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Chemistry in Physiology Chemistry Lecture Text Chapter 2 Atoms • smallest units of matter that can undergo chemical change • made up of three basic subatomic particles – protons – positively charged – neutrons – neutrally charged – electrons – negatively charged particles The Periodic Table of The Elements • Physiology requires some familiarity with basic chemistry – atomic and molecular structure – chemical bonds – pH – organic compounds (next week) The Nucleus • Nucleus = central body – Contains protons and neutrons • number of protons determines the element – Fundamental type of matter Atomic Number and Mass • Atomic number – number of protons in an atom • Atomic mass (weight) – the total number of protons and neutrons found within an atom – Isotopes = atoms of the same element with different atomic masses 1 Electrons • Revolve around the nucleus in certain volumes of space called orbitals • Several such orbitals: – innermost can hold two electrons – second layer can hold eight electrons – valence electrons = electrons in the outer shell Chemical Bonds • Atoms may give, take or share electrons in order to achieve full outer shell – link two or more atoms together through chemical bonds – molecules – structures consisting of atoms bound together by chemical bonds Covalent bonds • two or more atoms share their valence electrons • Nonpolar molecules – atoms share electrons equally • Polar molecules – Unequal sharing of electrons – Unequal charge between different regions of the molecule Electrons and the Periodic Table • Elements are arranged in columns by the # of valence electrons • atoms are most stable when the outermost orbital is full • most elements do not have full sets of valence electrons Types of Chemical Bonds 1. Covalent bonds 2. Ionic bonds 3. Hydrogen bonds Ionic Bonds • Between metal and nonmetal • One or more valence electrons completely transferred from one atom to another • Forms ions – atoms or molecules with unequal numbers of protons and electrons 2 Ionic Bonds Dissociation of Ionic Compounds • Cations – Positive charge – More protons than electrons – Metals • Anions – Negative charge – More electrons than protons – Non Metals • Attract each other • ionic bonds tend to be weak – Can dissociate in water – Water attracted electrostatically – forms hydration spheres around molecules – form ionic compound Water Solubility • Hydration sphere formation determines water solubility • Hydrophilic – Water soluble – Polar molecules and ions • Hydrophobic – Water insoluble – Nonpolar molecules Acidity and Alkalinity Hydrogen bonds • Polar molecules have weak electrostatic attraction for one another – Slight negative end to slight positive end • Responsible for water properties, protein shape, DNA structure, etc. Acidity and Alkalinity • In pure water, equal amounts of H+ and OH- are formed • Sometimes water molecules will split – Covalent bond between oxygen and a hydrogen will be broken – Form H+ (hydrogen ion) and OH- (hydroxide ion) – H2O ⇔ H+ + OH- – Generally, [H+] = 1 x 10 -7 M (= 0.0000001 M) – Neutral solution • Some solutes (acids) release H+ when mixed with water – ↑ [H+] above [OH-] – Acidic solution • Some solutes (bases) bind H+ or release OH- when mixed with water – ↓ [H+] below [OH-] – Alkaline or Basic solution 3 pH • Index of [H+] in a solution • Quantify acidity or alkalinity of a solution pH = log(1/[H+]) pH • Solutions w/ pH = 7.0 are neutral • Solutions w/ pH < 7 are acidic – [H+] > 1x10-7 M • Solutions w/ pH > 7 are alkaline • Example: for pure water [H+] = 1 x 10-7M – [H+] < 1x10-7 M pH = log (1/0.0000001) = log (10,000,000) =7 pH • pH can range from 0 to 14 • As pH increases, [H+] decreases • A difference of 1.0 in pH means a 10x difference in [H+] – A solution of pH 7 has 10x the [H+] of a pH 8 solution 4