Microbiology: A Systems Approach, 2 nd ed.

1. Microbiology: A Systems Approach, 2nd ed. Chapter 2: The Chemistry of Biology

2. 2.1 Atoms, Bonds, and Molecules: Fundamental Building Blocks • Matter: anything that occupies space and has mass • Can be liquid, solid, or gaseous state • Building blocks of matter- atoms • Subatomic particles of atoms- protons (p+), neutrons (n0), and electrons (e-) • Protons and neutrons make up the nucleus, electrons surround the nucleus • Held together by the attraction of positive protons to negative electrons

3. Figure 2.1

4. Different Types of Atoms: Elements and Their Properties Different numbers of protons, neutrons, and electrons in atoms create different elements Each element has a characteristic atomic structure and predictable chemical behavior Each assigned a distinctive name with an abbreviated shorthand symbol

5. The Major Elements of Life and Their Primary Characteristics Isotopes- variant forms of the same element that differ in the number of neutrons Radioactive isotopes used in research and medical applications and in dating fossils and ancient materials Electron orbitals and shells

6. Electron Orbitals and Shells An atom can be envisioned as a central nucleus surrounded by a “cloud” of electrons Electrons rotate about the nucleus in pathways called orbitals- volumes of space in which an electron is likely to be found Electrons occupy energy shells, from lower-energy to higher-energy as they move away from the nucleus Electrons fill the orbitals and shells in pairs starting with the shell nearest the nucleus Each element, then, has a unuiqe pattern of orbitals and shells

7. Figure 2.2

8. Bonds and Molecules Most elements do not exist naturally in pure form Molecule- the smallest particle of matter that can have independent existence; a distinct chemical substance that results from one atom of a noble gas (Ne) or the combination of two or more atoms (can be two atoms of the same element, such as O2). Compounds- are combinations of two or more different elements joined by chemical bonds. Chemical Bonds- When two or more atoms share, donate, or accept electrons Types of bonds formed and to which atoms and element bonds are determined by the atom’s valence

9. Figure 2.3

10. Covalent Bonds and Polarity: Molecules with Shared Electrons Covalent bonds- between atoms that share electrons (such as H2). The majority of molecules associated with living things are composed of single and double covalent bonds between C, H, O, N, S, and P.

11. Figure 2.4

12. Polar vs. Nonpolar Molecules • Some covalent bonds result in a polar molecule- an unequal distribution of charge (ex. H2O). • Polarity is a significant property of many large molecules, influencing both reactivity and structure. • An electrically neutral molecule is nonpolar • Van der Waals forces- weak attractions between molecules with low levels of polarity

13. Figure 2.5

14. Ionic Bonds: Electron Transfer Among Atoms • Electrons transferred completely from one atom to another, without sharing, results in an ionic bond (ex. NaCl) • Crystals with ionic bonds, when dissolved in a solvent, can separate in to charged particles called ions in a process called ionization • Cations- positively charged ions • Anions- negatively charged ions • These ionic molecules that dissolve to form ions are called electrolytes

15. Figure 2.6

16. Figure 2.7

17. Hydrogen Bonding Figure 2.8 • Weak bond between a H covalently bonded to one molecule and an O or N atom on the same or different molecule (such as between water molecules)

18. Chemical Shorthand: Formulas, Models, and Equations Molecular formula- gives atomic symbols and the number of atoms of the elements involved in subscript (H2o, C6H12O6). Molecular formulas might not be unique (glucose, galactose, and fructose, for example) Structural formulas illustrate the relationships of the atoms and the number and types of bonds

19. Figure 2.9

20. Chemical Equations • Equations are used to illustrate chemical reactions • Reactants- Molecules entering the reaction • Products- the substances left by a reaction

21. Types of Reactions • Synthesis: reactants bond together to form an entirely new molecule • A + B  AB • S + O2 SO2 • 2H2 + O2 2H2O (note that equations must be balanced) • Decomposition: bonds on a single reactant molecule are permanently broken to release two or more product molecules • AB  A + B • 2H2O2 2H2O + O2

22. Types of Reactions: • Exchange: The reactants trade places between each other and release products that are combinations of the two • AB + XY AX + BY (reversible reaction)

23. CATALYSTS • Catalysts- increase the rate of the reaction • Catalysts lower the energy required to get reactions started • Enzymes are biological catalysts • Most enzymes are proteins, but other substances, e.g. RNA can occasionally serve as enzymes

24. Solutions: Homogeneous Mixtures of Molecules • Solution- a mixture of one or more solutes uniformly dispersed in a solvent • The solute cannot be separated by filtration or settling • The rule of solubility- “like dissolves like” • Water- the most common solvent in natural systems because of its special characteristics • Hydrophilic molecules- attract water to their surface (polar) • Hydrophobic molecules- repel water (nonpolar) • Amphipathic molecules- have both hydrophilic and hydrophobic properties

25. Concentration of Solutions • Concentration- the amount of solute dissolved in a certain amount of solvent • In biological solutions, commonly expressed as molar concentration or molarity (M) • One mole dissolved in 1 L • One mole is the molecular weight of the compound in grams

26. Figure 2.11

27. Acidity, Alkalinity, and the pH Scale • Acidic solutions- when a component dissolved in water (acid) releases excess hydrogen ions (H+) • Basic solutions- when a component releases excess hydroxide ions (OH-) • pH scale- measures the acid and base concentrations of solutions • Ranges from 0 (most acidic) to 14 (most basic); 7 is neutral • pH = -log[H+]

28. Figure 2.12

30. Neutralization Reactions Neutralization reactions- occur in aqueous solutions containing both acids and bases Give rise to water and other neutral by-products HCl + NaOH  H2O + NaCl

31. The Chemistry of Carbon and Organic Compounds • Inorganic chemicals- usually do not contain both C and H (ex. NaCl, CaCO3) • Organic chemicals- Carbon compounds with a basic framework of the element carbon bonded to other atoms • Most of the chemical reactions and structures of living things involve organic chemicals

32. Carbon- the Fundamental Element of Life Valence makes it an ideal atomic building block Forms stable chains containing thousands of C atoms, with bonding sites available Can form linear, branched, or ringed chains Can form single, double, or triple bonds Most often associates with H, O, N, S, and P

33. Figure 2.13

34. Functional Groups of Organic Compounds • Special molecular groups or accessory molecules that bind to organic compounds- functional groups • Help define the chemical class of certain groups of organic compounds • Give organic compounds unique reactive properties • Reactions of an organic compound can be predicted by knowing the kind of functional group or groups it carries

36. 2.2 Macromolecules: Superstructures of Life • Biochemistry- study of the compounds of life • Biochemicals- organic compounds produced by (or components of) living things • Four main families- carbohydrates, lipids, proteins, and nucleic acids • Often very large, called macromolecules • All macromolecules except for lipids are formed by polymerization • Repeating subunits (monomers) are bound in to chains of various lengths (polymers)

37. Carbohydrates • Carbohydrates: Sugars and Polysaccharides • Most can be represented by the general formula (CH2O)n, where n = the number of units of this combination of atoms

38. Figure 2.14

39. Carbohydrates • Exist in a variety of configurations • Sugar (saccharide)- a simple carbohydrate with a sweet taste • Monosaccharide contains 3-7 carbons • Disaccharide contains two monosaccharides • Polysaccharide contains five or more monosaccharides • Monosaccharides and disaccharides are specified by combining a prefix that describes a characteristic of the sugar with the suffix –ose • Hexoses- six carbons • Pentoses- five carbons • Fructose- for fruit

40. The Nature of Carbohydrate Bonds Figure 2.15

41. The Functions of Polysacharides Figure 2.16a Structural support and protection Serve as nutrient and energy stores Cell walls in plants and many microscopic algae from cellulose

42. Other Important Polysaccharides Figure 2.16b Include agar, peptidoglycan, chitin, lipopolysaccharide, glycocalyx, and glycogen

43. Lipids: Fats, Phospholipids, and Waxes • Lipids- a variety of substances that are not soluble in polar substances • Will dissolve in nonpolar solvents • Main groups of lipids • Triglycerides-a single molecule of glycerol bound to three fatty acids • Includes fats and oils

44. Figure 2.17

45. Phospholipids • Phospholipids- Contain two fatty acids attached to the glycerol with a phosphate group on the third glycerol binding site • Important membrane molecules Figure 2.18

46. Miscellaneous Lipids • Steroids- complex ringed compounds commonly found in cell membranes and animal hormones • Best known- cholesterol • Waxes- esters formed between a long-chain alcohol and a saturated fatty acid

47. Figure 2.19

48. Proteins: Shapers of Life • Predominant organic molecules • Building blocks- amino acids • 20 different naturally occurring forms • Basic skeleton- a carbon (the α carbon) linked to an amino group (NH2), a carboxyl group (COOH), a hydrogen atom (H), and a variable R group • Peptide bond forms between the amino group on one amino acid and the carboxyl group on another.

50. Figure 2.20