The Chemistry of Biology

1. The Chemistry of Biology Inorganic Chemistry Organic Chemistry (Biochemistry) Analytical Chemistry

2. Chemistry: • The science that studies the structure and behavior of matter

3. Connections to LIFE: • Anything that takes up space and has mass is matter • Organisms take up space and have mass • Organisms are composed of matter • Matter exists in many forms—Rocks, metals, oils, gases, and humans

4. Essential Elements • Similar among organisms elements: C, H, N, O, P, S (+ trace) compounds: carbohydrates lipids proteins nucleic acids

5. Metabolic Processes/Chemical Reactions • Similar among all organisms: • Cellular respiration • Photosynthesis • Protein synthesis • DNA transcription and translation • Cellular transport mechanisms

6. Rocks and YOU! • The physical and chemical principles that govern nonliving things (remember physical science?) • Rocks • Also govern living things • YOU!

7. Types of Chemistry • Inorganic Chemistry: • studies non-C (carbon)-based chemicals • chemicals are small in size • most are found in non-living things *Examples: H2O, NaCl, CO2, CaCO3, H2SO4, HCl, HNO2

8. Organic Chemistry • studies chemicals based on C (carbon) • We call it Biochemistry • compounds are large in size • found in living things *4 classes of organic compounds 1. Carbohydrates 2. Lipids 3. Proteins 4. Nucleic Acids

9. Levels of Study - Chemistry • 1. Atoms: smallest unit of matter that still retain the properties of an element • 2. Element: group of the same type of atom • find information about elements on the periodic table • We symbolize atoms with the same abbreviation used for the element • Atoms and Elements are pure substances that cannot be broken down into smaller units

10. Elements important to living systems: C H N O P S …AND IN TINY AMOUNTS: Ca Na Cl K Fe I Mg

11. 3. Compounds: 2 or more elements bonded together • produce matter (new substance) with new properties (water, Table Salt) • 4. Mixtures: 2 or more elements thrown together, each retaining their elemental properties (Sand, Salsa)

12. Structure of the Atom • 3 particles: 1. Proton 2. Neutron 3. Electron • 2 regions: 1. Nucleus 2. Energy levels or Electron Cloud

15. Characteristics of Atomic Particles • 1. Proton • In the nucleus • Positive (+) charge • Mass of 1 amu • The number determines the atomic# of the atom/element • If this number changes, it’s a different atom/element!

16. 2. Neutron • In the nucleus • No charge (neutral) • Mass of 1 amu • Atomic# - Mass# • Altering the # of neutrons results in a new isotope and a change in the atomic mass

17. 3. Electron • Around the outside of the atom (e- cloud) • Negative (-) charge • Almost no mass! • The number and location determine the reactivity of an atom with other atoms • A change in number creates an ion or ionic compound

18. Neutral atom: p+ # = e- # • Ions • atom that has more p+ than e- (+ ion or cation) OR • Atom that has more e- than p+ (- ion or anion) • Isotope – atoms with the same number of protons but differing number of neutrons • Causes a change in the atomic mass number

19. Radioactive Isotopes • Defined: an isotope in which the nucleus is unstable and decays spontaneously giving off particles and/or energy to become stable • When the decay leads to a change in the number of protons, it transforms the atom to an atom of a different element • Types of Particles Released • Alpha - 2 protons + 2 neutrons (helium nucleus) • Beta – neutron decomposes to form p+ and e- • Gamma Rays – high energy radiation

20. Uses of Radioactive Isotopes • Fossil Dating • Tracers – Follow atoms through metabolism/Monitor bio processes • Diagnostic Tool – Medical Diagnosis • Used with imaging processes • Cancer Detection • Cancer Therapy

21. Diagraming an Atom • Start with the obvious! • Atomic number: from PT • Atomic number = # of p+ • # of p+ = # of e- • Atomic mass: from PT • # of p+ plus # of n◦ • Number of energy levels (Period # from PT) • Number of valence electrons (Group # from PT)

22. p+ and n go in the nucleus • e- placed in energy levels around the nucleus • The placement of electrons in the energy levels determines how and what atom it will react with to form bonds • 1st holds 2e-; 2nd holds 8e-; 3rd holds 8e-

23. The rule of 8 or octet rule • to be “stable” an atoms “wants” 8 electrons in its outermost energy level • An atom will gain, lose, or share electrons to become stable • In the process of gaining, losing, or sharing electrons, bonds (then compounds) are formed

24. Combining Atoms • 1. Mixing two or more atoms together - the mixture retains the properties of the mixed atoms. • 2. Bonding two or more atoms together - forms new compound/molecule with new properties

25. Pure or homogeneous substances – cannot be separated easily 1. Element 2. Compound • Heterogeneous substances – easily separated 3. Mixture

26. Compounds • Bonded elements are called compounds or molecules (if C is involved) 3 types of bonds 1. Ionic 2. Covalent 3. Hydrogen - an important covalent bond in living things

27. Ionic Bonds • Created when one element loses an e- and one element gains an e-; they are then attracted to each other by the opposite charges • Fairly strong bond • Important in salt compounds and other inorganic compounds • Form ions (charged particles) • K+, Cl-, Ca+, Na+

28. Covalent Bonds • Bonds between two atoms that share one or more electrons • The two shared electrons move about the nuclei of both atoms • Atoms in a molecule attract shared electrons to varying degrees – electronegativity • Non-polar (similar electronegativity/e- shared equally) • Polar (varried electronegativity/e- not shared equally) • Very strong bond • Very important in biochemistry and molecules making up living things

29. Hydrogen Bonds • Unique bond between a covalently bonded H atom and electronegative atom of another molecule. • Due to partial + charge of H • Hydrogen can donate 1 e- • Hydrogen ion H+ is called a proton donor • Weak bond • In living cells, the electronegative partners are usually O or N atoms

30. Mixtures - HomeworkDefine and list 3 examples of each • 1. solution • 2. suspension • 3. colloid • 4. acid • 5. base • 6. salt

31. Types of Mixtures • 1. Solution - one or more substances (solutes) are distributed evenly in another (solvent) • The particles of solute in solution cannot be seen with the naked eye • Examples of solvents: • Water – universal solvent • Alcohols, oils, acids, organic benzene

32. Examples of solutions: a. Koolaide b. Tea/coffee c. Pop d. Sweat e. Tears f. Urine

33. Types of Mixtures • 2. Suspension - a mixture with large size solutes that settle out over time • Examples: Salad dressings, any shakable, blood, barf/vomit, dust particles in air, mud, four in water

34. Types of Mixtures • 3. Colloid - a mixture of different size particles in a thick solvent so particles do not separate • Examples: Yogurt, pudding, mayo, cytoplasm, saliva, mucous,

35. Types of Mixtures • 4. Acids - a solution with an excessive amount of H ions H+ • Examples: HCl, H₂CO₃, H₂SO₄ • 5. Bases – a solution with an excessive amount of OH ions OH- • Examples: NaOH, NH₃+ (ammonia)

36. Types of Mixtures • 6. Salts - a mixture that will produce ions other than H+ and OH- • Examples: NaCl, CaCO₃, KCl • Environment affects the strength of ionic bonds – dissolving salt crystals in water weakens the ionic bond • Most drugs are manufactured as salts because they are quite stable when dry, but can dissociate (come apart) easily in water.

37. Chemical Reactions • Defined: the processes that • 1. make • 2. break up • 3. rearrange compounds • New structures means new functions • In organisms chemical reactions occur inside cells • Metabolism – sum of all chemical reactions that occur within an organism

38. Chemical Reactions and Bonds • Bonds are made • Bonds are broken • Bonds are rearranged - to produce new compounds • Replacement bonds - Atoms replace other atoms in a reaction

39. Energy and Reactions • Energy is needed for rxn to occur-endothermic or endergonic reaction • If energy of reactants is less than energy of products, energy is taken in (absorbed) • i.e. cooking food • Energy is given off during rxn-exothermic or exergonic reaction • If energy of reactants is greater than energy of products, energy is given out (released) during the reaction • i.e. light, gas, heat

40. Types of Reactions • Anabolic or Synthesis Rxn • Something is made – bonds are formed • Example: dehydration/condensation rxn – water is removed • Small sugar molecules join together to make disaccharides and water • Glucose and fructose  sucrose + water • C₆H₁₂O₆ + C₆H₁₂O₆  C₁₂H₂₂O₁₁ + H₂O • Anabolic rxn usually require energy

41. Types of Reactions • Catabolic or degradation Rxn • Something is broken down – bonds are broken • Example: hydrolysis rxn – water is added • Reverse of dehydration reaction • Water is recombined with the two hydroxyl groups and the disaccharide reverts back to being a monosaccharide • C₆H₁₂O₆ + C₆H₁₂O₆  C₁₂H₂₂O₁₁ + H₂O • Catabolic rxn usually release energy

42. Types of Reactions • Add phosphate to a compound • Phosphorylation – addition of a phosphate (PO₄³⁺) to a protein or other organic molecule • Example: photosynthesis and cell respiration

43. Types of Reactions • Oxidation/reduction rxn • e- are lost from one chemical (oxidation) and picked up by another (reduction) • The reaction is always coupled together • called a redox rxn • Example: photosynthesis and cell respiration

44. The Reaction Specialists • The Enzyme!! – A biological catalyst ~a protein ~needed to reduce temperature of activation in living systems (cold chemistry of life) ~speeds up a rxn, but is unchanged by the rxn ~ reuseable!!! ~needed in all living metabolic rxns

45. Factors that influence rxn rate • Temperature • Heat=faster/Cold=slower • Agitate/Stir • Poison – stops rxn • Heavy metals – stops rxn • Amount of substrate • Amount of reactants • Presence and amounts of enzymes