1. Basic Chemistry Chapter 2

2. Chemistry of Life All life processes involve chemical reactions (clinical day) –Ex. Ca ++ in muscle contraction Na +, K + in nerve impulses

3. anything that has mass and takes up space can you think of solids, liquids, and gases that might be found in the body? Matter

4. the capacity to do work No mass & does not take up space Types: –Potential: stored in bonds (chemical) –Kinetic: doing work (electrical, mechanical, radiant) Exergonic & endergonic reactions Energy

5. Composition of Matter 92 naturally occurring elements (112 known, 113- 118 are alleged) Living organisms require about 26 of these elements (table 2.1, p.28) About 96% (by mass) comes from Oxygen (O), Carbon (C), Hydrogen (H), and Nitrogen (N)

6. Atoms smallest complete unit of an element –Composed of dozens of subatomic particles, but we are only concerned with THREE! Subatomic Particle ChargeLocationWhat it tells you! Proton+NucleusIdentity of atom, mass Neutron0NucleusIsotope, mass Electron-Surrounds nucleus Properties of atom, negligible mass

7. Identifying Elements Atomic number equal to the number of protons in an atom (& electrons in neutral atom!) Atomic mass sum of the masses of all the protons & neutrons contained in nucleus

8. Isotopes atoms of same element with a different mass (due to neutrons) but same chemical properties Ex. C-12 and C-14 Radioactive isotopes used in many medical tests to tag biological molecules to be followed or traced i.e. PET scans, I-131 for thyroid activity, destroying localized cancers (Ra, Co, etc.)

9. Ions Charged particles Form ionic bonds Cations (+) Anions (-)

10. Find the Face (in the Beans)

11. Transfer or share electrons in order to fill their valence shell (stability) All atoms want 8 e- in their valence shell (except H & He) Ionic bond – transfer electrons Covalent bond – share electrons –Nonpolar: shares electrons equally –Polar: shares electron unequally Chemical Bonds

12. Chemical Bonding Due to electronegativity –How much an atom in a bond pulls electrons to itself –Ionic: >1.7 –Polar covalent: 0.4-1.7 –Covalent: <0.4

13. Hydrogen Bonding Weak bonds attraction of H to partial negative charge –Example: polar covalent bonds between oxygen and hydrogen

14. represents the numbers and types of atoms in a molecule –Ex. H 2 O, C 6 H 12 O 6 Molecular Formula

15. Chemical Reactions Metabolism= sum of all chemical reactions in the body –Synthesis (anabolism) A + BAB Energy absorbing i.e. growth, repair, protein synthesis –Decomposition (catabolism)AB A + B Energy releasing i.e. digestion of foods, breakdown of glycogen in liver to produce glucose –Single replacementAB + C AC + B –Double replacement AB + CD AD + CB

16. Rate of Chemical Reactions Temperature ( temp increases collisions) Concentration of reactants ( number = faster, more collisions) Particle size (smaller = faster, more collisions) Presence of catalysts –Affect rate of reaction without being changed by reaction –Biological catalysts: enzymes (proteins) –Shape matters! Like a puzzle piece

17. Biochemistry Inorganic compounds: lack carbon (with few exceptions) –Small, simple molecules –Water, salts, many acids & bases Organic compounds: carbon-containing compounds –Large covalently bonded molecules –Carbohydrates, lipids, proteins, nucleic acids

18. Inorganic Compounds Water –High heat capacity Absorbs & releases large amounts of heat Prevents sudden changes in body temperature (homeostasis!) –Polarity/solvent properties “universal solvent” Chemical reactions depend on solvent Transport/exchange medium Lubrication (synovial fluid in joints) –Chemical reactivity (hydrolysis reactions) –Cushioning Protective (CSF, amniotic fluid)

20. Oxygen –used to release energy from glucose Carbon dioxide –waste of metabolic processes Inorganic Compounds

21. Salts –Ionic compound containing cations other than H+ and anions other than OH- –Vital to body functions K+ & Na+ essential for nerve function, Fe2+ is essential for hemoglobin, Cl-, Ca++, Mg++, PO 4 -, CO 3 -, etc. –All salts are electrolytes (substances that conduct an electrical current in solvent) Release ions when dissolved in water –Functions in Table 2.1, page 28 Inorganic Compounds

22. Acids & Bases –Electrolytes Acids –Release H+ ions in solution –“proton donors” –HCL, acetic acid, carbonic acid Bases –Release OH- ions in solution –“proton acceptors” –HCO3- (important base in blood)

23. pH scale measures hydrogen ion concentration –pH 7 = neutral –pH >7= basics (more OH- than H+) –pH <7= acidic (more H+ than OH-) Normal blood pH for humans is 7.35 to 7.45 –If >, alkalosis –If <, acidosis Buffers- maintain pH

24. Organic Compounds

25. sugars, starches, glycogen, cellulose – 2-3% body weight –Plants- starches and cellulose (cannot digest) –Animals- source of energy- stored as glycogen Carbohydrates

26. Monosaccharides: 3 to 7 carbons –Ex. Glucose, fructose, galactose Carbohydrate utilized by the cell Many C 6 H 12 O 6 Carbohydrates

27. Disaccharides: 2 monosaccharides combine by dehydration synthesis (condensation) –Ex. Sucrose Broken apart by hydrolysis (add water) Carbohydrates

28. Polysacchride: 10-100s of monos –Ex. starch Carbohydrates

29. 18-25% in lean adults –Contain C, H, O - neutral –Fats- concentrated energy stored in adipose tissue Lipids

30. Triglycerides: Glycerol + 3 fatty acids Monounsaturated- one double bond Polyunsaturated- more than one double bond Saturated- no double bonds Lipids

31. Phospholipids- polar head and 2 non-polar tails (membrane) Lipids

32. Steroids- cholesterol, sex hormones, cortisol, etc. Lipids

33. 12-18% in lean adults –Structural and physiological enzymes –Made of amino acids (20)- held by peptide bonds –3D shape held by H-bonds (denatured with heat) Proteins

34. –Base + sugar + phosphate –DNA and RNA –ATP- provides energy for the cell Nucleic Acids

35. molecules with the same chemical formula and with the same kinds of bonds between atoms, but in which the atoms are arranged differently. share similar if not identical properties in most chemical contexts. Isomers