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Bio 232 Week 2

by: Adrienne Covington

Bio 232 Week 2 Bio 232

Adrienne Covington

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Human Anatomy and Physiology 1
Dr. Cockrell
Class Notes
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This 4 page Class Notes was uploaded by Adrienne Covington on Wednesday August 31, 2016. The Class Notes belongs to Bio 232 at Coastal Carolina University taught by Dr. Cockrell in Fall 2016. Since its upload, it has received 16 views. For similar materials see Human Anatomy and Physiology 1 in Biology at Coastal Carolina University.


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Date Created: 08/31/16
Biology 232 Lecture Notes Chapter 2. Chemistry Comes Alive.  Part 1: Basic chemistry I.  Definitions of matter and energy A.  What is matter? ­Anything that has mass and takes up space B.  States of matter 1. Solid­ has definite shape and volume 2. Liquid­ has definite volume, changeable shape 3. gas­ has changeable shape and volume 4. plasma­ an ionized gas state C.  What is energy? (Kinetic and potential energy)­The capacity to do work 1. Kinetic energy­ energy in action 2. Potential energy­ energy of position; stored (inactive) energy D. Forms of energy 1.  Chemical energy­ stored in the bonds of chemical compounds 2.  Electrical energy­ results from the movement of charged particles  3.  Mechanical energy­ directly involved in moving matter 4.  Radiant energy (electromagnetic energy) ­ energy traveling in waves II. The composition of matter A.  What are elements?  (Atoms and atomic symbols) ­unique substances that cannot be broken down by ordinary chemical means 1. Atoms­ more­or­less identical building blocks for each element 2. Atomic symbol­ one or two­letter chemical shorthand for each element B. Major, lesser and trace elements of the human body 1. Carbon, Oxygen, Hydrogen, Nitrogen. (Make up >96% of the body) 2. Calcium, Phosphorus, Potassium, Sulfur, Sodium, Chlorine, Magnesium,  Iodine, Iron (Lesser elements that make up 3.9% of the body 3. Trace elements make up less than .01% of the body. Required for enzyme  function in minute amounts C.  The structure of atoms 1.  The nucleus (protons, neutrons and atomic mass units) a. Protons­ have a positive charge and a mass of 1 atomic mass unit b. Neutrons­ have no charge and a mass of 1 atomic mass unit 2.  Electrons (planetary and orbital models) a. Electrons has a negative charge and 1/2000 the mass of a proton (0amu) b. Number of e­=p+ (atoms of elements have a neutral charge) C.  Identifying elements 1.  Atomic number­ equal to the number of protons (identifies an element) 2.  Mass number­ equal to the mass (number) of the protons and neutrons 3.  Isotopes­ atoms with the same number of protons but a different number of  neutrons 4.  Radioisotopes­ atoms that undergo spontaneous decay called radioactivity 5.  Atomic weight­ average of the mass numbers of all isotopes of an element as  they occur in nature III. How matter is combined A.  Molecules­ two or more atoms held together by chemical bonds B.  Compounds – two or more different kinds of atoms chemically bonded together C.  Mixtures (Solutions, colloids and suspensions), ­Two or more components physically intermixed (not chemically bonded). 1.  Solutions­ homogeneous mixtures of components a. Solvents­ substance present in the greatest amount (e.g. water) b. Solute­ substance(s) present in the smaller amounts c. Molarity­ moles of solute per liter of solution ­A mole of an element is equal to it’s atomic weight in grams 2.  Colloids­ heterogeneous mixtures whose solutes do not settle out 3.  Suspensions­ heterogeneous mixtures with visible solutes that tend to settle out 4.  Mixtures a. Components of mixtures are not chemically bonded b. Most mixtures can be separated by physical means c. Mixtures can be heterogeneous or homogeneous 5. Compounds a. Compounds are chemically bonded together b. Compounds cannot be separated by physical means c. All compounds are homogeneous D.  Chemical bonds 1.  Role of electrons a. Electron shells and energy levels­ surround the nucleus of an atom b. Valence shell­ The electrons in the outer levels involved in bonding c. Chemically inert and reactive elements­ Elements which have a full  valence shell are considered “inert” 2.  Types of chemical bonds a. Ionic bonds ­ Ions (anions and cations)­ formed by the transfer of  electrons from one atom to another. Ionic compounds dissociate into  charged ions in solution 1. Anions­ Negative charged ions that have gained electrons 2. Cations­ Positive charged ions that have lost electrons b. Covalent bonds (polar and nonpolar molecules)­ electrons are shared by atoms producing bonds and molecules. 1. Nonpolar molecules­ equal sharing of electrons between atoms 2. Polar molecules­ Unequal sharing of electrons between atoms c. Hydrogen bonds­ Weak attraction between a hydrogen atom on one  molecules and an electronegative atom of another (e.g. water). IV. Chemical reactions­ Occur when chemical bonds are formed, rearranged or broken A.  Chemical equations (reactants, products and molecular formulas) B.  Patterns of chemical reactions 1.  Synthesis reactions (e.g. dehydration synthesis)­ involve bond  formation forming larger molecules 2.  Decomposition reactions (e.g. hydrolysis)­ bonds are broken forming  smaller molecules 3.  Exchange reactions (e.g.. Oxidation­reduction reactions)­ Bonds are  both made and broken. a. Oxidation­reduction reactions (Redox) 1. Reactants losing electrons are oxidized 2. Reactants gaining electrons are reduced C.  Energy flow in chemical reactions  1. Exergonic reactions­ reactions that release energy 2. Endergonic reactions­ reactants that take in energy D.  Reversibility of chemical reactions (chemical equilibrium) ­All chemical reactions are theoretically reversible E.  Factors influencing chemical reactions 1.  Temperature­ chemical reactions proceed quicker at higher  temperatures  2.  Particle size­ the smaller the particle the faster the chemical reaction 3.  Concentration­ higher concentration of reacting particles produces  fasted reactions 4.  Catalysts (enzymes)­ increase the rate of a reaction without being  chemically changed 5. Enzymes­ biological catalysts


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