Biology 180 Chapter 2

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rosylyn
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200897
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Biology 180 Chapter 2
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2013-02-16 02:20:41
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Chapter 2 Notes
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  1. Matter
    Anything that takes up space and has mass.

    • Is composed of elements
    • – Elements cannot be broken down to substances with different chemical or physical properties
    • – 92 Naturally occurring elements
    • – 25 elements found in organisms

    • 11 Most common found in life

    • – 96% of life forms are carbon (C), oxygen (O) , hydrogen (H),nitrogen (N)
    • – ~4% of life forms are phosphorus (P), sulfur (S), calcium (Ca),potassium (K), Sodium (Na), Chlorine (Cl), Magnesium (Mg)

    • 14 remaining trace elements are only about 0.01% of the mass of life forms
  2. Atoms
    the smallest unit of matter that retains the properties of an element
  3. Atomic number
    – The number of protons in the element

    – Written as a subscript to the lower left of the atomic symbol
  4. Mass number
    – Approximation of the total mass of an atom

    – Nearly equals the mass of the protons plus the neutrons

    – Measured in amu (atomic mass units) or Daltons

    – Written as a superscript to the upper left of the atomic symbol
  5. Isotopes
    • Elements can have isotopes
    •    Atoms with the same number of protons but differ in the number of neutrons

    • Each isotope of an atom has a different atomic mass.
    •       Remember each neutron has a mass of approximately 1 Dalton
  6. Examples of isotopes
    • Carbon

    • – Carbon –12 is the most common isotope ofcarbon (99% of carbon)
    • • It has 6 protons, 6 neutrons, 6 electrons–

    • Carbon-13 (1% of carbon)
    • • It has 6 protons, 7 neutrons, 6 electrons

    • Carbon-14 is used for carbon dating youngfossils (<1% of carbon)
    • • It has 6 protons, 8 neutrons, 6 electrons
  7. Radioactive Isotopes
    • • The atomic nucleus of some isotopes can be unstable
    • • Release particles and energy as the nucleus decays
    • • Results in a change in proton number
    • – Therefore the isotope decays into a new element!
    • • Some useful isotopes in the field of biology
    • – 14C= Used to date fossils
    • – 32P= Used to radioactively label DNA
    • – 35S=Used to radioactively label proteins in vivo
    • – 125I=Used to radioactively label proteins in vitro
  8. Detecting radioactive isotopes
    Emissions expose photographic film

    • Scintillation counter
    • – Radioactive sample is placed in scintillation fluid

    – Energy released by isotope excites compounds in the fluid

    – Fluid flashes and the flashes are quantitated
  9. Chemical Bonds
    Chemical bonds

    – Interactions between valence electrons of atoms

    – Result in completing the valence shell

    • Examples

    • – Covalent- sharing valence electrons
    • – Ionic- transferring valence electrons
  10. Covalent bonds
    • • Nuclei share pairs of electrons
    • 1. Single Bond: atoms share 1 pair of electrons

    2. Double Bonds: atoms share 2 pairs of electrons

    3. Triple Bonds: atoms share 3 pairs of electrons

    • • No net charge results
    • • Valence shell of each atom completed
    • • Strongest bonds in living organisms
  11. Polar Covalent Bonds
    • • Electronegativity
    • – Sometimes one or more of the atoms shares itselectrons unequally
    • – Caused by the attraction of an atom for the electrons ofa covalent bond
    • – Results in one atom having the electrons in their spacemore often than the one(s) its sharing with

    • • Causes partial charges to form on a molecule
    • – Partial negative charge (delta-) occurs on the atom thatdraws the electrons
    • – Partial positive charge (delta+) occurs on the other•
  12. Polar covalent bonds in a water molecule
  13. Polar vs. Non-polar covalent bonds
    • • Non-polar covalent bonds
    •   – The atoms in the covalent bond equally share theelectrons
    •   – Equal electronegativity
    • • Polar covalent bonds
    • – The atoms in the covalent bond unequally share the electrons
    • – Unequal electronegativity
    • – A molecule that has polar covalent bonds is called a polar molecule

    Example: water is a polar molecule
  14. Ionic Bonds
    • • Transfer of electrons
    • – The difference in electronegativity of one atom results in the removal of electrons from the other atom
    • • Results in ions
    • – Cation is positive
    • – Anion is negative
    • • Ions are attracted to form ionic compounds
    • • Can form a crystal in the absence of water
    • – Responsible for strength of minerals like agate andmarble.
    • • Can ionize in water
    • -Relatively weak in water
  15. Hydrogen bonds
    • • Forms between molecules
    • • 2 electronegative atoms of different molecules are attracted to each other
    • – delta- is attracted to the delta+
    • •In living systems
    • – A delta+ hydrogen is attracted to a delta-
    • nitrogen or oxygen
  16. Van der Waals
    • Forms between molecules

    • Momentary positive and negative charges created by random placements of electrons

    – causes molecules to attract to each other

    • Occurs at very short distances between two atoms

    • Very weak chemical bond

    – But are additive and important when surfaces are brought very close together (remember the Geckos)
  17. Molecular shape
    • Weak interactions influence shape of large molecules in living systems

    • – Weak interactions in living systems
    • - hydrogen,ionic, and van der Waals

    • The functioning of a cell and communication between cells relies on“lock and key” interactions between molecules

    • – Ex. Receptor and ligands,
    • – Ex. Enzymes and substrates

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