BIOL 1020 - CHAPTER 2 LECTURE NOTES

Chapter 2: The chemical context of life

You must understand chemistry to understand life (and to pass this course)!

Overview: In many ways, life can be viewed as a complicated chemical reaction. Modern models of how life works at all levels typically have at least some aspect of chemistry as a major component or underpinning.

  1. Elements and Atoms
  2. elements – substances that cannot be further broken down into other substances (at least by ordinary chemical reactions)
  3. every element has a chemical symbol (H for hydrogen, O for oxygen, etc.); this is most familiar from the periodic table
  4. there are 92 naturally occurring elements, from hydrogen up to uranium
  5. 4 elements (oxygen, carbon, hydrogen, and nitrogen = O, C, H, N) make about 96% of the mass of most living things
  6. 8 others are consistently present in small amounts in living things (Ca, P, K, S, Na, Mg, Cl, Fe)
  7. several others are typically found only in trace amounts (trace elements); these tend to vary considerably in amount and even presence depending on the type of organism
  8. an atom is the smallest unit of an element that still retains the properties of that element
  9. atoms consist of subatomic particles
  10. electron - contributes no significant mass to the atom, but carries a (-1) electrical charge
  11. proton - contributes a mass of approximately 1 mass unit, and carries a (+1) electrical charge
  12. neutron - contributes a mass of approximately 1 mass unit, and carries no net electrical charge
  13. protons and neutrons are found in the nucleus (center) of an atom
  14. elements differ from each other because they contain different numbers of protons (all hydrogen atoms contain 1 proton, all carbon atoms contain 6 protons, all oxygen atoms contain 8 protons, etc.)
  15. atomic number = number of protons in the nucleus
  16. the periodic table has elements arranged largely according to atomic number
  17. protons + neutrons determine atomic mass
  18. each contribute ~1 atomic mass unit (amu, or Dalton)
  19. atoms that have the same number of protons but have different numbers of neutrons (therefore different masses) are referred to as isotopes
  20. atomic nuclei can undergo changes (decay)
  21. some elements are more stable than others
  22. some isotopes are more stable than others (most unstable = radioisotopes)
  23. decay rates are statistical averages, and are used for measuring time passage in many areas of science (carbon dating, etc.)
  24. the radiation emitted upon decay (alpha, beta, and/or gamma) can be used as a tool for experiments; can also be used medically; has other uses and dangers (nuclear power, nuclear bombs, radiation poisoning, etc.)
  25. radiation can cause mutations in DNA, can interfere with cell division
  26. electrons occupy orbitals surrounding the nucleus and move at the speed of light
  27. because ATOMS are electrically neutral the number of electrons an atom has always equals the number of protons
  28. electrons can exist at different energy levels, which correspond to orbitals
  29. the further away an orbital carries an electron from the nucleus, the higher the energy level of the electron
  30. electrons with similar energies make up an electron shell
  31. the outer electron(s) are known as the valence electron(s); collectively, they occupy the valence shell
  32. the chemical properties of an atom are largely determined by the valence electrons
  33. the science of chemistry mostly involves study of how electrons move about the nucleus, store energy, and determine chemical properties of substances as a result
  1. Describing Atomic Combinations
  2. atoms combine to form molecules and compounds
  3. molecule – two or more atoms held together by covalent bonds(defined later)
  4. may be composed of one or more elements (examples: O2, H2O)
  5. not all substances are molecular (NaCl, table salt, isn’t)
  6. if a substance is molecular, then an individual molecule is the smallest unit of the substance that exhibits the properties of the substance
  7. thus, a molecule differs in its physical and chemical properties from the elements that make it up
  8. compound - a specific combination of two or more different elements chemically combined in a fixed ratio
  9. compounds have unique physical and chemical properties that differ from those of the elements used to make it
  10. some compounds are held together by covalent bonds and are therefore molecular; some are held together by ionic bonds (defined later)
  11. chemists use two types of formulas to describe substances
  12. chemical formula - a shorthand formula showing the number of atoms of each element present in a molecule
  13. often called molecular formula if a molecule is involved; examples: H2O, CO2, O2, C6H12O6
  14. follows simplest ratio for ionic substances (NaCl, etc.)
  15. structural formula - shows the arrangement of atoms in a molecule
  16. examples:
  17. waterH─O─H
  18. carbon dioxideO═C═O
  19. molecular oxygenO═O
  20. the number of units of a substance are described using the mole
  21. molecular mass is the sum of the atomic masses of the atoms in the molecule
  22. since the actual mass of an atom is extremely small, it is convenient in the real world to work with a large number of atoms at the same time
  23. The amount of a substance that in grams has the same number as the atomic mass is a mole
  24. Thus, water has molecular mass 1+1+16 = 18; a mole of water has a mass of 18 g
  25. The mole is simply a conversion factor from the small scale of atomic mass units to the more familiar gram scale
  26. the factor represents the number of units (molecules or atoms) in a mole
  27. this factor, called Avogadro’s number, is 6.02 x 1023 atoms or molecules
  1. Chemical Bonds Hold Molecules Together and Store Energy
  2. recall that electrons in the outermost shell of an atom (valence electrons) determine the chemical behavior of the atom, i.e. what type and how many chemical bonds it can readily form
  3. most atoms in biological systems seek to have 8 electrons in their outermost shell (hydrogen seeks to have 0 or 2 electrons in its outermost shell)
  4. since atoms have the same number of electrons as protons, they meet this need to have a full valence shell by sharing, giving up, or acquiring electrons from other atoms; this forms chemical bonds
  5. a chemical bond is a reduced energy state
  6. bond energy is the amount of energy required to break a particular chemical bond
  7. there are two principle types of strong chemical bonds
  8. covalent bonds - electrons are shared between two atoms
  9. ionic bonds - one atom completely gives up an electron to another atom
  10. covalent bonds
  11. result in filled valence shells
  12. electrons are shared in pairs
  13. a single electron pair shared = a single covalent bond
  14. double and triple covalent bonds are also possible
  15. carbon forms 4 covalent bonds
  16. covalent bonds give molecules definite shapes
  17. the shared atomic orbitals require definite spatial arrangements that depend on the atoms involved in the bond
  18. covalent bonds can be nonpolar (equal sharing of electrons) or polar (unequal sharing of electrons)
  19. polar bonds result if one nucleus holds a stronger attraction on the electron pair
  20. molecules with polar bonds (polar molecules) have regions with partial charges
  21. ionic bonds
  22. when an atom gains or gives up one or more electrons, it is called an ion
  23. cations - ions that have lost one or more electrons; have a positive charge
  24. anions - ions that have gained one or more electrons; have a negative charge
  25. the suffix –ide indicates an anion
  26. polyatomic ions can also form
  27. covalently bound atoms that lose or gain electrons or protons
  28. only polyatomic ions can lose or gain protons
  29. polyatomic cations = positive charge; polyatomic anions = negative charge
  30. an ionic bond is formed by the attraction between a cation and an anion
  31. an ionic compound is a substance held together by ionic bonds
  32. ionic compounds dissociate into individual ions when dissolved in a polar substance, such as water
  33. hydration – surrounding the ions with the ends water molecules with the opposite (partial) charge
  34. hydrogen bonds
  35. weak interactions involving partially (+) charged hydrogen atoms
  36. the interaction is with another atom with a partial (-) charge
  37. can be within the same (large) molecule, or between molecules
  38. hydrogen bonds are common and important in living things
  39. water forms hydrogen bonds
  40. because they are weak, hydrogen bonds are relatively easy to manipulate
  41. collectively, hydrogen bonds can be very strong – they hold together the two strands of DNA, for example
  42. In aqueous systems (such as living organisms), the effective relative bond strengths are:

covalent bond > ionic bond > hydrogen bond

  1. Chemical Equations Describe Chemical Reactions
  2. Reactants are written on the left
  3. Products are written on the right
  4. an arrow () is used to show the direction the reaction proceeds

C6H12O6 + 6 O2 6 CO2 + 6 H2O + Energy

  1. double arrows of equal lengths () indicate equilibrium reactions (reactions proceeding simultaneously at equal rates in both directions)

N2 + 3 H2 2 NH3

  1. Sometimes, different lengths of double arrows are used to indicate which direction is favored

CO2 + H2O H2CO3

V.Oxidation-Reduction Reactions (redox reactions) Are Common in Biological Systems

  1. oxidation is a chemical process in which an atom, molecule, or ion loses an electron(s)
  2. reduction is the opposite – an electron is gained (charge is reduced)
  3. oxidation and reduction are always paired (hence redox reactions)
  4. example: rusting
  5. when iron rusts, iron oxide is formed by the oxidation of iron; this can be described by a chemical reaction as shown below:
  • 4 Fe + 3 O2 2 Fe2O3
  1. during the process iron atoms (Fe) become iron ions (Fe3+):
  • 4 Fe 4 Fe3+ + 12 e-
  1. therefore, we can say that iron atoms were oxidized to produce iron ions above
  2. on the flip side, the oxygen atoms gain electrons; we can say that the oxygen is reduced in the reaction:
  • 3 O2 + 12 e- 6 O2-
  1. oxygen is the most common oxidizing agent (hence the general term oxidation)
  2. in biological systems, typically molecules are oxidized and reduced
  3. very important in many processes such as photosynthesis, respiration
  4. electrons are less easily lost from molecules than from atoms
  5. molecules typically will lose the equivalent of a complete hydrogen atom when oxidized
  6. this means that both a proton and an electron are removed from the oxidized molecule and may be added to the reduced molecule

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