1. Chapter 2: The Chemical Context of Life

2. Matter  Anything that has mass and occupies space.  Smallest particle of an element (still retains elemental properties) Atom

3. Element  Matter made up of only one type of atom.  92 natural elements.  Each element has a symbol.

4. Compound  Compound - Elements combined in fixed ratios.  A compound has characteristics beyond those of its combined elements.

6. Question?  What elements are necessary for life?  Life requires about 25 chemical elements.

8. Macroelements  Elements needed in large amounts or quantities.  Macro = Large  Examples: C HOPKNS CaFe Mg NaCl

9. Microelements  Elements needed in very small quantities.  Also known as trace elements.  Micro = small  Examples: Cu, Co, Zn, Mo, I, Mn

10. Atomic Subparticles  Protons + charge, 1 Dalton mass  Neutrons no charge, 1 Dalton mass  Electrons - charge, essentially no mass

11. Atomic Model

12. Atomic Number  The number of protons in the nucleus.  Each element has its own atomic number. If you change the atomic number, you no longer have the same element.

13. Atomic Mass  The number of protons and neutrons in the nucleus.  The atomic mass can change.

14. Isotopes  Atoms of the same element with different atomic mass  Caused by changes in the number of neutrons  Used as “tracers”, used to kill cancer/bacteria cells, used to determine age of fossils/geological formations

15. Types of Isotopes  Radioactive Where the nucleus decays spontaneously, giving off particles and energy.  Heavy Has a stable nucleus, but masses more than the standard isotope for the element.

16. Energy  The ability to do work  THINK = ATP  ATP is how living organisms have the ability to do work They USE ATP!!!

17. Potential Energy  Is the energy that matter stores simply because of its position or location  Electrons have potential energy because of their position relative to the nucleus

19. Electron Energy Levels  Energy levels around the nucleus of an atom  1st level can have 2 electrons and has the lowest potential energy  Other levels can hold more than 2 electrons and have higher energy levels

20. Electron Orbitals  The three dimensional space where an electron is found 90% of the time.  Different orbitals have different shapes.  Each orbital can hold only 2 electrons.

21. Electron Orbitals

22. Chemical Behavior Of An Atom  Is determined by its electron configuration in the energy levels and orbitals  This determines who is can bond with (if anyone!)

23. Valence Electrons  The electrons in the outermost energy level  Electrons available chemical bonds  Atoms/Elements with same # of valence electrons will react similarly and will have similar characteristics

24. Octet Rule  The most stable condition is to have an outer level of 8 electrons  Exception - 1st level is stable with only 2 electrons  When stable - no chemical reactions will take place  Ex: Ne, He, (Noble gases)

25. Electrons of the first elements

26. Chemical Bonds  Forces that join atoms together to form molecules  Usually caused by sharing or transferring valence electrons

27. Bond Formation Depends On:  The number of valence electrons that must be gained, lost, or shared to reach the stable condition.

28. Chemical Bond Types  Nonpolar Covalent  Polar Covalent  Ionic  Hydrogen  van der Waals forces

29. Electronegativity  The attractiveness of a specific kind of atom towards e- in another atom Important in covalent bonds  Periodic Table trend: More electronegativity = stronger pull of e- He has highest EN Increases as you go right and up the table

31. Nonpolar Covalent  When electrons are shared equally between atoms  Very strong bond  Important in many molecules found in living things  Ex: carbon to hydrogen, hydrogen to hydrogen, oxygen to oxygen

32. Nonpolar Covalent  Can be single, double, or triple between two atoms  Each nonpolar covalent bond involves a pair of electrons

34. Polar Covalent  When electrons are shared unequally between atoms  Results in “polar” molecules that have charged areas  Use δ symbol  Ex: Water, H to O bonds

36. Ionic Bonds  Formed when electrons are transferred from one atom to another and ions are formed  Ex: NaCl  Why? Two atoms electronegativity are SO different that one atom gains e- completely

37. Two Types of Ions  Cations - have lost electrons (p+ > e-) giving them a positive charge. Ex: Sodium (Na+)  Anions - have gained electrons (p+ < e-) giving them a negative charge. Ex: Chlorine (Cl-)

38. Ionic Bonds  Formed when cations (+) and anions (-) attract each other  Weak chemical bond Why? Environment easily affects strength of this bond Ex: Salt  Solid in air/gas; Dissolves in liquid

39. Ionic

40. Hydrogen Bonds  When a hydrogen atom bonded to one molecule is attracted to the slightly negative area (often N or O) of another molecule  Very weak individual bond Can be a “strong” force if there are many H bonds.

41. Hydrogen bonds  Remember: H bonds occur BETWEEN MOLECULES (not b/t atoms within ONE molecule)  Ex: H bonds hold water molecules together

42. Hydrogen Bonds

43. Van der Waals  Result of e- ability to move at high speeds  Creates “spots” where there are “pools” of + and – charges  Weak chemical “bond”  Ex: gecko’s feet

44. Bond type B/t atoms w/in ONE molecule? B/t more than one different molecules? Weak or Strong Polar covalent YesNoStrong Nonpolar covalent YesNoStrong IonicYesNoWeak HydrogenNoYes (attraction of H in one to - atom in another) Weak (unless LARGE #) Van der Waals NoYesWeak

45. Molecular Shape  Each molecule on Earth has a characteristic shape Determined by the positions of the atom’s orbitals  Shape related to function

46. Molecular Shape  Molecular shape is crucial because it determines how most molecules of life recognize and respond to one another.  Ex: Viruses (reproduction), bacteria (reproduction), hormones/cell recognition

48. Chemical Reactions  The making and breaking of chemical bonds  Reactions do not destroy matter, they only rearrange it

49. Chemical Equations  A way to represent what is happening in a chemical reaction Ex: 2 H 2 + O 2 2 H 2 O

50. Parts of the Equation  Reactants: - the starting materials  Products: - the ending materials  Note - all atoms of the reactants must be accounted for in the products 2 H 2 + O 2 2 H 2 O

52. Chemical Equilibrium  When the conversion of reactants to products is balanced to the reverse reaction Ex: 3 H 2 + N 2 2 NH 3

53. Chemical equilibrium  Reversible rxn  When concentrations of react and prod STOPS changing Doesn’t necessarily mean concentrations are equal!!!!!  Rxn still continues

54. Summary  We will now put elements together to form molecules and build the next level in the hierarchy  Ch 3, 4, 5 (Properties of Water and Macromolecules)

55. Summary Continued  Recognize macro-elements and micro-elements and their roles in biological organisms.  Differentiate between elements and compounds.  Identify the basic principles of atomic structure and how they determine the behavior of an element.  Identify the main types of chemical bonds.  Discuss the relative strength of different types of chemical bonds.  Recognize that chemical reactions make and break chemical bonds.