Bio 232 Week 2
Bio 232 Week 2 Bio 232
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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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