1. Chemistry Lecture Text Chapter 2

2. Chemistry in Physiology Physiology requires some familiarity with basic chemistry –atomic and molecular structure –chemical bonds –pH –organic compounds (next week)

3. 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

4. The Nucleus Nucleus = central body –Contains protons and neutrons number of protons determines the element –Fundamental type of matter

5. The Periodic Table of The Elements

6. 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

7. 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

8. 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

9. 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

10. Types of Chemical Bonds 1.Covalent bonds 2.Ionic bonds 3.Hydrogen bonds

11. 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

12. Ionic Bonds Between metal and non- metal One or more valence electrons completely transferred from one atom to another Forms ions –atoms or molecules with unequal numbers of protons and electrons

13. Ionic Bonds Cations –Positive charge –More protons than electrons –Metals Anions –Negative charge –More electrons than protons –Non Metals Attract each other –form ionic compound

14. Dissociation of Ionic Compounds ionic bonds tend to be weak –Can dissociate in water –Water attracted electrostatically –forms hydration spheres around molecules

15. Water Solubility Hydration sphere formation determines water solubility Hydrophilic –Water soluble –Polar molecules and ions Hydrophobic –Water insoluble –Nonpolar molecules

16. 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.

17. Acidity and Alkalinity Sometimes water molecules will split –Covalent bond between oxygen and a hydrogen will be broken –Form H + (hydrogen ion) and OH - (hydroxide ion) –H 2 O  H + + OH -

18. Acidity and Alkalinity In pure water, equal amounts of H + and OH - are formed –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

19. pH Index of [H + ] in a solution Quantify acidity or alkalinity of a solution pH = log(1/[H + ]) Example: for pure water [H + ] = 1 x 10 -7 M pH = log (1/0.0000001) = log (10,000,000) = 7

20. pH Solutions w/ pH = 7.0 are neutral Solutions w/ pH < 7 are acidic –[H + ] > 1x10 -7 M Solutions w/ pH > 7 are alkaline –[H + ] < 1x10 -7 M

21. 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