The Chemistry of Life

The Chemistry of Life

Matter

  • Organisms are composed of matter.
  • Matter is anything that has mass and takes up space.
    • Weight = gravitational pull on matter
  • Matter is made of elements - substances that cannot be broken down to other substances by chemical reactions.

Naturally Occurring Elements

  • 92 naturally occurring elements
    • 73% not essential for life
  • 25 essential elements
    • Important to biology

Elements in the Human Body

  • Elements making up about 96% of human body weight:
    • Oxygen (O): 65.0%
    • Carbon (C): 18.5%
    • Hydrogen (H): 9.5%
    • Nitrogen (N): 3.3%
  • Elements making up about 4% of human body weight:
    • Calcium (Ca): 1.5%
    • Phosphorus (P): 1.0%
    • Potassium (K): 0.4%
    • Sulfur (S): 0.3%
    • Sodium (Na): 0.2%
    • Chlorine (Cl): 0.2%
    • Magnesium (Mg): 0.1%
  • Trace elements (less than 0.01% of human body weight):
    • Boron (B), chromium (Cr), cobalt (Co), copper (Cu), fluorine (F), iodine (I), iron (Fe), manganese (Mn), molybdenum (Mo), selenium (Se), silicon (Si), tin (Sn), vanadium (V), zinc (Zn)

Compounds

  • Substance of two or more elements in a fixed ratio
  • Have emergent properties not associated with elements

Emergent Properties of a Compound

  • Example: Sodium + Chloride = Sodium Chloride (Table Salt)

Atoms

  • Atoms are the smallest units of matter that retain the properties of the element.
  • Composed of:
    • Nucleus
      • Protons - mass of 1 and a + charge
      • Neutrons - mass of 1 and no charge
    • Electrons - negligible mass and a - charge
  • Atomic number is the number of protons.
  • Atomic weight is #protons + #neutrons (i.e., weight of nucleus)

Atomic Number and Mass

  • Elements vary in number of:
    • Protons
    • Electrons
    • Neutrons

Atomic Weight

  • Atomic mass or atomic weight = # of protons + # of neutrons
  • Example: Carbon
    • Atomic number: 6
    • 6 protons
    • 6 neutrons
    • Atomic weight: 12
    • Atomic symbol: C

Practice

  • Element example: ^{16}_{8}O
    • How many protons? 8
    • How many electrons? 8
    • How many neutrons? 8

Atoms (cont.)

  • Mass concentrated in nucleus
  • Nuclei of atoms never touch
  • Only electrons involved in chemical reactions

Electron Shells

  • Shell – energy level of electrons at an average distance from the nucleus of an atom

Energy

  • Energy = capacity to do work
    • Work = a force acting on an object
  • Potential energy = energy of matter because of location or structure

Energy Levels of an Atom’s Electrons

  • Electrons can move from one level to another only if the energy it gains or loses is exactly equal to the difference in energy between the two levels.

The Periodic Table of the Elements

  • Shows Electron Distributions
  • Includes:
    • Atomic number
    • Element symbol
    • Electron-distribution diagram
    • Atomic mass

Valence Electrons

  • Name given to electrons in outer-most shell
  • Properties of elements related to number of valence electrons (e.g., how stable or reactive they are, whether they are charged or neutral, etc.)
  • Inert (unreactive) elements have a full outer shell

Covalent Bonds

  • Covalent bonds form between atoms with incomplete valence shells; atoms share or exchange electrons.
  • Example: C6H12O6
    • 6 atoms of carbon
    • 12 atoms of hydrogen
    • 6 atoms of oxygen

Formation of a Covalent Bond

  • Atoms share electrons to fill their valence shells
  • Example: Hydrogen molecule (H2)

Covalent Bonding in Four Molecules

  • Examples:
    • Hydrogen (H2)
    • Oxygen (O2)
    • Water (H2O)
    • Methane (CH4)

Covalent Bonds (cont.)

  • Nonpolar bonds are between atoms with similar electronegativity (i.e., the ability to attract and hold electrons).
    • Ex.: N2, O2, CH3
  • Polar bonds are between atoms with somewhat different electronegativities. This results in an unequal sharing of electrons.
    • Ex.: H2O, CO2, NH4

Polar Covalent Bonds in a Water Molecule

  • Oxygen (O) is more electronegative than hydrogen (H), so shared electrons are pulled more toward oxygen.
  • This results in a partial negative charge on the oxygen (\delta-) and a partial positive charge on the hydrogens (\delta+
  • H2O

Ionic Bond

  • One atom strips another of electron
  • Changes charge of atoms
  • + and – attracted to each other
  • Example: Sodium chloride (NaCl)
    • Na (Sodium atom) --> Na+ (Sodium ion, a cation)
    • Cl (Chlorine atom) --> Cl- (Chloride ion, an anion)

Molecular Bonds

  • The two most important types of molecular bonds that form from biological standpoint are
    • Covalent bonds
    • Ionic bonds
  • Both of these involve chemical reactions among atoms and/or molecules that involve valence electrons changing their orbitals

Hydrogen Bonds

  • Hydrogen bonds are DIFFERENT from covalent and ionic bonds because they create interactions between two or more molecules—they do not cause chemical reactions to occur and they do not change the nature or shape of the molecules involved in hydrogen bond

Hydrogen Bonds (cont.)

  • A hydrogen bond results from the attraction between the partial positive charge on the hydrogen atom of a molecule and the partial negative charge on a more electronegative atom of another molecule.
  • Examples:
    • Water (H2O) and Ammonia (NH3)

Molecular Shape

  • Important in biology
  • Example: Morphine and Natural endorphin

Chemical Reactions

  • Reactants --> Products
  • Example:
    2H2 + O2 \rightarrow 2H_2O

Chemical Reactions: Making and Breaking Bonds

  • Example: Photosynthesis
    6CO2 + 6H2O \rightarrow C6H{12}O6 + 6O2
  • carbon dioxide + water --> glucose + oxygen