The Chemical Context of Life

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The Chemical Context of Life

• Chapter 2

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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.

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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.

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Nutrient Deficiencies
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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.

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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.

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Examples

• 23Mg Mass number ?? Atomic number ??
  
• 12
• 23
  12
• of protons ?? of electrons ??
  of neutrons ??
• 12 12
  11
• 14C Mass number ?? Atomic
  number ??
• 6
• 14
  6
• of protons ?? of electrons ??
  of neutrons ??
• 6 6
  8

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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.

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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.

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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.

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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.

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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 (H2O), methane (CH4).

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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 (H2 and O2).

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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.

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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).

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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.