Chemistry

1. Chemistry Chemistry for the biologist

2. Matter…Mass…Weight • Matter = Anything that takes up space and has mass • Mass = A measure of the amount of matter an object contains. • Weight is the measure of how strongly an object is pulled by the earth’s gravity. On earth weight can be used as a measure of mass.

3. Elements • Element = A substance that cannot be broken down into other substances by chemical reactions. There are 92 naturally occuring elements • 96% of living matter made up of C, O, H, and N • Ca, P, K, S , Na, Cl, and Mg are also common • Trace elements = elements require in very minute quantities. • B, Cr, Co, Cu, F, I, Fe, Mn, Mo, Se, Si, Sn, V, and Zn  • Compound = a pure substance composed of two or more elements combined in a fixed ratio. (e.g NaCl…CaCl2)

4. Atoms Atom = smallest unit of matter that retains the physical and chemical properties of an element.

5. Atomic Structure • Electrons – negatively charged subatomic particles circling a nucleus • Nucleus – contains neutrons and protons • Neutrons – uncharged particles • Protons – positively charged particles

6. Atomic Structure Atomic number = Number of protons in an atom of a particular element. In a neutral atom the # of protons = # of electrons11Na Mass number = Number of protons and neutrons in an atom. (close to but not exactly the same as atomic weight) 23Na 12 C 6 23 Na 11

7. Isotopes • Isotopes = atoms that have the same atomic number but different mass number. Same # of protons, but different # of neutrons. Some are radioactive • Radioactive isotope = Unstable isotope in which the nucleus spontaneously decays emitting sub-atomic particles and energy as radioactivity.

8. Energy • Potential energy = energy that matter stores because of its position or location. • Electrons have potential energy because of their position relative to the positively charged nucleus. • Electrons are arranged around the nucleus in electron shells which correspond to different energy levels. • Electrons with the lowest energy potential are located closest to the nucleus • Electrons with the highest energy are located further from the nucleus

9. (a) A ball bouncing down a flight of stairs provides an analogy for energy levels of electrons Third shell (highest energy level) Energy absorbed Second shell (higher energy level) First shell (lowest energy level) Energy lost Atomic nucleus (b)

10. Electron Configurations • Electron orbitals Orbital= The 3-D space where an electron will most likely be found 90% of the time. Electron configuration = The distribution of electrons in an atom’s electron shells. Valence electrons – electrons in outermost shell that interact with other atoms. Only the electrons of atoms interact, so they determine atom’s chemical behavior • Octet rule = A valence shell is complete when it contains 8 electrons (H and He are exceptions) • An atom with an incomplete valence shell is chemically reactive !!

11. Electron Configurations

12. Helium 2He Hydrogen 1H Atomic number 2 He 4.00 Atomic mass Element symbol First shell Electron- distribution diagram Lithium 3Li Beryllium 4Be Fluorine 9F Boron 5B Nitrogen 7N Neon 10Ne Carbon 6C Oxygen 8O Second shell Chlorine 17Cl Sodium 11Na Aluminum 13Al Silicon 14Si Argon 18Ar Magnesium 12Mg Phosphorus 15P Sulfur 16S Third shell

13. Chemical Bonds • Chemical bonds – when atoms combine by sharing or transferring valence electrons • Molecule – two or more atoms held together by chemical bonds • Compound – a molecule composed of more than one element

14. Chemical Bonds • Principal types of chemical bonds • Covalent bonds • Nonpolar covalent bonds • Polar covalent bonds • Ionic bonds • Hydrogen bonds – weak forces that combine with polar covalent bonds

15. Covalent Bonds • Covalent bond – sharing of a pair of electrons by two atoms Can be single, double or triple • Electronegativity – An atom’s ability to attract and hold electrons.

16. Nonpolar Covalent Bonds • Atoms with similar electronegativities • Shared electrons spend equal amount of time around each nucleus • No poles exist • Carbon atoms critical to life; forms four nonpolar covalent bonds with other atoms • Organic compounds contain carbon and hydrogen atoms

17. Nonpolar Covalent Bonds

18. Nonpolar Covalent Bonds

19. Polar Covalent Bonds • Unequal sharing of electrons due to significantly different electronegativities • Most important polar covalent bonds involve hydrogen • Allows for hydrogen bonding

20. Polar Covalent Bonds Water

21. Ionic Bonds Ion = a charged atom or molecule Anion = An atom that has gained one or more electrons from another atom and has become negatively charged. Cation = An atom that has lost one or more electrons and has become positively charged. Ionic bond = Bond formed by the electrostatic attraction after the complete transfer of an electron from a donor atom to an acceptor.

22. Hydrogen Bonds Hydrogen bond = Bond formed by the charge attraction when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom. • A weak bond that is 20X easier to break than a covalent bond • Occurs between oppositely charged portions of polar molecules • Biologically very important.!!!!! • Help stabilize the 3-D shape of large molecules (proteins, DNA, etc.) • They are temporary associations, forming and braking with relative ease.

23. Fig. 2-16  + Water (H2O) + Hydrogen bond  Ammonia (NH3) + + +

24. Chemical reactions Chemical reaction = process of making and breaking chemical bonds leading to changes in the composition of matter. • Involve reactant(s) and product(s) A + B  C + D Q + R  S T U + V • Matter is conserved • Reactions are reversible • products of the forward reaction become reactants for the reverse reaction