1. The Chemical Context of Life Chapter 2

2. Matter  Matter consists of chemical elements in pure form and in combinations called compounds; living organisms are made of matter.  Matter -- Anything that takes up space and has mass.  Element -- A substance that cannot be broken down into other substances by chemical reactions; all matter made of elements.  Life requires about 25 chemical elements  96% of living matter is composed of C, O, H, N.  Most of remaining 4% is P, S, Ca, K.  Trace element -- required by organisms in extremely small quantities: Cu, Fe, I, etc.

3. Matter cont.  Compound -- Pure substances made of two or more elements combined in a fixed ratio.  Have characterisitics different than the elements that make them up (emergent property).  Na and Cl have very different properties from NaCl.  Difference between mass and weight:  Mass -- measure of the amount of matter an object contains; constant.  Weight -- measure of how strongly an object is pulled by earth's gravity; varies.

4. Nutrient Deficiencies

5. Atomic structure determines the behavior of an element  Atom -- Smallest possible unit of matter that retains the physical and chemical properties of its element.  Subatomic Particles  1. Neutrons (no charge/neutral; found in nucleus; ~ 1 amu).  2. Protons (+1 charge; found in nucleus; ~ 1 amu).  3. Electrons (-1 charge; electron cloud; 1/2000 amu).  One amu approx equal to 1.7 x 10 -24 g.

6. Atomic Number and Atomic Weight  Atomic number = Number of protons in an atom of a particular element.  All atoms of an element have the same atomic number.  In a neutral atom, # protons = # electrons. Mass number -- Number of protons and neutrons in an atom; not the same as an element's atomic weight.

7. Examples  23 Mg Mass number ?? Atomic number ??  12  23 12  # of protons ?? # of electrons ?? # of neutrons ??  12 12 11  14 C Mass number ?? Atomic number ??  6 6 6 6  14 6  # of protons ?? # of electrons ?? # of neutrons ??  6 6 8

8. Isotopes Isotopes  Isotopes -- Atoms of an element that have the same atomic number but different mass number; different number of neutrons.  Half-life -- Time for 50% of radioactive atoms in a sample to decay.  Biological applications of radioactive isotopes include:  1. Dating geological strata and fossils.  Radioactive decay is at a fixed rate; by comparing the ratio of radioactive and stable isotope, age can be estimated. in a fossil with the  Ratio of Carbon-14 to Carbon-12 is used to date fossils less than 50,000 years old.

9. Isotopes cont.  2. Radioactive tracers  Chemicals labelled with radioactive isotopes are used to trace the steps of a biochemical reaction or to determine the location of a particular substance within an organism.  Isotopes of P, N and H were used to determine DNA structure.  Used to diagnose disease.  3. Treatment of cancer  Can be hazardous to cells.

11. Energy Levels  Electrons are directly involved in chemical reactions.  They have potential energy because of their position relative to the positively charged nucleus.  There is a natural tendency for matter to move to the lowest state of potential energy.  Different fixed potential energy states for electrons are called energy levels or electron shells.  Electrons with lowest potential energy are in energy levels closest to the nucleus.  Electrons with greater energy are in energy levels further from nucleus.  Electrons may move from one energy level to another.

12. Electron Configuration and Chemical Properties  Electron configuration -- Distribution of electrons in an atom's electron shells; determines its chemical behavior.  Chemical properties of an atom depend upon the number of valence electrons (electrons in the outermost energy level.  Octet rule -- A valence shell is complete when it contains 8 electrons (except H and He).  An atom with an incomplete valence shell is chemically reactive (tends to form chemical bonds until it has 8 electrons to fill the valence shell).  Atoms with the same number of valence electrons show similar chemical behavior.

13. Bonding in Molecules  Chemical bonds -- Attractions that hold molecules together.  Molecules --Two or more atoms held together by chemical bonds.  Covalent bond -- formed between atoms by sharing a pair of valence electrons; common in organic compounds.  Single covalent bond -- Bond between atoms formed by sharing a single pair of valence electrons.  Double bond -- share two pairs of valence electrons.  Triple bond -- share three pairs of valence electrons.  Compound = A pure substance composed of two or more elements combined in a fixed ratio.  For example: water (H 2 O), methane (CH 4 ).

15. Nonpolar Covalent Bonds   Electronegativity -- Atom's ability to attract and hold electrons.   The more electronegative an atom, the more strongly it attracts shared electrons.   Scale determined by Linus Pauling:   O = 3.5; N = 3.0; S and C = 2.5; P and H = 2.1.     Nonpolar bond -- Covalent bond formed by an equal sharing of electrons between atoms.   Occurs when electronegativity of both atoms is about the same.   Molecules made of one element usually have nonpolar covalent bonds (H 2 and O 2 ).

16. Polar Covalent Bonds   Polar bond -- Covalent bond formed by an unequal sharing of electrons between atoms.   Occurs when the atoms involved have different electronegativities.   In water, electrons spend more time around the oxygen than the hydrogens. This causes the oxygen atom to have a slight negative charge and the hydrogens to have a slight positive charge.

18. Ionic Bonds  Ion -- Charged atom or molecule.  Anion -- An atom that has gained one or more electrons from another atom; negatively charged.  Cation -- An atom that has lost one or more electrons; positively charged.  Ionic bond -- Bond formed by the electrostatic attraction after the complete transfer of an electron from a donor atom to an acceptor.  Strong bonds in crystals, but fragile bonds in water.  Ionic compounds are called salts (e.g. NaCl or table salt).

19. Biologically important weak bonds  Include: Hydrogen bonds; Ionic bonds in aqueous solutions; Van der Waals forces.  Hydrogen bond -- Bond formed by the charge attraction when a hydrogen atom covalently bonded to one electronegative atom is attracted to another electronegative atom.  Van der Waals -- charge attraction between oppositely charged portions of polar molecules.