Download CHAPTER 2: THE ATOMS AND MOLECULES OF ANCIENT EARTH

Chapter 2—The Atoms and Molecules of Ancient Earth
Lecture Outline
I.
Formation of Planets as Simulated by Computer Modeling
A. Dust and gases coalesce due to electrical and gravitational attractions.
B. Theory of Early Earth Conditions—Hot and violent
1. Immense heat due to accretion and volcanoes—Earth is molten.
2. Earth is bombarded by asteroids, one of which dislodges the moon.
3. Cooling causes condensation of H2O to form rain.
4. Volcanoes eject gases (CO2, N2, H2), forming the early atmosphere.
II.
Chemical Evolution Occurred Early in Earth's History
A. Radiometric dating technique determines age of rocks. (Fig. 2.1)
1. Atomic structure
a. Protons and neutrons in nucleus, electrons in surrounding orbitals.
b. Isotopes of an element—differ in the number of neutrons.
c. Radioactive isotopes:
(1) Unstable—decay by emitting either radiation or particles.
(2) Particle-emitting isotopes decay to daughter isotopes.
(3) Rate of decay for a specific isotope is constant and has a specific halflife.
2. Using radioactive isotopes to determine age of rocks. (Fig. 2.2)
a. Measure ratio of daughter isotopes to parent isotope in a rock sample.
b. Estimate the ratio that existed at the time the rock formed.
c. Calculate number of elapsed half-lives since the rock formed.
B. Estimating the Time when Chemical Evolution Occurred (Fig. 2.3)
1. Assume all components of solar system formed at same time, about 4.6 Ga
("giga ago").
2. Youngest known moon rocks are 3.8 Ga—correlates to slowing of asteroid
bombardment.
3. Oldest known cell fossils are found in rocks dated 3.5 Ga.
4. Conclusion—Chemical evolution occurred in the 300 million years between
cessation of asteroid bombardment (3.8 Ga) and age of oldest known cell fossils
(3.5 Ga).
III. Formation of Molecules from the Elements Available 3.8 Ga
A. Building Blocks of Chemical Evolution
1. 96% of every organism is composed of the elements C, H, O, N.
2. These elements were present in the forms of CO2, H2O, N2 and some CH4, H2,
and NH3.
B. Structure, Shape, and Bond Formation between Elements
1. Electrons occupy orbitals in energy shells around the nucleus.
a. Outermost orbital is the valence orbital. (Fig. 2.4b)
b. Electrons in outermost orbitals are farther from the protons of the nucleus
and will tend to move to an inner orbital if possible.
c. Atomic number = number of protons in nucleus. (Fig. 2.4a)
d. Mass number = number of protons plus neutrons in nucleus.
(1) Express in units of daltons, or atomic mass units (amu).
(2) 1 amu = 1.7 x 10–24 g, which is 1/12 the mass of a carbon atom that has
6 protons and 6 neutrons, hence the mass of one carbon atom = 12
amu. (Box 2.1)
(3) Mass of an electron is too small to affect the mass number.
2. Covalent Bonds and Molecular Structure
a. Occur between atoms that have unpaired electrons in valence orbitals.
b. A pair of electrons is shared between the valence orbitals, with each
electron attracted to the protons of each nucleus. (Fig. 2.5)
c. Single bonds share one pair of electrons, double bonds share two pairs, and
triple bonds share three pairs. (Fig. 2.6)
d. The orientation of the orbitals involved in electron sharing determines the
angle of the bond, which affects the overall shape of the molecule
(tetrahedral, pyramidal, flat planar). (Fig. 2.7)
e. Representing molecular structure: (Fig. 2.8)
(1) Molecular formula—shows types and numbers of atoms.
(2) Structural formula—shows linkages between atoms.
(3) Ball-and-stick model—3-D representation showing bond geometry.
(4) Space-filling model—most accurate 3-D spatial depiction.
f. Quantifying Molecules
(1) Mole = 6.022 x 1023 molecules (Avogadro's number)
(2) The mass of one mole of any molecule is the same as its molecular
weight in grams.
(3) Mol. wt. = sum of the mass numbers of the atoms in the molecule.
(4) Molarity = number of moles per liter of solution.
IV. Chemical Reactions between Atoms and Molecules
A. Reactants are converted to products; most reactions are reversible.
B. Chemical Equilibrium—Rate of forward reaction equals rate of reverse reaction:
1. Equilibrium is dynamic but also stable.
2. Equilibrium is disturbed by adding more reactant or product, or by altering the
temperature.
C. Energy = capacity to do work or supply heat.
1. Examples of different forms of energy:
a. Potential energy = stored energy
(1) Example—the potential energy of an electron in an outer shell.
(2) Example—the potential energy of H2O at the top of a waterfall.
b. Kinetic energy = the energy of motion
(1) Thermal energy—kinetic energy of molecular motion.
2.
(2) Temperature—a measure of thermal energy.
(3) Heat—thermal energy transferred from one object to another.
Energy is not created or destroyed; it just changes form.
a. Example—Potential energy of electron in outer shell becomes kinetic energy
when it moves to an inner shell. Potential energy of the electron is
decreased. (Fig. 2.9a,b)
b. Potential energy of H2O molecule at top of waterfall becomes kinetic energy
as it falls; mechanical, sound, and thermal energy when it strikes rocks at
bottom. (Fig. 2.9b)
D. Energy Transitions in Chemical Reactions
1. H = difference in potential energy between reactants and products
a. Exothermic reactions—Potential energy of products is less than that of
reactants.
(1) Heat is given off during the reaction.
(2) H = (–).
b. Endothermic reactions—Potential energy of products is greater than that of
reactants.
(1) Energy must be added to drive the reaction.
(2) H = (+).
2. S = difference in entropy (disorder) between reactants and products.
a. S is (+) if disorder increases—products less ordered than reactants.
b. S is (–) if products are more ordered than reactants.
3. Spontaneous chemical reactions
a. Sum total of energy transitions during the reaction must be toward
decreased potential energy and/or increased disorder.
b. G = H – TS; that is, the change in Gibbs free energy, or total energy, of
the reaction.
(1) Increase in temperature increases entropy
(a) Increase in temperature increases molecular motion.
(b) Molecules change state (solid to liquid to gas), resulting in rapid
jumps in entropy. (Fig. 2.10)
(2) Spontaneous reaction—G is (–).
(3) If G is (+), an input of energy is required for reaction to occur.
(4) If G = 0, the reaction is at equilibrium.
4. Chemical evolution—formation of products that have increased potential energy
and decreased entropy.
a. The equation G = H – TS will have a (+) value.
b. Energy must be added to drive the reaction.
V. Chemical Modeling of Chemical Evolution Reactions
A. Formaldehyde (H2CO) and hydrogen cyanide (HCN) are simple, carbon-containing
inorganic molecules that are key intermediates in forming larger organic molecules.
B. Can H2CO and/or HCN be produced by an input of energy to the types of molecules
available 3.8 Ga?
1. Pinto et al.—Hypothesis: CO2 + 2H2 + energy  H2CO + H2O
a.
b.
2.
Prepared computer model of all possible reactions.
Model used energy from photons of sunlight, which knock electrons from
valence orbitals and create free radicals. (Fig. 2.11)
c. Specified concentrations and temperature based on estimates of early Earth
conditions.
(1) Higher temperature increases number of collisions.
(2) Higher concentration of reactants increases collisions.
d. Specified reaction rates in the model based on measurements from
laboratory experiments under controlled conditions.
e. Model indicates appreciable quantities of formaldehyde are produced.
Zahnle develops a similar model showing HCN could also form.
VI. Importance of Redox Reactions to Chemical Evolution and Biology
A. Reduction-Oxidation Reactions
1. Gain of electron = reduction; loss of electron = oxidation.
a. Electrons can be transferred completely.
b. Electrons can just shift position to be closer to one atom than another. (Fig.
2.12)
2. Redox reactions are the most common chemical reactions in biology.
3. Reduction of carbon was a key step in chemical evolution.
a. Carbon is the most versatile molecule found in biological tissues.
(1) Each carbon atom can form four bonds with other molecules.
(2) Carbon atoms form the skeleton of organic molecules.
(a) Carbon atoms can be linked in many arrangements.
(b) A wide variety of molecular shapes is possible.
(c) Functional groups added to carbon skeleton impart a variety of
chemical reactivities to carbon molecules. (Table 2.1)
b. Reduction of CO2 by H2 forms H2CO, which is used as a building block to
form organic compounds (compounds containing at least one C–C bond).
(Fig. 2.13)
B. For carbon to be reduced, early atmosphere must have contained CH 4, H2, and NH3
(molecules that can give up electrons).
1. Volcanic ash is known to be rich in CO2, H2O, and N2.
2. Recent evidence indicates it also has small amounts of CH4, H2, and NH3.
C. Once organic molecules formed, only heat was needed to drive formation of more
complex organic molecules to complete chemical evolution. (Fig. 2.14)
1. Heat was widely available on early Earth as thermal energy.
2. Heat + potential chemical energy in bonds of organic compounds was sufficient
to drive formation of more complex organic compounds.
VII. Importance of Water to Chemical Evolution
A. Structure and Properties of Water
1. H2O is a polar molecule due to extreme electronegativity of oxygen.
a. Partial charges develop, resulting in slightly (+) and slightly (–) poles.
b.
2.
Hydrogen bonds link oppositely charged poles of H2O molecules in the liquid
state. (Fig. 2.15)
Unique Properties
a. Water is the most versatile solvent known.
(1) Any polar or charged molecule will dissolve in water.
(2) Substances in solution are more likely to collide than gases, which
increases the probability of chemical reactions.
b. Water expands as it forms a solid—Ice is less dense than liquid water.
(1) H2O molecules in ice form a regular crystal lattice more open than
irregular spacing of molecules in liquid. (Fig. 2.16)
(2) Flotation of ice likely prevented oceans from freezing solid before life
could form.
c. Water has a very high specific heat.
(1) Due to hydrogen bonding restricting molecular motion, a large increase
in temperature is required to raise the temperature of water by 1 oC.
(Table 2.2)
(2) Water insulates dissolved substances from large fluctuations in
temperature, thus facilitating the progress of chemical reactions.
B. pH and Acid-Base Reactions
1. Dissociation of water:
2H2O
H3O+ + OH–
2. Acids donate protons during a chemical reaction; bases take up protons.
3. pH is a measure of the concentration of hydrogen ions in solution.
a. Concentration of H+ in pure water at 25oC = 1 x 10–7M = pH 7.
b. pH scale based on antilogs of H+ concentration (pH = – log [H+]).
c. pH scale is logarithmic—A change of one unit constitutes a change in the
concentration of H+ by a factor of 10. (Fig. 2.17)
4. In solutions, the concentration of H+ is a major determinant of whether acid/base
reactions will occur. pH affects protonation state.
C. Conditions in the Oceans at the Time of Chemical Evolution
1. Earth cools; water condenses from atmosphere, forms rains.
2. Rainwater dissolves salts from rocks, forms oceans. (Table 2.3)
a. Salts are most abundant substances leached from rocks.
b. Salts are held by ionic bonds. (Box 2.2)
(1) Electrons are lost from valence shell of one atom and donated to the
valence shell of the other atom.
(2) The donor and recipient atoms each have filled valence shells.
(3) Donor atom carries a (+) charge (cation); recipient carries a (–) charge
(anion). (Fig. 2.18)
3. Chemical evolution may have begun in a salty ocean of pH 7.
VIII. Search for Extraterrestrial Life
A. Life probably requires water—Evidence of water recently found on Mars, Earth's
moon, and a moon of Jupiter.
B. Planets and moons that appear lifeless could have life below the surface—Bacteria
have recently been found living 1/2 mile below surface of Earth.
C. Life on Earth now known to exist in very extreme environments—hot springs, vents
on ocean floor, extremely salty lakes, and so forth.
D. Did chemical evolution occur only once in the universe, or many times?