounded by positive part of water and Na+ attracted and surrounded by negative part of water molecules - the resulting sphere of water molecules around each ion is called the hydration shell.

28. •A compound does not have to be an ion to be dissolved by water. Polar compounds are also water-soluble ( E.g. proteins, carbohydrates, nucleic acids).(refer to Fig 3.8: hydration of soluble protein) and become surrounded by a hydration shell as well. •Hydrophilic = any substance with an affinity for water (ions, polar molecules), even if that substance does not dissolve (E.g. cellulose) •Hydrophobic = any substance that neither dissolves nor has an affinity for water (nonpolar). E.g fats, waxes, etc...

29. Solute concentrations in Aqueous solutions Knowing concentrations is important in biology. This allows the combinations of substances in fixed ratios to make chemical solutions. •Molarity = Number of moles/Liter •A mole (mol) is a quantity = 6.02 x 1023 (Avogadro’s Number) molecules •Molecular weight(mass) = sum of the weight of all the atoms in a molecule. • Molar Mass = molecular weight(mass) (in grams) of a particular substance per 1 mole of that substance •e.g. 1 mole of sucrose has 6.02 x 1023 molecules and weighs 342 g. 1 mole of ethanol also has 6.02 x 10-23 molecules, but weighs only 46 g.

30. Acids and Bases •In pure water, [H+] = [OH-] = 10-7 •Acid = substance that increases the [H+] of a solution. A strong acid, such hydrochloric acid, dissociates completely when mixed with water. Base = A substance that reduces the [H+] in a solution. A strong base, such as sodium hydroxide, dissociates completely when mixed with water. Some bases, such as ammonia, reduce [H+] directly by accepting hydrogen ions Weak acids and bases do not dissociate completely. Ex: Carbonic Acid

31. The pH scale •pH = -log [H+] •For a neutral solution [H+] is 10-7 M, therefore -log 10-7 = -(-7) = 7 •Each pH unit represents a 10 fold change in concentration.

32. Buffers •Because biological systems are very sensitive to pH, they need to minimize changes in pH. They do this with buffers. •Buffers = compounds that resist changes to their own pH when acids or bases are introduced. •A buffer works by accepting hydrogen ions from the solution when they are in excess and donating hydrogen ions to the solution when they have been depleted. Most buffers are weak acids or bases. (acid-base pairs) Examples of physiological buffers: carbonate and phosphate buffers

33. Chapter 4: Carbon and the molecular diversity of life Biological diversity reflects molecular diversity. Of all chemical elements, carbon is unparalleled in its ability to form molecules that are large, complex, and diverse. This chapter focuses on the concepts of molecular architecture that highlight carbon's importance to life.

34. Carbon atoms are the most versatile building blocks of molecules •Valence of the major elements (Fig 4.4). •The chemical characteristics of an atom depend on its valence electrons, by determining the number of bonds an atom will form with other atoms. The tetravalence of carbon explains its versatility at making large, diverse, complex molecules possible

35. Variation in carbon skeletons contributes to diversity of organic molecules •Variation in carbon skeletons is important source of molecular complexity and diversity. •Hydrocarbons = organic molecules consisting of only carbon and hydrogen(Fig 4.5). •Major component of petroleum, a fossil fuel. •Nonpolar, therefore hydrophobic (Fig 4.6) •can vary in length•Position of double bonds can vary •may be branched or unbranched •can be arranged in rings

36. •Isomers = compounds with same molecular formula but different structures and hence properties. •Types of isomers (Fig 4.7): •1. Structural isomers - differ in covalent arrangement of their atoms•contributes significantly to their variation e.g. 18 types for C8H18, and 366,319 types for C20H42 •2. Geometric isomers - have same covalent relationships, but differ in spatial arrangements (cis- and trans-) •arise from inflexibility of double bonds •have distinct biological activities (ex: retinal) •3. Enantiomers - molecules that are mirror images of each other •asymmetric carbon attached to 4 different atoms or groups of atoms. •thalidomide was a mix of 2 enantiomers: one with sedative properties, the other caused birth defects •L-Dopa vs D-Dopa is another example (fig 4.8).

37. Functional groups •The distinctive properties of organic molecules depend not only on arrangement of carbon skeleton, but also on other molecular components attached to that skeleton. •Functional groups = groups of atoms attached to carbon skeleton which are commonly involved in chemical reactions. •All functional groups studied here are hydrophilic.

38. •Hydroxyl group (-OH) •compounds containing -OH are alcohols. Eg. ethanol, sugars, phenol •-OH group is polar. Therefore such compounds dissolve in water (sugars) •Carbonyl group (C=O) •aldehyde = carbonyl group on end carbon of chain •ketone =- carbonyl group attached to internal carbon •Variation in locations of functional groups along carbon chain is a source of variation. •Carboxyl group (-COOH) •compounds containing carboxyl group are called carboxylic acids. Eg. formic acid, acetic acid, amino acids •they are weak acids because they are a source of H+ ions •Amino group (-NH2) •compounds containing amino group are called amines. Eg. amino acids Note: amino acids are both amines and carboxylic acids. •Amino group can act as a base. •Sulfhydryl group (-SH) •organic compounds containing sulfhydryls are called thiols •important in stabilizing protein structure. •Phosphate group •Phosphate is an anion formed by dissociation of an inorganic acid called phosphoric acid (H3PO4) •Functions in energy transfer between organic molecules