Chemistry of Life

1. Chemistry of Life

2. Biological Chemistry I. Matter and Energy • A. Matter • Matter is anything that takes up space and has mass • Three states of matter exist (solid liquid and gas) • The fundamental units of matter are elements which are composed of atoms that are subdivided into subatomic particles.

3. B. Elements • Elements are substances that cannot be broken down into other substances by ordinary chemical means • Each element is represented by letters (e.g. H = hydrogen, C = carbon, etc...) • There are six elements that frequently occur in organic matter: CHNOPS = Carbon, Hydrogen, Nitrogen, Oxygen, Phosphorus, and Sulfur. • Elements are organized into compounds and molecules: • Molecules - two or more of the same element held together by chemical bonds. (e.g. O2) • Compounds - two or more different kinds of elements held together by chemical bonds. (e.g. NaCl)

5. C. Organization of Matter • Atoms are the smallest possible amount of an element. • Atoms of the same element share similar chemical properties. • Atoms are composed of subatomic particles: • Electrons (e-, negatively charged) high energy, low mass • Protons (p+, positively charged) low energy, high mass • Neutrons (n0, neutral charge) low energy, high mass

7. O2

8. Protons and neutrons are packed into a dense core called a nucleus • Positively charged protons are attracted to negatively charged electrons, but electrons have high amounts of energy, defying attraction to protons and spins around the nucleus. • The three dimensional space, where electrons are found around the nucleus, is called an orbital. • Carbon, Hydrogen, Oxygen, and Sulfur atoms

9. Elements are defined by the atomic number and mass number. • Atomic number = the number of protons in an atom • Mass number = the number of protons and neutrons in an atom • Examples - helium represented as He • Atomic number = 2 (two protons) • Mass number = 4 (two protons and 2 neutrons) • Isotopes - atoms of an element that have the same atomic number but different mass number. (therefore different number of neutrons)

11. Isotopes of Hydrogen

12. II. Energy and Energy Levels • Energy is the ability to do work (types = mechanical, chemical, thermal, and electrical) • Energy is defined as being either potential or kinetic: • Kinetic energy - energy of motion which is directly proportional to the speed of that motion. (e.g. electrons moving within an orbital) • Potential energy - energy stored by matter as a result of its location or spatial arrangement. Different states of potential energy of electrons in an atom is referred to as an energy level. The more an energy an electron possesses, the further away (thus high energy level) the electron will be from the nucleus

13. Element's Chemical Properties and Chemical Bonds • Chemical behavior of an atom is determined by the electron configuration of the outermost electron energy level. • Electron configuration is the distribution of electrons in each atom's energy level. • Electron configuration rules: • Electrons must first occupy lower electron levels before the higher levels can be occupied. • The first energy level of an atom has only two electrons and all higher energy levels have eight electrons. • If an atom doesn't have enough electrons to fill all energy levels, the outermost level will be the only one partially filled with valenceelectrons (electrons in outermost energy level). • Octet Rule - with the exception of the first energy level, the valence level is complete when it contains eight electrons.

14. O2 O=O

15. As a result of incomplete valence levels, atoms fill those levels by interacting with each other forming chemicalbonds (attractions that hold molecules together) • There are three general types of chemical bonds: ionic, covalent, and hydrogen. • Covalent bonds - chemical bond between atoms formed by sharing a pair of electrons. • Covalent bonds may be single, double, or triple. • example --> hydrogen gas • H2 (molecular formula = # and types of elements) • H-H (structural formula = # of elements & bonding)

16. H2 H-H

17. CH4 methane

18. Hydrochloric acid HCl

19. Ionic bonds - bond formed by the attraction after the complete transfer of an electron from a donor atom to an acceptor. • Such a relation ship forms an ion (charged atom). Clinically, we call these electrolytes. • There are two types of ions: • Anion - an atom that has gained one or more electrons from another atom and has become negatively charged. • Cation - an atom that has lost one or more electrons and has become more positively charged.

21. Figure 2.5: Formation of an ionic bond, p. 33. – + Na Cl Na Cl Sodium ion (Na+) Chloride ion (Cl–) Sodium atom (Na) (11p+; 12n0; 11e–) Chlorine atom (Cl) (17p+; 18n0; 17e–) Sodium chloride (NaCl) (a) CI– Na+ (b)

22. Figure 2.12: Dissociation of a salt in water, p. 40. Ions in solution Salt crystal Na+ Na+ Cl– Cl– + H Water molecule O + – H

23. Hydrogen bonds – weak bonds formed between a slightly positive hydrogen and a slightly negative atom.

24. IV. Chemical Reactions • Chemical Equation: Reactant + Reactant ---------- Product(s) • May be reversible • Tends toward equilibrium • Types of Reactions: • Synthesis reactions (A + B --> AB), usually anabolic, requires energy (endergonic) to build compounds. • Decomposition reaction (AB --> A + B), usually catabolic, releases energy (exergonic) to break down compounds. • Exchange/displacement reaction (AB + CD ---> AD + CB), may or may not require\release energy. • Redox reactions - compounds may gain or lose electrons: • oxidized - reactant loses an electron • reduced - reactant gains an electron • Chemical reactions are effected by particle size, temperature, concentration, catalysts, etc...

26. V. Inorganic and Organic Compounds Inorganic compounds • Compounds that contain no carbon or if containing carbon, may also contain elements other than HNOPS • Examples: water, salts, acids, and bases • Water and Its Properties: • High heat capacity - absorb/release large amounts of heat energy without changing in temperature itself. • High heat of vaporization - heat energy to cause transformation (disrupt hydrogen bonds) of water from liquid to gas. • Polarity - unequal distribution of electrons causing slightly positive hydrogens and slightly negative oxygens. • Solvent - water dissolves solutes (therefore compounds are dissociated in water). • Reactant - involved in hydrolysis reactions and dehydration synthesis reactions. • Cushion/shock absorber (e.g. joints and cerebral spinal fluid)

27. Water H2O

28. Salts are ionic compounds consisting of cations other than H+. The dissociation of salts with water forms electrolytes which are ions that conduct electrical current in solution. • Acids and bases... • Acids - hydrogen ion (H+ = proton) donors • Bases - hydrogen acceptors • pH - measure of protons in solution. (scale 0.0-14.0) • Neutralization - reacting acids with bases yielding a water and a salt.

29. Figure 2.12: Dissociation of a salt in water, p. 40. Ions in solution Salt crystal Na+ Na+ Cl– Cl– + H Water molecule O + – H

30. Figure 2.13: The pH scale and pH values of representative substances, p. 42. Concentration in moles/liter [OH–] [H+] pH Examples 0 100 10–14 1 10–1 10–13 2 Lemon juice; gastric juice (pH 2) 10–2 10–12 3 Grapefruit juice (pH 3) Sauerkraut (pH 3.5) 10–11 10–3 Increasing acidity 10–10 4 Tomato juice (pH 4.2) 10–4 5 Coffee (pH 5.0) 10–9 10–5 6 Urine (pH 5–8) Saliva; milk (pH 6.5) 10–8 10–6 7 Distilled water (pH 7) Human blood; semen (pH 7.4) 10–7 10–7 Neutral [H+] = [OH–] 8 Egg white (pH 8) Seawater (pH 8.4) 10–6 10–8 9 10–5 10–9 10 10–4 10–10 Milk of magnesia (pH 10.5) Increasing alkalinity (basicity) 11 10–3 10–11 Household ammonia (pH 11.5–11.9) 12 Household bleach (pH 12) 10–2 10–12 13 10–1 10–13 Oven cleaner (pH 13.5) 14 100 10–14

31. Organic compounds • Compounds containing carbon but may also contain hydrogen and oxygen • Biologically organic compounds may contain (in addition to C,H,O) nitrogen, phosphorus, and sulfur. • Types of organic compounds (Biological): • Carbohydrates - composed of units called saccharides • Lipids - composed of units called fatty acids • Proteins - composed of units called amino acids • Nucleic acids - composed of units called nucleotides

32. A. Carbohydrates • Compounds containing carbon, hydrogen, and oxygen in exact ratios (CnH2nOn) • Carbs are divided into two classes called simple sugars (monosaccharides and disaccharides) and complex sugars (oligosaccharides and polysaccharides). • Monosaccharides - one saccharide, made up of 5 (pentose) or 6 (hexose) carbons. • Ribose (pentose) is component of RNA and DNA • Glucose • Fructose • Galactose • all hexoses are biologically important in the production of energy.

34. Disaccharides - two saccharides formed from a synthesis (dehydration/synthesis) reaction. egs. sucrose (glu + Fru), lactose (glu + gala), and maltose (glu + glu). • Polysaccharides - starches (in plants) and glycogen (in animals), both composed of many glucoses. • Carbohydrates provide cellular fuel; glucose is oxidized in body cells and bond energy released during oxidation is transferred and trapped in the bonds of ATP molecules (adenosine triphosphate). ATP is then used in subsequent endergonic (energy requiring reactions).

35. Figure 2.14a-b: Carbohydrate molecules, p. 45. CH2OH CH2OH HOCH2 HOCH2 HOCH2 O O O O O H HO OH OH OH H H H H H HO H H H H H OH H OH OH HO CH2OH H H H H H H HO OH H H OH OH H OH OH H OH Glucose Fructose Galactose Deoxyribose Ribose (a) Monosaccharides CH2OH CH2OH Dehydration synthesis HOCH2 HOCH2 H2O O O O O H H H H H H H H + H HO H H HO OH H OH O OH HO HO CH2OH HO CH2OH Hydrolysis H2O OH OH H H OH H H OH Glucose Fructose Sucrose HOCH2 HOCH2 HOCH2 HOCH2 H H O O H O O H OH HO OH H H H H O OH H OH H OH H OH H O H HO H H H H OH H OH H OH OH H Glucose Glucose Galactose Glucose Maltose Lactose (b) Disaccharides

36. B. Lipids • Lipids are composed of fatty acids and glycerol. • Fatty acids are compounds containing long chains of carbons and glycerol is a compound containing a small chain of three carbons • Lipids are divided into three classes: Triglycerides, Phospholipids, and Sterols. • Triglycerides - considered the most usable form of energy in the body and is composed of three fatty acids bound to one glycerol by dehydration synthesis. Triglycerides may be saturated or unsaturated. • Phospholipids - component of cell membranes and is composed of one glycerol, two fatty acids. and a phosphate. • Sterols (steroids) - isoprene units (rings of carbon) egs. cholesterol and sex hormones.

37. C. Proteins • Proteins are composed of long chains of amino acids. • Peptide - short chain of amino acids (10-20?) • Polypeptide - long chain of amino acids • Structural Levels: • Primary = sequence of amino acids • Secondary = coiling of primary due to hydrogen bonding • Tertiary = folding of secondary due to hydrogen and sulfur bonds • Quaternary = many tertiary proteins bonded together

38. Figure 2.17: Amino acids are linked together by dehydration synthesis, p. 49. Peptide bond Dehydration synthesis H R O H R O H R O H R O H2O N C C + N C C N C C N C C OH H OH H OH H H H H H Hydrolysis H2O Amino acid Amino acid Dipeptide

39. Biological structures: Fibrous and Globular • Fibrous - strand-like appearance, mostly secondary structure, and referred to as structural proteins • Structural/mechanical - collagen, keratin, and elastin • Movement - actin and myosin in muscle • Globular - compact spherical tertiary proteins referred to as functional proteins • Functional proteins may denature • Example of globular proteins are: • catalysts - enzymes • transport - hemoglobin • pH regulation - plasma proteins • metabolism regulation - peptide and protein hormones • body defense - antibodies

40. Figure 2.18a,c: Levels of protein structure, p. 51. H C R C O O H N C H R C O O H N (c) Secondary structure (b-pleated sheet) C H R (a) Primary structure (polypeptide strand)

41. Enzymes - globular proteins that act as biological catalysts of reactions and are made up of protein and a cofactor/coenzyme (helpers of enzymes). • Mechanism of enzyme action: • Enzyme-substrate complex formation --> enzyme binds substance (substrate) on which it acts to a special site (active site) on the enzyme. • Enzyme-substrate complex undergoes an internal rearrangement that forms a product. • Enzyme releases the product of the reaction and now can catalyze another reaction.

42. Figure 2.18b,d,e: Levels of protein structure, p. 51. Heme group -helix (d) Tertiary structure (e) Quaternary structure (hemoglobin molecule) (b) Secondary structure (-helix)

43. Enzymes lower a reactions activation energy which is the energy required by compounds in order to react. Activation energy Activation energy Energy released by reaction Energy released by reaction Energy Energy (b) Enzyme-catalyzed reaction (a) Noncatalyzed reaction

44. D. Nucleic Acids • Nucleic Acids (DNA and RNA) are composed of nucleotides • Nucleotide are the basic building blocks of our genetic information (chromosomes) • Each nucleotide contains three components: • Pentose sugar (ribose) • Phosphate group • Nitrogenous base (adenine, guanine, cytosine, thymine, uracil)

45. Figure 2.22a-b: Structure of DNA, p. 56. Phosphate Sugar Adenine (A) Thymine (T) Sugar Phosphate O– H CH3 H OH N N H H O H P O CH2 O O H O– N H H H H N H N N O– H H O H N P H2C O O O H OH H O– Adenine nucleotide Thymine nucleotide A T C G A Sugar-phosphate backbone A Key: G Thymine (T) G Adenine (A) T A Cytosine (C) Guanine (G) A Deoxyribose sugar G C G C Phosphate T A Hydrogen bond