1. Functional Human Physiology Lab for the Exercise and Sport Sciences Basic Chemistry Overview Jennifer L. Doherty-Restrepo, MS, ATC, LAT Department of Health, Physical Education, and Recreation Florida International University

2. Basic Chemistry Overview As chemistry is not a prerequisite for this class, basic chemistry as an overview will be presented. This information will not be specifically covered on exams. Only chemistry as it relates to the function of the human body will be covered on the exams. Chemistry  The branch of science that studies the composition of substances and changes that take place in their composition.  The study of matter. Structure of Matter  Matter is anything that has mass (weight without gravity) and takes up space.  Includes solids, liquids, and gases.  Elements and Atoms 1)All mater is composed of fundamental substances called elements 2)Elements  One of a limited number of unique varieties of matter. At present, there are 109 unique chemical elements listed on the Periodic Table of Elements.

3. Elements Bulk elements:  Name/Symbol/Percentage in Body 1)Oxygen O 2 65.0% 2)Carbon C 4 18.5% 3)HydrogenH+9.5% 4)NitrogenN 4 3.2% 5)CalciumCa++ 1.5% 6)PhosphorusP-1.0% 7)PotassiumK+0.4%  Taken together, these 7 elements plus Sulfur (S), Chlorine (Cl-), Sodium (Na+), and Magnesium (Mg++) to comprise 99% of the body's material. Trace elements  Basic chemical substances required by the body.  Often play a role in enzyme structure.

4. Atoms The smallest particle of an element that exhibits the properties of that element. Smallest unit of matter Structure of Atoms  Nucleus  The central core of the atom containing protons and neutrons. 1)Protons  Positively charged particles found in the nucleus of an atom. The number of protons always stays the same for any given atom. It defines the element. 2)Neutrons  Particle found in the nucleus that has no electrical charge. It is neutral. The number of neutrons can vary between atoms of the same element.  Electrons 1)Negatively charged particles found orbiting outside the nucleus of an atom. 2)Electrons are located in shells or orbitals or energy levels. 3)The number of electrons varies in an atom and when the different than the number of protons, the atom is an ion. 4)Electrons in the outer orbitals (valence electrons) determine the reactivity of the atom.

5. Ions An atom that has either a negative or positive charge.  Cation 1)Positively charged ion resulting from more protons than electrons. 2)Cations are attracted or move toward anything negative.  As in a negative ion or a negative electrode (anode)  Anions 1)Negatively charged ions resulting from more electrons than protons. 2)Anions are attracted or move toward anything positive.  As in a positive ion or a positive electrode (cathode) Isotopes  An atom of an element that has more or fewer neutrons in its nucleus (than protons) giving it a larger or smaller atomic weight or mass.

6. Molecules The smallest unit of matter that retains the properties of that material.  Example 1)A molecule of sugar is the smallest fragment of sugar that can still be recognized as sugar.  When all the atoms in a molecule are the same, they form a molecule of that element.  When some of the atoms in a molecule are different, they form a chemical compound. Molecular formula  Method of representing the elements and the number of atoms of each element in a molecule.  Examples 1)One molecule of water has the molecular formula of H 2 O, meaning it has two hydrogen atoms and one oxygen atom. 2)A molecule of sugar or glucose is C 6 H 12 O 6 meaning it has 6 carbon, 12 hydrogen, and 6 oxygen molecules.

7. Polarity of Molecules Refers to an unbalance of electrical charge. Polar molecule  Occurs when part of the molecule has a partial charge  The molecule is asymmetric or unbalanced  Soluble in water and not soluble in fats (lipids)  Example 1)Water is a polar molecule and allows water molecules to interact or loosely bond to each other. This quality give water its liquid properties. Non-polar molecule  Electrically balanced and symmetrical molecule  Usually soluble in fats (lipids) and not soluble in water

8. Compounds One or more molecules composed of two or more different atoms, ions, or elements.  Example 1)Table salt is NaCl. 2)It is composed of one sodium ion and a chlorine ion. The charges of the ions hold or bond the ions of each element together.

9. Bonding of Atoms Chemical bonds are the forces of attraction that hold the atoms in a molecule together. Requires energy to make and break chemical bonds or to cause a chemical reaction. Chemical reactions  Making and breaking of bonds between atoms that either require or produce energy.  Occurs when atoms combine with each other or break apart from each other.  All life processes depend on chemical reactions.  Chemical reactions depend on valence electrons.

10. Bonding of Atoms Valence  Describes the ability of an atom to combine with other atoms.  It is a function the number of electrons in the outer energy level or orbitals of an atom. Orbitals or Shells  Concentric rings or layers where electrons are found circling a nucleus of an atom.  The closer the orbital to the nucleus, the tighter the electrons are held to the nucleus and the less like the atom is to lose that electron (ionize) or form chemical bonds.

11. Types of Chemical Bonds Hydrogen bonds  Simplest and weakest bonds.  Occurs when a molecule containing hydrogen, the positive part of a polar (as in water) molecule, temporarily becomes attracted to the negative part of another polar molecule. Ionic bonds  Bonding that results when one atom gives up an electron (becomes a positive ion) and becomes attracted to another ion that has gained the electron (becomes a negative ion).  The oppositely charged ions are attracted to each other and hold the atoms together. Covalent bonds  Occurs when atoms fill the outer orbitals of another atom by sharing electrons.  Electrons are not gained or lost.

12. Energy The capacity to do work. Energy exists in several forms and is often converted from one form to another. Types of Energy  Electrical energy 1)Results from the flow of electrons or ions (or protons). 2)It is the force that causes particles to move when opposites attract. 3)Opposites attract, like repels.  Chemical energy 1)The energy held in chemical bonds. 2)Concentration is considered chemical energy during certain chemical processes like diffusion and osmosis. 3)Particles always move from areas of higher concentration to lower concentration. 4)Must spread out or move away from other particles.  Gradient or Potential 1)Differences in concentration or electrical charge.

13.  Mechanical energy 1)The energy of movement as when the body parts are moved.  Heat energy 1)A product of many chemical reactions and some of the heat released in these reactions is used to maintain temperature homeostasis. All body activities depend on the availability of chemical energy. Most of the chemical energy in the cell is provided by ATP, Adenosine Triphosphate, which we will discuss later. Electrochemical reactions Combination of chemical and electrical energy Basis for the functioning of the nervous system via nerve impulse conduction, and of the muscle tissues, via muscle cell contraction.

14. Electrolytes A substance that ionizes in water or a molecule that becomes an ion when added to a solution.  Example 1)NaCl ↔ Na+ + Cl- The solution that results will conduct an electrical current

15. Acids  A substance that releases hydrogen ions (H+ or protons) when added to water.  Example 1)HCl ↔ H+ + Cl-  The hydrogen ion is called a proton (H+ ions)  Acids are called proton donors because they produce H+ ions.

16. Bases  A substance that releases OH-or hydroxyl ions when added to water OR an ion that combines with H+ ions.  Example 1)NaOH ↔ Na+ + OH-  Bases produce negatively charged OH- or hydroxyl ions.  Basic solutions are also called alkaline.  Bases are called proton acceptors because they take up hydrogen ions.  When this occurs water is formed. 1)OH- + H+ ↔ H2O

17. Salts A compound produced by a reaction between an acid and a base.  Example 1)HCl + NaOH ↔ H2O+ NaCl Also, a salt is an electrolyte that dissociates into cations (+ ions) and anions (- ions), neither of which is OH- + H+.  Example 1)NaCl ↔ Na+ + Cl-

18. pH Scale Acid and Base Concentrations are measured on the pH scale  A measure of the hydrogen ion concentration of a solution.  Ranges from 0 (most hydrogen ions) strong acid or very acidic to 14 (no hydrogen ions) most alkaline or basic.  The pH scale is a logarithmic scale. 1)This means that every increment represents a 10-fold increase. 2)Example  A pH of 4 represents a concentration of hydrogen ions that is 10 times greater than a pH of 5 and 100 times greater than a pH of 6.

19. Buffers Substances that resist large and/or sudden changes in the pH of a solution by reacting with a strong acid or base to form a weaker acid or base. Deviations from the normal pH range are controlled by buffers. The function of a buffer is to convert strong acids, which are relatively unstable and ionize easily (completely dissociate providing lots of H+), into between acids which are relatively stable and do not ionize easily (do not completely dissociate - provide only a few H+).

20. Most buffers in the human body consist of a weak acids and the salt of that acid that functions as a weak base.  Example 1)The carbonic acid-bicarbonate buffering system 2)Used to reduce the acidic effects of CO2 in body fluids. 3)Based on the bicarbonate ion (HCO3-), a weak base and carbonic acid (H2CO3), a weak acid.  CO 2 + H 2 O ↔ H 2 CO 3 ↔ HCO 3 - + H+ When pH drops (acidity sharply increases), the bicarbonate ion weak base binds the H+ and raises the pH back to normal. When pH rises (alkalinity sharply increases), the carbonic acid buffer releases H+ to reduce the pH back to normal. Objective is to always keep body fluids within normal homeostatic pH ranges.