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Lecture 3 : How Matter is Organized

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  1. Lecture 3 : How Matter is Organized • Chemistry is the science of the structure and interactions of matter. • all living things consist of matter. • Matter is anything that occupies space. • mass is the amount of matter in any object. • weight is the force of gravity acting on matter. • In outer space, weight is close to zero, but mass remains the same as on Earth.

  2. Chemical Elements • Elements are substances that can not be split into simpler substances by ordinary means. • 92 elements exist in nature • 26 of naturally occurring elements are in the body • represented by chemical symbols ( first 1-2 letters of name )NaNatrium=Sodium • 4 elements form 96 % of the body’s mass • hydrogen, oxygen, carbon and nitrogen • Trace elements are present in tiny amounts • such as copper,selenium & zinc

  3. Structure of Atoms • Atoms are the smallest units of matter that retain the properties of an element • Atoms consist of 3 types of subatomic particles • protons, neutrons and electrons • Nucleus contains protons (p+)&neutrons (neutral charge) • Electrons (e-) surround the nucleus as a cloud (electron shells are designated regions of the cloud)

  4. Electron Shells • Most likely region of the electroncloud in which to find electrons orbiting the nucleus • Each electron shell can hold onlya limited number of electrons • first shell can hold only 2 electrons • 2nd shell can hold 8 electrons • 3rd shell can hold 18 electrons • higher shells (up to 7) hold many more electrons • Number of electrons = number of protons • Each atom is electrically neutral; charge = 0

  5. Atomic Number & Mass Number • Atomic number is number of protons in the nucleus. • Mass number is the sum of its protons and neutrons.

  6. Ions, Molecules, & Compounds • Ions are formed by ionization • an atom that gave up or gained an electron • written with its chemical symbol and (+) or (-) • Molecule • molecules = when 2 of the same atoms share electrons • compound = when 2 or more different atoms share electrons • written as molecular formula showing the number of atoms of each element (H2O)

  7. Chemical Bonds • Bonds hold together the atoms in molecules and compounds • An atom with a full outer electron shell is stable and unlikely to form a bond with another atom • Octet rule states that biologically important elements interact to produce chemically stable arrangements of 8 electrons in the valence shell. • Whether electrons are shared, lost or gained determines the types of bonds formed

  8. Free Radicals • Atom with an unpaired electron in its outmost shell • Unstable and highly reactive • Can become stable • by giving up electron • taking one off another molecule (breaking apart important body molecules)

  9. Free Radicals & Your Health • Produced in your body by absorption of energy in ultraviolet light in sunlight, x-rays, by breakdown of harmful substances, & during normal metabolic reactions, fried foods, smoking, drugs, exercise. • Linked to many diseases -- cancer, diabetes, Alzheimer, atherosclerosis and arthritis • Damage may be slowed with antioxidants such as vitamins C and E, selenium & beta-carotene (precursor tovitamin A). Increased serum thiol levels with spinal adjustments.

  10. Atoms achieve stability by gaining, losing or sharing electrons to fill their outermost energy level. The interaction often involves the formation of chemical bonds which hold the participating atoms together once the reaction has ended. 3 basic types of chemical bonds • Ionic • Covalent • Hydrogen

  11. Bonds continued • When chemical bonding occurs, the result is the creation on NEW CHEMICAL ENTITIES called MOLECULES and COMPOUNDS. • Molecule refers to any chemical structure consisting of atoms held together by a COVALENT BOND. • Compound is any chemical substance made of atoms or 2 or more elements regardless of the bond. • Not all compounds consist of molecules. Salt (NaCl) held together by ionic bond. • Water=H2O is a compound. It contains 2 elements Hydrogen and Oxygen held by covalent bond

  12. Ionic Bonds • Positively (cations) and negatively charged ions (anions) attract each other to form an ionic bond • Atoms become ions by losing or gaining electrons • We assign + to charge on proton and - to charge on electron. • In the body, ionic bonds are found mainly in teeth and bones • An ionic compound that dissociates in water into + and - ions is called an electrolyte( Na+Cl-)salt

  13. The Ionic Bond in Sodium Chloride • Sodium loses an electron to become Na+ (cation) • Chlorine gains an electron to become Cl- (anion) • Na+ and Cl- are attracted to each other to form the compound sodium chloride (NaCl) -- table salt • Ionic compounds generally exist as solids

  14. Covalent Bonds • Some atoms can complete there outer shells not by gaining or losing electrons but by SHARING. • Individual Hydrogen atoms DON’T exist in nature. Instead we find Hmolecules. Molecular H consists of a PAIR of H atoms. • In chemical shorthand Molecular hydrogen = H2. H is symbol for Hydrogen and the subscript 2 indicates the number of atoms.

  15. Covalent Bonds • Atoms share electrons to form covalent bonds • Electrons spend most of the time between the 2 atomic nuclei • single bond = share 1pair • double bone = share 2 pair • triple bond = share 3 pair • Polar covalent bonds share electrons unequally between the atoms involved

  16. Polar Covalent Bonds • Unequal sharing of electrons between atoms. • In a water molecule, oxygen attracts the hydrogen electrons more strongly so the electrons spend more time orbiting the oxygen nucleus than orbiting the hydrogen nucleus • Oxygen has greater electronegativity as indicated by the negative Greek delta sign.

  17. Hydrogen Bonds • Covalent & ionic bonds tie atoms together to form molecules and compounds. Other WEAK forces also act between adjacent molecules and even between atoms within a large molecule. The MOST IMPORTANT OF THESE WEAK ATTRACTIVE FORCES is the HYDROGEN BOND. • A HYDROGEN BOND is the attraction between + on the HYDROGEN atom of a polar covalent bond and a – on an oxygen or nitrogen atom of another polar covalent bond. • Hydrogen bonds are to weak to create molecules but can change molecular shape. Surface tension( insects )

  18. States of Matter • Matter exists as either SOLID, LIQUID or GAS. • Solids maintain their volume and their shape at ordinary temperature and pressure. • Liquids have a constant volume but no fixed shape. • Gas has neither a constant volume nor fixed shape. • WATER is the only substance that can exist as ALL 3.

  19. Lecture 4 BASIC ENERGY CONCEPTS • To understand the relationship between MATTER and ENERGY it’s essential to discuss chemical reactions. • WORK- Movement of an object or change in the physical structure of matter ( walking or running) ( converting water to water vapor) • ENERGY- Capacity to perform work. Need energy to do work. • Kinetic energy- energy in motion. Falling off ladder, it’s kinetic energy that does damage. • Potential energy- stored energy that has potential to do work. • The potential energy in batteries is converted into kinetic energy that vibrates sound producing membranes of headset of I-POD.

  20. Potential  Kinetic Energy • Cells perform work as they synthesize complex molecules into, out of, and within the cell. • Cells of skeletal muscle at REST have potential energy in the form of protein filaments and the covalent bonds between molecules inside the cells. • When the muscle contracts it performs work. Potential energy is converted into kinetic energy and HEAT IS RELEASED. • As a result when you exercise, your body temperature rises.

  21. Thermodynamic laws • Energy can’t be created or destroyed • Energy can only be converted from one form to another. • The conversion is never 100% efficient. Each time energy exchange occurs some energy is released into non useable form.

  22. TYPES of chemical reactions • 3 types of chemical reactions. • Decomposition • Synthesis • Reaction ( exchange/reversible )

  23. Decomposition Reactions--Catabolism • Large molecules are split into smaller atoms, ions or molecules • All decomposition reactions occurring together in the body are known as catabolism • Usually are exergonic since they release more energy than they absorb

  24. Catabolism • AB A + B • When you eat a meal of Proteins, Fats, and Sugars which are too complex to be absorbed, decomposition reactions in the digestive tract break down these molecules. • Decomposition involving water is important in breakdown of complex molecules. HYDROLYIS is taking: H20  H + OH • When a covalent bond is broken, it releases kinetic energy that can perform the work. • When the energy is released cells perform vital functions.

  25. Synthesis Reactions--Anabolism • Two or more atoms, ions or molecules combine to form new & larger molecules • All the synthesis reactions in the body together are called anabolism • Usually are endergonic because they absorb more energy than they release • Ex: The formation of water from hydrogen and oxygen molecules. • DEHYDRATION SYNTHESIS or condensation is formation of a complex molecule by the REMOVAL of WATER. The opposite of hydrolysis.

  26. Exchange Reactions • Substances exchange atoms • consist of both synthesis and decomposition reactions • Example • HCl + NaHCO3  H2CO3 + NaCl • ions have been exchanged between substances

  27. Reversible Reactions • Chemical reactions can be reversible. • Reactants can become products or products can revert to the original reactants • Indicated by the 2 arrows pointing in opposite directions between the reactants and the products • AB A + B

  28. Reversible Reactions • This is important with acid & bases to balance the pH. • If you take carbonic acid and add OH then your left with bicarbonate ( base). H2CO3  HCO3 • If we take Bicarcbonate and add H then we end up with carbonic acid. HCO3 H2CO3 • Below is a type of response to pH in body fluids. • C02 + H2O  H2CO3 HCO3 + H • 

  29. Inorganic Compounds & Solvents • Most of the chemicals in the body are compounds • Inorganic compounds • usually lack carbon& hydrogen are structurally simple • water, carbon dioxide (a by-product of cell metabolism), oxygen (important in metabolic reaction,) acids, bases and salts (compounds held together by IONIC BONDS ) • Organic compounds • contain carbon & hydrogen • always have covalent bonds

  30. WATER PROPERTIES • We will be concerned primarily with water and how its properties are essential for life. Most of the other inorganic molecules exist in association with water. Both carbon dioxide and oxygen for example are gas molecules that are transported in body fluids. All inorganic acids, bases and salts dissolve in body fluids as well. Water is known as the universal SOLVENT

  31. Water • Water makes up about 2/3 of total body weight. • The bonds in water are oriented in such a way that the hydrogen atoms are close together. As a result the water molecule has a (+) and (–) poles. Water is called POLAR MOLECULE. • Many inorganic compounds are held together by ionic bonds. In water these COMPOUNDS undergo IONIZATION = dissociation. • Aqueous solution containing +/- ions will conduct electrical current. + & - move toward each other. • Many inorganic/organic molecules will dissolve in water=solution. The medium in which other atoms, ions molecules are dispersed is called the solvent. The dispersed substances are the solute.

  32. Water Functions • Heat capacity is high. It can absorb a large amount of heat with only a small increase in its own temperature. • large number of hydrogen bonds in water • bonds are broken as heat is absorbed instead of increasing temperature of water. • The cohesive nature of water molecules allows blood and body fluids to flow smoothly through vessels and tissue spaces and allows for uniform distribution of solutes. • Major component of lubricating fluids within the body • mucus in respiratory and digestive systems • synovial fluid in joints • serous fluids in chest and abdominal cavities • organs slide past one another

  33. Electrolytes • Soluble inorganic molecules whose ions conduct an electrical current in solution are electrolytes. • NaCl is an electrolyte. • Changes in the concentration of electrolytes in body will effect every vital function. Ex: decreased potassium levels will lead to general muscular paralysis and increased levels will cause weak or irregular hear beats. • The balance(homeostasis) is regulated by Kidneys ( ion excretion) the digestive tract (ion absorption) and skeletal system (ion storage or release)

  34. Electrolytes • + IONS (Cations)- IONS (Anions) • Na - Sodium Cl • K - Potassium HCO3 – Bicarbonate • Ca - Calcium HPO4 – Phosphate • Mg - Magnesium SO4 – Sulfate • PROTEINS • These are used in acid/base balance and action potentials.

  35. Hydrophilic & Hydrophobic Compounds • Some organic molecules contain polar covalent bonds which also attract water molecules. The hydration spheres that form may then carry these molecules into SOLUTION. • Molecules such as GLUCOSE ( an important soluble sugar) that interact readily with water is called HYDROPHILIC. (Water- Loving)

  36. Hydrophobic • When organic molecules lack polar covalent bonds or have few. These molecules don’t have + or – terminals and are said to be non polar. • The most familiar hydrophobic molecule are fats and oils of all kinds. • Think of what happens if you spill gasoline or oil in ocean. They form slicks instead of dissolving.

  37. Inorganic Acids, Bases & Salts • Acids, bases and salts always dissociate into ions if they are dissolved in water • acids dissociate into H+and one or more anions • bases dissociate into OH-and one or more cations • salts dissociate into anions and cations, none of whichare either H+ or OH- • Acid & bases react in the body to form salts • Electrolytes are important salts in the body that carry electric current (in nerve or muscle)

  38. ACID and BASES • The human body can only remain healthy if it’s fluids exist within a narrow band of pH. The pH scale measures the relative concentration of HYDROGEN H+ and HYDROXIDE OH-. • The pH scale goes from 0-14 • 0-6.9 is ACID: stomach HCL, beer, wine, tomatoes saliva ( contains more Hydrogen than Hydroxide) • pH of 7 = neutral ( equal # of Hydrogen & Hydroxide) • pH above 7 = Base ( more Hydroxide than Hydrogen) blood, ocean water, bleach, ammonia (7.35-7.45 blood) Should eat more alkaline FORMING foods to remain healthier. Diet should have little acid forming foods. 80% alkaline foods and 20% acidic foods

  39. Acid/ Base • The strength of acid/base is determined by the ability of the substances to dissociate into individual ions. • HCL( hydrochloric acid) ( muriatic acid)  H+ Cl- • Can lower the pH of pools • NaOH (sodium hydroxide) Strong base----- Na+ + OH- • Strong bases have industrial uses like drano. • If acid or base is swallowed the UNIVERSAL ANTIDOTE is RAW EGG WHITES, WATER & MILK. Then take to hospital for further evaluation. Never vomit the acids up. You will damage the tissues a second time.

  40. Buffer Systems of the Body • Body fluids vary in pH but the range of each is limited and is maintained by a variety of buffering systems. • gastric juice 1.2 to 3.0; saliva 6.35 to 6.85; bile 7.6 to 8.6 and blood 7.35 to 7.45 • Buffers stabilize the pH of a solution by removing or replacing hydrogen ions. They involve a WEAK ACID and its related salt. • carbonic acid - bicarbonate buffer system NaHCO3 otherwise known as baking soda. Alka-Seltzer can neutralize the acid in the stomach. • HCL + NaOH - H2O + NaCl • This reaction produces water and salt. By adding HCL you neutralize both the strong acid and strong base.

  41. Carbonic Acid Bicarbonate Buffering System • Carbonic Acid H2CO3 • Bicarbonate HCO3 • If you add H which is an ACID it will go to the BASE to make it less basic or more acidic. If there is HCO3 in a flask of water and you add H, the result is H2CO3 • ADD H • HCO3 H2CO3 A base became an acid • ADD OH • H2CO3 HCO3 A acid became a base

  42. Lecture 5 -Organic Compounds • Always contain carbon and hydrogen & generally Oxygen • Usually long chain of Carbon atoms linked by covalent bonds • The carbon atoms typically form additional bonds with Hydrogen or Oxygen. • Make up 40% of body mass. • 4 Main Classes: Carbs, Lipids, Proteins, and Nucleic Acids.

  43. Functional groups of Organic Compounds • Carboxyl Group –COOH : Acts as an acid, releasing H bonds to become R…--COO - Examples are Fatty acids and Amino Acids • Amino Group –NH2Can accept or release H bonds depending on the pH. An example are Amino acids • Hydroxyl Group –OHStrong base dissociate to release hydroxide ions ( OH-) Examples Carbohydrates, amino acids, and fatty acids. • Phosphate Group –PO4 Links other molecules to form larger structures. Stores energy in high NRG bonds. Example: phospholipids, nucleic acids.

  44. Isomers • Isomers have same molecular formulas but different structures (glucose & fructose are both C6H12O6 • STRUCTURALFORMULA OFGLUCOSE

  45. Carbohydrates • Sugars and Starches that make up roughly ½ of U.S. diet. • Important for energy sources that are catabolized rather than stored. • 3 sizes of carbohydrate molecules • monosaccharides • disaccharides • polysaccharides

  46. Carbohydrates • Diverse group of substances formed from C, H, and O • ratio of one carbon atom for each water molecule (carbohydrates means “watered carbon”) • glucose is 6 carbon atoms and 6 water molecules (H20) • Main function is source of energy for ATP formation • Forms only 1 % of total body weight • glycogen is storage form of glucose in liver and muscle tissue • sugar building blocks of DNA & RNA(deoxyribose & ribose sugars) C6H12O6 - Glucose

  47. Monosaccharides • Called simple sugars • Contain 3 to 7 carbon atoms • We can absorb only 3 simple sugars without further digestion in our small intestine • glucose found in syrup or honey (metabolic fuel) • fructose found in fruit • galactose found in dairy products

  48. Disaccharides • Formed by combining 2 monosaccharides by dehydration synthesis (releases a water molecule) • sucrose (table sugar)= glucose & fructose • Maltose(beer) = glucose & glucose • Lactose(dairy) = glucose & galactose (lactose intolerance) • ALL carbs except monosaccharides need to be broken down by hydrolysis before they can provide useful energy. • Sugars are stored as fat for energy. ( artificial sweeteners can’t be broken down)

  49. Polysaccharides- 3 Complex Carbohydrates • Contain 10 or 100’s of monosaccharides joined by dehydration synthesis. • STARCH: The storage form of glucose in plants. High concentrations of glucose in plant or animal cells is TOXIC. Surplus stores of glucose is converted to starch for safety. • Glycogen is the storage form of glucose in animals. Glycogen is stored in the liver (hepatocytes) and in muscle cells. If serum glucose levels fall, glycogen is converted into glucose • Cellulose: can only digest small amounts via bacterial action. Cellulose is called dietary fiber. Water soluble found in fruits veggies. LOWER CHOLESTEROL. Insoluble fiber is in nuts and grains.

  50. Lipids = fats • Formed from C, H and O • includes Fats, oils, waxes, • 18-25% of body weight • Hydrophobic • insoluble in polar solvents like water • BUTcan combines with proteins for transport in blood (lipoproteins LDL’s, HDL’s ) • They form the essential structural component of all cells. • They provide twice the energy as carbs gram for gram when broken down in the body. • 5 Classes: Fatty acids, eicosanoids, glycerides, steroids, phospholipids.