BIOCHEMISTRY

1. BIOCHEMISTRY The chemical basis of life

2. ATOMS • Basic unit of matter • Two regions • Nucleus • Electron cloud

3. Subatomic Particles • Protons  • Positively charged particles • Located in the nucleus • Neutrons  • Neutral particles • Located in the nucleus • Electrons  • Negatively charged particles • Located in the electron cloud • These are the particles involved when atoms bond with other atoms

4. Electrons orbit the nucleus. • An atom is only about 0.0000000001 meters big. It would take 10 BILLIONS atoms lying side by side to equal 1 meter!

6. Atomic Number • The number of protons distinguishes an atom of one type from another. • All atoms of the same element have the same number of protons. • This unique number is called theatomic number.

7. Atomic Mass • The atomic mass is equal to the total number of protons plus neutrons in an atom. Atomic Mass = P + N

8. Mass Number • The mass number can be written as a superscript above the symbol and the atomic number as a subscript below the symbol.

9. Elements & Isotopes • Elements: Simplest Pure Substance • Elements of Life • 96%  Carbon (C), Hydrogen (H), Oxygen (O), and Nitrogen (N) • 3%  P, S, Ca, K, Na, Mg, Fe, Cl • 1%  other trace elements • Isotopes • Atoms of the same element that contain a different number of neutrons • Radioactive isotopes will breakdown at a specific rate and are used in determining the age of various things (i.e. fossils & rocks)

10. Isotopes • Atoms of the same element can have different numbers of neutrons. • These different forms of the same element are called isotopes. • The atomic mass is the average mass of all the known isotopes of the element.

12. Atoms • All atoms are neutral; they have the same number of electrons as protons. Example: An atom of 42He has an atomic number of 2 and a mass of 4. Therefore, it has 2 protons and 2 neutrons in its nucleus. Since it has 2 positive protons (neutrons are neutral) it must have 2 negative electrons to make the total charge neutral.

13. How are electrons arranged? • Electrons are located in different energy levels. • The farther away from the nucleus the electron is found, the higher the energy. • As electrons move from a lower level to a higher level energy is absorbed. • As electrons move from a higher level to a lower level energy is release in the form of light.

15. Periodic Table Group • The elements are listed in order of increasing atomic number. • By looking at the row (period) number you can determined how many energy levels an atom has. • By looking at the column (group) number you can determine how many electrons are in the outermost level. Period

16. Comparing Atoms

17. How does one kind of atom differ from another? • Number of protons determines an element. • Even if atoms bond or break apart, the number of protons will always be the same.

18. Chemical Compound • Pure substance formed by two or more elements chemically combined. • Ex: water: H2O, sodium chloride NaCl

19. Types of Bonds • Ionic: The TRANSFER of electrons between a metal and a nonmetal • Covalent: The SHARRING of electrons between two or more nonmetals • Two Types: Polar and Nonpolar • Metallic: A “sea of electrons” around two or more metals • Animations: http://ithacasciencezone.com/chemzone/lessons/03bonding/mleebonding/covalent_bonds.htm

20. Bonding by Analogy • Sometimes it helps to think of bonds (which you can't see) in terms of familiar things you can see.  This is called an analogy.  Let's use the natural attraction between dogs and bones as an analogy to the attraction between opposite charges and atomic or intramolecular bonds. http://ithacasciencezone.com/chemzone/lessons/03bonding/dogbonds.htm#Ionic%20Bonding

21. A neutral atom contains an equal number of positive and negative charges.  • In a sense, the atoms fight over the available electrons in much the same way two or more dogs will fight over bones.  • The bone and the electron are very similar.  • The Dog - Bone analogy works quite well for three of the four types of atomic bonds.  (van der Waal's forces are the only one which cannot be represented with this analogy) • http://ithacasciencezone.com/chemzone/lessons/03bonding/dogbonds.htm#Ionic%20Bonding

22. Ionic bonds: One big greedy thief dog! • Covalent bonds:  Dogs of equal strength. • Polar Covalent bonds: Unevenly matched but willing to share. • Metallic bonds:  Mellow dogs with plenty of bones to go around.

23. Ionic Bonds

24. Covalent Bonds • Sharing of electrons • These are stronger bonds than either of the other two types because the electrons are shared. • Your body is based upon carbon bonding. • So the covalent bond is considered the most important bond with regards to life. • Interestingly, Si, just above C in the periodic table, with its covalent bonding, is the basis for the computer industry.

25. Metallic Bonds

26. Van der Waals Bonds • When molecules are close together an attraction can develop between oppositely charged regions of nearby molecules • Example: water molecules

27. Geckos and Van der Waals • Geckos can stick to so many surfaces in a seemingly impossible manner. • Specifically, the tiny hairs on the gecko's feet (called setae) are split at the microscopic level into as many as 1,000 branches. As a result, even though the Van der Waals forces acting on an individual tip is small, the adhesion of a billion or so tips adds up to enough force to let the gecko stick to basically anything.

28. Water (H2O) • Most abundant compound in organisms • Water is polar • Unequally shares electrons between hydrogen & oxygen atoms • Makes it possible for other compounds to dissolve in water

29. Water has Hydrogen Bonds • Adhesion – the attraction of unlike molecules to one another • Cohesion – the attraction of like molecules to one another Cohesion causes water to form drops, surface tension causes them to be nearly spherical, and adhesion keeps the drops in place.

31. Properties of Water • Solutions and suspensions • Mixtures are composed of 2 or more elements or compounds that are physically mixed, but not chemically combined • There are two types of mixtures that can be made with water • Solutions • suspensions • Solutions – made of a solute and a solvent

32. Properties of Water • Solutions • Made of a solute and a solvent • Usually water acts as a solvent • Polar water molecules are able to pull apart the solute to form a solution • Water is known as the “universal solvent” • Many biological fluids are solutions • Solutions are also known as homogeneous mixtures

33. Properties of Water • Suspensions • Particles do not dissolve in H2O and remain suspended • Colloids are a type of suspension • Examples of suspensions: smoke, fog, jello, blood

34. pH scale and buffers • The pH of substances ranges from 0-14 • 7 = neutral • 0 - 6.9 = acid • 7.1 - 14 = base • Buffers • Prevent sharp, sudden changes in pH so that the body can maintain homeostasis • pH of most fluids in the body = 6.5 - 7.5

35. Acidic (acid) 0 - 7 Basic (base) 7 - 14 Neutral

36. Acids, Bases and the pH scale • [H+] = concentration of hydrogen ions • [OH-] = concentration of hydroxide ions • The pH scale measures the concentration of H+ ions (how acidic something is) • Ranges from 0-14 • At 7: H+ ions and OH- ions are equal so it is neutral • 0-7: acidic, has more H+ ions • 7-14: basic, has more OH- ions

37. ACIDS: release H+ when mixed with water • Sour, corrosive • Ex: HCl, H2SO4 • Always have H at front of formula • BASES: release OH- when mixed with water • Bitter, slippery, usually in cleaners • NaOH, CaOH • Always have OH at end of formula

38. What is a buffer? • Buffers are weak acids or bases that react with strong acids and bases to prevent sharp changes in pH • Helps to neutralize • Help to control pH in blood, digestive tract, etc. to maintain homeostasis • Ex: Antacids buffer the stomach from the Hydrochloric Acid (HCl)

39. Organic vs. Inorganic • All compounds can be separated into two groups: • Inorganic • Does not contain carbon • Non-living (never alive) • Examples: Oxygen gas, metals, rocks, water • Organic • Contains carbon • Living (or dead – once was alive) • Examples: wood, grass, diamonds, petroleum

40. Inorganic Compounds • Usually do not contain CARBON (except good old CO2) • WATER-- a very curious material • salts, compounds in our bones, etc. but none as numerous as the ORGANIC compounds in living things

41. Why Carbon Compounds? • Carbon (C) forms strong, stable COVALENT bonds • Carbon forms almost infinite chains when bonded to other C atoms • Chains may form as ring structures with single or double bonds • Ex: Polymerization

42. We eat polymers! • Hey, come on over here and have a big slice of POLYMER pizza. It's not as strange as it sounds. French fries are loaded with a polymer called starch, which your body digests into sugar to use as fuel.

43. Polymerization • Monomers (small) • One unit of a compound • Polymers • Many monomers combine to make a polymer • Macromolecules (huge) • Many large molecules combined • Polymers are everywhere: http://pslc.ws/macrog/paul/

44. Containers • Fast food often comes in boxes made of polystyrene foam. • Cup lids are made of polystyrene, but in plastic form instead of foam. • Napkins are made of paper, which is made from wood pulp, and that wood pulp has an awful lot of the polymer cellulose. • The trays are made of polyethylene. Most of the prizes in the kids' meals are made from polystyrene and polyethylene or polyvinyl chloride.

45. Clothing • The polymers in clothes can be everything from plant materials, to synthetics, to proteins like silk and wool. • Sweaters are also made out of acrylics, like polyacrylonitrile. • Spandex is a special kind of polyurethane that's very stretchy. Spandex is also used in bicycle pants, swim suits, and other items of stretchwear.

46. Carbohydrates • Made of C, H, & O • Functions • Main energy source in organisms • Structural component in plants • Types • Sugars • gives off energy when broken down • Cellulose  twisted chain of sugars, not digestible by humans • Chitin  hard cellulose found in the exoskeletons of invertebrates • Ex. Sucrose, fructose, glucose • Starches • used as a storage molecule for sugars • Many athletes eat these before events • Ex. Bread, rice, pasta, corn

47. Lipids • Made of C, H, O • in the form of glycerol and fatty acid chains • Commonly called fats, oils, & waxes • Functions • Storage of energy • Parts of biological membranes • Water proof coverings • Chemical messengers (steroids) • Insoluble in water • Ex. Lard, butter, oil, hormones, steroids

48. Steriods • Steroids occur in animals in something called hormones. The basis of a steroid molecule is a four-ring structure, one with five carbons and three with six carbons in the rings. • Many body builders and athletes use anabolic steroids to build muscle mass. The steroids make their body want to add more muscle than they normally would be able to. • The body builders wind up stronger and bulkier (but not faster). Never take drugs to enhance your body. Those body builders are actually hurting their bodies. They can't see it because it is slowly destroying their internal organs and not the muscles. When they get older, they can have kidney and liver problems. Some even die!

49. Lipids... • Saturated fats • all Carbons attached by single bonds with the maximum H atoms • meats, dairy • Unsaturated fats • C atoms joined by double bond, not with the maximum H atoms (more double bonds=polyunsaturated) • liquid fats at room temp.--sesame, peanut, canola oils

50. Fats • There are two kinds of fats, saturated and unsaturated. • Unsaturated fats have at least one double bond in one of the fatty acids. • A double bond happens when two electrons are shared or exchanged in a bond. They are much stronger than single bonds. • Saturated fats have no double bonds. Fats have a lot of energy stored up in their molecular bonds. That's why the human body stores fat as an energy source. When it needs extra fuel, your body breaks down the fat and uses the energy. One molecule of sugar only gives a small amount of energy, a fat molecule gives off many times more.