1. Chemistry Elements, Atoms and Molecules

2. Why Chemistry? Nature is not neatly packaged into the individual life sciences. Biology is a multidisciplinary science, drawing on the insights from other sciences.

3. Essential Elements  96.3% of living matter O = 65%, C = 18.5%, H = 9.5%, N = 3.3%  3.7% Ca, P, K, S, Na, Cl, Mg  trace elements (<0.01% each) required in minute amounts B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, Zn

4. Some trace elements, like iron (Fe) are required by all organisms. other trace elements are required only by some species. for example, a daily intake of 0.15 milligrams of iodine is required for normal activity of the human thyroid gland.

5. Atoms Subatomic particles: atomic nucleus  protons (p +)  neutrons (n) cloud around the nucleus  electrons (e -)

6. Electrostatic Charges The attractions between the positive charges in the nucleus and the negative charges of the electrons keep the electrons in the vicinity of the nucleus.

7. Mass of Subatomic Particles 1 dalton = 1.67 X 10 -24 g p = 1 dalton n = 1 dalton e = 1/2000 dalton * Dalton = a.m.u. (atomic mass unit)

8. Relationships Atomic N° = # of protons 2 mass N° = # p + # n 4  Atomic weight = a measure of its mass, can be approximated by the mass number.  For example, He has a mass number of 4 and an estimated atomic weight of 4 daltons.  More precisely, its atomic weight is 4.003 daltons

9. Radioactive Isotope  Most isotopes are stable; they do not tend to lose particles. Both 12 C and 13 C are stable isotopes.  Some isotopes are unstable and decay spontaneously, emitting particles and energy. 14 C is a one of these unstable or radioactive isotopes.  In its decay, an neutron is converted to a proton and electron. This converts 14 C to 14 N, changing the identity of that atom. Half life 14 C = 5600 yrs. Half life 40K = 1.3 billion yrs.

10. Radioactive Isotope Uses Used for  geological dating  Treatment to trace atoms in metabolism Diagnosis research

11. Energy Levels or Electron Shells Electrons of an atom may vary in the amount of energy that they possess. The farther electrons are from the nucleus, the more potential energy they have.

12. Electron Shells Outermost shell (valence shell) determines behavior of atom  Valence electrons electrons in the outermost shell (valence shell)  Atom’s valence number of covalent bonds that it can form  Octet Rule  valence shell complete with 8 electrons (except H, He)

13. Electron Shells

14. Electron Orbitals and Shells  Electron orbitals space around nucleus where electrons are most likely to be found (90% of the time) can be shaped differently  Shells -energy levels within an orbital the first shell, closest to the nucleus, has the lowest potential energy electrons in outer shells have more potential energy electrons can only change their position if they absorb or release a quantity of energy that matches the difference in potential energy between the two levels

15. Electron Orbitals and Shells Superimposed orbitals

16. Covalent Bond  Sharing of electron pairs single bonds double bonds triple bonds  covalent bonds nonpolar covalent bonds  electrons shared equally polar covalent bonds  electrons shared unequally

17. Single Covalent Bond the sharing of a pair of valence electrons by two atoms Molecular formula Electron Structural Space Model distribution formula

18. Double Covalent Bond sharing two pairs of valence electrons Molecular formula Electron Structural Space distribution formula model

19. Nonpolar Covalent Bond Carbon and hydrogen have similar electronegativities Molecular formula Electron Structural Space distribution formula mod el

20. Polar Covalent Bond  oxygen has a much higher electronegativity than does hydrogen.  Compounds with a polar covalent bond have regions that have a partial negative charge near the strongly electronegative atom and a partial positive charge near the weakly electronegative atom.

21. Ions Atoms that have gained or lost electrons Cations and Anions  lost electron = cation  gain electron = anion Cl - = 17p + 18e + 17nNa+ = 11p + 10e + 11n 11p + 11e + 11n 17p + 17e + 17n

22. Activity: atoms 1. Match the following….isotope, cation, or anion When an atom loses an electron you form… When an atom gains an electron you form… When an atom changes the number of neutrons you form…

23. When an atom loses an electron? ---cation When an atom gains an electron? ---anion When an atom changes the number of neutrons you form… ---isotope

24. Activity: atoms 39 K with atomic number 19 What is the proton number? Neutron number? Number of electrons? Mass?

25. 39 K with atomic number 19 What is the proton number? 19 Neutron number? 20 Number of electrons? 19 Mass? 39

26. Hydrogen Bond A slightly positive H atom of a polar covalent bond in one molecule is attracted to a slightly negative atom of a polar covalent bond in another molecule For example, ammonia molecules and water molecules link together with weak hydrogen bonds

27. Van der Waals Interactions Even molecules with nonpolar covalent bonds can have partially negative and positive regions.  Because electrons are constantly in motion, there can be periods when they accumulate by chance in one area of a molecule  This created ever-changing regions of negative and positive charge within a molecule. Molecules or atoms in close proximity can be attracted by these fleeting charge differences, creating van der Waals interactions

28. Chemical Reactions Reactants products 2 H 2 O2O2 += 2 H 2 O

29. Redox Reaction (topic under cellular respiration and photosynthesis) Chemical reaction involving the transfer of 1 or more electrons from one reactant to another Redox = oxidation-reduction reaction Red = reduction, gain of electrons by a substance ox = oxidation, loss of electrons by a substance

30. Redox Reaction (cellular respiration) C 6 H 12 O 6 + 6O 2 6CO 2 + 6H 2 O Reducing agent Oxidizing agent Oxidized molecule Reduced molecule Chemical equilibrium The End