1. Human Anatomy & Physiology, Sixth Edition Elaine N. Marieb 2 Chemistry Comes Alive Part A

2. Matter  Has mass & takes up space  States of  Solid – definite shape & volume  Liquid –definite volume, changeable shape  Gas – changeable shape & volume

3. Energy  Capacity to do work  Types of energy  Kinetic – motion/action  Potential – position; stored energy

4. Forms of Energy  Chemical – atomic bonds  Electrical – movement of charged particles  Mechanical – moving matter  Radiant – energy traveling in waves  Conversion between forms

5. Composition of Matter  Subatomic particles  Electrons (e - )  Neutrons  Protons  Atoms  Unique arrangements of subatomic particles  Can’t break down chemically  Elements  Matter of a single type of atom  Atomic symbols

6. Properties of Elements  Periodic table  Elements grouped by properties  Unique physical & chemical properties  Physical properties –detected by senses  Chemical properties –way atoms interact

7. Elemental Composition of the Human Body  Major Constituents  O ~ 65%  C ~ 18.5%  H ~ 9.5%  N ~ 3.2 %  Remaining ~ 3.8%  Ca, P, K, S, Na, Cl, Mg, I, Fe  Trace elements  Zn, Mn, Cu

8. Atomic Structure  Nucleus  Neutrons  no charge  mass = 1 atomic mass unit (amu)  Protons  +1 charge  mass = 1 amu  Electron orbitals  Electrons  orbit nucleus in energy levels (orbitals)  -1 charge  mass = 0.0005 amu

9. Atomic Models Planetary Model Electrons move around the nucleus in fixed, circular orbits Orbital Model Regions around the nucleus in which electrons are most likely to be found

10. Characteristics of Atoms & Elements  Atomic number –  # protons  Atomic mass –  # protons + # neutrons  Isotope –  neutron # can vary  different # of neutrons  Atomic weight – average mass of all isotopes  11 C, 12 C, 13 C, 14 C or 235 U, 238 U  Radioisotopes – atoms that undergo spontaneous decay called radioactivity  131 I, 99m Tc, 60 Co, 14 C, or 235 U

11. Subatomic Configuration of Elements Figure 2.2

12. Isotopes of Elements Figure 2.3

13. Chemically Inert Elements  Inert elements have full outer e - orbitals  Full = 8 e - (except for He) Figure 2.4a

14. Chemically Reactive Elements  Reactive elements have unfilled outer orbitals Figure 2.4b

15. Molecules & Compounds  Molecule – ≥ 2 atoms bonded  Compound – ≥ 2 different kinds of atoms bonded  ALL Compounds Are Molecules  BUT  All Molecules NOT Compounds

16. Chemical Bonds  Chemical bonds formed by e - in outer orbitals (valence shell)  The Octet rule  Atoms interact to have 8 e - s in valence shells  (except H only 2 e - s)  lose, gain or share  Types of Bonds  Ionic – loss or gain e - s  Covalent – sharing e - s

17. Ionic Bonds  Ions  charged atoms  Anions - charge = gained e -  Cations + charge = lost e -  Example: NaCl (sodium chloride)

18. Formation of an Ionic Bond  Ionic compounds form a crystal structures

19. Covalent Bonds  Sharing e - s fill outer orbitals Single bond each atom donates 1 e -

20. Double & Triple Covalent Bonds  Atoms share 2 or 3 e - s

21. Polar & Nonpolar Bonds  Polar bonds  Unequal e - sharing  Atoms w/ 6 -7 valence shell e - s = electronegative  Atoms w/ 1-2 valence shell e - s = electropositive  Electronegative & electropositive atoms form ionic bonds  Nonpolar bonds  Equal sharing of e - s  Atoms with 3-5 valence electrons form covalent bonds

22. Comparison of Ionic, Polar Covalent, & Nonpolar Covalent Bonds

23. Hydrogen Bonds  Very important type of bond for life functions  Allows reversible interactions between molecules  Due to unequal sharing of H’s electron with N or O  Responsible for properties of H 2 O  Very important bonds between large macromolecules (ie proteins & nucleic acids)

24. Hydrogen Bonds in H 2 O Figure 2.9

25. Properties of Water  High heat capacity – absorbs & releases large amounts of heat before changing temperature  High heat of vaporization – changing from a liquid to a gas requires large amounts of heat  Polar solvent properties – dissolves ionic substances, forms hydration layers around large charged molecules, & serves as the body’s major transport medium  Reactivity – is an important part of hydrolysis & dehydration synthesis reactions

26. Mixtures & Solutions  Mixtures – two or more components physically intermixed but not chemically bonded  Solutions – homogeneous mixtures of components  Solvent – substance present in greatest amount  Solute – substance(s) present in smaller amounts  Colloids - heterogeneous mixtures whose solutes do not settle out  Suspensions - heterogeneous mixtures with visible solutes that settle out

27. Concept of Concentration  Concentration refers to the amount of a substance in a given volume  A critical concept to master

28. Units of Concentration  Percent – parts per hundred  PPM – parts per million  PPB – parts per billion  Molarity – moles per liter  Molality – moles particles per kg solvent  Mass/volume  g/ml (gram per milliliter)  mg/ml (milligram per milliliter )   g/ml (microgram per milliliter)  Mass/Mass  mg/kg body mass   g/kg body mass

29. Molecular Weight, Moles & Molarity  Molecular weight  The mass of a mole of atoms or molecules  Has units of g/mole  = atomic weight of an atom in grams  1 mole of C = 12 g OR C is 12g/mole  Molecule’s molecular weight = sum of the atomic weights of its atoms  1 mole of H 2 O = 1g + 1g + 16g = 18 g  MW of H 2 O = 18g/mole

30. Molecular Weight, Moles & Molarity  What’s a mole??  The number of molecules in the gram molecular weight of that molecule  Always 6.02 x 10 23  This magical number is Avogadro’s Number  MW of C = 12  there are 12g C/mole C  there are 6.02 x 10 23 C atoms in 12g of C  MW of H 2 O = 18  there are 18g H 2 O/mole H 2 O  there are 6.02 x 10 23 H 2 O molecules in 18g of H 2 O

31. Molarity – The Standard of Chemical Concentration Terms  Molarity = moles of solute per liter (L) of solvent  Abbreviated with M  A 5M NaCl solution contains 5 moles of NaCl molecules per liter of solution  So 1 L of a 5 M NaCl solution contains 3.01 x 10 24 molecules of NaCl

32. Calculating Molarity  Chemical formula for glucose is C 6 H 12 O 6  MW = 180 g/mole  What is concentration of glucose in a can of coke?  42g of glucose/355ml H 2 O  How many moles of glucose?  42g / 180g/mole = 0.233 moles  How many liters of coke?  355 ml / 1000ml/L = 0.355 L  How many moles/liter?  0.233 moles/0.355 L = 0.656 M

33. Examples using Molarity  Ion concentrations in cells & body fluids  [Na] = 0.15M outside cells & 0.015M inside cells  [K] = 0.005M outside cells & 0.15M inside cells  Molar is often converted to millimolar (mM)  M = moles/L  mM = millimoles/L  0.15M = 150mM  Simply multiply M by 1000 to convert M to mM

34. Mass/Volume & Percent Concentration Terms  Many molecules in the body are measured in mass/volume  Normal glucose = 100mg/dL  or 100mg/100ml  or 1mg/ml or 1g/L  or 0.001g/ml  or 0.1g/100ml  Cholesterol should be below 200mg/dL  or 0.2g/dL  or 0.2g/100ml  Many substances are described in percentages  Percentage is g/100g or g/100ml  Blood [glucose] would = 0.1%  Blood [cholesterol] would = 0.2%

35. Chemical Reactions  Forming or breaking chemical bonds  Chemical equations show:  Reactants & products & their relative amounts

36. Patterns of Chemical Reactions  Combination reactions: Synthesis reactions which always involve bond formation A + B  AB  Decomposition reactions: Molecules are broken down into smaller molecules AB  A + B  Exchange reactions: Bonds are both made & broken AB + C  AC + B

37. Oxidation-Reduction (Redox) Reactions  Reactants losing electrons are electron donors & are oxidized  Reactants taking up electrons are electron acceptors & become reduced C 6 H 12 O 6 + 6O 2 6CO 2 + 6H 2 O glucose carbon dioxide

38. Energy Flow in Chemical Reactions  Exergonic reactions – reactions that release energy  Endergonic reactions – reactions whose products contain more potential energy than did its reactants

39. Rates & Reversibility of Chemical Reactions  Chemical reactions proceed with measurable rates  All reactions are theoretically reversible  Equilibrium (dynamic)  Forward & reverse reactions proceed at same rate A + B AB

40. Factors Influencing Rate of Chemical Reactions  Temperature  higher temperatures increase rates  Concentration  higher [reactant] increases rates  Catalysts  Molecules that increase reaction rates  Enzymes  Biological catalysts with high specificities