Chemistry Comes Alive

1. Chemistry Comes Alive

2. The Chemistry of Life • Atoms, Ions and Molecules • Water and Mixtures • Energy and Chemical Reactions • Organic compounds

3. Matter • The “stuff” of the universe • Anything that has mass and takes up space • States of matter • Solid – has definite shape and volume • Liquid – has definite volume, changeable shape • Gas – has changeable shape and volume

4. The Chemical Elements • Element • simplest form of matter with unique chemical properties • Each element has unique physical and chemical properties • Physical properties – those detected with our senses • Chemical properties – pertain to the way atoms interact with one another

5. Major Elements of the Human Body • 98.5% of body weight consists of • Oxygen (O) • Carbon (C) • Hydrogen (H) • Nitrogen (N)

6. Lesser and Trace Elements of the Human Body • Lesser elements make up 3.9% of the body and include: • Calcium (Ca), phosphorus (P), potassium (K), sulfur (S), sodium (Na), chlorine (Cl), magnesium (Mg), iodine (I), and iron (Fe) • Trace elements make up less than 0.01% of the body • They are required in minute amounts, and are found as part of enzymes

8. Periodic Table of the Elements • Atomic number of each element • number of protons in its nucleus • Periodic table • letter symbols of elements arranged by atomic number http://pearl1.lanl.gov/periodic/default.htm

9. Atomic Structure • Nucleus - center of atom contains • protons: positive charge, mass of 1 amu • neutrons: neutral charge, mass of 1 amu • atomic mass = total # of protons + neutrons • Electron shells • electrons: negative charge • # of electrons = # of protons, atoms have neutral charge • electrons further from nucleus have higher energy • valence electrons are in the outermost shell • interact with other atoms • determine chemical behavior • octet rule - atoms react to obtain a stable number of 8 valence electrons

10. Bohr Planetary Model of an Atom

11. Models of Some Elements p+ represents protons, no represents neutrons

12. Isotopes and Radioactivity • Isotopes • elements that differ in the number of neutrons • 1H, 2H, 3H • extra neutrons result in increased atomic weight • “heavy water” • have no change in chemical behavior • same valence electrons • Atomic weight • Average atomic mass of the mixture of isotopes of an element found in a sample

13. Isotopes of Hydrogen • radioisotopes decay to stable isotopes releasing radiation • radioisotopes decay to stable isotopes releasing radiation Figure 2.3

14. Radioisotopes and Radioactivity • Isotopes • same chemical behavior, differ in physical behavior • Radioisotopes • unstable isotopes • Radioactivity • radioisotopes decay to stable isotopes releasing radiation Marie Curie

15. Ionizing Radiation • High energy • Ejects electrons from atoms • Destroys molecules and produces free radicals • sources include: • UV light, X rays, nuclear decay (, , ) •  particle - • 2 protons + 2 neutrons can’t penetrate skin •  particle - • free electron - penetrates skin a few millimeters •  particle - • high energy, penetrating; very dangerous

16. Ionizing Radiation 2 • Physical half-life • time for 50% of atoms to decay • 90Sr - 28 yr. • 40K - 1.3 billion years • Biological half-life • time for 50% of atoms to disappear from the body • function of decay and physiological clearance • Cesium 137 - physical half-life -- 30 years - biological half-life -- 17 days • Radiation exposure • background radiation • radon gas from decay of uranium in granite • cosmic rays

17. Molecules and Chemical Bonds • Molecules • two or more atoms of same element covalently bonded • Compounds • two or more atoms of different elements covalently bonded • Molecular formula • itemizes each element present and its quantity • Structural formula • shows arrangement of atoms • needed to show structural isomers

18. Concentration of Solutions • Percent, or parts per 100 parts • Molarity, or moles per liter (M) • Mole – Avagadro’s number of molecules • 6.02 X 1023 • A mole of an element or compound is equal to its atomic or molecular weight (sum of atomic weights) in grams

19. Types of Chemical Bonds • Ionic • Covalent • Hydrogen

20. Chemical Bonds • Electron shells, or energy levels, surround the nucleus of an atom • Bonds are formed using the electrons in the outermost energy level • Valence shell – outermost energy level containing chemically active electrons • Octet rule – except for the first shell which is full with two electrons, atoms interact in a manner to have eight electrons in their valence shell

21. Chemically Inert Elements • Inert elements have their outermost energy level fully occupied by electrons Figure 2.4a

22. Chemically Reactive Elements • Reactive elements do not have their outermost energy level fully occupied by electrons Figure 2.4b

23. Ions • Ions - carry a charge, unequal numbers of protons and electrons • Ionization - transfer of electrons from one atom to another ( stability of valence shell)

24. Anions and Cations • Anion - atom gained electron, net negative charge • Cation - atom lost an electron, net positive charge

25. Ionic Bonds • Attraction of oppositely charged ions to each other forms an ionic bond - no sharing of electrons • Ionic bonds are weak and dissociate in water • These compounds tend to form crystals...

26. Formation of an Ionic Bond Figure 2.5a

27. Covalent Bonds • Formed by sharing valence electrons • Types of covalent bonds • single covalent bond • double covalent bond • Triple covalent bond

28. Polar and Nonpolar Molecules • Electrons shared equally between atoms produce nonpolar molecules • Unequal sharing of electrons produces polar molecules • Atoms with six or seven valence shell electrons are electronegative • Atoms with one or two valence shell electrons are electropositive

29. Comparison of Ionic, Polar Covalent, and Nonpolar Covalent Bonds Figure 2.8

30. Hydrogen Bonds • Weakest of the bonds • Attraction between polar molecules – no sharing of electrons • Greatest physiological importance • properties of water • shapes of complex molecules • proteins, DNA

31. Hydrogen Bonding in Water

32. The Chemistry of Life • Atoms, Ions and Molecules • Water and Mixtures • Energy and Chemical Reactions • Organic compounds

33. Adhesion and Cohesion • Adhesion - attraction between one substance and another substance • Cohesion - attraction between one substance and itself • water is very cohesive due to hydrogen bonds • Surface tension • elastic surface film caused by the attraction of molecules at the surface from those below

34. Thermal Stability of Water • Heat capacity: amount of heat required to raise the temperature of 1g of a substance by 1°C • Calorie: amount of heat required to raise the temperature of 1g of water by 1°C • Water stabilizes internal temperature of the body • high heat capacity • its hydrogen bonds inhibit increased temperature (molecular motion) caused by increased heat • effective coolant • 1 ml of perspiration removes 500 calories from the body

35. Properties of Water • Reactivity – is an important part of hydrolysis and dehydration synthesis reactions • Cushioning – resilient cushion around certain body organs InterActive Physiology®: Fluid, Electrolyte, and Acid/Base Balance: Introduction to Body Fluids PLAY

36. Mixtures and Solutions • Mixtures – two or more components physically intermixed (not chemically bonded) • Solutions – homogeneous mixtures of components • Solvent – substance present in greatest amount • Solute – substance(s) present in smaller amounts

37. Solvency • Solvency - ability to dissolve matter • Hydrophilic - charged substances that dissolve easily in water • Hydrophobic - neutral substances that do not easily dissolve in water • Water is the universal solvent, important for metabolic reactions and transport of substances

38. Water as a Solvent • Water molecules overpower the ionic bond above between Na+Cl- by forming hydration spheres. • Note orientation of water molecules: negative pole faces Na+, positive pole faces Cl-

39. Mixtures Substances that are physically blended but not chemically combined • Solutions • Colloids • Suspensions

40. Solutions • Solute < 1nm • pass through membranes • Transparent • e.g. copper sulfate solution

41. Colloids • Particles 1 to 100nm • to large to pass through membranes • Cloudy • e.g. milk protein

42. Suspensions • Particles >100nm • Cloudy or opaque • Separate on standing • e.g. blood cells

43. Measures of Concentration • Weight per Volume • weight of solute in a given volume of solution • e.g. IV saline contains 8.5 g/L NaCl • Percentages • weight or volume of solute in solution • e.g. IV D5W (5% w/v dextrose in distilled water) • 5 grams of dextrose in add 100ml water • Molarity • number of moles of solute/liter in solution • physiologic effects of a chemical based on the number of molecules in solution

44. Salts • Inorganic compounds • Contain cations other than H+ and anions other than OH– • Are electrolytes; they conduct electrical currents

45. Acids and Bases • Acids release H+ and are therefore proton donors HCl  H+ + Cl – • Bases release OH– and are proton acceptors NaOH  Na+ + OH–

46. Acid-Base Concentration (pH) • Acidic solutions have higher H+ concentration and therefore a lower pH • Alkaline solutions have lower H+ concentration and therefore a higher pH • Neutral solutions have equal H+ and OH– concentrations

47. pH • pH - based on the molarity of H+ on a logarithmic scale • pH = -log [H+] • for molarity of H+ = 100,10-1,10-2,etc. • pH = - log [100] = 0, - log [10-1] = 1, etc. • a change of one number on the pH scale therefore represents a 10 fold change in H+ concentration • Our body uses buffers to resist any change in pH

48. pH Scale

49. Buffers • Systems that resist abrupt and large swings in the pH of body fluids • Carbonic acid-bicarbonate system • Carbonic acid dissociates, reversibly releasing bicarbonate ions and protons • The chemical equilibrium between carbonic acid and bicarbonate resists pH changes in the blood InterActive Physiology®: Fluid, Electrolyte, and Acid/Base Balance: Acid/Base Homeostasis PLAY

50. The Chemistry of Life • Atoms, Ions and Molecules • Water and Mixtures • Energy and Chemical Reactions • Organic compounds